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THE   CHEMISTRY 

of  the 

COORDINATION 
COMPOUNDS 

1 
»  * 

Edited   by 

JOHN  C.  BAILAR,  JR. 

University  of  Illinois 
Urbana,  Illinois 


DARYLE  H.  BUSCH 

Editorial  Assistant 


American  Chemical  Society 
Monograph  Series 


REINHOLD       PUBLISHING       CORPORATlbN 

NEW  YORK 
CHAPMAN  &  HALL,  LTD.,  LONDON 

1956 


QD 

53 


Copyright  1956  by 


REINHOLD  PUBLISHING  CORPORATION 


All  rights  reserved 


Library  of  Congress  Catalog  Card  Number  56-6686 


REINHOLD  PUBLISHING  CORPORATION 

Publishers  of  Chemical  Engineering  Catalog,  Chemical  Materials 
Catalog,  "Automatic  Control,"  "Materials  &  Methods";  Ad- 
vertising   Management    of     the     American     Chemical     Society 


Printed  in  the  U.S.A.  by 
The  Waverly  Press,  Inc.,  Baltimore,  Md. 


Fred  Basolo 
B.  P.  Block 


Contributors 


X 


Robert  C.  Brasted 
Clayton  F.  Callis 
Leallyn   B.  ('LAPP 
William    E.  Coolly 
Bodie  E.  Douglas 

GUNTHER    L.    ElCHHORN 

Stanley  J.  (  Iill 
Roi    1).  Johnson 

Hans    H.   JoNASSEN 


Raymond  \\  Keller 
Stanley  Kirschner 
Ernest  H.   Lyons,  Jr. 
J.  A.  Mattern 
Niels  C.  Nielsen 
Thomas    1).   O'Brien 
Robert  \\\  Parry 
James  V.  Quagliano 
R.  L.  Rebertus 
Carl  L.  Rollinson 
Donald   II.   WlLKlNS 


Digitized  by  the  Internet  Archive 

in  2012  with  funding  from 

LYRASIS  Members  and  Sloan  Foundation 


http://archive.org/details/chemicoorOObail 


General   Introduction 

American  Chemical  Society's  Series  of  Chemical  Monographs 

By  arrangement  with  the  Interallied  Conference  of  Pure  and  Applied 
Chemistry,  which  met  in  London  and  Brussels  in  July,  1919,  the  American 
Chemical  Society  was  to  undertake  the  production  and  publication  of 
Scientific  and  Technologic  Monographs  on  chemical  subjects.  At  the  same 
time  it  was  agreed  that  the  National  Research  Council,  in  cooperation 
with  the  American  Chemical  Society  and  the  American  Physical  Society, 
should  undertake  the  production  and  publication  of  Critical  Tables  of 
Chemical  and  Physical  Constants.  The  American  Chemical  Society  and 
the  National  Research  Council  mutually  agreed  to  care  for  these  two  fields 
of  chemical  progress.  The  American  Chemical  Society  named  as  Trustees, 
to  make  the  necessary  arrangements  of  the  publication  of  the  Monographs, 
Charles  L.  Parsons,  secretary  of  the  Society,  Washington,  D.  C;  the  late 
John  E.  Teeple,  then  treasurer  of  the  Society,  New  York;  and  the  late  Pro- 
fessor Gellert  Alleman  of  Swarthmore  College.  The  trustees  arranged  for 
the  publication  of  the  ACS  Series  of  (a)  Scientific  and  (b)  Technological 
Monographs  by  the  Chemical  Catalog  Company,  Inc.  (Reinhold  Publish- 
ing Corporation,  successor)  of  New  York. 

The  Council  of  the  American  Chemical  Society,  acting  through  its  Com- 
mittee on  National  Policy,  appointed  editors  (the  present  list  of  whom 
appears  at  the  close  of  this  sketch)  to  select  authors  of  competent  authority 
ID  their  respective  fields  and  to  consider  critically  the  manuscripts  sub- 
mitted. 

The  first  Monograph  of  the  Series  appeared  in  1921.  After  twenty-three 
years  of  experience  certain  modifications  of  general  policy  were  indicated. 
In  the  beginning  there  still  remained  from  the  preceding  five  decades  a 
distinct  though  arbitrary  differentiation  between  so-called  "pure  science" 
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knowledge.  Furthermore,  most  workers  in  the  sciences  were  coming  to  see 
the  artificiality  of  the  separation.  The  methods  of  both  groups  of  workers 
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edge for  the  service  of  man.  The  officers  of  the  Society  therefore  combined 
the  two  editorial  Boards  in  a  single  Board  of  twelve  representative  members. 

Also  in  the  beginning  of  the  Series,  it  seemed  expedient  to  construe 


VI  GENERAL  INTRODUCTION 

rather  broadly  the  definition  of  a  Monograph.  Needs  of  workers  had  to  be 
recognized.  Consequently  among  the  first  hundred  Monographs  appeared 
works  in  the  form  of  treatises  covering  in  some  instances  rather  broad  areas. 
Because  such  necessary  works  do  not  now  want  for  publishers,  it  is  con- 
sidered advisable  to  hew  more  strictly  to  the  line  of  the  Monograph  char- 
acter, which  means  more  complete  and  critical  treatment  of  relatively 
restricted  areas,  and,  where  a  broader  field  needs  coverage,  to  subdivide  it 
into  logical  subareas.  The  prodigious  expansion  of  new  knowledge  makes 
such  a  change  desirable. 

These  Monographs  are  intended  to  serve  two  principal  purposes:  first, 
to  make  available  to  chemists  a  thorough  treatment  of  a  selected  area  in 
form  usable  by  persons  working  in  more  or  less  unrelated  fields  to  the  end 
that  they  may  correlate  their  own  work  with  a  larger  area  of  physical 
science  discipline;  second,  to  stimulate  further  research  in  the  specific  field 
treated.  To  implement  this  purpose  the  authors  of  Monographs  are  ex- 
pected to  give  extended  references  to  the  literature.  Where  the  literature 
is  of  such  volume  that  a  complete  bibliography  is  impracticable,  the 
authors  are  expected  to  append  a  list  of  references  critically  selected  on  the 
basis  of  their  relative  importance  and  significance. 

AMERICAN  CHEMICAL  SOCIETY 

BOARD    OF   EDITORS 

William  A.  Hamor,  Editor  of  Monographs 
Associates 

L.  W.  Bass  C.  H.  Mathewson 

T.  II.  Chilton  Laurence  L.  Quill 

Norman  Hackerman  W.  T.  Read 

J.  Bennett  Hill  Arthur  Roe 

C.  G.  King  Walter   A.   Schmidt 

S.  C.  Lind  E.  R.  Weidlein 


Prefa 


ce 


Werner's  coordination  theory  has  been  a  guiding  principle  in  inorganic 
chemistry  and  in  the  theory  of  valence  since  its  publication  sixty  years  ago. 
Indeed  :t  might  be  said  to  underlie  our  modern  concepts  of  molecular 
structure.  The  current  theories  of  acidity,  basicity,  amphoterism,  and 
hydrolysis  grew  directly  from  it,  and  the  assumption  of  the  complete  ioniza- 
tion of  solid  salts  is  implicit  in  it.  In  recent  years,  the  coordination  theory 
has  found  increasing  application  in  many  types  of  chemical  work.  For 
example,  its  usefulness  in  the  selection  of  organic  precipitants  for  metallic 
ions  and  in  explaining  biological  phenomena  are  well  known.  It  is  also  the 
basis  for  our  understanding  of  the  role  of  metal  ions  in  leather  tanning,  in 
the  dyeing  of  cloth,  and  in  regulating  plant  growth.  Coordinating  agents 
are  used  in  winning  metals  from  their  ores,  in  electroplating,  in  catalyzing 
reactions  and  in  obviating  the  effects  of  undesirable  catalyses,  in  precipi- 
tating metallic  ions  and  in  preventing  their  precipitation,  and  in  many 
other  ways.  Still  other  uses  await  study  and  exploration. 

So  much  interest  has  developed  in  the  theory  of  coordination  and  in 
coordination  compounds  in  recent  years  that  a  need  has  arisen  for  a  book 
describing  them.  I  began  the  preparation  of  such  a  book  several  years  ago, 
but  the  literature  on  the  coordination  compounds  is  so  vast,  and  is  growing 
so  rapidly,  that  it  soon  became  apparent  that  the  task  is  too  great  for  one 
person.  I  have  therefore  asked  some  of  my  students  and  former  students  to 
help  me  with  it.  I  am  grateful  to  them  for  their  help,  and  proud  to  present 
their  work. 

No  attempt  has  been  made  to  cover  the  chemistry  of  coordination  com- 
pound.^ completely  to  do  so  would  require  many  volumes.  Rather,  we 
have  attempted  to  select  ideas  which  are  fundamental  and  stimulating  and 
applications  which  are  both  illustrative  and  useful.  Even  so,  it  has  been 
necessary  to  omit  extensive  discussion  of  such  important  topics  as  the  use 
of  complex  ions  as  catalysts,  metal  ion  deactivators,  methods  of  preparing 
complex  ions,  and  the  details  of  many  physical  methods  which  are  used  in 
the  study  of  coordination  compound.-. 

In  the  interest  of  saving  space,  we  have  often  used  a  single  reference 
number  lor  several  related  articles.  When  one  of  these,  articles  is  referred  to 
later,  it  is  designated  by  the  original  number,  followed  by  a  letter  of  the 
alphabet  which  show-  its  position  in  the  list. 

Our  thanks  are  due  to  Prof.  \.  J.  Leonard.  Prof.  C.  S.  Vestling,  Prof. 

vii 


viii  PREFACE 

II.  A.  Laitinen  and  Dr.  Eleanora  C.  Gyarf as  who  have  read  portions  of  the 
manuscript,  and  have  made  valuable  suggestions  concerning  them. 

In  addition  to  serving  as  a  coauthor,  Dr.  Daryle  II.  Busch  has  assisted 
a  greal  deal  with  the  editorial  work,  and  I  wish  to  express  special  gratitude 
to  him.  Without  his  excellent  help,  it  is  doubtful  if  the  work  could  have  been 
completed. 

A  person  who  has  never  written  a  book  may  wonder  why  authors  so  fre- 
quently acknowledge  the  patience  and  understanding  of  their  wives.  These 
are,  indeed,  not  idle  words.  Many  of  the  hours  which  went  into  the  working; 
on  this  book  were  taken  from  evenings  which  would  otherwise  have  been 
spent  with  my  family  or  from  time  which  might  have  been  spent  in  doing 
the  many  odd  jobs  that  tall  to  the  lot  of  every  householder.  My  wife  has 
not  only  borne  this  with  patience  and  understanding,  but  has  lent  valuable 
advice  and  encouragement .  To  her  goes  my  most  grateful  acknowledgment. 

John  C.  Bailar,  Jr. 

I'rbana,  Illinois 
June,  1950 


Contents 

Preface 

1.  A  General  Survey  ok  the  Coordination  Compounds,  John 

( ' .  Bailor,  Jr.,  and  Daryle  II .  Busch 1 

2.  The    Early    Development   ok   the   Coordination    Theory, 

John  C  Bailor,  Jr 100 

.*;.  Modern  Developments    The  Electrostatic  Theory  of  Co- 
ordination Compounds,  Robert  IT.  Parr;/  (Did  Raymond  N. 

Keller 119 

4.  Modern  Developments    The  Electron  Pair  Bond  and  the 

Structure  ok  Coordination  Compounds,  Raymond  X .  Keller 

ami  Robi  it  II'.  Parry 1">7 

5.  Chelation  and  the  Theory  of  Heterocyclic  Ring  Forma- 

tion [nvolving  Metal  Ions,  Robert  W.  Parry 220 


(i.  Large  Rings,  Thomas  I).  O'Brien 


)o 


,.  General    [somerism    ok   Complex    Compounds,    Thomas    I). 

O'Brien 261 

8.  Stereoisomerism  ok  Hexacovalent  Atoms,  Fred  Basolo 274 

'.».  Stereochemistry    ok   Coordination    Number    Four,    B.    R. 

Block 354 

10.  Stereochemistry  and  Occurrence  ok  Compounds  Involving 

the    Less   Common    Coordination    Numbers,    Thomas   b. 

O'Brien :)82 

11.  Stabilization  ok  Valence  States  Through  Coordination, 

James  V .  Quagliano  and  R.  L.  R<l><  rtus  398 

12.  Theories   oi    Acids,    Bases,    Amphoteric    Hydroxides   and 

Basic  Salt-,  as  Applied  to  the  Chemistry  ok  Complex 

ifPOUNDS,  Fred  Basolo 4  Hi 

!:;.  Olation  and  Related  Chemical  Processes,  Carl  L.  Rollinson       ws 

14.  The  Poly  A<  ids,  Hans  B.  Jonasst  n  mat  Stanley  Kirschner  472 

15.  Coordination  Compounds  of  Metal  Ion-  with  Olefins  and 

Olefin-Like  Substances,  Bodu  E.  Hour/las  487 

ix 


\  CONTESTS 

.    16.  Metal  Carbonyls  and  Nitrosyls,  ./.  A.  Mattern  and  Stanley 

J.  Gill 509 

17.  Organic  Molecular  Compounds,  Leallyn  B.  Clapp 547 

-^  IX.  Physical    Methods   i\   Coordination  Chemistry,  Robert  C. 

Brasted  and  William  E.  Cooley 563 

19.  Coordination  Compounds  i.\  Electrodeposition,  Robert  W. 

Parry  <ni<l  Erru  Bt  //.  Lyons,  Jr 625 

20.  'I'm-:  Use  of  Coordination  Compounds  in  Analytical  Chem- 

istry, James  V .  Quagliano  and  Donald  H.  Wilkins 672 

21.  Coordination  Compounds  i\  Natural  Products,  Gunther  L. 

Eichhorn 698 

22.  Dyes  and  Pigments,  Roy  I).  Johnson  and  Niels  ('.  Nielsen  743 

23.  Water  Softening   Through   Complex    Formation,   Roy   I). 

Johnson  and  Clayton  F.  Cattis 7(>8 

Subject  Index 785 


1 


A  General  Survey  of  the  Coordination 
Compounds 

John  C.  Bailor  Jr.,  and  Daryle  H.  Busch* 

University  of  Illinois,  Urbano,  Illinois 

Since  coordination  compounds  differ  greatly  in  nature  and  stability, 
chemists  are  not  completely  agreed  on  a  simple  definition  of  the  term. 
Marly  workers  in  the  field  had  few  of  the  modern  physical-chemical  tools 
at  their  disposal,  and  if  a  material  satisfied  the  law  of  definite  proportions, 
they  were  inclined  to  consider  it  a  compound.  For  example,  crystals  of  the 
composition  (XHOaZnOo  are  readily  obtained  from  an  aqueous  solution 
containing  zinc  chloride  and  ammonium  chloride.  These  can  be  recrystal- 
lized  without  change  of  composition,  and  the  substance  was  long  considered 
to  be  a  complex  compound  in  which  zinc  shows  a  coordination  number  of 
five.  X-ray  analysis  has  shown,  however,  that  only  four  of  the  chlorine 
atoms  are  close  to  the  zinc  while  the  fifth  is  much  more  distant.  Similarly, 
the  clathrates  were  once  believed  to  be  coordination  compounds. 

According  to  the  theory  of  Sidgwick  and  Lowry,  a  coordinate  bond  (and 
hence,  a  coordination  compound)  can  be  formed  between  any  atom  or  ion 
which  can  accept  a  share  in  a  pair  of  electrons  (the  acceptor)  and  any  atom 
or  ion  which  can  furnish  a  pair  of  electrons  (the  donor).  The  donor  is  non- 
metallic  it  may  be  part  of  a  neutral  molecule,  like  CO,  IU>,  or  XTI::  .  or 
part  of  an  ion,  like  CI  ,  COjT  or  NH2CH2COO  .  Ordinarily,  an  acceptor 
requires  several  donors,  which  may  be  alike  or  different.  The  resulting 
complex  may  be  a  positive  ion,  a  negative  ion,  or  a  neutral  molecule. 

Even  if  we  accept  the  idea  that  a  coordinate  bond  consists  of  a  shared 
pair  (or  pairs)  of  electrons,  a  question  remain-  as  to  the  nature  and  the 
necessary  degree  of  such  sharing.  In  many  cases,  the  donor  and  acceptor 
are  bound  in  such  a  way  ih;it  the  reaction  of  formation  is  noi  reversible  t<» 
any  detectable  degree.  In  aqueous  solution,  the  hexamminecobalt(III) 
ion  [Co(NH3)e]+++,  shows  no  detectable  dissocial  ion1  and  the  analogous 
tri8(ethylenediamine)cobalt(III)  ion,  [Co  en.,]*  *  \  retains  its  optical  ac 
tivity  in  solution  for  many  week.-  at  ordinary  temperatures.  Both  of  these 

Now  :it  Ohio  State  University,  Columbus,  Ohio. 

1.  Flagg.  ./.  Am.  I  63.  057     L941 

1 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

ions  arc  stable  in  concentrated  hydrochloric  acid,  and  react  only  slowly 
with  hydrogen  sulfide  and  with  sodium  hydroxide. 

The  copper(II)  tetrammine  ion,  [Cu(NH3)4]++,  can  be  easily  detected 
in  solution  by  its  deep  blue  color,  and  its  salts  can  he  crystallized  from 
solution.  It  is  of  a  different  order  of  stability  from  the  cobalt  (III)  hexam- 
niine  ion,  however,  as  it  is  readily  destroyed  by  acids  or  by  heating.  In 
solution,  it  exists  in  equilibrium  with  [Cu(H20)4]++  and  ammonia.  The 
fact  that  the  formation  of  the  complex  is  accompanied  by  a  color  change, 
by  a  change  in  oxidation-reduction  potential  of  the  copper(II),  and  by 
other  changes  in  properties  clearly  indicates  that  there  is  a  true  chemical 
bond  between  the  copper  ion  and  the  ammonia  molecules. 

Sodium  chloride  absorbs  ammonia  when  under  pressure,  but  liberates  it 
when  the  pressure  is  released'-'.  No  doubt  there  are  attractive  or  adsorptive 
tones  which  tend  to  hold  the  two  substances  together,  but  they  are  weak 
and  poorly  characterized. 

In  general,  the  small,  highly  charged  cations  form  the  most  stable  co- 
ordinate bonds,  and  it  is  often  mistakenly  supposed  that  the  ability  to 
form  complexes  is  limited  to  the  transition  metals.  This  is  far  from  being 
so,  as  is  seen  from  the  fact  that  the  beryllium  derivative  of  acetylacetone 
can  be  distilled  without  decomposition  at  270°C. 

Even  the  alkali  metal  ions  form  complexes,  as  shown  by  the  work  of 
Sidgwick  and  Brewer3.  They  found  that  sodium  benzoyl  acetone  has  the 
properties  of  a  salt;  it  is  insoluble  in  nonpolar  solvents,  and  upon  heating 
it  chars  instead  of  melting.  If  recrystallized  from  95  per  cent  ethanol,  it 
takes  up  two  molecules  of  water  from  the  solvent,  yielding  a  dihydrate 
that  melts  at  1 15°C  and  is  appreciably  soluble  in  toluene.  It  is  evident  that 
the  dihydrate  is  a  chelated  compound. 

CH, 

C  — Ox  H20 

<  X 

c=o'      xh2o 
/ 

CH3 

Salicylaldehyde  (and  similar  compounds)  also  forms  sodium  chelates3' 4' h-  6. 
The  nature  of  the  electron  sharing  is  discussed  in  Chapters  3  and  4. 

2.  Clark,  .1///.  ./.  8ci.t  7,  1    (1924). 

l^wick  and  Brewer,  ./.  Chem.  Soc,  127,  2379  (1925);  Brewer,  J.  Chem.  Soc, 
1931,  361. 
\.  Hantssch,  />'-     .39,  3089  (1906). 

5  Weygand  and  Porkel,  ./.  prakt.  Chem.,  116,  293  (1927). 

6  Brady  and  Bodger,  ./.  Chem.  Soc.,  1932,952. 


GENERAL  SURVEY  3 

Suffice  it  to  Bay  here  that  stability  depends  upon  many  factors  and  cannot 
be  directly  correlated  with  bondPtype.  Among  the  many  other  factors  that 
are  important  in  determining  stability  are  charge  on  the  acceptor  atom, 
nature  of  the  donor  atom  and  of  the  molecule  of  which  it  is  a  part,  chela- 
tion, cationic,  anionic,  or  neutral  nature  of  the  complex,  and  the  nature  of 
the  ion  with  which  it  is  associated  (if  the  complex  is  an  ion). 

The  relationship  between  the  donor  and  acceptor  atoms  is  especially 
interesting.  Nearly  all  of  the  complexes  of  the  light  metals  (Periodic  groups 
IA,  IIA,  IIIB,  IVB)  contain  oxygen  as  the  donor  atom.  It  may  be  furnished 
in  the  form  of  water,  hydroxide  ion,  an  oxyanion,  an  alcohol,  ether,  ketone, 
or  in  a  variety  of  other  ways.  These  light  metals  seldom  coordinate  with 
molecules  containing  nitrogen,  sulfur,  carbon,  or  the  halogens.  Vanadium, 
at  the  head  of  group  VB,  is  a  powerful  oxygen  coordinator,  but  also  shows 
some  ability  to  form  ammines  and  complex  cyanides.  Proceeding  across  the 
periodic  table  toward  the  right  from  vanadium,  we  encounter  elements 
which  easily  coordinate  with  nitrogen.  Thus,  chromium  forms  a  large  num- 
ber of  ammines,  most  of  which  are  slowly  destroyed  in  water  solution.  The 
ammines  of  manganese  are  still  less  stable,  and  neither  iron(II)  or  iron(III) 
ion  reacts  with  ammonia  in  water  solution  to  give  ammines.  These  ions 
coordinate  instead  with  hydroxy!  ions  generated  in  the  water  by  the  addi- 
tion of  ammonia.  With  cobalt,  nickel,  copper,  and  zinc,  however,  stable 
ammines  are  formed.  The  ions  of  these  metals  retain  the  ability  to  co- 
ordinate with  oxygen  in  even  greater  degree  than  do  the  ions  of  the  lighter 
metals,  but  the  tendency  to  form  links  with  nitrogen  is  still  more  pro- 
nounced. Starting  with  vanadium,  too,  we  see  an  increasing  tendency  to 
coordinate  with  carbon — all  the  elements  from  vanadium  to  zinc  form 
stable  cyanides,  those  from  chromium  to  nickel  form  carbonyls,  and  copper, 
at  least,  forms  compounds  with  olefinic  substances.  The  ability  of  the 
metals  in  this  series  to  combine  with  sulfur  also  increases  toward  copper. 
Vanadium,  chromium  and  manganese  occur  in  nature  in  oxide  ores,  iron 
both  in  oxide  and  sulfide  ores,  and  cobalt,  nickel,  copper  and  zinc  largely 
as  sulfide  ores. 

In  the  fifth  and  sixth  series  of  the  periodic  table,  there  is  an  increased 
tendency  to  form  stable  complexes  with  halides.  This  is  present  in  the 
fourth  series  to  some  degree,  but  is  increasingly  important  in  the  later 
series,  as  is  illustrated  by  the  solubility  of  silver  chloride  in  hydrochloric 
acid  and  the  reaction  of  platinum  and  gold  with  chlorine  water  and  aqua 
regia  to  form  [PtCle]"  and  [AuCl4]~. 

The  elements  of  Periodic  groups  VIII,  IB,  and  IIP  are  of  special  interest.. 
All  of  them  form  complex  cyanides,  but  only  palladium,  silver,  platinum, 
rhodium,  and  mercury  are  known  to  form  compounds  with  the  ethylenic  dou- 
ble bond.  All  of  them  form  ammines  (the  ammines  of  mercury  readily  Lose 


4  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

protons,  hut  the  metal-nil rogen  bond  remains),  but  the  platinum  metals 
and  gold  form  few  complexes  containing  a  metal-oxygen  bond.  This  does 
not  mean  that  such  a  bond  is  not  stable,  but  only  that  the  metal-halide 
and  metal-sulfur  bonds  are  more  stable. 

The  metals  of  periodic  groups  II I  A,  EVA,  and  \  A  form  many  complexes 
in  which  the  donor  atom  is  oxygen,  sulfur,  or  a  halogen.  Compounds  in 
which  the  donor  is  carbon  or  nitrogen  are  much  less  common. 

The  Donob  Properties  of  the  Halogens 

The  halide  ions  often  coordinate  strongly,  and  halo-  complexes  are  well 
known;  fluorosilicates,  bromoplatinates,  and  iodomercurates  are  familiar. 
These  ions  are  often  thought  of  as  substituted  oxy-  anions,  but  this  has 
arisen  through  pedagogic  convenience  rather  than  strict  parallelism,  for 
while  a  halide  ion  occupies  one  coordination  position,  just  as  an  oxide  "ion" 
does,  its  elect rovalcnce  is  1  instead  of  2.  Thus  the  statement  that  Na2SiF6 
is  analogous  to  Na^iOg  is  somewhat  misleading,  for  in  solid  sodium  sili- 
cate, the  silicate  ions  are  linked  together  through  oxygen  atoms  in  such  a 
way  that  each  silicon  is  surrounded  by  four  oxygens,  while  in  the  fluoro- 
silicate,  each  silicon  is  surrounded  by  six  fluorines.  A  much  closer  analogy 
exists  between  the  halide  ions  and  the  hydroxyl  ion,  as  is  shown  by  the  series 
II,|PtCl6];  H2[PtCl5(OH)];  H2[PtCl4(OH)2];  H2[PtCl3(OH)8];  H2[PtCl2(OH)4]; 
H2[PtCl(OH)5];  H2[Pt(OH)6],  all  of  the  members  of  which  are  known  ex- 
cept the  fourth.  These  acids,  or  their  alkali  salts,  can  be  obtained  from  the 
chloro-platinate  by  stepwise  substitution  of  hydroxo-  groups  for  chloro- 
groups7  •  8  • 9  •  10. 

For  convenience,  the  complexes  formed  by  halide  ions  may  be  considered 
to  be  of  two  general  types;  those  containing  only  halide  ions  as  ligands  (with 
the  possible  exception  of  solvent  molecules)  and  those  containing  halide 
ions  as  a  less  abundant  donor  species,  as  is  the  case  among  the  halopen- 
tammines  of  cobalt(III)  and  chromium(III).  Although  the  stabilities  of 
complexes  is  generally  dependent  both  on  the  nature  of  the  central  metal 
ion  and  on  the  nature  of  the  donor  group,  these  complexes  may  be  grossly 
divided  into  two  major  stability  groups;  i.e.,  those  very  stable  complexes 
of  the  heavy  metals,  such  as  the  platinum  group  metals  and  mercury,  which 
give  only  a  faint  test  for  halide  ion  in  water  solution,  and  those  relatively 
labile  halide  complexes  of  the  type  formed  by  the  elements  of  the  first 
transition  group  and,  in  general,  the  more  electropositive  metals.  These 

7  Miolati  and  Bellucci,  /.  anorg,  Chem.,  26,  209  (1001). 

B  Miolati,  /.  anorg.  Chem.,  22,  145  (1900) 

9  Miolati,  Z.  anorg.  Chem.,  88,  261  (1903). 

in  Bellucci,  /.  anorg.  Chem.,  44,  168  (1906 


GENERAL  8URVE1  5 

two  major  stability  groups  correspond  to  the  penetration  and  normal  com- 
plexes discussed  in  Chapter  4. 

Many  of  the  reported  halide  complexes  of  metallic  elements  are  char- 
acterized solely  by  the  composition  of  solids  obtained  from  solutions  of 
mixed  halides.  The  weakness  of  this  type  of  evidence  as  a  criterion  for 

complex  formation  is  exemplified  by  the  fact  that  the  compound  written 
as  KjCuCl4*2H20  has  been  shown  by  x-ray  means  to  exist  as  copper(II) 

chloride  2-hydrate  admixed  with  potassium  chloride  in  the  crystal  lattice". 

Occurrence  and  Nature  of  the  "Strictly"  Halide  Complexes 

In  order  to  facilitate  an  understanding  of  the  extent  of  the  occurrence  of 
halo-  complexes,  and  to  illustrate  the  trends  occurring  among  the  families 
and  periods  of  the  periodic  system  of  elements,  a  brief  discussion  of  the 
halide  complexes  follows. 

Family  II A.  In  group  IIA,  only  tetrafiuoroberyllate  ion,  [BeF4]=,  is  well 
characterized.  Its  salts  bear  marked  resemblance  to  sulfates12.  This  is  not 
unexpected  since  tetrafiuoroberyllate  ion  is  isoelectronic  and  isosteric  with 
sulfate  and  also  approximately  the  same  size13.  Mitra14  reports  that  mono- 
hydroxytrifluoroberyllate  resembles  sulfate  even  more  closely,  citing  such 
evidence  as  the  isomorphism  of  the  salts.  The  corresponding  chloro-  complex 
is  much  less  stable,  evidence  for  its  existence  being  confined  to  freezing  point 
behavior  of  beryllium  chloride-alkali  chloride  mixtures15.  Double  fluorides 
of  magnesium  with  alkali  metal  ions  have  been  reported;  however,  their 
complexity  is  unlikely  since  the  crystal  structure  of  KMgF3  is  close-packed 
and  does  not  show  discrete  anionic  complexes16. 

Family  II B.  Complexes  with  all  four  halide  ions  are  reported  for  zinc 
and  cadmium.  In  the  solid  state,  the  complexes  seem  to  vary  from  [ZnX3]_ 
and  [CdX3]~  to  [ZnX5]-  and  [CdXfi]l~.  However,  it  seems  probable  that 
[ZnX4]=  represents  the  maximum  ratio  of  halide  to  zinc  in  true  combina- 
tion (see  page   1).  Studies  of  complex   halides   of  cadmium17,    zinc18,    and 

11.  Hendricks  and  Dickinson,  ./.  Am.  Chem.  Soc.,  49,  2149  (1927). 

12.  Kruss  and  Moroht,  Ann.,  260,  161  (1890);  Hay,  et  «/..  Z.  anorg.  Chem.,  201,  289 

(1931);  205,  257  (1932);  206,  209  (1936);  227,  32,  103  (1936);  241,  165  (1939). 

13.  Ray  and  Sarkar.  ./ .  I nd .  Chem.  Soc.  6,  987    1929);  Ghosh,  Mitra,  and  Ray:  ./. 

Ind.  Chem.  Soc.,  30,  221   .1953). 

14.  Mitra,  Science  and  Culture,  18,  393  (1963 

15.  Schmidt.  Ann.  ehim.,  [X]  11,  351  (1929);  O'Daniel  and  Tscheischwile,  Z.  Krist., 

104,  124  (1942). 

16.  Wells,  "Structural  Inorganic  Chemistry,"  p.  89,  London,  Oxford   University 

Press,  191s. 

^17.  Leden,  Z.  phys.  Chem.,  188,  100  (1941);  Ermolevka  and  Makkaveeva,  Zhur. 
Obschchei  Kkim.,  22,  1741  (1952  ;  Markman  and  Tur'yan;  Zhur.  Obschchei 
Khun.,  22,  1926  (1952);  Btrocchi,  Gazz.  ehim.  Hal..  80,  231    I960 


6  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

mercury(Il  lu  in  solution  support  the  possibility  that  the  most  characteris- 
tic species  arc  [MX*]  and  [MXJ".  The  order  of  stability  of  the  cadmium 
and  mercury  complexes  is  I  >  Br  >  CI  (There  is  some  doubt  that  fluoride1 
ion  form  complexes  with  these  two  metals  in  solution). 

Family  III  A.  The  halide  complexes  of  group  MA  illustrate  the  in- 
version in  relative  stability  of  the  [MXW]("~3)_  anions  upon  descending  the 
scries.  The  fhioro-  complexes  of  aluminum  are  by  far  the  best  characterized 
and  most  stable  of  all  the  haloaluminates.  The  anion  [Al  F6]=  is  remarkable 
in  a  number  of  ways.  It  represents  the  only  6-coordinate  haloaluminate, 
the  only  class  of  haloaluminates  which  may  be  prepared  in  water20,  the 
only  haloaluminates  occurring  in  nature,  and  it  is  apparently  the  monomelic 
parent  unit  of  a  family  of  condensed  fluoroaluminates  all  of  which  contain 
hexafluoroaluminate  units  in  their  solid  structures21.  However,  some  doubt 
remains  concerning  the  nature  of  the  complex  species  existing  in  solutions 
of  aluminum  ions  and  fluoride  ions22.  Chloride  and  bromide  form  complexes, 
M[AlXi],  with  the  corresponding  simple  aluminum(III)  halides  in  organic 
solvents20  or  from  melts  of  the  mixed  halides23.  The  tetrahedral  A1X4  unit 
also  exists  in  the  liquid  and  vapor  states  of  the  aluminum  (III)  halides,  which 
arc  dimeric24. 

The  halide  complexes  of  gallium(III)  are  relatively  rare,  the  best  known 
species  being  the  fluorides25,  [GaF6p  and  [GaF5(H20)]=.  There  is  little 
indication  that  the  remaining  halides  have  any  great  tendency  to  form 
complexes  with  gallium(III)  ions.  In  contrast  to  this  behavior,  and  to  the 
behavior  of  aluminum,  indium(III)  and  thallium(III)  form  well  charac- 
terized complexes  with  chloride  and  bromide  (and  iodide  in  the  case  of 

l'.i  Sherrill,  Z.phys.C hem.,  43,  705  (1903);  47, 103  (1904);  Garrett,/.  Am.  Chem. Soc., 
61,  2744  (1939);  Nayar,  Srivastava,  and  Nyar, ./.  Ind.  Chem.  Soc.,  29,  241,  248, 
250  (1952);  Kazi  and  Desai,  Current  Set.,  (India),  22,  15  (1953);  Ellendt  and 
Cruse,  Z.  physik.  Chem.,  201,  130  (1952). 

20    Malquori,  Atti  R.,  (GJ  5,  510  (1927);  [61  7,  745  (1928). 

21.  Thilo,  Naiurwiss.,  26,  529  (1938);  Brosset,  Z  anorg.  Chem.,  235,  139  (1937). 

22.  Bavchenko  and  Tananaev,  ./.  Gen.  Chem.,  U.S.S.R.,  21,  2505  (1951);  cf.  Chem. 

Mis.,  47,  5836o  (1953);  Tananaev  and  Nekhamkina,  Trudy  Komissii  Anal- 
Khim.,Akad.  Nauk.  S.S.S.R.,3,  89  (1951);  cf.  Chem.  Abstracts,  A7,58S5e  (1953)- 

23.  Kendall,  Crittenden,  and  Miller,  J.  Am.  Chem.Soc,  45, 969  (1923);  Plot  nikov  and 

Gorenbein,  ./ .  Gen.  Chem,  Rues.,  5,  1108  (1935). 

24.  Harris,  Wood,  and  Hitler../.  Am.  Chem.  Nor.,  73,  3151  (1961);  Gerding  and  Smit, 

/..  physik.  Chem.,  50B,  171  (1941);  Deville  and  Troast,  Compt.  rend.,  45,  821 
1857);  Palmer  and  Elliott,/.  Am.  Chem.  Soc.,  60,  1862  (1938);  Smits,  Meter- 
ing, and  Kamermans,  Proc.  Acad.  Sci.,  (Amsterdam),  34,  1327  (1931);  Smita 
and  Meijering,  Z  physik.  Chem..  41B,  98  (1938 
_'."»  Hannebahn  and  Klemm,  Z  anorg.  Chem.,  229,  341  (1936);  Pugh,  ./.  Chem.  Soc., 
1937,  1046,  1969 


GENERAL  SURVEY  7 

thallium)26  -7  2S.  They  apparently  form  no  fluoro-  complexes.  The  most 
typical  species  is  [MXJ",  although  the  enneachlorodithallate(III)  ion, 
[TljCW  has  been  studied  extensively29. 


From  such  observations  it  is  commonly  suggested  that  the  more  electro- 
positive cations;  i.e.,  A1+++  and  Ga+++,  tend  to  form  electrostatically  bound 
complexes  and,  in  consequence,  show  their  greatest  affinities  for  the  most 
electronegative  halogens.  On  the  other  hand,  the  relatively  less  electroposi- 
tive ions,  In+++  and  Tl+++,  show  a  much  greater  tendency  to  form  covalent 
bonds,  and  for  that  reason  are  most  susceptible  to  complexation  with  the 
larger,  more  easily  polarized  halide  ions. 

Family  IV A.  Similar  behavior  is  observed  among  the  elements  of  group 
IVA  (excluding  carbon).  Only  the  octahedral30  hexafluorosilicate  exists  in 
the  case  of  silicon,  while  germanium(IV)  forms  the  analogous  [GeFc]=  ion31 
and  the  relatively  unstable  hexachlorogermanate32.  The  complexes  [SnX6]= 
are  reported  for  all  four  of  the  halides33.  That  fewer  halogen  complexes  are 
formed  by  lead  (IV)  is  a  direct  result  of  the  strongly  oxidizing  nature  of  the 
ion. 

Family  VA.  Tripositive  arsenic  and  antimony  are  almost  unique  in  their 
ability  to  exist  either  as  the  central  atom  in  a  complex  species  or  as  the  do- 
nor atom  in  complexing  with  another  metal  ion  (a  property  which  is  probably 
shared  only  by  selenium  and  tellurium).  The  latter  role  will  be  discussed  at 

26.  Hoard  and  Goldstein,  ./.  ('hem.  Phys.,  3,  645  (1935). 

27.  Klug  and  Alexander,  J.  Am.  Chem.  Soc,  70,  3064  (1948). 

28.  Benoit,  Bull.  soc.  chim.,  France,  1949,  518. 

29.  Hoard  and  Goldstein,  ./.  Chem.  Phys.,  3,  199  (1935);  Powell  and  Wells, ./.  Chem. 

Soc,  1935,  1008. 

30.  Ketelaar,  Z.  Krist.,  92,  155  (1935);  Hoard  and  Vincent,  ./.  .1///.  Chem.  Soc,  62, 

3126  (1940). 

31.  Miiller, ./.  Am.  Chem.  Soc.,  43,  1087  (1921);  Wykoff  and  Muller,  Am.  J.  Sci.,  [5\ 

13,  346  (1927). 

32.  Laubengayer,  Billings,  and  Xewkirk,  ./.  .1///.  Chem.  Soc,  62,  546  (1946). 

33.  Skrabal  and  Gruber,  Monats.,  38,  1«.)  (1917);  Briggs,  /.  anorg.  Chem.,  82,  441 

L913);  Casey  and  Wyckoff ,  Z.  Krist.,  89,  469  (1934);  Dickinson,/.  Am.  Chem. 
Nor..  44,  276  (1922);  Ketelaar,  Rietdyk,  and  Stoverer,  Ree.  txav.  chin,.,  56,  '.hi; 

1937);  Goeteanu,£er.,60, 1312  (1927) ;  Seubert, fler., 20, 793  (1887);Brauner,  J. 
Chem.  Soc,  65,  393  (1894). 


s 


CHEMISTRY  OF  THE  COORDIN ATION  COMPOUNDS 


some  length  later  (page  78).  Species  of  the  types  [MX4]~  and  [MX6]=  have 
beenreported  (for  M  -  As,  X  =  CI  or  Br84;  for  M  =  Sb  or  Bi,  X  =  F  or 
(  1  ').  Bismuth  (II  I)  and  antimony  (III)  also  form  hexahalo-anions.  Recent 
x-ray  investigations  of  complex  antimony(III)  fluorides86  have  been  inter- 
preted  as  showing  that  the  pair  of  "s"  electrons  of  the  antimony  are  stereo- 
chemically  active.  Thus,  K2SbF5 ,  which  contains  discrete  SbF5  units,  is 
nol  strictly  5-coordinate  hut  is  octahedral 


F<^ 


Similarly,  the  ion  [Sb2F7]_,  in  its  cesium  salt,  is  probably  made  up  of  two 
trigonal  bipyramids  sharing  a  fluoride  ion  at  a  common  apex  and  with  one 
corner  of  each  equatorial  plane  occupied  by  an  electron  pair. 


Sb/ 
F 


T 


The  only  halide  complex  of  arsenic(V)  is  [AsF6]-37.  The  anions  [SbX6]~ 
have  been  reported  for  X  =  F,  CI,  or  Br.  The  bromide  complexes  differ 
from  the1  chloro-  and  fluoro-  species  in  being  highly  colored  and  readily 
bydrolyzed.  They  may  be  polybromides  of  antimony(III)37d.  Bismuth(V) 
does  nol  form  the  fluoro-  complex  corresponding  to  that  of  antimony,  but 
gives  [BiOF6]- instead88. 

First  Transition  Series.  By  far  the  most  interesting  halide  complexes 
occurring  among  the  metals  of  the  first    transition  series  are  the  fluoride 

34.  Petzold,  /.  anorg.  Chem.,  214,  355,  365  (1933);  Dehr,  ./.  Am.  Chem.  Soc,  48,  275 
L926). 
rutbier  and  Muller,  /.  anorg.  Chem.,  128,  137  (1023);  Ephriam  and  Masimann, 

54,  396  i  L923  I, 
Bystrom  and  Wilhelmi,  Arkiv  Kemi,  3,  373,  461  (1052);  Bystrom,  Nature,  167, 

0     I  '.i51). 
3chrewelius,  Z.  anorg.  Chem., 223,  1035  (1035);  Weinland  and  Feige,  ttrr.,36,244, 
L903     Petzold,  Z.  anorg.  Chem.,  215,  92  (1033). 
38    Ruff,  '/.  anorg.  Chem.,  57,  220  (1908  , 


GENERAL  SURVEY  9 

complexes.  Some  of  these  are  uniquely  stable  toward  hydrolysis  while  others 
may  support  unusually  high  oxidation  states  for  the  metal  ions.  The  rela- 
tive resistance  of  some  of  the  fluoro-  complexes  to  dissociation  or  hydrolysis 
in  aqueous  medium,  as  compared  to  the  remaining  halo-  complexes,  is  an 

indication  of  the  relative  affinities  of  the  transition  ions  for  these  donors. 
It  is  obvious  that  the  affinity  lor  fluoride  ion  in  these  cases  must  exceed 
that  for  the  oxygen  donor  species  of  the  solvent  water,  and  it  is  likely  that 
the  affinity  for  oxygen  donors  is  greater  than  that  for  chloride  or  bromide, 
although  our  picture  is  greatly  distorted  in  this  latter  case  by  the  omni- 
presence of  water  as  the  solvent.  The  extreme  difficulty  with  which  fluoride 
ion  is  oxidized  apparently  makes  the  existence  of  strongly  oxidizing  metal 
fluoride  complexes  possible;  however,  it  is  not  true  that  the  highest  known 
elect rovalences  of  a  given  metal  invariably  occur  in  fluoride  complexes. 
Figure  1.1  illustrates  this  point  by  comparing  oxy-  complexes  of  the  ele- 
ments of  the  first  transition  series  with  the  corresponding  fluoro-  complexes. 
The  general  character  of  the  fluoride  complexes  of  these  metals  may  be 
judged  from  the  fact  that  most  of  the  complexes  containing  higher  valence 
states,  such  as  heptafluorocobaltate(IV),  are  decomposed  by  water39.  Some 
of  the  complexes  of  the  more  common  oxidation  states  are  much  more 
stable. 

7- 

UJ  o 
OD 

25-| 

z> 
24 

I3 

*2 
Q 

X  H 

o 


—I — I — I — I — I — I — I — I — I 
Ti     v    cr  Mn  Fe  CO  Nl    Qj  zn 

Fig.  1.1.  Maximum  valencies  of  the  elements  of  the  first  transition  series. 
o    =  Maximum  valencies  found  in  oxy-  complexes. 
X  =  Maximum  valencies  found  in  fluro-  complexes. 

The  fluoro-  complexes40  of  iron  (III)  are  noteworthy  because  of  their  im- 
portance in  analytical  chemistry.  Iron  (III)  also  forms  relatively  stable 
complexes  with  chloride  ion  as  indicated  by  their  extractability  from 
aqueous  hydrochloric  acid  with  ether41. 

Cobalt (II)  forms  a  number  of  complex  fluorides  and  chlorides42.  Physico- 

39.  Klemm  and  Huss,  Z.  anorg.  allgem.  Chem.,  258,  221  (1949). 

40.  Remy  and  Busch,  Ber.,  66,  961  (1933). 

41.  Dodson,  Forney,  and  Swift, ./.  Am.  Chem.  S<><\,  58,  2573  (1936);  Lindquist,  Arkiv 

Kemi  Min.  Geol.,  24A,  No.  1  (1947). 

42.  Gmelin,  "Handbuch  der  Anorganisen  Chemie,"  Vol.  58A,  pp.  398-461,  Berlin, 

Verlag  Chimie  G.m.b.,  1932. 


10  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

chemical  studies43  on  solutions  of  cobalt(II)  halides  in  the  presence  of 
excess  halide  ion  indicate  the  existence  of  [CoX4]=,  the  stability  of  the  com- 
plexes decreasing  in  the  order  Cl~  >  Br-  >  I-.  Even  the  chloro  complex  is 
not  very  stable,  its  formation  being  detectable  spectrophotometrically  only 
in  hydrochloric  acid  which  is  at  least  2N.  A  fluoro-  complex  of  tetrapositive 
cobalt,  K3C0F7 ,  has  been  prepared39  by  fluorination  of  mixtures  of  potas- 
sium chloride  and  cobalt(II)  chloride.  It  is  fairly  stable  toward  reduction, 
l>n t  at  lf)0°  is  slowly  converted  by  hydrogen  to  potassium  hexafluoroco- 
baltate(III). 

The  halide  complexes  of  dipositive  nickel  are  poorly  characterized,  the 
fluoride  compounds  being  best  known.  When  treated  with  elemental  fluorine 
at  elevated  temperatures,  mixtures  of  potassium  chloride  and  nickel  chlo- 
ride yield  potassium  hexafluoronickelate(IV)39,  which  is  readily  hydrolyzed 
and  may  be  reduced  to  K2NiF4 . 

The  composition  of  K2MnF6  coupled  with  the  presence  of  manganese (IV) 
in  a  soluble  compound  justifies  the  assumption  that  the  substance  is  a  true 
complex39- 44, 45.  Manganese(III)  forms  fluoro-  and  chloro-  complexes  having 
five  halogen  atoms  and,  presumably,  one  water  molecule  attached  to  each 
manganese46. 

Complex  titanium  halides  of  the  form  [TiXfi]=,  where  X  =  F,  CI,  or  Br, 
have  been  characterized47- 48.  Of  these,  the  fluoro-  complex  is  the  most  stable. 

The  halo-  complexes  of  vanadium  are  best  characterized  for  the  triposi- 
tive  oxidation  state  of  the  metal  ion,  higher  valent  vanadium  tending  to 
form  oxy-  and  hydroxyhalo-  complexes.  The  hexafluorovanadates(III)  and 
pentafluoroaquovanadates  have  been  identified49,  as  have  complex  chlorides 
of  the  type  M2[VCl5(H20)]50.  Tripositive  chromium  also  forms  halo-  com- 
plexes of  the  type  M2[CrX5H20]44  and  hexanuorochromates(III)51. 

43.  Barvinok,  Zhur.fiz.  Khim.,  U.S.S.R.,  22,  1100  (1948);  Zhur.  Obshchei  Khim.,  19, 

612,  1028  (1949);  Varadi,  Acta  Univ.  Szeged.,  chim.  et  phys.,  2,  175  (1949);  3,  62 
(1950) . 

44.  Weinland  and  Laurenstein,  Z.  anorg.  allgem.  Chem.,  20,  40  (1899);  Jenssen  and 

Bardte,  Angew.  Chem.,  65,  304  (1953). 

45.  Bode  and  Wendt,  Z.  anorg.  Chem.,  269,  165  (1952);  Cox  and  Sharpe,  J.  Chem.  Soc. 

1953,  1783. 

46.  Weinland  and  Dinkelacker,  Z.  anorg.  Chem.,  60,  173  (1908). 

47.  Ruff  and  Ipsen,  Ber.,  36,  1777  (1903);  Rumpf,  Compt.  rend.,  202,  950  (1936); 

Rosenheim  and  Schutte,  Z.  anorg.  Chem.,  26,  239  (1901). 

48.  Cox  and  Sharpe, ./.  Chem.  Soc,  1953, 1783;  Wernet,  Z.  anorg.  allgem.  Chem.,  272, 

279  (1953). 
19.  Neumann,  Ann.,  244, 336  (1888);  Werner  and  Gubser,£er.,  34, 1579  (1901);  Chris- 

tensen,  •/.  prakt.  Chem.,  [2]  35,  161  (1887);  Schulter,  Compt.  rend.,  152,  1107, 

1261  (1911  I 
60    Stahler,  Ber.,  37.  nil  (1904  , 
51.  Fabris.  Oazz.  chim.  ital.,  20,  582  (1890);  Helmolt,  Z.  anorg.  Chem.,  3,  125  (1898). 


GENERAL  SURVEY  1  1 

Scandium  forms  several  complex  halides,  among  which  arc  the  fluoro- 
complexes  [ScF4]~,  [ScFJ",  and  [ScF8]".  There  is  Borne  evidence  thai  fche 
remaining  elements  of  periodic  family  1 1  IB  also  form  fluoro-  complexes, 
although  these  arc  noi  so  well  characterized  as  those  of  the  other  transition 
elements81.  The  complexity  of  KI.aF,  is  unlikely  since  the  crystal  structure 
indicates  the  presence  of  no  finite  [LaFJ    groups68. 

Although  COpper(I)  complexes  are  known''1  with  chloride,  bromide,  and 
iodide  ions,  no  fluoride  complexes  appear  to  exist.  A  great  variety  of  com- 
plex halides  has  been  reported  for  eopper(II).  The  complexity  of  some  of 
the  double  salts  formed  by  copper(II)  chloride  and  copper(II)  bromide 
with  alkali  halides  is  in  doubt  since  x-ray  data  show  that  K2CuCl4-2H20 
and  (NH^sCuBfi^HsO  exist  as  lattice  compounds  of  the  simple  salts. 
However,  physical  evidence  indicates  that  [CuCl§]~  and  [CuCl4]=  do  exist55. 
The  latter  is  reported  to  be  a  distorted  tetrahedron56.  Copper(III)  has  been 
reported  in  K3CuF639. 

The  relatively  greater  tendency  of  the  metallic  ions  of  the  first  transition 
series  to  form  complex  ions  with  fluoride  and  chloride  rather  than  with 
bromide  and  iodide  and  the  general  tendency  of  the  complexes  to  dissociate 
or  hydrolyze  in  solution  appears  to  justify  the  supposition  that  the  binding 
force  involved  is  essentially  electrostatic.  This  suggestion  is  supported  by 
the  considerable  stability  of  hexafluoroferrate(III)  and  hexafluorotitan- 
ate(IV)  which  involve  electronic  states  normally  associated  with  unusually 
stable  gaseous  ions  (Chapter  3). 

Second  and  Third  Transition  Series  and  Family  IB.  In  contrast  to  the 
elements  of  subgroups  IIIA,  IVA,  VA,  and  VIA,  the  elements  of  the  three 
transition  series  show  a  marked  increase  in  the  importance  of  their  higher 
oxidation  states  as  the  atomic  weight  of  the  metal  increases.  This  is  related 
to  the  types  of  compounds  formed  by  each  element,  since  high  oxidation 
states  ions  usually  exist  in  covalent  compounds.  The  halide  complexes  of 
the  platinum  metals  include  some  of  the  most  widely  known  complex  ions. 
This  is  doubtless  a  consequence  of  the  fact  that  their  simple  compounds  are 
for  the  most  part  "simple"  in  name  only  (for  example,  platinum(II)  chlo- 
ride is  not  salt-like  but  exists  as  bridged,  covalent,  giant  molecules). 

Complexes  of  the  type  [PtX6]=  have  been  characterized  for  all  four  of  the 

52.  Dergunov,  Doklady  Akad.  Navk,  S.S.S.R.,  85,  1025  (1952);  cf.  Chem.  Abs.  47, 

1524b  (1953). 

53.  Ref.  16,  p.  290. 

54.  Szabo  and  Szabo,  Z.  physik.  Chem.,  166,  288  (1933);  Fontana,  Gorin,  Kidder,  and 

Meredith,  Ind.  Eng.  Chem.,  44,  363  (1952);  Harris,  ./.  Proc.  Roy.  Soc.,  N.S. 
Wales,  85,  138  (1952). 

55.  Rossi  and  Strocchi,  Gazz.  chim.  Hal.,  78,  725  (1948).  (see  Ref.  72c) 

56.  Helmholz  and  Kruh.  ./.  .1///.  Chem.  Soc.,  74,  1176  (1952). 


L2  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

common  halides,  the  chloride  and  bromide  being  the  easiest  to  prepare57. 
The  iodo  complex  tends  to  liberate  iodine  with  the  reduction  of  the  plati- 
num to  the  dipositive  slate58,  while  salts  of  hexafluoroplatinate(IV)  readily 
hydrolyze.  The  complex  fluorides  have  been  prepared  by  heating  the  addi- 
tion product  of  the  ehloroplatinate  and  bromine  trifluoride59.  They  are 
diamagnetic,  indicating  drsp*  hybridization  and  covalent  bonding  (despite 
the  high  electronegativity  of  the  fluorine).  The  Pt — F  bond  distance  is 
greater  t  han  that  expected  for  a  covalent  link,  which  indicates  that  the  bond 
hasa  considerable  degree  of  ionic  character60.  Mixed  halo-  complexes,  such  as 
IPtrhBr^,  have  been  prepared61,  as  well  as  the  series  of  hydroxychloro- 
anions  [PtCl„(OH)6-nl=  (page  4).  The  planar  tetrahalide  complexes  of 
platinum(II)  have  been  prepared  with  chloride,  bromide,  and  iodide.  Salts 
of  these  anions  are  generally  obtained  by  reduction  of  the  corresponding 
hexahaloplatinate(IV)  salts  with  sulfur  dioxide62,  potassium  oxalate6213  ■ 
62c'  6i,  potassium  hydrogen  sulfite64,  hydrogen  sulfide65,  potassium  hypophos- 
phite66,  or  hydrazine  salts67.  Grinberg68  has  suggested  that  reduction  by 
hydrazine  salts  proceeds  in  two  steps: 

K2PtClfi  +  N.>H4-2HC1  ->  Pt°  +  N2  +  2KC1  +  6HC1 

2KC1  +  K>PtCl6  +  Pt°  ->  2K,PtCl4 

In  support  of  this  argument,  Grinberg  has  sho\ui  that  hexachloroplatinate 
ion  is  reduced  to  tetrachloroplatinate(II)  by  platinum  black  which  has  been 
freshly  prepared  by  the  reduction  of  hexachloroplatinate (IV)  with  hydra- 
zine sulfate.  Exchange  experiments  have  shown  that  halide  ions  of  plati- 
num(II)  complexes  are  labile,  the  bromide  of  [PtBr4]=  being  subject  to 
complete  exchange;  however,  the  central  platinum  atom  does  not  undergo 

57.  Weber,  J.  Am.  Chem.  Soc.,30,29  (1908);  Rudnick  and  Cooke,  J.  Am.  Chem.  Soc., 

39,  633  (1917);  Bielmann  and  Arduson,  Ber.,  36,  1365  (1903);  Gutbier  ami 
Bauriedel,  Ber.,  41,  4243  (1908). 

58.  Datta,  ./.  Chem.  Soc,  103,  426  (1913). 

59.  Sharpe,  ./.  Chem.  Soc,  1950,  3444;  1953,  197;  Schlesinger  and  Tapley,  ./.  Am. 

Chem.  Soc,  46,  276  i  L924  i. 

60.  Mellor,  Report   of  the  Brisbane  meeting  of  the  Australian  and  New  Zealand 

IlBSoc.  for  the  Advancement  of  Science,  Vol.  XXVIII,  131,  May  1951. 

61.  Klement,  /.  anorg.  Chem.,  164,  I!).")  d<)27). 

62.  Claua,  Ann.,  107,  137  (1868);  Klason, Ber., 37, 1360  (1904);  Vezea,Bull. soc. chim., 

|3]  19,  879  (1898). 
63    Mikhelis,   Zhur.    priklad.    Khim.,   26,   221    (1953);   cf.   Chem.  Abs.,  47,   11060i 

195 
64.  Lea,  .1///.  ./.  S«\,  [3]  48,  398,  loo  (1894). 
65    Bottger,  J.  prakt.  Chem.,  |1]  91,  251  (1863) 
1///.  ./.  Set.,  [3]  48.  :VM  (1894 
■  1'degershel  and  Shagesultanova,  Zhur.  priklad.  Khim.,  26,  222  (1953);  Cooley 
and  Busch,  unpublished  experiments  (1954). 
I  irinberg,  //////•.  priklad.  Khim.,  26,  224  (1953). 


GENERAL  SURVEY  L3 

exchange*.  The  rates  of  exchange  vary  in  the  order  CN  >  I  >  Br  >  ('1  . 
It  is,  at  first  thought,  paradoxical  thai  the  complexes  having  the  greater 
thermodynamic  stabilities  exchange  most  rapidly  (AF, ,„,,,.  :  [PtClJ  ,  —21.8; 
[PtBrJ™,  —24.5).  This  ease  of  "self -displacement"  may  be  a  peculiarity  of 
planar  complexes  since  ferrocyanide  ion  docs  not  exchange  with  cyanide 
ion  in  water70.  The  diammine  Pt  (^N II ; )vHr:  which  was  once  thought  to  con- 
tain tripositive,  5-coordinate  platinum  has  been  shown  rather  to  exist  as  a 
molecular  compound  of  lIV^XII^Br,!  and  (Pt  lv(\H:;>,Br,]71. 

In  contrast  to  platinum,  the  tet  rapositive  oxidation  state  of  palladium  is 
rather  unstable.  The  hexachloro-  and  hexabromopalladate(IV)  anion-  may 
he  prepared78  in  much  the  same  way  as  are  the  platinum  complexes;  how- 
ever, their  solutions  are  unstable  toward  evolution  of  the4  halogen  and  they 
both  react  with  aqueous  ammonia  to  liberate  nitrogen.  The  hexafluoro- 
palladate(IV)  has  recently  been  prepared  by  Sharpe73.  Its  salts  are  yellow; 
they  darken  rapidly  in  air  and  are  immediately  hydrolyzed  in  cold  water. 
Salts  of  the  planar  tetrahalopalladate(II),  [PdX4]=,are  known71  for  X  =  CI, 
Br.  and  I. 

The  great  affinity  of  palladium(II)  for  halide  ions  may  be  seen  from  the 
dissociation  constant75  of  [PdCl4]=  (Kd  =  b  X  10-14).  The  supposed  pal- 
ladiumflll)  complex,  MjPd111^76  probably  contains  both  palladium(II) 
and  palladium (IV), 

The  tendency  for  higher  oxidation  states  to  become  more  stable  with 
increasing  atomic  weight  of  the  metal  is  illustrated  by  cobalt,  rhodium,  and 
iridium.  The  only  strictly  halogen  complexes  in  which  cobalt  has  an  oxida- 
tion number  greater  than  two  are  the  fluoro-  complexes.  Dipositive  rho- 
dium, on  the  other  hand,  forms  no  complexes.  Rhodium  is  tripositive  in  all 
of  its  halogen  complexes  except  the  recently  reported  rhodium(IV)  fluoro- 

69.  Grinberg and Filinov,  Compt.  rend.  acad.  sci.,  U.R.8.S.,  23,  912  (1939);  cf.  Chem. 
Abe.  34,  12462  (1940);  31,  453  (1941);  cf.  Chem.  Abs.,  37,  5719  (1943) ;  Grinberg, 
Bull. acad. 8ci.,U.RS.S.,Ser.phy8., 349    1940);  cf.  Chem.  Abs.  35,  3895»  (1941  . 

7().  Grinberg  and  Nikol'skaya,  Zhur.  priklad.  Khim.,  24,  893  (1951);  cf.  Chem.  Aba. 
47,4709a    19.53). 

71.  Cohen  and  Davidson.  ./ .  .1///.  Chem.  Sue.  73,  1965    1951    :  Brossett,  Arkiv  Kemi 

Mir,.  Geol.,  25A,  No.  19    1948). 

72.  Puche,  (  •/..  200,  1206  (1935);  208,  656    1939  ;  Rosenheim  and  Maas,  / 

l  hem.,  18,  331  (1898);  Gutbier  and  Krell,  Ber.,  38,  2385    L905  , 
Sharpe,  ./.  (  hem.  Soc,  1953,  L97. 
74.  Gutbier,  Ber.,  38,  2107    1905);  Gutbier  and  Krell,  Ber.,  38,  3969    L90S  ;  Gutbier, 
Krell  and  Janssen,  /  anorg.  Chem.,  47,  23,  1292  (1906);  Gutbier  and  Woernle, 
47,271ti    L906  ;  Gutbier  and  Fe\\neT,Z.  anorg.  Chem.,  96, 129    1916  ;  Dickinson, 
J.  Am.  Chem.  Soc.,  44,2404    L922  ;  Cox  and  Preston,  J   Chem   Soc,  1988, 1089; 
Theilacker,  /   anorg.  Chem.,  234,  161     l" 
75    Templeton,  Watt,  and  Garner,  ./.  Am.  Chem.  Soc.,  65,  1608    L943  . 
Wohler  and  Martin,  /  anorg.  Chi  m.,  57,  398    L908 


14  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

complexes77.  Three  formulations  are  reported  for  the  halorhodiates(III), 
M2KhX5 ,  M3RhX6 ,  and  M2Rh2X9  .  All  three  types  are  known  for  bromide 
and  chloride78,  but  the  only  fluoro-  complex  is  the  ion  [RhF6]~.  The  struc- 
tures of  most  of  these  compounds  are  still  open  to  question. 

Both  tripositive  and  tetrapositive  iridium  form  complexes  with  chloride 
and  bromide.  The  iridium(III)  complexes  are  of  the  types  [IrX6]-  and 
[lrX5(ll20)]=79,  whereas  iridium (IV)  is  found  in  the  anion  [IrXe]=,  (X  = 
Br,  CI,  or  F).  The  hexabromo  compound  is  unstable  toward  evolution  of 
bromine80. 

Ruthenium(III)  and  ruthenium (IV)  form  a  variety  of  complex  halides 
and  aquohalo-  or  hydroxohalo-  complexes.  Ruthenium  trichloride  appar- 
ently exists  in  several  hydrated  forms,  analogous  to  the  hydrate  isomers  of 
chromium  (III)  (see  Chapter  7)81.  Some  of  the  probable  "hydrate  isomers" 
are  Ru(H20)Cl3 ,  which  contains  no  ionizable  chloride,  and  the  reported 
cis  and  trans  forms  of  [RuCl2(H20)2]Cl.  Dwyer  and  Backhouse81  suggest 
that  the  ruthenium  is  6-coordinate  in  all  of  these  complexes.  As  compared 
to  the  similar  platinum  compounds,  halide  complexes  of  ruthenium  show  a 
marked  tendency  to  hydrolyze  and  to  retain  water  in  their  coordination 
spheres.  As  with  the  platinum  analogues,  [RuBr6]=  is  less  easily  hydrolyzed 
than  [RuCl6]=81.  Ruthenium(III)  forms  two  types  of  anionic  chloro-  com- 
plexes [RuCl6]=  and  [RuCl5(H20)]=,  while  ruthenium(IV)  forms  the  com- 
plexes formulated  as  [RuCl6]=  and  [RuCl5(OH)]-82- 83-  84.  It  has  been  shown 
that  [RuCl5(OH)]=  is  actually  dimeric  in  the  crystalline  state,  having  the 
structure  [Cl5Ru — 0 — RuCl5]4~  (see  p.  167).  Fluorination  of  hexachloro- 
ruthenate(IV)  yields  a  white  crystalline  compound  of  the  composition 
K2RuF8 ,  which  hydrolyzes  readily  and  darkens  on  standing85.  It  is  possible 
that  ruthenium (VI)  is  present,  and  that  it  is  octacoordinate. 

77.  WeiseandKlemm,Z.  anorg.  allgem.  Chon.,  272,  211  (1953);  Sharpe,/.  Chem.  Soc, 

1950,  3444. 

78.  Delepine,  Bull.  soc.  chim.,  Belg.,  36,  108  (1927);  Gut  bier  and  Bertsch,  Z.  anorg. 

Chem.,  129,  67  (1923);  Meyer  and  Hoehne,  Z.  anorg.  Chem.,  231,  372  (1937); 

Meyer,  Kawkzyk,  and  Hoehne,  232,  410;  Poulenc,  Compt.  rend.,  190,  639  (1930) ; 

Ann.  chim.,  [Xi]  4,  567  (1935). 
79    Delepine,  Bull.  soc.  chim.,  [4]  3,  901  (1908);Delepine-Tard,  Ann.  chim.  phys.,  [10] 

4,  2S2  (1935 
B0.   Delepine,  .1////.  chim.  phys.,  [9]  7,  277  (1917);  Schlesinger  and  Topley,  J.  Am. 

Chem.  Soc,  46,  276  (1924);  Dobroborskaya,  Zhur.  priklad.  Khim.,  26,  223 

(1953);  cf.  Chem.  Abe.,  47,  U061g  (1953). 
81.  Dwyer  and  Backhouse,  J    Proe.  Roy.  Soc,  X.S.  Wales,  83,  138  (1949). 
$2.  Gutbierand  Niemann, Z.  anorg.  ('hem .,  141,  312  (4924) ;  Howe, ./.  Am.  Chem.  Soc, 

49,  2389  (1927);  Charonnat,  .1////.  chim.,  [10]  16,  72  (1931);  Compt.  rend.,  181, 
;    L925 
Howe,  ./.  .1///.  Chem.  Soc.,  26,  942    L904 
84.  Charonnat,  Compt.  rend.,  180,  1271  (1925). 
85     \vnsley:  Peacock,  and  Robinson,  Chem.  hid.,  1952,  1002. 


GENERAL  SURYIA  L5 

The  hexahalo-  salts  M  I  >-.Y  and  MiOsXa  are  reported  where  X  =  CI  or 
Br  in  the  first  case86  and  for  X  =  F,  CI,  Br,  or  I,  in  the  latter"7.  Recrystal- 
lization  of  the  hexachloro-  and  hexabromoosmiate(IV)  salts  from  dilute 
halogen   acid   leads   to   hydrolysis.    Mixed   halogen   complexes,   such   as 

[OsClsBr]"  and  [OsCl;J3r3]=,  and  hydroxohalo-  complexes,  such  as 
[OsX.s((  )H  )] '",  are  also  reported88.  Osmium  also  forms  halo-  complexes  in 
its  higher  oxidation  states.  Osmium(YI)  exists  in  the  tet rahaloosmyl  com- 
plexes [( )s(  )-j\;!~  Vl,  and  the  oxydihaloosmyl  complexes  (( )s(  ):iX2]=90.  X-ray 
data  show  that  the  salts  Mo[Os02X4]  are  similar  in  crystal  structure  to  po- 
tassium hexachloroplatinate(IV)91.  Fluoride  ion  combines  with  osmium 
(Mil)  fluoride  to  produce  a  white  solid  that  may  be  a  9-  or  10-coordinate 
complex9'2;  the  material  has  not  been  analyzed.  Dissolution  of  osmium(VIII) 
oxide  in  fluoride  solution  leads  to  the  formation  of  unstable  compounds 
which  presumably  contain  complex  anions,  such  as  [Os04F2]=  93. 

The  halo-  complexes  of  rhenium  are  intermediate  in  character  between 
those  of  the  platinum  metals  and  those  of  the  remaining  transition  ele- 
ments. Thus,  rhenium (IV)  forms  complexes  of  the  type  [ReX6]=  with  fluo- 
ride (like  the  IVB,  VB,  and  VIB  metals)  and  also  with  the  other  halogens, 
even  iodide  (a  behavior  more  to  be  expected  of  the  platinum  metals)94. 
An  interesting  similarity  is  found  between  some  rhenium  (IV)  and 
rhenium(V)  chloro-  complexes  and  those  of  ruthenium(III)  and  ruthen- 
ium (IV).  In  addition  to  hexachlororhenate(IV),  the  pentachlororhenium 
complexes  [ReIVCl5(OH)]=    [RevCl50]=   and  [ReIV2Cli0O]4-  also  exist94f. 

Molybdenum  and  tungsten  form  complex  halides  or  oxyhalides  in  their 
di-,  tri-,  penta-,  and  hexavalent  states.  Tripositive  molybdenum  forms 
fluoro-  and  chloro- complexes  of  the  types  [MoX5(H20)]=  and  [MoX6]s95. 

86.  Claus  and  Jacob}', J .  prakt.  Chem.,  90,  78  (1863) ;  Crowell,  Brenton,  and  Evenson, 

J.  Am.  Chem.  Soc.,  60,  1105  (1938). 

87.  Ruff  and  Tscherch,  Ber.,  46,  932  (1913);  Dwyer  and  Gibson,  Nature,  165,  1012 

(1950;;  Wintrebert,  Ann.  ekim.  phys.,  [7]  28,  133  (1903). 

88.  Krauss  and  Wilkin,  Z.  anorg.  Chem.,  137,  360  (1924). 

89.  Wintrebert,  Ann.  chim.  phys.,  [7]  28,  54,  86  (1903). 

90.  Wintrebert,  Ann.  chim.  phys.,  [7]  28,  114  (1901). 

91.  Hoard  and  Grenko,  Z.  Krist.,  87,  100  (1934). 

92.  Ruff  and  Tscherch,  Ber.,  46,  929  (1913). 

93.  Tschugaev,  Compt.  rend..  167,  162  (1918);  Krauss  and  Wilkin,  Z.  anorg.  Chem., 

145,  151  (1925). 
(.)4.  Ruff  and  Kwasnik,Z.  anorg.  CAem.,219,76  (1934);  Schniid.  Z.  anorg.  CAem.,212, 

187  (1933);H6lemann,Z.  anorg.  Chem.,  211,  195  (1933);  Nod. lack  and  Nod. lack. 

Z.  anorg.  Chem. ,216,  120  (1933) ; Briscoe, Roderson,  and  Etudge,  J.  Chei 

1931,  3218;  Jezowska-Trzebiatowska,  Trav.  soc  sd.  et  lettres  Wroclaw,  Ber.  B, 

39,  5  (1953). 
95.  Rosenheim  and  Braun,  Z.  anorg.  ('htm.,  46,  ^>2<)  (1905);  Foerster  and  Fricke,  Z. 

angew.  Chem.,  36,  458  (1923). 


L6  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Eowever,  only  the  dimeric  anion  [W2C19]-  is  known  for  tungsten(III)96  (for 
structure',  see  page  7).  It  seems  likely  that  a  tungsten-tungsten  bond  is 
pi  (Hi it  in  this  anion  since  the  substance  is  diamagnetic97.  The  most  stable 
oxyhalo-  complexes  of  molybdenum  and  tungsten  in  their  penta-  and  hexa- 
positive  states  are  fluoro-  complexes,  such  as  [MoVI02F4]=,  [WVI02F4]=,  and 
[MovOFb]-,  all  of  which  are  isomorphous  with  [NbvOF5]=.  The  affinity  of 
fluoride  ion  for  hexavalent  molybdenum  and  tungsten  may  be  illustrated 
by  the  fact  that  most  of  the  precipitation  and  color  reactions  of  molybdate 
and  tungstate  ions  are  masked  by  the  presence  of  fluoride  ion98. 

An  interesting  feature  of  the  halogen  complexes  of  niobium  and  tantalum 
is  the  occurrence  of  high  coordination  numbers  (see  Chapter  10).  This  is 
undoubtedly  associated  with  the  fact  that  the  only  significant  strictly  halo- 
geD  complexes  of  these  metals  are  those  of  the  fluoride  ion.  Both  of  these 
elements  form  heptafluoro-  anions  of  the  type  [MVF?]=.  Their  structures  are 
discussed  on  page  393.  In  addition,  tantalum (V)  forms  an  8-coordinate 
fluoro-  complex  [TaF8]=  which  exists  in  the  form  of  a  tetragonal  antiprism". 
Six-coordinate  hexafluoroniobate(V)  is  also  known,  as  is  its  tantalum  ana- 
log100. The  heptafluorotantalate(V)  is  somewhat  more  stable  than  the  nio- 
bium^) compound  which  hydrolyzes  to  [NbOF5]=,  and  this  difference  has 
served  in  helping  to  separate  the  two  metals.  Oxyhalo-  complexes  are 
formed  by  both  metals,  the  oxyfluorides  being  the  most  stable. 

The  same  trends  are  observable  among  the  halogen  complexes  of  zir- 
conium and  hafnium,  the  outstanding  characteristics  being  variable  coor- 
dination number  and  decreasing  stability  of  the  complexes  with  increasing 
atomic  weight  of  the  halide.  The  latter  point  is  illustrated  by  the  fact  that 
zirconium  dioxide  is  dissolved  by  hydrofluoric  acid  and  that  only  the  fluoro- 
complexes  are  stable  in  aqueous  media101.  The  chloro-  and  bromo-  complexes 
are  prepared  in  alcohol102.  The  complexes  are  of  the  types  [MX5]~, 
[MX5(H20)]=,  [MX6]=,  [MXJS  (see  Chapter  10).  The  structure  of  the  sup- 
posed 5-coordinate  species  is  still  open  to  question103.  The  fluoro-  complexes 
are  used  in  the  separation  of  hafnium  and  zirconium104. 

96.  Olsson,  Ber.,  46,  566  (1913);  Olsson,  Collenberg,  and  Sandved,  Z.  anorg.  chem., 

130,  16  (1923). 

97.  Brossett,  Nature,  135,  824  (1935);  Pauling,  Chem.  Eng.  News,  1947,  2970. 
its.  Feigl,  .1/.//.  Chem.  Aria,  2,  397  (1948). 

99    Hoard,  ./.  -1///.  Chem.  Sac,  61,  1252  (1939);  64,  633  (1942);  dc  Marigroc,  Compt. 

rend.,  63,  85  (1866);  Board,  Paper  presented  at  6th  annual  symposium,  Div. 

Phys.  and  [norg.  Chem.,  Columbus,  Ohio,  Dec.,  1941. 
KM).  Halm  and  Putter,  '/. .  anorg.  Chem.,  120,  71  (1922). 
nil.  Connick  and  McVey,  •/.  .1/,/.  Chem.  Soc,  71,  3182  (1949). 

Schwarz  and  Giese,  /.  anorg.  Chem.,  176,209  (1928);  Rosenheim  and  Frank,  Ber., 

38,  812    L905  , 
Haendler  and  Robinson,/.  .1///.  Chem.  Soc. ,75,  3846  (1953);  Haendler,  Wheeler, 

and  Robinson,  ./.  Am.  Chem.  Soc,  74,  2352  (1952). 
lot.  Larsen,  Fernelius,  and  Quill,  //,</.  Eng.  Chem.,  Anal.  Ed.,  15,  512  (1943);  Schultz 


GENERAL  SURVE1  17 

The  solubilities  of  the  silver  halides  increase  sharply  as  the  concentration 
of  excess  halide  ion  is  increased106.  The  study  of  this  solubility  dependence 
indicates  the  format  ion  of  a  scries  of  complexes  ranging  from  [AgjX]H  !  to 
[AgXJ  ,  and  possibly  [Ag»Xe]4  m.  Theorderof  stability  of  both  silver  and 
gold  halide  complexes  is  I  >  Br  >  CI  (as  is  also  commonly  observed  among 
the  platinum  metals).  The  silver  complexes  best  known  in  the  solid  .state 
are  of  the  types  [AgXJ  and  [AgXg]  ln7.  I'nipositive  gold  normally  forms 
2-coordinate,  linear  complexes  of  the  type  [A11CI2]  "'s,  while  gold(III)  forms 
L-coordinate,  planar  complexes  of  the  type  |AnX,|  '"''.  Gold  forms  many 
bridged  halogen  compounds  (page  19).  The  substance  having  the  em- 
pirical formula  CsAuCla  should  be  formulated  as  Cs2AuIAuIIICl6 ,  contain- 
ing equivalent  amounts  of  gold(I)  and  gold(III)  (see  Chapter  9). 

Complexes  Containing  Halide  Groups  as  a  Less  Abundant  Donor 
Species 

Many  metals,  especially  those  of  the  platinum  group,  form  halo-  com- 
plexes containing  three1,  four,  or  five  halide  groups;  however,  with  the  ex- 
ception of  the  hexafluorocobaltate(III),  cobalt(III)  complexes  are  not 
known  with  more  than  three  halide  groups.  Indeed,  the  mixed  complexes 
which  have  been  most  significant  in  the  development  of  the  coordination 
theory  are  those  which  contain  one,  two,  or  three  coordinated  halides  and 
five,  four  or  three  neutral  groups.  Chloropentamminecobalt(III)  chloride, 
[Co(XH3)5Cl]Cl2 ,  is  one  of  the  longest  known  cobalt  (III)  ammines  and  is 
the  chief  product  obtained  by  atmospheric  oxidation  of  solutions  containing 
cobalt (II)  chloride,  ammonium  chloride,  and  ammonium  hydroxide.  The 
coordinated  chloride  is  only  slowly  removed  by  the  action  of  silver  nitrate, 
even  when  heated.  The  salt  serves,  however,  as  a  starting  material  for  the 
preparation  of  many  other  cobalt  (III)  ammines,  not  only  by  replacement  of 
the  chloride,  but  also  by  replacement  of  one  of  the  ammonia  molecules. 
Heating  with  ammonium  carbonate,  for  example,  gives  carbonatotetram- 
minecobalt(III)  chloride,  [Co^Hs^CCyCl.  It  has  also  been  utilized    in 

and  Larsen,  J.  Am.  Chem.  Soc,  72,  3610  (1950);  Huffman  and  Lilly,  •/      t 
m.  Soc,  73,  2902  (1951). 

105.  Eber  and  Schuhly,  J.  prakt.  Chem.,  158,  176  (1941);  Z.  anorg.  allgem.  ('Ik  ///.,  248, 

32  (1941). 

106.  Bern  and  Leden,  Svensk.  Kern.  Tidskr.,  65,  88  (1953);  Z.  Naturforsch.,  89,  719 

(1953);  Chateau  and  Pounadiev,  Science  et  indus.  phot.,  23,  225  (1952);  Y..t 
simirskii,  Doklady  Akad.  Nauk.,  S.S.S.R.,  77,  819  (1951);  cf.  Chem.  Abs.,  45, 
7102  (1951). 

107.  Forbes  and  Cole,  /.  Am.  Chem.  Soc.,  48,2492    L921  \;  Harris  and  Schafer,/.  P 

8oe.,  X.s.  Wales,  85,  148  (1952);  Harris,  •/.  Proc.  Roy.  Soc.,  N.8.  Wales 
85,  142  (1952). 

108.  Lengfield,  Am.  Chem.  J.,  26,  324    L901). 

109.  Cox  and  Webster,  J.  Chem.  Soc,  1936,  1635. 


18  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

studies  directed  a1  elucidation  of  the  mechanism  of  substitution  reactions 
of  6-coordinate  complexes110. 

The  two  forms  of  dichlorobis(ethylenediamine)cobalt(III)  chloride, 
[Co  ei^CyCl,  are  used  in  the  preparation  of  other  ethylenediamine  cobalt 
Baits.  .Both  the  cis  and  trans  forms  of  this  complex  are  readily  prepared, 
and  are  stable  in  water  solution  for  some  time,  though  the  change  in  color 
of  the  solution  indicates  aquation;  the  w's-dibromobis(ethylenediamine)co- 
balt(III)  ion  rearranges  with  extreme  ease  to  the  trans  form,  and  both  iso- 
mers aquate  rapidly;  the  corresponding  iodo  compounds  are  not  known. 

Chloropentamminechromium(III)  chloride,  [Cr(NH3)5Cl]Cl2 ,  is  ob- 
tained, together  with  the  hexammine,  by  the  action  of  liquid  ammonia  on 
anhydrous  chromium(III)  chloride.  Once  formed,  the  pentammine  is  con- 
verted to  the  hexammine  with  extreme  slowness,  which  may  be  due,  how- 
ever, to  the  very  slight  solubility  of  the  pentammine  in  liquid  ammonia. 

ns-Diehlorobis  (ethylenediamine)  chromium  (III)  chloride  is  most  easily 
obtained  by  the  thermal  decomposition  of  tris (ethylenediamine) chrom- 
ium (III)  chloride.  The  reverse  reaction  takes  place  very  slowly  when  the 
dichloro-  salt  is  suspended  in  ethylenediamine.  Complexes  of  very  similar 
type  are  also  encountered  in  the  chemistries  of  the  platinum  metals. 

In  general,  the  complexes  containing  halo-  groups  as  less  abundant  donor 
species  may  be  grouped  according  to  the  same  classification  as  that  given 
for  the  strictly  halide  complexes;  i.e.,  those  which  show  little  tendency  to 
dissociate  in  solution  (penetration  complexes),  and  those  which  change 
upon  dissolution  in  a  polar  solvent  as  a  result  of  displacement  by  solvent 
molecules  (normal  complexes).  Only  the  first  class  of  compounds  is  of  great 
significance  here  since  the  more  labile  species  cannot  experience  a  change  in 
the  state  of  aggregation  without  extensive  change  in  their  natures.  Thus, 
[Fe(XH3)2Cl2]  cannot  be  dissolved  in  water  and  subsequently  recovered, 
while  many  strictly  halide  complexes  may  dissociate  in  solution  but  still  be 
recoverable  in  the  original  form  upon  removal  of  the  solvent. 

(  omplexes  Involving  Halogen  Bridges 

The  halide  ions  sometimes  donate  pairs  of  electrons  to  two  metallic  ions 
simultaneously,  forming  a  "bridge."  Aluminum  chloride  (page  6)  and 
rhenium(III)  chloride111  have  been  shown  to  have  the  structures 

CI  CI  CI 

\  /  \  / 

M  M 

/    \     /    \ 
CI  CI  CI 

1  lii    Br0nsted,  '/..  phyaik.  Chem.,  102,  169  (1922);  Garrick,  Trans.  Faraday  Soc,  33, 
L937);  Lamb  and  Fairball,  ./.  Am.  Chctn.  Soc,  45,  378  (1923);  Lamb  and 
Maiden,./.  .1///.  Chem.  Soc.  ,33,  1873  (1911) ;  Adell,  Z.  anorg.  allgem.  Chem.,  249, 
251  (1942). 


GENER  I/.  SURVE] 


L9 


and  other  volatile  metal  halides  are  probably  similar.  The  dimeric  tertiary 
phosphine  and  arsine  compounds  also  contain  double  halide  bridges  (page 
81 )  and  a  number  of  olefine  complexes  and  thio  ether  complexes  have  hern 
formulated  in  the  same  way  (see  page  83).  Alky]  derivatives  <>f  gold  bro- 
mide arc  dimeric  and  probably  have  the  structure112 

H  Br  R 


Au 


An 


R 


R 


The  presence  of  double  bridges  in  platinum(II)  chloride  results  in  the  forma- 
tion of  an  infinite  chain  of  PtCl4  groups. 

In  addition  to  double  halogen  bridges,  triple  or  single  bridges  may  be 
formed.  The  triple  bridge  is  illustrated  by  ions  of  the  type  [Mni2X9]=  (see 
page  7),  while  single  halogen  bridges  occur  in  such  species  as  [A1F5]= 
(page  389).  The  compounds 


CI 


Ag 


/ 


Co(XH,)4 


CI 


S04 


and 


Ag 


CI 


CI 


Co(NH3)3(H20) 


S04 


may  also  exemplify  single  bridges.  When  silver  ion  is  added  to  a  solution 
of  the  dichlorotetramminecobalt(III)  ion,  silver  chloride  does  not  precipi- 
tate at  once,  but  the  silver  ions  lose  their  ionic  property  through  coordina- 
tion with  the  chloride  of  the  cobalt  (III)  complex.  The  ion  so  formed  is  not 
stable,  however,  and  slowly  precipitates  silver  chloride. 

The  phenomenon  of  "interaction  absorption"  is  often  observed  in  bridged 
halogen  complexes.  When  the  halides  (cyanides,  or  oxides)  of  a  metal  in 
two  different  oxidation  states  are  associated  in  a  single  molecule  or  ion  (or 
possibly  in'  such  relatively  less  intimate  admixture  as  crystal  compounds  or 
solutions — the  point  is  not  clear),  a  high  degree  of  color  is  developed.  Thus, 
CuCl,CuCl;  SbClgSbCl* ;  Sn('l,-Sn('l,  ;  ( •>,Au,Au,,<  'I,  ;  and 
[Pd(XH3)2Br2]-lPd(\Il;;)-jBr4]  are  all  highly  colored60.  In  none  of  these  casee 


111.  Wriggee  and  Biltz,  Z.  anorg.  allgem.  Chem.,  228,  :>7J    r.  ».;•■, 
1 1J .  Gibson  and  Simonsen,  J.  Chem.  Soc.,  1930,  2531 ;  Buroway,  ei  al.,  .1   { ' h<  m 
1937,  1090. 

113.  Werner,  Z.  anorg.  Chem.,  14,  31  (1897). 

114.  Werner,  Z.  anorg.  Chem.,  15,  155  (1897). 


Soc, 


20  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

has  conclusive  evidence  for  an  intermediate  oxidation  state  of  the  metal 
been  obtained;  indeed,  .strong  evidence  indicates  the  nonexistence  of  such 
states.  In  the  first  example  the  ridiculous  assumption  of  the  ion  Cu1,5+ 
would  be  necessary,  while  in  the  case  of  the  diammino  palladium  compound, 
x-ray  data  and  magnetic  behavior  definitely  preclude  the  existence  of  the 
intermediate  state.  Nonetheless,  a  resonance  between  the  two  oxidation 
states  produces  high  color  and  probably  renders  the  two  metal  atoms  indis- 
tinguishable.  The  probablity  that  a  halogen  (or  similar)  bridge  is  necessary 
for  this  phenomenon  is  supported  by  the  fact  that  rapid  electron  exchange 
occurs  between  the  coordinately  saturated  complexes,  [OsIIdipy3]++  and 
[Osmdipy3]+++,  without  the  development  of  high  color115. 

The  Donor  Properties  of  Oxygen 
Hydrate  Formation 

All  metallic  ions  apparently  form  hydrates  in  aqueous  solution,  frequently 
surrounding  themselves  with  large  numbers  of  molecules  of  water.  Part  of 
this  water  is  held  by  van  der  Waals  forces  only,  but  it  is  difficult  to  escape 
the  conclusion  that  in  every  case  a  few  molecules  at  least  are  coordinated 
to  the  metallic  ion.  In  many  cases,  of  course,  the  hydrates  can  be  crystal- 
lized from  the  solution.*  These  usually  retain  only  enough  molecules  of 
water  t  o  satisfy  the  coordination  number  of  the  metallic  ion,  but  sometimes, 
as  with  the  alums,  stable  hydrates  contain  more  than  this  amount.  To 
account  for  these  we  may  assume  that  (a)  the  excess  water  is  not  chem- 
ically combined,  but  is  held  in  place  by  the  demands  of  the  lattice  structure, 
(b)  the  coordination  number  of  the  metal  is  abnormal,  (c)  second  and  even 
third  coordination  spheres  are  formed  (d)  the  molecules  of  water  are  poly- 
meric or  (e)  part  of  the  water  is  combined  with  anion.  It  is  often  assumed 
that  water  of  hydration  which  is  not  lost  at  100°C  must  be  chemically  com- 
bined, but  this  does  not  necessarily  follow,  for  lattice  compounds  sometimes 
show  considerable  stability.  On  the  other  hand,  chemically  combined  water 
may  escape  from  salts  at  low  temperatures — even  at  room  temperature — if 
the  anion  is  one  which  readily  coordinates  with  the  cation,  thus  displacing 
the  water  from  the  coordination  sphere. 

Werner  recognized  that  water  molecules  are  sometimes  held  by  feeble, 
Qonchemical  forces  in  writing  formulas  such  as  [Co(NH3)5Cl]Cl2-H20.  The 
water  may  be  removed  from  this  compound  without  changing  its  properties 
except  for  disruption  of  the  crystal  lattice,  while  dehydration  of  the  isomeric 
[Co(N  1 1;;  >.  I  M  )|( '!:,  is  accompanied  by  change  in  color  and  solubility,  and  by 
Loss  of  ionic  function  of  one  chloride1  ion. 

LIS.  Dwyer,  Mellor,  and  Gyarfas,  Nature,Mt  176  (1950). 

Man)   anions  also  have  the  power  of  combining  with  water — this  union  takes 
place  through  hydrogen  bonding. 


GENERAL  SURVEY  21 

In  his  early  papers,  Werner111  also  gave  expression  to  the  though.1  thai 
Beveral  coordination  spheres  can  form  around  a  positive  ion.  I  [e  argued  that 
when  water  molecules  form  a  coordination  sphere  around  a  positive  ion,  a 
negative  charge  is  induced  on  the  inner  surface  of  the  sphere,  80  that  the 

outer  surface  hears  a  positive  charge,  just  as  the  metal  ion  itself  does.  This 
enables  it  to  attract  another  sphere  of  water  molecules,  which  will  likewise 
hear  an  induced  charge.  The  process  may  he  repeated  several  time-. 

Closely  related  to  this  hypothesis  was  the  thought  that  water  exists  in 
hydrates  in  the  polymeric  form.  In  view  of  the  fact  that  water  as  such  ifl 
iated,  this  is  not  an  unreasonable  assumption,  though  Werner  had 
little  experimental  evidence  on  which  to  support  it.  The  fact  that  many 
salts  contain  exactly  twice  as  many  water  molecules  as  can  he  explained  by 
the  coordination  theory  made  it  an  easy  assumption.  Such  an  explanation 
seems  naive,  hut  the  fact  that  "multiple  coordination  spheres"  do  exist  in 
solution  cannot  be  denied.  Their  existence  has  been  demonstrated  by  the 
diffusion  studies  of  Brintzinger  (Chapter  18)  and  by  the  polarographic  work 
of  Laitinen  and  his  co-workers117. 

As  is  to  be  expected,  the  ease  with  which  metallic  ions  form  hydrates  in- 
creases with  increasing  charge  and  with  decreasing  radius.  The  ions  of  the 
alkali  metals  except  lithium  and  sodium  are  seldom  hydrated  in  the  solid 
state,  and  the  hydrates  of  these  two  are  unstable;  divalent  ions  of  the  lighter 
metals  are  usually  hydrated  (unless  they  exist  in  highly  insoluble  com- 
pounds) and  trivalent  ions,  nearly  always  so.  In  any  periodic  group  the 
stability  of  the  hydrates  is  greatest  for  the  smallest  ions,  while  the  number 
of  water  molecules  normally  held  is  greatest  for  the  large  ions.  Even  in 
complexes  in  which  water  molecules  undoubtedly  occupy  positions  in  a  true 
coordination  sphere,  the  nature  of  the  oxygen-metal  bond  varies  a  great 
deal.  Hunt  and  Taubells  showed  that  the  water  in  the  hydrated  forms  of 
Al  ,  Ga  and  Th4+  exchange  with  the  solvent  water  in  about  three 
minutes,  so  the  metal-oxygen  bond  must  have  a  considerable  degree  of 
ionic  character.  The  hydrated  chromium(III)  ion,  on  the  other  hand,  ex- 
changes very  slowly,  the  halftime  being  about  forty  hours.  They  made  the 
observation,  also,  that  all  of  the  cations  studied  show  a  greater  affinity  for 
H2018  than  for  H2016.  The  hydrated  cobalt  (III)  ion  exchanges  rapidly.  This 
is  probably  not  due  to  a  lack  of  covalenl  bonding,  but  to  a  rapid  electron 
exchange  between  the  hydrated  cobalt(III)  and  cobalt (II)  ions,  and  a  rapid 
exchange  between  the  latter  and  the  solvent  water119. 

116.  Werner,  Z.  anorg.  Chem.,  3,  267    1S93). 

117.  Laitinen,  Bailar,  Holtaclaw,  and  Quagliano,  ./.  .1///.  Qk  m.  Soc  ,  70,  2999    1948  ; 

Laitinen,  Frank  and  Kivalo,  ./.  .1///.  Chem.  8oe.,  75,  2866    1"~ 

118.  Bunt  and Taube, /.  Chem.  Phyt  .  18,  757    1950  ;  19, 602    1951   . 

119.  Friedman,  Taube,  and  Hunt,  ./.  Chem.  Phys.,  18,  759    I960  ■ 


99 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Hydroxy  1  Coordination 

The  hydroxide  ion  has  a  strong  coordinating  tendency,  partly  because  it 
has  i hive  pairs  of  unshared  electrons,  but  chiefly  because  of  its  negative 
charge.  The  hydrates  of  highly  charged  metallic  ions  readily  lose  protons 
with  the  formation  of  hydroxo  complexes: 

[A1(H20)6]+++  -»  H+  +  [Al(H,0)5OH]++  ->  H+  +  [Al(H20)4(OH)2]+,  etc. 

The  aquo  ammine  complexes  undergo  the  same  type  of  reaction: 
[Co(NH8)6H20]+++  ^±  [Co(NH3)5OH]++  +  H+.  The  phenomenon  underlies 
our  present  theories  of  acidity,  hydrolysis  and  amphoterism,  and  is  discussed 
in  Chapter  12. 

The  hydroxide  group  can  act  as  a  bridging  group  between  two  metallic 
ions,  under  which  conditions  it  is  almost  entirely  devoid  of  basic  properties. 
This  bridge  forming  ability  may  extend  to  great  lengths  and  an  interesting 
theory  of  colloidal  oxides  has  been  based  upon  it  (Chapter  13). 

Werner's  postulate  that  basic  salts  are  polynuclear  complexes  held  to- 
gether by  hydroxy  1  groups120  has  been  shown,  by  x-ray  studies,  to  be  un- 
tenable in  most  cases.  The  basic  chlorates  and  perchlorates  of  lead  have  not 
been  studied  by  x-ray  analysis,  but  the  conductivities  and  other  properties 
of  their  solutions  indicate  that  they  have  the  structures 


X2 


una 


•(  M"  bridges  are  common  in  the  polynuclear  cobalt  complexes.  The  chief 
constituenl  of  Vbrtmann's  sulfate,  which  is  obtained  by  oxidation  of  an 

120.  Werner,  Ber.,  40,  I  HI  (1907). 

121.  Weinland  and  Stroh,  Ber.}  55,  2210,  2706  (1922). 

122    Weinland  and  Paul,  Z.  anorg.  Chem.,  129,  243  (1923). 


(ihWhlx'AL  SI  7,'lA'l 


23 


ammoniacaJ  solution  of  a  cobalt  salt,  is 

/     \ 
(NH,)4Co  Co(NH,)< 

\       / 
OH 


(S04)2 


Such  ions  as 


oil 

/      \ 
(XH3)4Co  Co(NH3)4 

\       / 

OH 

/       \ 

(NH,)8Co— OH— Co(NH,), 

\       / 
OH 


and 


;co(nh3). 


have  been  known  for  many  years.  The  hexol  salt  is  of  special  interest,  as  it 
was  the  first  strictly  inorganic  compound  to  be  resolved  into  optical  anti- 
podes1'24. Adamson,  Ogata,  Grossman,  and  Newbury125  have  come  to  the 
conclusion  that  Durrant's  salt  has  the  bimolecular  structure 


K. 


Alcohol*  and  Kthers 


OH 

/       \ 

(C204)2Co  Co(C204)2 

\       / 
OH 


The  organic  derivatives  of  water,  the  alcohols  and  ethers,  show  much  less 
tendency  to  form  coordination  compounds  than  docs  water;  nevertheless,  a 

123.  Werner,  Ber  .  40,  4609  (1907). 

124.  Werner,  Ber.,  47,  3087    L914 

l_'.V  Adamson,  Ogata,  Grossman,  and  Newbury,  0  \  K    Contract  23809,  Technical 
Report,  M.uch  L954. 


L 


24  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

large  number  of  such  compounds  is  known.  The  compounds  of  the  alcohols 
are  more  stable  than  those  of  the  ethers,  the  stability  in  each  series  de- 
creasing as  the  size  of  the  organic  group  increases.  Because  of  the  chelation 
effect,  the  polyhydric  alcohols  form  somewhat  more  stable  compounds  than 
do  the  monohydric  alcohols.  Glycol  is  able  to  displace  water  from  hydrates 
of  heavymetals,  each  alcoholic  hydroxyl  group  taking  the  place  of  one  mole- 
cule of  water  in  the  coordination  sphere126.  Glycerol  ordinarily  behaves  as  a 
bidentate  donor,  also,  adjacent  hydroxyl  groups  coordinating.  The  third 
hydroxyl  group  is  prevented  from  combination  by  steric  factors.  The  di- 
valent ions  of  the  alkaline  earths127,  and  of  cobalt,  nickel,  copper,  and  zinc, 
all  form  compounds  in  this  way,  those  of  the  heavy  metals  being  rather 
unstable.  Other  poly  hydroxy  alcohols  and  even  the  sugars  form  coordina- 
tion compounds,  the  tendency  to  combine  with  the  ions  of  the  alkaline  earths 
being  particularly  noticeable.  The  purification  of  sugar  through  the  precipi- 
tation of  calcium  and  strontium  "saccharates"  is  of  interest  in  this  connec- 
tion. The  structure  of  these  compounds  has  not  been  studied  in  detail,  but 
they  are  evidently  coordination  compounds  rather  than  salts. 

In  the  presence  of  polyhydric  alcohols  such  as  mannitol  and  sorbitol, 
sodium  hydroxide  does  not  precipitate  iron  (III)  ion128.  Addition  of  barium 
chloride  to  such  basic  solutions  gives  pale  yellow,  crystalline  products  con- 
taining the  alcohol,  iron,  and  barium  in  a  1:1:1  ratio.  Traube  and  Kuhbier 
write  the  formula  of  this  product  as 

CH,— CH— CH— CH— CH CH2 

I  I  I  I  I  I 

O  O         O         OH     O  O 

/  X    / 

7  Ba 

Fe 

but  they  cite  no  evidence  to  support  such  a  formulation.  Scale  models  indi- 
cate that  it  is  improbable  that  three  consecutive  hydroxyl  groups  are  co- 
ordinated to  the  iron.  According  to  Traube  and  Kuhbier,  treatment  of  this 
product  with  sodium  sulfate  gives  Na[FeC6Hio06]-3H20,  in  which  there 
must  be  two  uncoordinated  hydroxyl  groups.  Several  similar  compounds 
containing  sugars  or  polyhydroxy  acids  and  a  variety  of  metal  ions  have 
been  prepared  and  analyzed,  but  their  structures  have  not  been  deter- 
mined129. Some  of  these  oxidize  in  the  air  to  formic  acid,  carbon  dioxide, 
and  similar  compounds180. 

126.  Cum  and  Bockisch,  Ber.,  41,  3465  (1908);  Griin  and  Boedecker,  Ber.,  43,  1051 

(1910). 

127.  Grttu  and  Husmann,  Ber.,  43,  1291  (1910). 
Us    Traube  and  Kuhbier,  Ber.,  65,  187  (1932). 

129.  Traube  and  Kuhbier,  Ber.,  66,  1545  (1933);  69,  2655  (1936). 
13G.  Traube  and  Kuhbier,  Ber.,  65,  190  (1932);  69,  2664  (1936). 


GENERAL  SURVEY 


25 


The  ethanolamines  can  coordinate  through  cither  oxygen  or  nitrogen. 
Tettamanzi  and  Carliul  found  that  triethanolamine  tonus  addition  com- 
pounds of  the  type  M.\  _  \  (',11,011);;  (where M  isCo,  Ni,Cu,Cd,  PI),  Ca, 
Mgj  or  Sr  .  some  of  the  compounds  being  hydrated.  No  Btudy  of  the  struc- 

Ures  of  these  compounds  lias  been   made,   bul    in   view    of    the    Structural 

similarity  of  triethanolamine  and  nit rilotriacet ic  acid,  one  may  assume  the 
presence  oi  chelate  rings,  their  number  depending  upon  the  coordination 
number  of  the  metal  ion: 


(ho-ch2-ch2)3_x      (ch2ch2  ohj3_x 


Ethers  form  addition  compounds  with  a  wide  variety  of  compounds. 
Confirmation  of  this  is  found  in  the  high  solubility  of  the  heteropolyacids, 
of  uranyl  nitrate,  and  of  magnesium  iodide,  in  ethers.  The  best  known  of 
the  ether  coordination  compounds  are  those  formed  with  the  Grignard 
reagent.  Spacu132  has  prepared  some  interesting  compounds  in  which  ether 
and  pyridine  share  the  coordination  sphere:  [Mg  py4(ether)2]Br2  and 
[Mg  pyi  ether]I2  . 

The  formation  of  a  deep  color  in  the  well  known  iron  (III)  chloride  test 
for  phenols  indicates  that  phenols  form  compounds  with  the  heavy  metals. 
In  the  thermometric,  conductometric,  and  spectrophotometric  titration  of 
phenol  with  iron(III)  chloride,  Banerjee  and  Haldar133  find  breaks  at  molar 
ratios  of  1:3  and  1:6.  Upon  electrolysis,  the  iron(III)  ion  goes  to  the  anode. 
These  findings  suggest  the  reactions 

Fe^+_>  [Fe(OC6H5)3]°^  [Fe(OC,H,).]- 

(  atechol,  because  of  the  effect  of  chelation,  forms  stable  complexes  with 
the  heavy  metals: 


K*[MC  /C«H')»}XH*0 


131.  Tettamanzi  and  Carli,  Gazz.  chim.  Hal.,  63,  566  (1933);  64,  315  (1934);  AM  accad. 
sci.  7  'asse  sci.  fis.,  mat.  nat.,  68,  500  (1933);  Garelli,  AM  accad.  sci. 

fit.,  mat.  not.,  68,  398  (1933). 
-  Cluj,  1,  72  (1921). 

Banerjee  and  Baldar,  Natun  .  165,  1012  (1950). 
134.  Weinland  and  Binder,  Ber.t  45,  148,  1113  (1912);  46,  874    1913);  Weinland  and 
Walther,  Z.  anorg.  Chem.,  126,  Ml  (191 


26  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

If  the  phenolic  group  can  take  part  in  the  formation  of  a  chelate  ring 
with  souk1  other  strongly  coordinating  group,  very  stable  complexes  may 
be  formed.  Thus,  naphthazarin  reacts  quantitatively  with  beryllium  ion  to 
give  the  complexes 

HO    O 


AND 


HO    O 
in  which  the  coordinated  oxygen  atoms  are  doubtless  equivalent135. 

Peroxide  Coordination 

Many  salts  have  been  shown  to  crystallize  with  hydrogen  peroxide  "of 
crystallization"  136.  In  some  cases,  at  least,  this  may  be  chemically  com- 
bined with  the  salt,  as  is  shown  by  cryoscopic  measurements137. 

The  peroxo  group  may  serve  as  a  bridge  between  two  cobalt  ions.  When 
an  ammoniacal  cobalt (II)  solution  is  allowed  to  stand  in  the  air,  the  first 
product  formed  is  a  brown  decammine-/x-peroxo-dieobalt(III)  salt, 
[(NH3)5Co — 02 — Co(NH3)o]X4138,  which  upon  further  oxidation  is  converted 
to  the  deep  green  [(NH3)6Co — 02 — Co(NH3)6]X5  in  which  one  of  the  cobalt 
atoms  seems  to  have  achieved  a  valence  of  4+.  The  dicobalt(III)  salts 
are  reduced  to  cobalt(II)  by  four  equivalents  of  arsenic(III)  oxide  (one 
equivalent  for  each  cobalt  and  two  for  the  peroxo  group)  while  the  co- 
balt(III)-cobalt(IV)  salts  require  five  equivalents  of  reducing  agent.  The 
brown  dicobalt(III)  salt  is  diamagnetic,  whereas  the  cobalt(III)-cobalt(IV) 
sail  is  paramagnetic139. 

135.  Underwood,  Toribara,  and  Neuman,  J.  Am.  Chem.  Soc.,  72,  5597  (1950). 

136.  Tanatar,2?er.,32,  1544  (1899);  Z.  anorg.  Chem. ,2%,  255  (1901);  Rudenko,  J.  Russ. 

Phys.  Chem.  Soc,  44,  1209  (1912);  Kazanetzkii,  ./.  Russ.  Phys.  Chem.  Soc.,  46, 
1110  (1914). 

137.  Jones  and  Murraj  .  Am.  Chem.  ./.,  30,  205  (1903);  Maass  and  Hatcher,  J.  Am. 

Chem.  Soc.,  44,  2472  (1922). 
L38    \  ortmann,  Monatshefte,  6,  404  (1885);  Werner  and  Mylius,  Z.  anorg.  Chem.,  16, 
246  (1898  ;  Werner,  .1////.,  375,  1  (1910). 
' .  1'u  and  Rehm,  Z.  anorg.  allgem.  Chem.,  237,  79  (1938). 


GENERAL  SURVEY 


27 


'Vortman's  sulfate"  is  a  mixture  of  materials,  containing  the  sulfates  of 


in, 

(NH,)4Co 


o, 


Ml 


III 
Co(NH,)< 


(A)         and 


o2 

III/     \IV 
(XH3)4Co  Co(NH3)4 

\     / 
XII. 


(B) 


Compound  B,  on  wanning  with  sulfuric  acid,  liberates  one  and  a  half  atoms 
of  oxygen,  and  on  further  heating,  two-thirds  of  an  atom  of  nitrogen,  leav- 
ing t he  cobalt  in  the  dipositive  state.  These  reactions,  again,  confirm  the 
tetravalency  of  one  cobalt  atom.  The  surprising  stability  of  these  com- 
pounds is  illustrated  by  the  reaction 


/     \ 
(XH3)4Co  Co(NH3)4 

\     / 
NH, 


X4  +  en 


en2Co 


05 


XH, 


Coen< 


X4 


Compound  B  and  its  ethylenediamine  analog  are  both  paramagnetic140. 
The  ethylenediamine  compound  can  be  reduced  to  the  dicobalt(III)  state 
by  nitrite,  hydrazine,  ferrocyanide,  arsenite  or  thiosulfate,  but  not  by  hy- 
droxylamine,  hydrogen  peroxide  or  mercury(I)  ion.  The  product  of  the 
reduction  can  then  be  reoxidized  to  the  Co(III)-Co(IV)  state  by  treatment 
with  permanganate,  hypochlorite,  bromine,  bromate,  or  nitric  acid,  but  not 
by  dichromate,  peroxide,  or  mercury(II),  iron(III)  or  silver  ions.  These 
reactions  establish  the  reduction  potential  at  about  one  volt. 

The  peroxo-  group  in  compound  B  can  be  replaced  by  other  groups  with 
reduction  of  the  cobalt  to  the  3+  condition.  Thus 


02 

/     \ 
(XH3)4Co  Co(XH3)4 

\     / 
NH, 


+  SO: 


(XH3)4Co 


S04 

/       \ 
>  ( 

\       / 

XII: 


Co(XH3)4 


Among  the  other  doubly  bridged  cobalt  (III),  cobalt  (IV)   compounds 
described  by  Werner  the  triply  bridged  compound 


XII., 
Ill/     \IV 
(NH  )  Co-OH-Co(XTI3)f 

\     / 
02 


CI 


k  worthy  of  note1 


140.  Malatesta,  Gazz.  chim.  ital.,  72,  287  (1942), 

141.  Werner,  Ann.,  375,  104  (1910). 


28 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Brimm142  has  pointed  out  that  most  of  the  results  which  have  been  inter- 
preted  to  show  the  presence  of  tetrapositive  cobalt  in  these  compounds  can 
be  explained  on  the  assumption  that  they  contain  the  superoxide  group. 

Connick  and  McVey143  have  identified  two  peroxo  complexes  of  plu- 
tonium(IV)  in  aqueous  solution.  While  the  structures  of  these  are  not 
lii i ally  proved,  they  seem  to  contain  the  rings 


Pu 


OH 


02 


Pu 


and 


Pu 


02 


<> 


Pu 


Metallic  Oxide  Coordination 

Metallic  oxides  frequently  coordinate  with  metallic  ions,  as  is  evidenced 
by  the  increased  solubility  of  such  oxides  in  salt  solutions.  Beryllium  oxide, 
for  example,  dissolves  readily  in  saturated  beryllium  sulfate  solution,  at 
the  same  time  increasing  the  solubility  of  the  sulfate  itself144.  The  solubility 
relations  indicate  that  each  beryllium  ion  combines  with  four  beryllium 
oxide  molecules.  The  compound  [Be(BeO)4]S04  is  more  soluble  than  its 
analog  [Be(H20)4]S04 .  The  structure  given  for  the  complex  is  supported 
by  the  lowering  of  the  freezing  point,  which  indicates  that  addition  of 
beryllium  oxide  to  a  solution  of  beryllium  sulfate  does  not  increase  the 
number  of  ions  in  solution.  Beryllium  selenate  gives  the  same  result  as  the 
sulfate.  A  related  situation  is  found  in  the  anion  commonly  described  as 
[RuCl5OH]=,  but  which  is  shown  by  crystal  analysis  to  be  the  oxo  complex 
[Cl5Ru— 0— RuCl5]4-  145. 

Oxyanion  Coordination 

The  anions  of  all  oxyacids  have  donor  properties,  but  in  very  different 
degree.  It  is  sometimes  said  that  the  nitrate  and  perchlorate  ions  do  not 
enter  into  complex  formation,  but  this  is  not  true.  Nitratopentammineco- 
balt(III)  salts  were  prepared  by  some  of  the  earliest  investigators,  and  were 
described  in  detail  by  Jorgensen146.  Later  investigations  have  led  to  the 
preparation    of    [Co(NH3)3(N03)3],147    [Co(NH3)4(N03)2]N03-H2(V48    and 

142.  Brimm,  private  communication.  Quoted  in  Kleinberg  "Unfamiliar  Oxidation 

States,"  p.  100,  University  of  Kansas  Press,  1950. 

143.  Connick  and  McVey,  National  Nuclear  Energy  Series,  Vol.  14B  (The  Transura- 

nium Elements),  p.  445,  1949. 

144.  Sidgwick  and  Lewis,  J.  Chem.  Soc,  1926,  1287. 

145.  Mellor,  Report  of  the  Brisbane  meeting  of  the  Australian  and  New  Zealand  As- 

sociation for  the  Advancement  of  Science,  28,  137  (1951). 

146.  Jorgensen,  J.  prakt.  Chem.,  [2]  23,  227  (1881). 

147.  Jorgensen,  Z.  anorg.  Chem.,  5,  185  (1894). 

148.  Birk,  Z.  anorg.  allgem.  Chem.,  164,  241  (1927). 


GENERAL  SURVEY  29 

[Co  imi.(\();;)2]N03-H20149.  The  last  two  are  shown  to  be  dinitrate  salts 
rather  than  aquo  nitrato  salts  by  the  fact  tliat  the  loss  of  water  does  not 
change  the  properties  greatly. 

Transference  measurements  on  solutions  of  plutonium(IV)  in  \M  HN03 
indicate  the  existence  of  the  complex  [Pu(N03)]+,  which  coordinates  with 
more  nitrate  ions  as  the  concentration  of  IIX03  is  increased.  In  bM  acid, 
the  bright  green  ion  [Pu(X03)6]=  is  present,  and  (XH4)2[Pu(N03)6]  can  be 
crystallized  from  the  solution.  Thorium  shows  a  similar  behavior,  giving  a 
salt  which  is  isomorphous  with  the  plutonium(IV)  and  cerium(IV)  com- 
pounds150. 

G.  F.  Smith  and  his  students  have  demonstrated  the  existence  of  both 
nitrate  and  perchlorate  cerium(IV)  ions151  but  the  exact  structure  of  the 
ions  is  not  yet  clear.  The  oxidation-reduction  potential  of  the  cerium (III)- 
eerium(IY)  couple  varies  greatly  with  the  nature  of  the  acid  present.  In 
IN  acid,  the  electrode  potentials  (referred  to  the  normal  hydrogen  elec- 
trode) are  HC104 ,  1.70  volts;  HN03 ,  1.61  volts;  H2S04 ,  1.44  volts;  HC1, 
1.28  volts.  This  variation  indicates  that  either  the  Ce(III)  or  the  Ce(IV) 
or  both,  combine  with  the  anion  of  the  acid.  Duval152  has  reported  pentam- 
minecobalt  complexes  in  which  chlorate,  bromate,  iodate  and  perchlorate 
groups  occupy  the  sixth  coordination  position. 

The  sulfate  ion  can  occupy  either  one  coordination  position,  or  two.  In 
either  event,  of  course,  it  contributes  a  charge  of  minus  two  to  the  ion  of 
which  it  becomes  a  part.  The  first  type  of  compound  is  illustrated  by  sul- 
fatopentamminecobalt(III)  bromide,  [Co(XH3)5S04]Br153,  which  is  pre- 
pared by  heating  the  chloropentammine  chloride  with  concentrated  sul- 
furic acid.  The  sulfate  group  in  the  coordination  sphere  is  not  readily 
replaced,  but  is  precipitated  by  boiling  with  barium  salts.  The  ion  slowly 
aquates  on  standing  in  solution: 

[Co(NH3)5S04]+  +  H20  -+  [Co(NH3)5(H20)]+++  +  SOr 

Sulfato-aquo  complexes  of  several  types  evidently  exist  in  aqueous  solutions 
of  chromium  (III)  sulfate154. 

Cases  in  which  the  sulfate  group  occupies  two  positions  in  the  same  co- 
ordination sphere  are  not  as  well  known.  The  double  sulfates  of  iron,  chrom- 

149.  Schramm,  Z.  anorg.  allgem.  Chem.,  180,  170  (1929). 

150.  Hindman,  National  Nuclear  Energy  Series,  Vol.  14B  (The  Transuranium  Ele- 

ments), p.  388,  1949. 

151.  Smith,  Sullivan,  and  Frank,  Ind.  Eng.  Chem.,  Anal.  Ed.}  8,  449  (1936) ;  Smith  and 

Getz,  Ind.  Eng.  Chem.,  Anql.  Ed.,  10,  191  (1938);  Kott,  thesis,  University  of 
Illinois,  1940. 

152.  Duval,  Ann.  Chim.,  18,  241  (1932). 

153.  Jorgensen,  J.  prakt.  Chem.,  [2]  31,  270  (1885). 

154.  Enlmann,  Angew.  Chem.,  64,  500  (1952). 


:*() 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


ium  and  the  rare  earths  may  contain  the  anions  [M(S04)3]=,  but  they  are 
too  unstable  to  exist  in  solution.  The  case  of  potassium  iridium  sulfate, 
3KaS04-Ir2(S04)8-2H20  or  K3[Ir(S04)3]-H20,  is  perhaps  a  little  more  cer- 
tain, for  this  salt  does  not  give  the  characteristic  tests  for  sulfate  ion155.  Wein- 
land  and  Sierp168  have  prepared  alkaloid  salts  of  the  acids  H3[Fe(S04) (€204)2] 
and  H:!( Fe(S<  VMC^  )*)],  in  which  the  sulfate  group  is  evidently  doubly 
coordinated.  Duff157  claims  to  have  prepared  [Co  en2S04]Br-H20,  but  Job158 
and  Ephraim  and  Flugel159  believe  the  salt  to  be  [Co  en2(H20)S04]Br,  in 
which  the  sulfato  group  occupies  only  one  coordination  position.  In  any 
event j  the  sulfate  group  is  not  held  very  tenaciously,  for  in  solution  the 
complex  ion  is  rapidly  converted  to  [Co  en2(H20)2]+++. 

Several  cases  are  known  in  which  the  sulfato  group  acts  as  a  bridge  be- 
tween two  metal  atoms,  but  in  every  case  it  must  evidently  be  accompanied 
by  some  other  bridging  group.  When  octammine-ju-amino-ol-dicobalt(III) 
chloride, 


NH2 

/      \ 
(NH3)4Co  Co(NH3)4 

\       / 
OH 


Cl4, 


is  heated  with  sulfuric  acid,  the  "ol"  bridge  is  replaced  by  a  sulfato  bridge: 

NH2 


(NH3)4Co 


Co(NH3)4 


\        / 

oso 
o2 


The  sulfato  bridge  is  eliminated  by  heating  with  concentrated  hydrochloric 
acid;  chloroaquo-octammine-/x-amino-dicobalt(III)  chloride 


Cl3  160. 


[CI  H20  "1 

1  I 

(NH3)4Co— NH2— Co(NH3)JC!4 


results.  The  /x-amino-sulfato  compounds  are  also  obtained161  by  the  action 
of  sulfur  dioxide  upon  salts  of  the  /x-amino-peroxo  series  (see  page  27). 

155.  Delepine,  Compt.  rend.,  142,  1525  (1906). 

156.  Wcin land  and  Sierp,  Z.  anorg.  Chem.,  117,  59  (1921). 
L57    Duff,  •/    Chem.  Sue,  121,  450  (1922). 

L68  Job,  Bull.  80C.  chim.,  |4]  33,  15  (1923). 

159  Ephraim  and  Flugel,  Helv.  chim.  Acta,!,  727  (1924). 

L60  Werner,  Beddow,  Baselli,  and  Steiniteer,  Z.  anorg.  Chem.,  16,  109  (1898). 

163  Werner,  .1/,//.,  375,  15  (1910 


GENERAL  SURVEY  3] 

Gibson  and  his  co-workers"1'-  have  studied  a  case  of  a  very  different  type 
of  sulfate  bridging.  The  substance  iCjII.-.hAn-jSO.,  was  Pound  to  be  a  dimer 
in  acetone,  and  probably  has  the  structure 


Foss  and  Gibson163  have  reported  a  similar  compound  in  which  the  phenyl 
phosphate  group,  C6H5OP03=,  replaces  the  sulfate. 

The  sulfate  ion  has  the  rather  unusual  ability  to  form  hydrates;  metallic 
sulfates  usually  crystallize  from  solution  with  one  molecule  of  water  more 
than  other  salts  containing  the  same  metallic  ion.  Thus  the  vitriols  of  the 
divalent  ions  of  magnesium,  zinc,  cadmium,  vanadium,  chromium,  man- 
ganese, cobalt,  and  nickel  are  heptahydrates  and  that  of  copper  is  a  penta- 
hvdrate.  In  these  complexes,  two  oxygens  of  the  sulfate  ion  are  hydrogen 
bonded  to  the  water. 

The  tellurate  and  iodate  ions  are  remarkable  in  that  when  they  co- 
ordinate with  copper,  they  stabilize  the  trivalent  state,  forming  such  com- 
plexes as  [Cu(Te06)2]9-  and  [Cu(I06)2]7- 164-  165. 

The  bleaching  of  solutions  of  iron  (III)  chloride  by  addition  of  phosphate 
ion  indicates  the  existence  of  phosphate  complexes166.  Ricci167  advanced 
evidence  for  the  existence  of  H3[FeCb,P04]  and  H3[FeCl3As04],  but  later 
work  indicates  that  the  complexes  probably  contain  no  chlorine.  Jensen168 
found  the  solubility  of  FeP04  and  A1P04  to  rise  with  increasing  phosphate 
ion  concentration,  but  to  be  independent  of  the  chloride  ion  concentration. 

162.  Gibson  and  Weller,  ./.  Chem.  Soc.,  1941,  102;  Evens  and  Gibson,  ./.  Cht  m.  Soc., 

1941,  Hi!». 

163.  Foss  and  Gibson,  ./.  Chem.  Soc,  1949,  3075. 

164.  Malatesta,  Gozz.  chim.  itol.,  71,  407,  580  (1941  I. 

165.  Lister,  Can.  ./.  Chem.,  31,  638    1953). 

166.  Weinland  and  Ensgraber,  '/.    anorg.  Chem.,  84,  340    L91 1  , 

L67.  Ricci  and  Meduri,  Gazz.  chim.  itol.,  64,  235    1934);  Ricci  and  Lamonica,  G 

(■hint,  itol.,  64,  294  (1934   ;  Ricci  and  Saraceno,  thesis,  University  of  Messina, 
1929. 

168.  Jensen,  Z.  anorg.  aUgem.  Chem.,  221,  1  (1934). 


32  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

I)i-,  tri-  and  polyphosphates  all  show  a  remarkable  ability  to  form  stable 
complexes,  even  with  the  alkaline  earth  ions,  so  some  of  them  have  found 
wide  use  industrially  (Chapter  23).  Pyrophosphate  complexes  of  many 
metals  have  been  studied  in  solution  by  a  variety  of  physical  methods. 
For  example,  Haldar189  has  studied  the  pyrophosphate  complexes  of  Cu++, 
Ni++,  and  Co++  by  thermometric  and  conductometric  titrations,  and  by 
magnetic,  cryoscopic,  and  transport  measurements.  He  finds  evidence  for 
i  he  existence  of  two  series  of  complexes,  [M(P207)]=  and  [M(P207)2]6_. 
Watters  and  Aaron170  report,  in  addition,  copper  complexes  with  Cu:P2074_ 
ra1  ios  of  2: 1  and  4:1,  which,  however,  exist  only  in  dilute  solutions. 

The  carbonate  ion  forms  coordinate  bonds  easily,  as  witnessed  by  its 
strong  tendency  to  unite  with  hydrogen  ions.  In  the  metal  amminessuchas 
|(  <)(\II;;)4C03]+  it  seems  to  occupy  two  coordination  positions.  In  view  of 
i he  fact  that  this  coordination  entails  the  formation  of  a  four-membered 
ring,  it  is  surprisingly  stable.  Because  the  pentammine  [Co(XH:05CO3]Cl- 
1 1-<  I  gives  an  alkaline  reaction,  and  because  he  thought  that  the  molecule  of 
water  could  not  be  removed  without  destruction  of  the  complex,  Werner 
was  of  the  opinion  that  the  formula  of  the  salt  should  be  written 
i(,(»(\H3)5HC03]Cl(OH)171.  Lamb  and  Mysels172,  however,  found  that  all 
of  the  water  can  be  removed  without  destruction  of  the  complex,  so  it  is 
evidently  not  essential  to  the  constitution  of  the  complex.  On  the  other 
hand,  the  carbonato  complex  does  undergo  aquation  in  water  solution,  first 
yielding  [Co(NH3)5HC03]++  and  then  [Co(NH3)5(H20)]+++  173.  The  anala- 
gous  ion,  [Co(NH3)4C03]+,  aquates  to  fCo(NH3)4(HC03)H20]++,  and  then 
to  [Co(XH3)4(H20)2]+++ 174.  Stranks  and  Harris175  studied  the  exchange  in 
solution  of  C-labelled  carbonate  with  the  carbonate  in  [Co(NH3)4C03]+ 
and  Yankwich  and  McXamara176  did  the  same  with  [Co  en2C03]+.  The 
exchange  takes  place  through  the  intermediate  formation  of  a  bicarbonate 
complex. 

By  using  labeled  oxygen,  Taube  and  his  students  demonstrated  that  in 
the  cases  of  [Co(NH3)6C03]+  and  [Co(NH2)4C08]+  exchange  does  not  in- 
volve rupture  of  the  cobalt-oxygen  link,  but  rather,  of  the  carbon-oxygen 
bond177. 

169.  Haldar,  Smnr,  and  Culture,  14,  340-1  (1949);  Nature,  166,  744  (1950). 

170.  Watters  and  Aaron,  ./.  .1///.  Chem.  S<>c,  75,  611  (1953). 

171.  Werner,  Ber.  40,  4101   (1907 

172    Lamb  and  Mysels,/.  .1///.  Chem.  Soc.,  67,  468  (1945). 
17::    I. ami)  and  Stevens,  •/.  .1///.  Chem.  So,-.,  61,  3229  (1939). 
17  1    ll.ii  lis  ;in(l  Si  tanks.  Trans.  Faraday  Soc.  48,  137  (1952). 
177,    Stranks  and  Harris.  ./ .  Chem.  Phye.,  19,  267  ^  1951). 
L76    Vankwich  and  McNamara,/.  Chem.  Phys., 20,  1325  (1952 
177.  Hunt.  Rutenberg,  and  Taube,  ./.  Am.  Chem.  Soc,  74,  268  (1952);  Posey  and 
Taube, ./.  .1///.  Chem.  Soc.  75,  4099    i" 


GENERAL  SURVEY 


33 


McCutcheon  and  Schuele178  bave  recently  isolated  the  interesting  ion 
[Co(C03)i]"  as  the  hexamminecobalt(III)  salt;  its  existence  clearly  indi- 
cates thai  tin1  carbonate  ion  can  fill  two  coordination  positions. 

Organic  inion  Coordination 

Many  organic  anions  form  stable  coordination  compounds.  Formate  and 
acetate  ions  form  strong  bonds,  but  monocarboxylic  acids  with  Longer  chains 
show  a  rapidly  decreasing  ability  to  coordinate.  Formate  and  acetate  often 
bind  two  metal  atoms  together,  each  oxygen  of  the  carboxy]  group  linking 

to  a  different  metal  atom. 

R 


M— OC=0— M. 

When  the  carboxy]  group  is  attached  to  only  one  metal  atom,  however,  it 
tills  but  one  position  in  the  coordination  sphere.  Complexes  of  the  types 
[Co(NH;  *OOCCH,]++  m  and  [Co(XH3)5OOCH]++  180  are  well  known  and 
easily  prepared.  The  solubilities1"1  and  stabilities182  of  several  similar  com- 
plexes containing  a  variety  of  aliphatic  anions  have  been  studied. 

•  and  Bailar183  were  able  to  effect  a  partial  resolution  of  a-chloropro- 
pionic  and  a-bromopropionic  acids  through  the  formation  of  stable  cobalt 
<•(  tmplexes  containing  levo-propylenediamine,  [Co  ?-pn2(OOC  •  CHX  •  CH:j)o]+. 
The  solubility  of  lead  sulfate  in  solutions  of  sodium  acetate  has  inspired 
much  research,  and  many  formulas  have  been  postulated  for  the  complexes 
which  are  formed184.  Weinland  and  his  students121  report  the  isolation  of  the 
polynuclear  complex  ions 


/    \ 


Ph 


PI 


\ 


PI 


peva  and  Batyrshine186,  however,  report  only  the  formation  of  [Pbac]+, 
pPbacj]-,  and  [PbacJ",  the  last  being  the  most  important  in  analytical 
work. 

178.  McCutcheon  and  Schuele,  /.  Am.  Chem.  Soc,  75,  1845    H»53). 

179.  1  4,  171     1953). 

180    \  atsimirekii,  ./.  Gen.  Chem.    I    S.S  R.),  20,  140s    I960  . 

181.  Linhard  and  Rau,  Z.  anorg.  cUlgem.  Chem.,  271,  121  '1952). 

182.  Bunton  and  Llewellyn,  J.  CI         -        1953,  L6  - 

tnd  Bailar,  ./.  Am.  Chem.  Soc.,  74,  1820    L952). 
184.  Weinland,  "Einfuhrung  in  die  Chemie  der  Komplexverbindungen,"  Becond  Edi- 
tion, pp.  391    100,  Enke,  Stuttgart,  1924. 
Is-.").  Toropova  and  Batyrshina,  Zkur.  Anal.  Kkim.,  4,  337    194 


:;i  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

The  "basic  acetate"  method  of  separating  the  ions  of  the  trivalent  metals 
in  qualitative  analysis  involves  the  formation  of  acetate  complexes.  Wein- 
land  and  his  students  studied  many  of  these184  and  isolated  some  very 
complex  materials  which  they  thought  were  true  chemical  entities. 

Among  the  examples  in  which  the  carhoxyl  group  forms  a  bridge  between 
two  metal  atoms  are  the  "basic"  beryllium  salts,  Be40(OOC-R)6 ,  in  which 
R  represents  (II:;.  (YII5,  etc.  These  compounds  are  readily  formed  and 
are  stable,  volatile,  and  soluble  in  nonpolar  solvents.  Structural  studies186 
indicate  the  presence  of  a  central  oxygen  surrounded  tetrahedrally  by  four 
beryllium  ions.  Each  edge  of  the  tetrahedron  is  composed  of  the  grouping 

R 
Be — O — C — O — Be.     Similar     compounds     of     zinc187     and     zirconium, 
I  XrO)40(OOCR)6 ,188  are  known. 

The  oxalate  ion  forms  a  great  many  stable  coordinate  compounds,  usually 
acting  as  bidentate  group.  The  best  known  are  those  of  the  types 
!M"()x;J4-,  [MmOx3h  and  [MmOx2]-.  The  tris-(oxalato)  complexes 
have  been  studied  extensively,  especially  in  regard  to  their  stereochemistry. 
(Chapter  8).  The  oxalate  group  can  share  the  coordination  sphere  with 
ammonia,  ethylenediamine,  water,  or  other  groups.  Oxalatobis(ethylene- 
diammine)  cobalt(III)  chloride,  [Co  en2Ox]Cl,  is  readily  obtained  by  the 
action  of  an  alkali  oxalate  upon  the  dichloro  salt189;  the  corresponding 
chromium  salt  is  prepared  by  the  action  of  ethylenediamine  upon  the  tris- 
(oxalato)  salt190.  Hamm  and  Davis191  have  studied  the  formation  of  these 
ions  by  the  reaction  of  [Cr(H20)6]+++  and  oxalate  ion,  and  Hamm192  has 
followed  the  rate  of  isomerization  of  [Cr(H20)20x2]~  in  water  solution.  He 
postulates  that  upon  collision  with  the  ion,  a  water  molecule  knocks  one  end 
of  an  oxalate  group  away  from  the  chromium  and  takes  its  place;  on  return 
of  the  oxalate,  either  the  cis-  or  trans-  isomer  may  be  formed,  depending 
upon  which  molecule  of  water  is  eliminated.  A  small  amount  of  alkali  con- 
verts the  diaquo  compounds  to  hydroxoaquo-  compounds,  the  cis  isomer 

1S6.  Bragg  and  Morgan,  Proc.  Roy.  Soc.  London,  A104,  437  (1923);  Morgan  and  Ast- 

bury,  Proc.  Hoy.  Soc.  London,  A112,  441  (1926),  Pauling  and  Sherman,  Proc. 

Natl.  Acad.  Sri.,  20,  340  (1934). 
is?.  Auger  and  Robin,  Compt.  rend.,  178,  1546  (1924);  Wyart,  Bull.  Soc.  Fr.  Min.,  49, 

1  is  (1026). 
188.  Tanatar  and  Kurowski,  Chem.  Centralblatt,  1908  (1)  1523. 
L80.  Werner  and  Vilmos,  Z.  anorg.  Chem.,  21,  153  (1899);  Price  and  Brazier, ./.  Chem. 

Soc,  107,  1376,  1726  (1915). 
l'Mi    Werner  and  Schwarz,  .1////.,  405,  222  (191  \-. 
I'M.  Hamm  and  Davie, ./.  .1///.  Chem.  Sue.  75,  3085  (1953). 
192    Hamm,  ./.  Am.  Chem.  Soc,  75,  609  (1953). 


a i:\f-: ual  sritVEY 


35 


of  which  is  converted  upon  heating  into  the  tetrakis(oxalato)-M-diol-salt, 

OH 


M4 


Ox,Cr  CrOx, 

\       / 

oil 


Larger  amounts  of  alkali  change  the  diaquo  salts  to  dihydroxo  salts,  still 
without  breaking  the  chromium-oxalate  linkage 

Weinland  and  Paul1--  have  isolated  several  compounds  of  the  ion 
[Pr>Ox]+~,  in  which  all  four  of  the  oxygen  atoms  are  probably  bonded  to 
the  metal: 


/ 


0— c=o 


Pb 


Pb 


O— C^O 


Solubility  studies193  have  indicated  the  existence  of  analagous  ions  of  zinc 
and  cadmium. 

The  stability  of  the  oxalato  complexes  is  largely  due,  no  doubt,  to  the 
formation  of  five-membered  rings.  Compounds  are  known,  however,  in 
which  rings  are  not  formed.  Griinberg's  method  of  determining  the  con- 
figuration of  cis-trans  isomers  of  the  type  [Pt(XH3)2X2]194  is  based  upon  the 
inability  of  the  trans-isomer  to  yield  a  chelate  oxalato  derivative.  (See 
Chapter  9). 

The  oxalate  ion,  like  the  sulfate  ion,  forms  hydrates.  Werner  has  pointed 
out195  that  a  large  number  of  compounds  containing  complex  oxalate  anions 
crystallize  with  water,  even  if  the  cation  is  one  which  is  usually  anhydrous. 

The  malonate  ion  coordinates  with  metallic  ions  to  give  a  six-membered 
ring,  which  is  not  as  stable  as  the  five-membered  ring  formed  from  the 
oxalate  ion.  Schramm  has  studied  the  formation  of  malonatotetrammine- 
cobalt(III)  compounds  in  some  detail196.  Anions  of  other  dibasic  organic 
acids  form  cations  of  the  type  [Co  en2A]+,  but  seem  unable  to  form  anionic 
complexes  like  those  formed  by  oxalates  and  malonates.  Complexes  of  some 
difunctional  acids  are  discussed  in  Chapter  6. 

a-Hydroxy  acids  often  coordinate  readily,  the  hydroxy]  and  carboxy] 
group  both  coordinating,  and  the  chelation  effect  enhancing  the  stability 


193.  Vosburgh  and  Beckman,  ./.  Am.  Chan.  Soc,  62,  1028  (1940). 

194.  Gr&nberg,  Helv.  ckim.  Acta,  14,  455  (1931). 

195.  Werner,  "New  Ideas  on  Inorganic  Chemistry,"  Translated  by 

London,  Longmans,  Green  &  Co.,  1911. 

196.  Ref.  140,  p.  161. 


Hedley,  p.  113, 


36  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

of  the  compounds  formed.  The  hydrogen  of  the  hydroxyl  group  may  be  lost 
simultaneously,  so  that  the  organic  group  contributes  a  charge  of  minus 
two  to  the  complex.  Thus,  coordination  with  the  copper(II)  ion  gives 


The  copper  complexes  containing  glycollic  and  lactic  acids  are  not  very 
stable197  but  those  containing  the  stronger  salicylic  and  mandelic  acids  are 
easily  isolated198.  Boron  forms  stable  compounds  even  with  the  simpler 
a-hydroxy  acids199,  and  Boesken  and  his  co-workers  were  able  to  resolve  the 
bis-(a-hydroxybutyro)borate  ion200  as  well  as  the  bis(salicylato)borate  ion201. 
The  work  of  Jantsch202  on  the  rare  earth  glycolates  and  lactates  indicates 
that  some  chelation  takes  place.  His  values  for  the  equivalent  conductances 
of  various  lanthanum  salts  are  as  follows: 

v  X 

acetate  1024  89.5 

phenylacetate  1200  91.2 

glycolate  1200  70.3 

lactate  1024  54.1 

Salicylate  ion  differs  from  its  meta-  and  para-  isomers  in  being  able  to 
form  chelate  rings,  which  greatly  stabilizes  its  coordination203.  Many  recent 
studies  have  been  made  on  solutions  of  metal  ions  and  a-hydroxy  acids,  such 
as  salicylic,  lactic,  citric,  glycollic,  and  tartaric;  these  studies  lead  to  a 
knowledge  of  the  compositions  and  stabilities  of  the  complexes  formed,  but 
do  not  give  information  on  their  structures.  The  work  of  Bertin-Batsch  and 
of  Bobtelsky  and  his  collaborators204  is  typical. 

The  compounds  of  the  a-amino  acids  are  of  great  stability,  and  have  re- 
ceived extensive  study.  Ley205  and  Bruni  and  Fornara206  suggested  that 

197.  Wark,  J.  Chem.  Soc,  123,  1815  (1923). 

198.  Wark,  J.  Chem.  Soc,  1927,  1753. 

L99.  Rosenheim  and  Vermehren,  Ber.,  57,  1337  (1924). 

200.  Boeseken,  Muller,  and  Japhongjouw,  Rec.  trav.  chim.,  45,  919  (1926). 

201.  Boeseken  and  Meulenhoff,  Proc.  Acad.  Set.  Amsterdam,  27,  174  (1924). 

202.  Jantsch,  Z.  anorg.  allgem.  Chem.,  153,  9  (1926);  Jantsch  and  Griinkraut,  Z.  anorg. 

allgem.  Chem.,  79,  305  (1913). 

203.  Bertin-Batsch,  Ann.  chim.,  7,  481  (1952). 

Jin  Bobtelsky  and  Eeitner,  Bull.  soc.  chim.  France,  1951,  494;  Bobtelsky  and  Graus, 
J.  A  m  <  "h<  »i .  Soc,  75,  4172  (1953) ;  Bobtelsky  and  Bar-Gadda,  Bull.  soc.  chim. 
Franc*  ,  1953,  276,  687. 

205.  Ley,  Z.  Elektrochem.,  10,  954  (1904). 

206.  Bruni  and  Fornara,  Aiti  accad.  Lincei,  [5]  13,  II,  26  (1904);  Bruni,  Z.  Elektro- 

chem., 11,  93  (1905). 


GENERAL  SURVEY  M 

copper  glycine  is  an  inner  complex.  The  deep  blue  color  of  the  compound 
indicates  copper-nitrogen  linkages,  and  the  possibility  of  the  formula 
CiuXIICIU'ooiu  is  eliminated  by  the  facl  thai  N,N-diethylglycine 
gives  an  analagous  compound.  The  compound  is  a  nonelectrolyte,  and  i1  is 
evident  thai  the  copper  is  coordinately  saturated,  for  it  absorbs  ammonia 
only  very  slowly.  Finally,  the  properties  of  copper  glycine  are  very  similar 
to  those  of  diamminecopper(II)  acetate  [Cu(OOCCH3)2(NH3)2],  which 
seems  to  justify  the  formula 


The  copper(II)  compounds  of  a-amino  acids  are  so  stable  that  they  do  not 
respond  to  most  of  the  usual  tests  for  copper(II)  ion.  Hydrogen  sulfide  de- 
posits copper  sulfide,  and  boiling  alkalies  precipitate  copper  oxide,  but  both 
reactions  take  place  slowly.  The  opening  of  the  ring  by  ammonia  to  give 
[Cu(XH3)2(OOCCH2XH2)2]207  is  an  interesting  reaction.  The  remarkable 
stability  of  the  copper  chelate  of  the  a-amino  acid  group  is  illustrated  by 
the  work  of  Kurtz208  who  studied  several  acids  of  the  type 

XHo— (CH2)Z— CH— COOH, 

I 
NHs 

where  X  =  2,  3,  or  4  (a , 7-diaminobutyric  acid,  ornithine,  and  lysine).  In 
each  case  the  usual  properties  of  the  carboxyl  group  and  the  adjacent 
amino  group  are  completely  masked,  but  the  other  amino  group  retains  its 
characteristic  behavior,  and  Kurtz  was  able  to  carry  out  reactions  on  it, 
without  affecting  the  coordinated  amino  group. 

The  cobalt  complexes  of  the  a-amino  acids,  [Coamac3],  exist  in  two  stereo- 
isomeric  forms  (see  page  283),  both  of  which  are  remarkably  stable,  being 
unat tacked  by  50  per  cent  sulfuric  acid.  Elliott209  has  utilized  this  stability 
in  the  preparation  of  highly  insoluble  and  stable  "super  complexes"  by  the 
reaction  of  cobalt  (III)  hydroxide  with 

IK  ><  >c— CH— (CH2)n— CH— COOH 
I  I 

NHS  XII: 

Chromium(III)  forms  inner  complexes  which  are  similar  but  of  less  sta- 
bility; they  are  .-lowly  decomposed  by  hot  acids,  by  sodium  hydroxide,  and 

Ley,  Ber.,  42,  354  (1909). 

208.  Kurtz.  ./.  Biol.  Chem.,  122,  177  (1937-8);  180,  1253  (1949). 

209.  Elliott,  thesis,  University  of  Illinois,  1943. 


38 


CHEMISTRY  OF  THE  ('OOEI)I XATIOX  COMPOUNDS 


to  ;i  degree,  by  foiling  water.  Keller210  has  studied  the  reactions  of  a  large 
number  of  a-amino  acids  with  chromium  (III)  hydroxide  and  chromam- 
mines  in  boiling  water.  In  all  cases  compounds  of  the  formula  [Cr(amac)3] 
aeem  t<>  form,  l>ut  are  quickly  hydrolyzed  to 

OH 

/       \ 
amac2Cr  Cramac2 

\       / 

which  in  turn  hydrolyze  slowly  to 

OH    OH    OH 

/       \l/       \ 
amacoCr  Cr  Cramac2 

\       /l\       / 

OH    OH    OH 

and  more  complex  products.  Cobalt  amino  acid  compounds  undergo  the 
same  reactions,  but  much  more  slowly. 

Platinum  does  not  readily  coordinate  with  oxygen,  but  the  coordinating 
tendency  of  the  a-amino  acids  is  so  great  that  such  compounds  as 


K 


PtCl; 


0 ( 

:=o" 

and 

NH2— ( 

}H2    _ 

O  C=0> 


CH- 


can  be  formed211,212,213.  Even  a-amino  acids  containing  tertiary  nitrogen 
atoms  will  coordinate  Avith  platinum  strongly,  as  is  shown  by  the  optical 
resolution  of  the  ion 


(N02)2Pt 


\ 

CHj     C2H5 


c=o 

I 

CH, 


Heterocyclic  acids  having  a  carboxyl  group  in  the  a-position  to  the  ring 
nitrogen  (picolinic,  quinolinic,  quinaldinic,  etc.)  form  inner  complexes.  The 
compounds  with  iron(II),  which  arc  deeply  colored,  have  been  studied  by 

210.  Keller,  thesis.  University  of  Illinois,  L940. 

211.  Ley  and  Picken,  />'</•.,  45,  377  (1912). 

212    I  Irinberg  and  Ptitzuin,  .1////.  inst.  platine,  No.  9,  55  (1932). 
Grinbergand  Ptitzuin,  Am,,  inst.  platine,  No.  n.  77  (1933). 
_•]  l    Kueblerand  Bailar,  J.  Am.  Ckem.  Sac,  74,  3535  (1952). 


GENERAL  SURVEY 


39 


Ley  and  his  co-workers-1'.  The  corresponding  copper(II)  compounds  are 
light  in  color,  and  are  probably  not  coordination  compounds. 

The  fi-amino  acids  also  form  inner  complexes  with  the  transition  metals, 
hut  these  are  less  stable  than  those  of  the  a-acids.  Hearn218  has  shown  that 
a-amino  acids  can  be  distinguished  from  the  0-aeids  by  the  fact  that  the 
former  react  with  cobalt  (III)  hydroxide  to  give  colored  complexes,  while 
the  latter  do  not. 

The  y-,  5-,  and  e-amino  acids  do  not  form  chelate  rings  with  metals,  so 
form  normal  salts217. 

Among  the  amino  acids,  the  derivatives  of  acetic  acid  are  particularly 
noteworthy  for  their  chelating  ability.  The  tridentate  iminodiacetic  acid 
gives  many  complexes,  which  in  general  are  more  stable  than  those  of  gly- 
cine. For  example,  the  first  and  second  stability  constants  of  the  zinc  com- 
plex of  glycine  are  4.8  and  4.1,  while  for  the  zinc  complex  of  iminodiacetic 
acid  they  are  7.8  and  5.7218.  Nitrilotriacetic  acid  forms  still  more  stable 
complexes,  the  two  dissociation  constants  for  the  zinc  complex  being  10.5 
and  3.0219.  The  great  difference  between  the  two  values  in  the  case  of  the 
triacetic  acid  doubtless  reflects  the  fact  that  the  zinc  ion  cannot  accept  all 
of  the  possible  donor  groups  in  two  of  the  donor  anions.  The  complex  which 
is  formed  in  this  case220  is 


-i4  — 


OOC-CH2—  N 


The  most  remarkable  of  the  acetic  acid  derivatives,  however,  is  ethylene- 
diaminetetraacetic  acid  (often  abbreviated  EDTA  or  H4Y).  This  substance 
is  potentially  hexadentate,  but  complexes  in  which  only  four  or  five  groups 
are  coordinating  are  well  known.  The  complexes  of  EDTA  are  remarkably 
stable,  so  have  been  investigated  extensively  from  the  industrial  point  of 

215.  Ley,  Schwarte,  and  Miinnich,  Ber.,  57,  349  (1924). 

216.  Hearn,  thesis,  University  of  Illinois,  1951. 

217.  Tschugaeff  and  Serbin,  Compt.  rend.,  151,  1361  (1910);  Pfeiffer  and  Lubbe,  ./. 

prakt.  Chem.,  [2]  136,  321  (1933). 

218.  Flood  and  Loras,  Tids.  Kjemi,  Bergsvesen  Met.,  6,  83  (1945). 
JIM.  Schwarzenbach,  Chimin,  3,  1  (1949). 

220.  Schwarzenbach  and  Biedeimann,  Eelv.  Ckitn.  Ada,  31,  331  (1948). 


in 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


\  i.w  .  More  than  four  hundred  and  fifty  articles  were  published  during  1952 
describing  uses  of  this  reagenl  or  stability  constants  of  its  metal  derivatives. 
1 1  has  been  used  in  water  softening  (Chapter  23),  electroplating,  controlling 
the  metal  contenl  of  dye  baths,  in  removing  lead  and  other  heavy  metals 
from  the  human  Bystem,  in  the  treatment  of  chlorosis  in  plants,  and  in  many 
other  ways, 

The  stability  of  the  EDTA  complexes  is  illustrated  by  the  fact  that 
neither  the  copper(II)  or  the  nickel  compound  is  destroyed  by  sodium  or 
ammonium  hydroxide.  The  nickel  compound  is  not  attacked  by  dimethyl- 
glyoxime  or  hydrogen  sulfide,  but  is  destroyed  by  potassium  cyanide.  The 
copper  compound  gives  the  usual  reactions  of  Cu++  when  treated  with  potas- 
sium cyanide,  hydrogen  sulfide,  or  potassium  ferrocyanide221. 

The  ability  of  ethylenediaminetetraacetic  acid  to  form  stable  complexes 
depends  upon  the  fact  that  when  it  coordinates  it  forms  multiple  fused  five- 
m< -inhered  chelate  rings.  Pfeiffer  and  Simons222  compared  the  calcium  deri- 
vatives  of  methylaminediacetic  acid 


CHoCOO^ 


CH8— N 


\ 


Ca 


CHoCOO/  2 
and  ethylenediaminetetraacetic  acid, 

cir-cocr 


ff. 


-f-CH2— N 


Ca 


CH2COO 


Hs 


which  differ  only  in  that  the  two  nitrogen  atoms  in  the  latter  are  linked 
together  through  the  ethylene  bridge.  The  methylamine  complex  reacts 
slowly  with  oxalate  ion  to  precipitate  calcium  oxalate,  but  the  ethyl- 
enediamine  complex  does  not.  Pfeiffer  and  Simons  came  to  the  conclu- 
sion that  these  complexes  are  hexadentate,  for  the  structurally  similar 
I  K  M  )CCH(CH3)NIICH2CH2NCH(CH3)COOH  does  not  form  a  stable  cal- 
cium complex. 

Several  studies  have  been  made  of  the  effect  of  ring  size  on  the  stability 

of  complexes  of  this  type.  Schwarzenbach  and  Ackermann223  investigated 

i  iee    1 1« »«  x  JCH2)2N(CH2)nN(CH2COOH)2 ,  where  n  varies  from  two 

to  five.  In  general,  the  stability  of  the  alkaline  earth  compounds  decreases 

increases.  When  "n"  is  4  or  ~>,  the  two  ends  of  the  molecule  seem 

221.  Brintzinger  and  Hesse,  /.  anorg.  allgem.  Chem.,  249,  113  (1942). 
Pfeiffer  and  Simons,  Ber.,76B,847  (1943). 
3<  bwaraenbach  and   Ackermann,  Help.  Ckim.  Ada,  31,  1029  (1948). 


GENERAL  SURVEY  41 

able  to  act  independently,  for  complexes  of  the  type  M-Y  can  be  formed. 
Chaberek  and  Mart  el  l-'-'1  found  the  stabilities  of  the  complexes  of  ethylene- 
diaininediacetic-dipropionic  acid  to  be  considerably  less  than  those  of  the 
tetraacetic  acid. 

Some  ca>es  are  known  in  which  EDTAdoes  not  act  as  a  hexadentate  co- 
ordinator, even  though  six  positions  are  open  to  it.  Thus,  Schwarzenbach228 
prepared  the  compounds  [CoHYBr]-  and  [CoHY(N02)]~.  Removal  of  the 
bromide  or  oitro  group  allows  the  unattached  carboxyl  group  to  coordinate 
with  the  cobalt  to  form  [CoY]~.  Busch'-"-6  has  shown  that  the  palladium(II) 
chelate  has  the  structure 

CH2-CH2\ 
HpC7/  Pd  /    CH2 

_  o=c — o' 'O  —  c=o 

The  stereochemistry  of  the  EDTA  complexes  is  discussed  in  Chapter  8. 
Carbonyl  Coordination 

The  carbonyl  group  of  aldehydes  has  rather  weak  donor  properties,  but 
addition  compounds  of  aldehydes  with  several  of  the  light  metals,  such  as 
magnesium227,  and  with  the  wreakly  basic  elements,  such  as  tin  and  anti- 
mony,*28 are  known.  The  carbonyl  group  of  esters  also  forms  rather  weak 
coordinate  links  with  these  metals229.  Simple  aliphatic  ketones  show  similar 
behavior. 

The  1,3-dicarbonyl  compounds,  through  their  ability  to  enolize,  form 
stable  chelate  rings  with  a  large  number  of  metals.  In  many  cases  the  com- 
pounds so  obtained  are  nonionic,  insoluble  in  water,  soluble  in  nonpolar 
solvents,  and  volatile.  Acetylacetone  has  received  the  most  attention  in 
this  regard,  but  dibenzoyl  methane,  benzoylacetone,  acetoacetic  ester,  sali- 
cylaldehyde,  benzoyl  pyruvic  acid,  and  o-hydroxyacetone  are  important. 
Thenoyltrifluoroacetone  (TTA), 

O  O 


"C      CHo — C — CF; 


. 


224.  Chaberek  and  Martell,  J.  Am.  Chem.  Soc,  74,  6228  (1952). 

225.  Schwarzenbach,  Helv.  Chim.  Acta,  32,  839  (1949). 

226.  Busch  and  Bailer,  J.  Am.  Chem.  Soc,  in  press,  1956. 

227.  Menschutkin,  Izvest.  St.  Petersburg  Polyttch.  Inst.,  6,  39  (1906). 

228.  Menschutkin,  ./.  Russ.  Phys.  Chem.  Soc.,  44,  1929  (1912);  Rosenheim  and  Soil 

man,  Ber.,  34,  3377  (1901);  PfeifTer,  Ann.,  376,  296  (1910). 

229.  Menschutkin,  Izvest.  St.  Petersburg  Polytech.  Inst.,  4,  101  (1906);  6,  L01     L906 

Lewy,  J.  prakt.  Chem.,  37,  480  (1846). 


42 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


has  received  much  attention  because  of  the  great  stability  of  its  compounds. 
The  classic  paper  of  Morgan  and  Moss  on  the  acetylacetone  compounds230 
reviews  the  Literature  up  to  1914  and  describes  the  preparation  of  many 
compounds.  Metallic  ions  having  a  coordination  number  twice  the  ionic 
charge  give  nonelectrolytic  complexes: 


"C—G 

CH3 
M=Be,Cu,Ni,ETC. 


CH3 
M=AI,Cr,r%CcvETC 


CH3 
M=Th.  Zr,Hf,Ce,Pu,E-rc . 


Many  of  these  compounds  show  exceptional  stability,  the  beryllium  com- 
plex, for  example,  boiling  without  decomposition  at  270°C  at  atmospheric 
pressure.  Molecular  weight  determinations  indicate  that  these  compounds 
are  monomeric.  Wilkins  and  Wittbecker231  have  utilized  this  stability  in 
the  preparation  of  beryllium  containing  polymers.  They  report  that  tet- 
rake tones  form  linear  polymers  of  the  types 


R  R 

B  C_Y_C 

'     o  —  c/  x  —  o 

R  R 


Be: 


•o=c 


R 


>  — 


c>=°>e/ 


AND 


x  y°=c  — 

Be  ,CH 


/     \ 


o— c; 


— /C 
R 


Be  CH 

■o'  Nj-c; 


—  c=0\  / 

HC  •     Be 

*C  —  o'    \ 


where  Y  is  any  one  of  a  variety  of  organic  groups. 

There  is,  however,  some  popular  misconception  as  to  the  stability  of  the 
diketone  chelates.  The  statements  that  the  rare  earths  can  be  separated 
through  the  volatility  of  their  acetylacetonates232,  and  that  the  molecular 
weights  of  the  rare  earth  acetylacetonates  can  be  determined  by  their  vapor 


230.  Morgan  and  Moss,  ./.  Chem.  Soc,  105,  189  (1914). 

231.  Wilkins  and  Wittbecker,  U.  S.  Patent  2,659,711  (Nov.  17,  1953). 

232    Bphraim,  "Inorganic  Chemistry,"  English  Edition  by  Thome,  London,  Gurney 
and  Jackson,  L926. 


<,i:\i:i;.\l  si'HVK) 


\:\ 


densities18'  are  incorrect     scandium  acetylacetonate  is  readily  volatile280,  **, 

hut  those  oi  the  true  rare  earths  decompose  on  heating288,288,  Brimm288 
found  that  the  rare  earth  compounds  of  dibenzoylmethane  and  benzoylace- 
tone  are  readily  decomposed  by  traces  of  moisture  with  the  formation  of 

[M(dik(it<>nrM(  HI  )<  II-( ))],  These  compounds  are  soluble  in  organic  solv- 
ents, hut  are  not  volatile. 

When  the  coordination  Dumber  of  the  central  ion  is  less  than  twice  the 
elect rovalence,  cat  ionic  compounds  are  formed,  as  illustrated  by  the  com- 
pounds containing  boron,  silicon  and  titanium287 


M=Si,Ti 

These  compounds  are  of  special  interest  because  of  their  stereochemical 
possibilil  ies  and  because  they  show  typical  metalloid  elements  in  the  role  of 
cations.  Similar  compounds  of  other  1 ,3-diketones  have  been  described238. 
If,  on  the  other  hand,  the  coordination  number  of  the  central  atom  is 
more  than  twice  the  electrovalence,  the  coordination  sphere  will  tend  to  fill 
itself  with  other  neutral  groups237.  Iron(II)  forms  the  compounds 


Y=NH3,pq,£en,(t)NHNH2, 

PIPERIDfNE,  NICOTINE 


all  of  which  are  soluble  in  organic  solvents,  insoluble  in  water,  and  deeply 
colored-'.  On  heating  in  vacuo  the  ammonia  compound  is  converted  to 
dibenzoylmethane  iron. 

233.  Hein,  "Chemische  Koordinationtheorie,"  p.  153,  Zurich,  Hirzel  Verlag,  1050. 

234.  Meyer  and  Winter,  Z.  anorg.  Chem.,  67,  414  (1910). 

235.  [Jrbain,  Ann.  ckim.,  [7]  19,  212  (1900). 

236.  Brimm,  thesis,  University  of  Illinois,  1940. 

237.  Dilthey,  Ber.,  36,  923  (1003);  37,  588  (1904);  .1////.,  344,  300  .1905). 

238.  Dilthe:     /;■       36.  1595   3207  (1903);  ./.  prdkt.  Chem.,  [2]  111,  147  (1925). 

239.  Emmerl  and  Gsottschneider.  Ber..  66,  L871   (1933). 


I  }  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

2,  t-Pentanediono-dimethyl  thallium 


/CH3 
CH3  ^0=C 

> 


civ  o— c 


XCH3 


has  unusual  properties240.  It  is  soluble  in  benzene,  has  a  low  melting  point, 
and  sublimes  readily.  On  the  other  hand,  it  is  also  soluble  in  water,  giving 
an  alkaline  solution.  This  solution  shows  the  usual  properties  of  the  di- 
methyl thallium  ion,  so  it  appears  that  the  coordinate  bonds  are  broken  by 
water. 

The  diketone  compounds  which  are  soluble  in  organic  compounds  have 
achieved  considerable  importance  as  agents  for  the  separation  of  metal  ions 
through  the  techniques  of  solvent  extraction.  If  two  metals  in  aqueous  solu- 
tion, are  in  equilibrium  with  a  diketone,  if  the  equilibrium  constants  are 
different  and  if  the  complexes  are  soluble  in  a  solvent  immiscible  with 
water,  the  metals  can  be  separated  by  liquid-liquid  extraction241.  Since  the 
extent  of  dissociation  of  the  complex  of  any  metal  can  be  changed  by  chang- 
ing the  pH  of  the  solution,  the  method  is  widely  applicable.  If  a  speci- 
fied metal  is  to  be  separated  from  several  others,  the  pH  is  adjusted  so 
that  that  metal  (and  those  with  smaller  dissociation  constants)  will  be 
extracted  into  the  organic  layer.  This  is  then  extracted  with  water,  the 
pH  of  which  is  adjusted  to  allow  only  the  extraction  of  the  metal  in  ques- 
tion, since  its  complex  has  the  largest  dissociation  constant  of  those  now 
present.  Bolomey  and  Wish242  used  this  technique  to  separate  radioberyl- 
lium  from  the  other  metals  obtained  with  it  by  cyclotron  bombardment. 
Huffman  and  Beaufait243  employed  the  method  to  separate  zirconium  and 
hafnium,  using  thenoyltrifluoroacetone  as  the  complex  former.  The  dis- 
tribution  coefficient  of  the  zirconium  complex  is  about  twenty  times  that 
of  i  lie  hafnium  complex,  so  excellent  separation  was  achieved. 

This  extraction  technique  can  also  be  used  to  determine  the  formulas  of 
complexes  and  the  degree  of  hydrolysis  of  metal  ions  in  aqueous  solution, 
as  was  shown  by  ( Jonnick  and  McYey  in  their  study  of  the  zirconium  ion244. 
By  determining  the  extraction  coefficient  of  the  zirconium  complex  of  then- 

240.  Menziee,  Sidgwick,  Fox,  and  Cutliffe,  ./.  Chem.  Soc,  1928,  1288. 

241.  ( lalvin,  Manhattan  Project  Report  CN-2486,  December  1944;  Experientia,  6, 135 

(1950). 

242.  Bolomey  and  Wish,./.  .1///.  Chem.  Soc,  72,  4483  (1950). 
213.  Huffman  and  Beaufait,/.  .1///.  Chem.  Soc,  71,  3179  (1949). 
244.  Connick  and  McVey,  J,  A»,.  Chem.  Soc,  71,3182  (1949). 


GENERAL  SURVEY  45 

oyltrifluoroacetone  between  benzene  and  water  as  a  function  of  the  TTA 
activity  in  benzene,  they  were  able  to  establish  the  composition  of  the  che- 
late as  [Zr(TTA).i].  By  measuring  the  distribution  of  the  zirconium  between 
the  benzene  and  water  phases  as  a  function  of  pH,  they  then  demonstrated 

that  in  the  pi  I  range  —0.4  to  2.0,  the  zirconium  ion  exists  largely  as  a  mix- 
ture of  Zr*+  and  Zr(OH)+++ 

Steinbach  and  Preiser248  have  suggested  that  the  complexing  agenl 
(acetylacetone,  in  their  example)  can  serve  also  as  the  solvent  for  the 
complex.  Using  this  technique,  they  have  effected  the  analytical  separation 
of  zinc  and  copper  ions. 

Oxygen  Carrying  Chelates 

Hemoglobin  and  hemocyanin  were  long  considered  to  be  unique  in  their 
ability  to  absorb  and  release  oxygen,  but  several  types  of  synthetic  com- 
pounds are  now  known  which  possess  this  property.  Their  behavior  is  illus- 
trated by  a  simple  experiment:  If  cobalt  nitrate  solution  is  treated  with 
ammonium  chloride  and  ammonium  hydroxide  in  the  absence  of  air,  a  pink 
precipitate  forms.  When  air  is  bubbled  through  the  suspension,  a  brown 
color  develops,  but  when  nitrogen  is  substituted  for  the  air,  the  pink  color 
returns.  This  cycle  can  be  repeated  many  times.  Interestingly  enough,  the 
experiment  fails  if  ethylenediamine  is  substituted  for  ammonia. 

Pfeiffer,  Breith,  Lubbe,  and  Tsumaki246  reported  that  bis-(salicylal)ethyl- 
enediiminecobalt(II) 

Qv 

CH=lsT       XN 

I         I 

CH2-CH2 

(A) 

darkens  in  air.  Tsumaki-'47  found  that  this  is  due  to  absorption  of  oxygen 
and  thai  the  process  is  reversible.  It  has  since  been  found  that  other  cobalt 
chelates  also  show  this  property.  ( !alvin  and  his  students  and  Diehl  and  his 
students  have  studied  compound  (A)  and  many  derivatives  of  it.  Diehl248 
reports  thai  the  parent  compound  contains  one-half  mole  of  water  per  co- 
balt atom,  and  believes  thai  two  molecules  of  the  chelate  are  held  together 

245.  Steinbach  and  Preiser,  Anal.  Chem.,25,  881  (1053). 

246.  Pfeiffer,  Breith,  Lubbe,  and  Tsumaki,  Ann.,  503,  si  (1933). 

247.  Tsumaki,  Bull.  Ch*  Japan,  13,  252    L938 

248.  Diehl  and  co  workers    Bach,  Harrison,  Liggett,  Chao,  Brouns,  Curtis,  Bensel- 

meir,  Schwandl ,  Mathews  .  Iowa  Sim,  Coll.  J.  Sri.,  21,  271,  278,  287,  311,  316, 
326,  335  (1047);  22,  91,  110,  126,  129,  141,  150,  165  (1948);  23,  27:;    1949 


46  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

by  an  aquo  bridge.  This  is  a  unique  situation,  for  no  other  cases  of  aquo 
bridges  are  known.  Calvin  and  his  group249  have  studied  compound  (A)  and 
some  of  its  derivatives  from  the  structural  point  of  view.  Both  Calvin  and 
Diehl  report  that  most  of  these  compounds  exist  in  several  different  isomeric 
forms,  only  one  of  which  (for  each  compound)  is  active  toward  oxygen. 
Compound  (A)  is  paramagnetic,  apparently  having  one  unpaired  electron 
per  cobalt  atom.  Diehl  reports  that  it  does  not  absorb  carbon  monoxide  or 
nitrous  oxide,  but  that  it  absorbs  nitric  oxide  and  nitrogen  dioxide.  He  is  of 
the  opinion  that  it  will  absorb  other  paramagnetic  gases,  but  not  diamag- 
netic  ones. 

When  put  under  pressure  of  oxygen,  these  materials,  either  in  the  solid 
state  or  in  solution  in  quinoline  or  similar  solvents,  absorb  one  mole  of  oxy- 
gen for  each  two  moles  of  chelate,  and  release  it  again  when  the  pressure  is 
decreased.  In  each  repetition  of  the  cycle,  however,  there  is  a  small  amount 
of  irreversible  oxidation,  so  the  ability  to  absorb  oxygen  gradually  de- 
creases. 

Calvin's  group  also  prepared  compound  (B) 

CK<-p 

(CH2)3-NH  — (CH2)3 
(B) 

and  several  analogs  of  it.  Compound  (B)  has  three  unpaired  electrons  per 
cobalt  atom,  and  reversibly  absorbs  one  mole  of  oxygen  per  atom  of  co- 
balt250. 

Calvin's  x-ray  studies  on  compound  (A)  show  that  it  crystallizes  in  layers, 
with  holes  running  through  the  layers.  These  holes  are  big  enough  to  con- 
tain oxygen  molecules,  and  the  passages  between  them,  while  smaller,  are 
sufficiently  large  to  allow  such  molecules  to  go  through  without  great  diffi- 
culty. 

Cobalt (II)  histidine  chelates  in  water  solution  will  absorb  oxygen 
reversibly261.  Histidine  compounds  of  iron  are  oxidized  irreversibly,  while 
those  of  nickel  and  copper  are  not  oxidized  at  all.  The  unoxygenated  cobalt 
histidine  complex  is  paramagnetic  to  the  extent  of  three  unpaired  electrons 
per  cobalt  atom,  while  the  oxygenated  compound  is  diamagnetic.  Hearon  is 
of  the  opinion  that  cobalt  is  four  covalent  in  this  compound,  and  that  the 

249.  Calvin  and  co-workera  (Bailee,  Wilmarth,  Barkelew,  Aranoff,  Hughes),  J.  .1///. 

Ch  m.  Snr.,  68,  2254,  2257,  2263,  2267,  2273  (1946). 
260.  Harle  and  Calvin,  J.  Am.  Chem.  Soc. ,  68,  2612  (1946). 


GENERAL  SURVEY  47 

amino  acid  is  coordinated  to  the  metal  only  through  nitrogen  atoms 


Two  molecules  of  this  chelate  absorb  one  molecule  of  oxygen.  It  does  not 
combine  with  carbon  monoxide.  According  to  Hearon251d- e,  the  oxygenated 
molecule  has  either  the  structure 


(g  is  a  molecule  of  water 
or  some  other  neutral 
group) 


OR 


Michaelis252  has  also  measured  the  magnetic  susceptibility  of  the  cobalt 
histidine  compounds. 

The  properties  of  hemoglobin  and  its  oxygen  carrying  capacity  are  dis- 
cussed in  Chapter  21.  Like  the  other  oxygen  carrying  chelates,  it  is  para- 
magnetic when  deoxygenated,  but  diamagnetic  in  the  oxygenated  form253. 
A-  is  well  known,  it  combines  with  carbon  monoxide  more  firmly  than  with 
oxygen,  and  with  cyanide  ion  or  pyridine  still  more  firmly. 

The  Doxor  Properties  of  Sulfur 

The  donor  properties  of  sulfur  are  quite  different  from  those  of  oxygen. 
In  general,  they  are  somewhat  more  restricted  as  regards  the  nature  of  the 
acceptor  atom,  but   in  some  types  of  compounds,  they  are  exceptionally 

251.  Burk.  Bearon,  Caroline,  and  Schade, ./.  Biol.  Chem.,  165,  723  (1946);  Burke,  I! 

ron,  Levy,  and  Schade,  Federation  Proc.,  6,  212  (1947  ;  Hearon,  Federation 
.  6,  256  260   L947  :./.  Nat.  Cancer  Inst.,  9,  1     L94S  ;  Hearon,  Burk,  and 
Schade,/.  Natl.  Cancer  Inst.,  9,  :>>:>>:    1049). 
Michaelis,  Arch.  Biochem.,  14,  17  (1942). 
253.  Pauling  and  Coryell,  Proc.  Natl.  Acad.  Set.,  22,  159,  210    L936). 


48  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

strong.  The  thioethers,  for  example,  form  much  more  stable  compounds 
than  the  corresponding  oxyethers.  The  coordination  of  sulfide  (or  hydro- 
sulfide)  ion  with  the  sulfides  of  arsenic,  antimony,  tin,  copper,  and  mercury 
is  well  known  and  is  of  great  importance  in  qualitative  analysis.  Similarly, 
the  preferential  coordination  of  sulfide  ion  plays  an  important  part  in  the 
metallurgies  of  copper  and  nickel.  The  Orford  process  exploits  the  ampho- 
teric behavior  of  copper  and  iron  toward  sulfide  in  the  separation  of  these 
metals  from  nickel.  The  separation  is  not  quantitative,  but  repetition  of 
the  process  gives  further  separation. 

Thiohydrate  Formation 

Liquid  hydrogen  sulfide  shows  little  resemblance  to  water  in  its  solvent 
properties254,  although  some  inorganic  salts  dissolve  in  it.  A  few  thiohy- 
drates  have  been  isolated255  • 256  ■ 257  and  thiohydrolysis  probably  takes  place 
through  the  formation  of  unstable  thiohydrates.  Morgan  and  Ledbury258 
concluded  that  organic  sulfides  coordinate  readily  with  those  metals  which 
occur  as  sulfides  in  nature,  or  which  form  very  stable  sulfides.  They  also 
found  that  the  reactions  of  metal  ions  with  dimethyldithiolethylene  show 
analogies  to  their  reactions  with  hydrogen  sulfide.  Thus,  copper(II)  and 
gold  (III)  chlorides,  which  are  readily  reduced  by  hydrogen  sulfide,  form 
the  compounds 

CH3  CH3 

/5-CH2  /S-f* 

CI2Cu  AND  CI3Au 

XS CH2  S  —  CH2 

CH3  CH3 

which  readily  revert  to  copper(I)  and  gold  (I)  compounds.  TschugaefT259 
found  that  of  the  dithioethers,  RS(CH2)nSR  (n  =  0,  1,  2,  3,  5),  only  the 
compounds  having  n  =  2  formed  stable,  well-characterized  chelates. 

Dithiane,  C4H8S2 ,  forms  complexes  with  the  ions  of  the  coinage  metals, 
platinum,  mercury,  and  cadmium260.  The  ratio  of  dithiane  to  metal  varies 

254.  Antony  and  Magri,  Gazz.  chim.  ital.,  35,  206  (1905). 

266.  Plotnikov,  ./.  Ruse.  Phys.  Chem.  Soc,  45,  1162  (1913). 

256.  Hill/,  and  Keunecke,  Z.  anorg.  allgem.  Chem.,  147,  171  (1925). 

267    Ralston  and  Wilkinson,  ./.  Am.  Chem.  Soc,  50,  258  (1928). 

268.  Morgan  and  Ledbury, ./.  Chem.  Soc,  121,  2882  (1922). 

260.  Tschugaeff,  Ber.,  41,  2222  (1908);  TschugaefT  and  Kobljanski,  Z.  anorg.  Chem., 

83,  8    L913); Tschugaeff,  Compi.  rend.,  154,  33  (1912);  Tschugaeff  and  Subbo- 

tin,  Ber.t  43,  1200  (1910). 
280    Bouknighl  and  Smith   /.  Am.  Chem.  Soc.,  81,  28  (1939). 


GENERAL  SURVEY 


49 


from  two  to  one,  as  in  2AgN<  I    C4H8S2,  toonetotwo,  as  in  AgNi ).  -2(  ,1 1  > 
The  cation  in  the  former  may  have  the  bridge  structure 


CIU'Il 


\ 


Ag— S 


S— Ag 


(II  (II 


Thioethers  and  Thiols 

Pfeiffer881  has  pointed  out  that  the  thioethers  show  a  strong  tendency  to 
unite  with  salts  of  such  metals  as  nickel,  copper,  and  zinc,  and,  especially 
with  those  of  platinum  and  palladium.  Diethyl  sulfide  reacts  with  plati- 
num(II)  chloride  to  give  three  compounds  of  the  empirical  formula 
Pt(SEt-_. (jClj  .  the  yellow  a-  and  0-isomers  being  the  trans  and  cis  com- 
pounds, respectively*1,  and  the  y-isomer  being  the  dimer  [Pt(SEt2)J 
[PtCl4]'263.  The  a-  and  /3-forms  are  easily  converted  into  each  other  by  crys- 
tallization from  suitable  solvents.  The  differences  between  these  a-  and 
£-  forms  are  so  much  greater  than  is  usually  shown  by  cis-trans  isomers  that 
Angell,  Drew,  and  Wardlaw  concluded  that  the  isomerism  is  structural 
rather  than  spatial264a.  They  proposed  the  formulas 


(«) 


Et2S 


...CI 
SEt2 

yPt  AND       (/S)     PL 


Et2SN     /CI 


CI 


'SEta 

''CI 


but  Drew  and  Wyatt2Wb  later  concluded  that  the  a-salt  has  the  trans  struc- 
ture : 


CI 


Pt 
/     \ 


CI 


B]  • 


The  great  differences  in  the  two  isomers  may  be  explained  on  the  basis  of 
the  strong  trans  influence  of  the  coordinated  sulfur. 

261.  Pfeiffer,  "Organische  Molekulverbindungen,"  p.  159,  Second  Edition,  Stuttgart, 

Enke,  1927. 

262.  Jensen,  Z.  anorg.  allgem.  Chem.,  225,  97,  115  (1935). 

263.  Tschugaeff  and  Benewolensky,  Z.  anorg.  Chem.,  82,  120  (1913);  Drew,  Preston, 

Wardlaw,  and  Wyatt,  ./.  Chem.  80c. ,  1933,  1294;  Cox,  Saenger  and  Wardlaw, 

./.  '  .  1934,  182. 

264a.  Angell,  Drew,  and  Wardlaw,  ./.  Cfo  m.  Soc.,  1930,  349 
264b.  Drew  and  Wyatt,  ./.  Chem.  8oe.t  1934,  56. 


50 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


The  ion  1 1 >t  (SEt2)4]++  i-s  unstable,  and  its  salts  with  simpler  anions  have 
not  been  isolated  in  the  solid  state.  The  iodide  apparently  cannot  exist  even 
in  solution286.  With  ions  such  as  [PtCl4]=  [PtCl6]=  and  [Pt(N02)4]=,  how- 
ever, it  forms  stable,  insoluble  salts.  Upon  heating  or  solution,  chloro- 
platinites  of  this  type  frequently  rearrange  to  a  mixture  of  the  a-  and  0- 
monomeric  forms: 


[Pt(SMe2)4][PtCl4]  -»  2[Pt(SMe2)2Cl2]263b'c 

The  chloroplatinate  decomposes  on  heating  to  give  a  mixture  of 
[Pt(SEt2)2Cl2]  and  [Pt(Et2S)2Cl4]263. 

Several  tetrahalides  of  the  type  [Pt(R2S)2X4]  are  known264, 265- 266.  Several 
of  them  have  been  shown  to  exist  in  a-  and  /3-forms,  which  are  readily 
interconvertible. 

Disulfides  behave  similarly,  but  occupy  two  positions  in  the  coordination 
sphere.  The  compound 


Et 

1 

1 
S— C 

/ 
Pt 

\ 

s— c 

1 

}H2 

yii.2 

Cl2 

1 

Et 

which  may  serve  as  an  example,  cannot  exist  in  a  trans  form,  but  /?-  and 
7-  forms  analagous  to  those  described  above  have  been  prepared.  The 
/3-form  reacts  with  ethylenediamine  to  give  the  rather  unstable  mixed  com- 
pound [Pt  es  en]Cl2265.  Bennett,  Mosses,  and  Statham267  were  of  the  opinion 
that  dithioether  complexes  of  the  type  [Pt  es  X2]  should  exist  in  racemic 
and  meso  forms  because  of  the  asymmetry  of  the  donor  atoms,  but  they 
were  unable  to  isolate  the  two  geometrical  isomers.  Mann,  however,268 
resolved  a  compound  containing  coordinated  sulfur  as  its  center  of  asym- 
metry  (see  page  325). 

The  dibenzylsulfide  complex  [Au{S(C7H7)2}Cl2]  is  noteworthy  because  its 
simplest  formula  suggests  the  possibility  that  it  may  contain  gold(II)269. 

266.  Tflchugaefl  and  Fraenkel,  Compt.  rend.,  164,  33  (1912). 

266  Blomstrand  and  Weibull,  J.  prakt.  Chem.,  [2]  38,  352  (1888);  Blomstrand  and 

Enebuske,  ./.  prakt.  Chem.,  [2]  38,  3G5  (1888);  Blomstrand  and  Rudelius,  J. 
prakt.  Chun.,  [2]  38,  508  (1888);  Blomstrand  and  Londahl, ./.  pmkt.  Chem.,  [2] 
38,  515  (1888). 

267  Bennett,  Mosses,  and  Statham,  J,  Chem.  Soc,  1930,  1668. 
.v-    Mann,  ./.  Chem.  8oc,  1930,  1746. 

269    Herman,  />'<  r  .  38,  2813  (1905) ;  Raj  and  Sen,  ./.  Tnd.  Chew.  Soc,  7,  67  (1930). 


GENERAL  si  RVEY 


5] 


Such  is  not  the  case,  however,  as  the  substance  is  diamagnetic270.  Prom  the 
molecular   weight,   electrical   conductivity,   magnetic   susceptibility,   and 

crystallographic  data  it  is  concluded  that  the  substance  is  a  Lattice  com- 
pound containing  equivalent  amounts  of  goldi  1 1  and  goldi  111),  IAihSRoCI]- 
[Au(SR2)Cl8]270. 

[ridium(III)271  and  rhodium(  1 1 1 )-"-  form  the  species  |M  (SK,»:,( 'l:;|.  The 
iridium  complex  has  been  separated  into  its  isomeric  forms.  The  anionic 
complex  jlnSR jU'l,]  has  also  been  prepared*78.  Surprisingly,  treatment  of 
these  complexes  with  amines  results  in  the  replacement  of  the  thioether 
groups  first27*. 

Livingstone  and  Plowman274  have  prepared  soma  halogen  bridged  com- 
plexes of  0-methylmercaptobenzoic  acid  which  contain  different  metal  ions. 


(M  =  Hg^rCu11). 

Most  of  the  remarkable  hexadentate  chelating  agents  of  Dwyer  and 
Lions  (Chapter  8)  contain  two  coordinating  sulfur  atoms.  A  fine  demonstra- 
tion of  the  much  greater  affinity  of  cobalt  (III)  for  ether-type  sulfur  than  for 
ether-type  oxygen  is  found  in  the  fact  that  so  long  as  one  sulfur  atom  is 
present,  the  complexes  are  resolvable  into  optical  isomers,  while  substitu- 
tion of  oxygen  atoms  for  both  sulfurs  leads  to  cobalt  (III)  complexes  which 
are  too  unstable  to  resolve275. 

Gonick,  Fernelius,  and  Douglas276  determined  the  formation  constants  of 
ties  of  sulfur  and  nitrogen  containing  chelating  agents  with  the  ions  of 
copper,  nickel,  cobalt,  zinc,  and  silver.  A  comparison  of  the  data  with  similar 
data  for  a  series  of  analagous  polyamines  indicated  that  nitrogen  is  prob- 
ably a  stronger  donor  for  the  metals  studied,  except  silver.  However, 
2-aminoethanethiol,  which  coordinates  as  a  negative  ion,  forms  the  most 
-table  complexes  of  the  entire  group. 


27(1 
971 


Brain,  Gibson,  Jarvis,  Phillips,.  Powell,  and  Tyabji,  ./.  ('hem.  Soc,  1952,  30S6. 
Ray  and  Adhikari, ./.  Ind.  Chem.  Soc,  9,  251  (1932);  Ray,  Adhikari,  and  Ghosh, 

./.  Ind.  Chem.  Soc.,  10,  279    1933). 
Dwyer  and  Nyholm, ./.  Proc.  Roy.  Soc.  N.S.  Wales,  78,  67    194  l  . 
Ray  and  Ghosh,  ./.  Ind.  Chem.  Soc.,  13,  138  (1936);  Ray,  Adhikari.  and  Ghosh, 
./.  Ind.  Chem.  Soc.,  10,  27.",    L933  ;  Ray  and  Adhikari,  ./.  Ind.  Chem.  Soc.,  11, 
•",17    L934  ;  Lebedinskii  and  Gurin,  Compt.  "ml.  aca4.  set.  U.R.S.S.,  40,  322 
L943  . 
Livingstone  and  Plowman,  J.  Proc   !:■■     Sen      V.S.  Wales,  86,  116    1962). 
275.  Dwyer,  Gill,  Gyarfas,  and  Lions.  ./.  Am.  Chem.  Soc.,  75,  1526    1963 

.'.nick,  Fernelius,  and  Douglas,  Technical  Report  to  O.N.R.,  Oct.  16,  1963. 


1,1 
273 


274. 


52 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


In  addition  to  the  marked  stability  of  complexes  containing  the  negative 
mercapt  ide  ion  toward  dissociation  in  solution,  stabilization  of  the  ligand  or 
of  a  high  oxidation  state  of  the  metal  may  occur.  Thus,  the  complex 


Co 


P=C — NH-/         V 


WS  —  CH; 


stabilizes  the  ligand  toward  oxidation  (when  uncomplexed  it  is  rapidly  oxi- 
dized by  air  to  the  disulfide)  and  at  the  same  time  stabilizes  the  strongly 
oxidizing  cobalt  (III)  species277.  The  great  specificity  of  the  metal  ion  in  this 
behavior  is  illustrated  by  comparison  with  the  reaction  of  thioglycolic  acid 
and  iron  ions.  In  air-free  alkaline  solution,  the  complex 


is  formed.  Air  oxidizes  the  iron(II)  ion  to  iron(III)  which  in  turn  catalyzes 
the  oxidation  of  the  ligand  to  the  corresponding  disulfide278. 

Gold  complexes  of  a-thiol  fatty  acids  may  prove  useful  in  the  treatment 
of  such  maladies  as  tuberculosis  and  leprosy279.  The  complex  formed  from 
diethylgold  monobromide  and  2-aminoethanethiol  is  also  of  interest.  From 
its  molecular  weight  it  is  assigned  the  structure 


Et 


Et 


Au 


NH2CH2 


CH, 


However,  the  compound  is  remarkable  in  that  the  coordinated  sulfur  atom 
is  quite  reactive.  The  compound  reacts  explosively  with  methyl  iodide  and 
more  moderately  with  ethyl  bromide.  The  picrate  salt  of  the  product  of 
treat  incut  with  ethyl  bromide  was  shown  to  be  identical  with  the  complex 
prepared  from  S-ethyl-2-aininoethanethior280. 

J77  Feigl,  Nature,  161,  435  (1948);  Anal.  Chem.,  21,  1298  (1949). 

278.  Leussing  and  Kolthoff,  •/.  Am.  Chem.  Soc,  75,  3904  (1953). 

279  Kundu,  J.  Ind.  Chem.  Nor.,  29,  592  (1952). 

280  Ewenfl  and  Gibeon, ./.  Chem.  Soc,  1949,  431. 


QE  VERAL  SURY/)  53 

The  ability  of  seleno-  and  telluromercaptides  and  ethers  to  form  com- 
plexes similar  to  those  of  the  sulfur  analogs  is  illustrated  by  the  mercury(II) 
halide  complexes  used  in  the  characterization  of  these  donor  molecules281. 
Gould  and  McCulloughlM  have  expressed  the  opinion  thai  diarylselenoxides 

coordinate  to  mercury <  1 1 1  through  the  selenium  atom. 

Thio<arl>oii\  I  (  .001  dimit  ion 

Many  thiocarbony]  compounds  show  strong  donor  properties.  Among  the 

simplest  of  these  is  thiourea,  which  coordinates  through  the  sulfur  rather 
than  through  nitrogen,  thus  occupying  only  one  coordination  position. 
Thiourea  coordinates  with  salts  of  almost  all  of  the  heavy  metals.  With 
compounds  of  tripositive  iridium288  and  rhodium284  it  forms  whole  series  of 
compounds,  such  as  [Ir  tu3Cl3],  [Ir  tu4Cl2]Cl,  [Ir  tu5Cl]Cl2  ,  and  [Ir  tu6]Cl3  . 
Thiourea  reacts  with  the  cis  and  trans  isomers  of  platinum(II)  eompounds 
of  the  type  [Pt  a>X2]  yielding  different  products  and  in  so  doing  serves  as 
the  basis  of  Kurnakov's  test  which  is  widely  used  to  distinguish  I  jet  ween 
the  isomers285.  (Chapter  9)  Another  interesting  application  of  thiourea  to 
the  chemical  determination  of  structure  is  found  in  the  work  of  Gent  and 
Gibson288  with  the  dimeric  [Et-jAu  SCX]2 .  The  failure  of  the  complex  to 
react  with  such  nitrogen  bases  as  ammonia,  dipyridyl,  and  ethrylenediamine, 
and  its  reaction  with  thiourea  to  produce  [Et2Au(SCX)tu]  is  interpreted  to 
mean  that  the  thiocyanate  is  coordinated  through  the  sulfur,  and  that  the 
original  compound  has  the  structure 

CN 

I 
E1  S  Et 

\     /\     / 
Au  Au 

/     \/     \ 

Et  S  Et 

I 

CN 

Jensen  has  studied  the  compounds  formed  between  thiosemicarbazide 

281.  Morgan  and  Burst  all,  ./.  Chen  .  8oe.t  1929,  1096;  1930,  1497;  1931,  173;  Carr  and 
rson,  /.  Chem.  Soc.,  1988,  282;  Kraffl  and  Lyons,  Ber.,  27, 176]  (1894  , 
Gould  and  McCullough,  J.Am.  Chem.  Soc,  73.  3195  [1961). 
Lebedinskii,  Shapiro,  and  Kasatkina,  Ann.  inst.  platine,  t'.s.s.ir,  No.  12,  93 
L935). 
284.  Lebedinskii  and  Volkov,  An,,,  inst.  plain,,  .  U.S.S.R.,  No.  12,  79    I" 

Kurnakov,  •/.  Ruse.  Phye.  Chem.  Soc.,  25,  565    1893  ;  <■<.  Chem.  Centr.t  65,  I.  WW 
18Q 
•  nt  and  Gibson,  ./.  Chen    8oe.t  1949,  L835. 


54 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


and  platinum(II)287,  palladium (II)287,  and  nickel288  ions.  The  thiosemicar- 
bazide  molecule  occupies  two  coordination  positions,  evidently  coordinating 
thus 


/ 


NHS 


M 


/ 


-NH 

I 
C— NH2 


S 


Upon  the  addition  of  thiosemicarbazide,  potassium  chloropalladate(II)  first 
gives  [Pd  thio2][PdCl4]  and  then  [Pd  thio2]Cl2 .  If  this  latter  compound  is 
heated  in  weakly  acid  solution,  it  changes  to  the  insoluble  inner  complex 


NH— NH2         NH2— NH 

/  \     /  \ 

HN=C  Pd  O 


=NH 


There  is  evidence  that  this  exists  in  two,  presumably  cis-trans-,  forms.  The 
platinum  and  nickel  compounds  behave  similarly. 

Diketonedithiosemicarbazone  (thiazone)  and  its  homologs 


S  S 


NH2— C— NHN=CR—  C  R'=NNH— C- 


-NHj 


act  as  tetradentate  ligands  forming  inner  complexes  with  copper(II)  and 
nickel(II)  ions 


V 


/ 


N  N 

HN  M       XNH 

I     A    I 

HN=C— S       S— C=NH 


These  complexes  are  quite  stable,  dissolving  in  strong  acid  as  the  soluble 
Baits,  [M(thiazone)]X 

Ammonium  dithiocarbazide  reacts  with  platinum(II)  in  a  manner  com- 
parable both  to  thiourea  and  to  thiosemicarbazide.  With  /rfl/^-[Pt(NH3)2Cl2] 

_'s7    Jensen,  /.  anorg.  allgem.  Chem.,  221,  6  (1934 

288.  Jensen,  /.  anorg.  allgem.  Chem.,  221,  11  (1934). 

289    B&hr  and  Hess,  /   anorg.  allgem.  Chem.,  268,  351  (1952). 


ai:\/:ir\L  survey 


r>-) 


the  reaction  is 

S 

II 
2S— ('— XIINH.     +    frans-[Pt(NH,),Cl,] 


NH,NHC— S  NH, 


Pt 


S— C— XI  I. Mi- 
ll 
S 


+    2C1" 


+    2NH3  +  2C1- 


X  1 1 


while  the  cis-isomer  undergoes  the  reaction 

S 

II 
2S— C— XHXH7  c7S-[Pt(XH3)2Cl2]     -> 

S  S 

II  II 

cs  sc 

/  \  /  \ 

HX  Pt  NH 

\      /     \      / 

NH2  NH2 

From  these  reactions  and  the  fact  that  tetrammineplatinum(II)  ion  is  not 
attacked  by  the  dithiocarbazide  ion,  it  is  concluded  that  the  sulfur  groups 
may  displace  chloride  rapidly  but  that  the  ammonia  is  displaced  only  as  a 
consequence  of  the  trans  influence  of  the  coordinated  sulfur290. 

Inner  complexes  are  formed  by  eobalt(II),  cobalt  (III),  nickel(II),  and 
palladium(II)  with  thiodicyandiamidine  (guanyl  thiourea), 

NH  S 

II  II 

H,X( '—  XII-  CXH2. 

Copper  differs  by  forming  a  complex  of  the  type  [Cu(thicy)  SO,|.  The  co- 
balt, copper,  and  palladium  complexes  decompose  in  warm  alkali,  deposit- 
ing insoluble'  metal  sulfides,  thus  providing  evidence  for  the  participation  of 

290.  Chernyaev  and  Mashentsev,  Izvest.  Sektora  Platiny  i  Drugikh  Blagorod.  Metal., 
Inst.  Obshchei  i  Neorg.  Khun.,  Akad.  Nauk  8.S.S.R.,  23,  72  1949);  cf.  Chem. 
Abs.  45,  2812d     L951);  Mashentsev  and  Chernyaev,  Doklady  Akad.  Nauk 

-   S  S  R  .  79,  803    1951);  cf.  Chrm.  Ah*.  46,  2!»40g  (1052). 


56 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


the  sulfur  atom  in  coordination  (A).  The  nickel  complex  fails  to  give  this 
tesl  and  is  thought  to  have  structure  B291. 


s\  /NH 

HN  Ni  NH 

X  —  UH^     XNH— c( 


HN                                                  ^NH 

>-S           /NH2-CX 

HN                        Pd                        .NH 

>-nh/    xs-< 

HNC                                              XNH 

<w 


(B) 


The  occurrence  of  nickel (IV)  in  sulfur  complexes  testifies  to  the  great 
tendency  of  that  donor  to  form  strong  covalent  bonds.  Hieber  and  Briick292 
found  that  air  oxidation  of  a  strongly  alkaline  suspension  of  the  nickel(II) 
complex  of  o-aminothiophenol  produces  the  deep  blue  complex 


CEh^^'^XD 


A  similar  bridged  disulfo  compound  is  formed  by  dithiobenzoic  acid 


Dithiooxamide  (rubeanic  acid)  forms  insoluble  complexes  with  nickel,  and 
copper  ions293.  These  substances  have  the  properties  of  inner  salts,  and  be- 
cause of  the  steric  requirements  of  the  ligand,  they  exist  as  bridged  polymers 


>     x    /\    i.   /\ 


'NH 


NH*  ^S        NH'  NS  x 


Anion  and  Kane-""  have  used  the  linear  nature  and  the  light  absorption  of 
this  polymer  in  the  manufacture  of  a  device  for  the  polarization  of  light.  A 
sheet  of  plastic  is  soaked  in  a  solution  of  dithiooxamide,  which  causes  the 
precipitation  of  the  complex  within  the  plastic.  When  the  plastic  sheet  is 

291.  Ray  and  Chaudhury,  J.  Ind.  Chem.  Soc.,  27,673  (1950);  Poddar  and  Ray,  /.  Ind. 
Chem.  Soc.,  29,  279  (1952). 

Hieber  and  Bruck,  Naturwias.,  36,  312    L949). 

Jensen,  Z.  anorg.  Chem.,  262,  •_,--,7  (1944  ;  Ray,  Z.  anal.  Chem.,  79,  95  (1929). 
294.  Anion  and  Kane,  U.S.  Patent  2  505  085,  April  25,  1950. 


GENERAL  SURVE1  57 

Stretched  ill  one  direction,  the  polymer  chains  arc  oriented  parallel  to  each 
other.  The  bridging  ability  of  tins  donor  molecule  is  also  illustrated  in  the 
dimeric  derivative  of  diethylgold  monobromide 


E^A- 
Au 

1 

Au 

Bt^     X 

H 

^Et 

Other  Sulfur  Donors 

The  thioeyanate  ion  has  unshared  pairs  of  electrons  on  both  the  sulfur 
and  the  nitrogen.  Werner  at  one  time296  supposed  the  two  isomers  of 
[Co  enj(NCS)2]+  to  be  structurally  different,  one  having  a  cobalt-nitrogen 
link  and  the  other  a  cobalt-sulfur  link.  This  hypothesis  was  based  upon  the 
fact  that  the  thioeyanate  group  of  one  of  the  isomers  is  destroyed  by 
chlorine,  leaving  the  nitrogen  (in  the  form  of  ammonia)  in  union  with  the 
metal,  while  the  thioeyanate  group  of  the  other  isomer  is  completely  elim- 
inated by  this  treatment.  Werner  later  found'297,  however,  that  the  two  com- 
pounds are  stereoisomers,  and  that  the  thioeyanate  group  is  attached  to 
the  metal  through  the  nitrogen  in  both  cases.  The  sulfur  of  the  thioeyanate 
group  probably  does  have  strong  donor  properties,  however,  and  in  the 
case  of  gold  it  is  the  sulfur  atom  which  preferentially  coordinates  Werner 
reported  that  silver  nitrate  does  not  precipitate  silver  thioeyanate  from  a 
solution  of  [Co(XH3)s(XCS)]++  or  similar  complexes  but  the  silver  loses 
its  ionic  character.  He  supposed  that  the  silver  coordinates  with  the  sul- 
fur2'^. Waggener,  Mattern,  and  Cartlcdge299  however,  have  found  the 
stability  of  these  dinuclear  complexes  to  be  much  less  than  reported  by 
Werner. 

The  sulfite  ^roup  evidently  occupies  only  one  coordination  position  in 
most  cases,  and  from  the  fact  that  salts  of  the  ion  [Co(XH3)4(S03)2]_  are 
yellow  or  brown,  it  may  be  inferred  that  these  compounds  contain  a  sulfur- 
cobalt  link. 

The  action  of  sulfite  ion  on  platinum(II)  complexes  is  also  most  easily 
explained  on  the  basis  of  a  metal-sulfur  bond.  Sulfite  acts  differently800  on 

396.  I . .'  era  and  ( iibaon,  •/ .  ( 'hi  n  .  Soc.s  1949,  131 . 

Werner  and  Braunlieh,  Z.  anorg.  Chetn.,  22,  91,  123  (1900). 

297.  Werner.  .1„„..386,  1     L912). 

298.  Werner,  Ann.,  386,  50  (1912). 

Waggener,  Mattern.  and  Cart  ledge,  al>st  racl  8,  [22nd  meeting,  American  Chemi- 
Sepl .  1962. 
300.  Gurin,  Doklady  Akad.  Nauk  S.S.S./r.  50,  Jul    L946). 


58  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

the  cis-  and  trans-  isomers  of  dichlorodiammineplatinum(II) 

cis  [Pt(NH,)2CU]  +  l\a,S();;  ->  Xa6[Pt(S08)4]  +  2NaCl  +  2NH3 
«rans-[Pt(NH,)8Cla]  +  2Na2S08  -*  *mns-Na2[Pt(NH3)2(S03)2]  +  2NaCl 

This  behavior  is  quite  similar  to  the  reaction  of  these  isomers  with  thiourea. 

Aside  from  the  complexes  with  aromatic  nitrogen  molecules,  ruthe- 
nium^ I)  is  besl  known  in  its  very  unusual  sulfite  complexes.  Treat- 
ment of  cWoropentammineruthenium(III)  ion  with  sodium  bisulfite 
produces  the  two  complex  compounds,  [RuIT(NH3)4(S03H)2]  and 
\a  ,|  Etun<  X  HsMSOaMSOsH)*,]  -6H20.  The  dipositive  oxidation  state 
of  the  ruthenium  was  verified  by  analysis  and  magnetic  measurements301. 
Upon  dissolution  in  acid,  [RuII(NH3)4(S03H)2]  is  converted  to 
[RuII(NH3)4(S02)X]X.  The  action  of  ammonium  hydroxide  on  the  dibisul- 
ntotetrammine  produces  the  nonelectrolyte,  [RuII(NH3)5(S03)]302.  This 
compound  is  also  sensitive  to  acid,  transforming  to  [RuII(NH3)5(S02)]++. 

Rhodium(III)  and  iridium(III)  form  complexes  of  the  type 
MI3[MIII(NH3)3(S03)3]303.  Iridium  also  forms  a  compound  in  which  the 
sulfite  group  is  reported  to  be  bidentate,  [Ir(S03)3Cl2]5~,  but  the  alternate 
possibility  of  halogen  bridging  has  not  been  disproved. 

Riley304  has  prepared  salts  of  the  dark  red  selenitopentamminecobalt(III) 
ion,  [Co(XH3)5(Se03)]+,  but  his  experiments  did  not  show  whether  the 
selenite  group  is  attached  through  the  selenium  or  through  the  oxygen. 
Several  selenite  complexes  of  nickel,  copper,  and  cobalt  have  been  obtained 
by  Ray  and  Ghosh305,  who  found  them  to  be  less  stable  than  the  correspond- 
ing sulfite  compounds. 

The  thiosulfate  group,  with  unshared  electrons  on  both  oxygen  and  sul- 
fur, could  conceivably  coordinate  through  either  or  both.  When  it  occupies 
but  one  coordination  position,  union  with  the  metal  evidently  takes  place 
through  the  oxygen,  for  the  ion  [Co(NH3)5S203]+  is  red306.  This  ion  is  very 
stable,  for  it  is  formed  when  [Co(NH3)5Cl]S203  or  [Co(NH3)5Br]S203  is 
allowed  to  stand  at  35  to  40°C307.  The  stability  of  the  thiosulfate-cobalt 
bond  is  further  attested  by  the  reaction  of  [Co(NH3)5S203]+  with  potassium 
cyanide,  which  yields  K4[Co(S203)(CN)5].308  Duff167  reported  the  prepara- 

301.  Gleu,  Breuel,  and  Rehm,  Z.  anorg.  allgem.  Chem.,  235,  201  (1938). 

302.  Gleu  and  Breuel,  Z.  anorg.  allgem.  Chem.,  235,  211  (1938). 

303    Lebedinskii  and  Shenderetskaya,  Izvest.  Sektora.  Platiny  i  Drugikh  Blagorod. 

Metal.,  hist.  Obehchei  i  Neorg.  Khim.  Akad.  Nauk  S.S.S.R.,  21,  164  (1948);  cf. 

Chem.  Abe.,  44,  L0566a  (1950);  Gurin,  Doklady  Akad.  Nauk  S.S.S.R.,  56,  217 
1936);  of.  Chem.  Abe.  43,  1676a  (1949). 
304.  Riley,  ./.  Chem.  Soc,  1928,  2985. 

306  Raj  and  Ghosh,  ./.  Indian  Chem.  Soc.,  13,  494  (1936). 
Ray,  •/.  Indian  Chem.  Soc.,  4,  64  (1927). 

307  Sarkar  and  Daa  Gupta,  J.  Ind.  Chun.  Soc,  7,  835  (1930). 

308  Ray, ./.  Ind.  Chem.  Soc,  4,  325  (1927). 


GENERAL  SURVEY  59 

tion  of  [Co  enj  S.j().;]Br-3II2(),  which  he  thought  contained  a  doubly  co- 
ordinated thiosulfate  group.  The  evidence  for  this  is  Blight,  however,  and 

the  correct  formula  may  well  be  [Co  en2  (H,(  ))S,(  ):;]Br-L>I  I,( ).  YVeinlaiaP'' 
has  suggested,  hut  without  experimental  evidence,  that  the  double  potas- 
sium bismuth  thiosulfate  is 


K3 


jH20 


in  which  coordination  takes  place  through  both  oxygen  and  sulfur.  The 
fixation  process  in  photography  depends  upon  the  formation  of  thiosulfato- 
silver  anions,  of  which  several  have  been  reported310.  If  coordination  takes 
place  through  the  oxygen,  sulfates  should  give  analagous  compounds. 

The  Doxor  Properties  of  Nitrogen 

The  solvent  properties  of  ammonia  closely  resemble  those  of  water,  and 
solvation  is  as  important  in  ammonia  solutions  as  it  is  in  aqueous  solutions. 
The  donor  properties  of  nitrogen  are  as  strong,  or  stronger,  than  those  of 
oxygen,  and  some  of  the  metal-ammonia  compounds  show  remarkable 
stability.  Many  of  them  (including  those  of  cobalt,  chromium  and  the 
platinum  metals)  do  not  lose  ammonia  when  heated  above  200°C  or  wrhen 
treated  with  sodium  hydroxide  or  hydrochloric  acid.  The  ammines  of  copper, 
silver,  zinc,  and  several  other  metals  are  equally  well  known,  but  are  much 
less  stable,  and  are  decomposed  by  dilute  acids  or  bases.  Ammines  of  the 
alkali  and  alkaline  earth  metals  are  completely  decomposed  by  water,  and 
some  of  them  are  stable  only  at  low  temperatures. 

Ammines 

The  hydrates,  especially  those  of  the  highly  charged  metallic  ions,  readily 
liberate  hydrogen  ions,  with  the  formation  of  aquohydroxo  complexes.  An 
analagous  reaction  takes  place  with  ammines,  but  it  is  less  pronounced  than 
with  hydrates.  From  a  study  of  the  ammines  of  rhodium,  Griinberg  and 
rmanir11  concluded  that  the  acid  dissociation  of  coordinated  water  is 
105  times  as  great  as  that  of  coordinated  ammonia.  The  loss  of  protons  by 
ammines  is  particularly  noticeable  with  the  complexes  of  the  very  heavy 
metal>,  as  is  illustrated  by  the  formation  of  HgXH-jCl  when  mercuric  chlo- 
ride is  treated  with  ammonia.  Other  illustrations  involve  the  ammines  of 

309.  Weinland,  "Einfuhrung  in  die  Chemie  der  Komplexverbindungen,"  Second  Ed., 

p.  148,  Stuttgart,  Enke,  1924. 

310.  Bassett  and  Lemon,./.  Chem.  Soc.j  1933,  112:;;  Aflhihara  and  Mateuda,  Kogaku 

8huho,  Kyushu  Univ.  (Technological  Reports,  Kyushu  Univ.),  25,  11  (1952);  cf. 
Chem.  Abs.,  47,  12075g  (1953). 


00 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Table  1.1.  Colors  of  Some  Anhydrous  Salts,  and  Their  Hydrates 

and  ammonates 


CoCl2 

[Co(H20)6]Cl2 

[Co(NH3)6]Cl2 

Blue 

Red 

Rose 

CuCl2 

[Cu(H20)4]Cl2 

[Cu(NH3)4]Cl2 

Brown 

Blue 

Deep  Blue 

NiCl2 

[Ni(H20)6]Cl2 

[Ni(NH3)6]Cl2 

Light  brown 

Green 

Blue 

CrCl3 

[Cr(H20)6]Cl3 

[Cr(NH3)6]Cl3 

Violet 

Gray-violet 

Yellow 

[Cr(H20)5Cl]Cl2 

H20 

[Cr(NH3)5Cl]Cl2 

Green 

Rose  red 

[Cr(H20)4Cl2]Cl 

•2H20 

[Cr(NH3)4Cl2]Cl 

Green 

cis-  violet 
trans-  green 

platinum311, 312,  gold313,  and  osmium314.  Ammines  of  the  lighter  elements 
also  lose  protons  to  some  extent,  as  is  indicated  by  the  fact  that  the  hydro- 
gen atoms  in  such  complexes  as  [Co(NH3)6]"HH"  are  readily  exchanged  for 
deuterium  when  placed  in  heavy  water315. 

Water  and  ammonia,  coordinated  to  ions  of  the  same  metal,  do  not  al- 
ways stabilize  the  same  valence  state  (Chapter  11).  For  example,  hydrated 
cobalt  (III)  compounds  are  very  strong  oxidizing  agents,  while  ammoniated 
cobalt (II)  compounds  are  strong  reducing  agents.  The  hydrates  and  am- 
mines often  show  similar  colors,  but  this  is  by  no  means  a  general  rule. 
Table  1.1  summarizes  a  few  examples.  Peters316  made  the  first  systematic 
and  extended  study  of  the  stability  of  ammines.  He  subjected  ninety  seven 
salts  to  the  action  of  dry  ammonia  gas  at  atmospheric  pressure  and  by 
measuring  the  volume  of  ammonia  absorbed  in  each  case,  calculated  the 
formulas  of  the  ammines  obtained.  Following  Peters,  Ephraim317,  W. 
Biltz318,  Clark319, 320,  and  others  studied  the  reactions  of  salts  with  anhydrous 

311.  Grunberg,  Z.  anorg.  Chem.,  138,  333  (1924);  Griinberg  and  Faermann,  ibid.,  193, 

193  (1930). 

312.  Tschugaeff,  Z.  anorg.  Chem.,  137,  1  (1924). 

313.  Block  and  Bailar,  J.  Am.  Chem.  Soc.,  73,  4722  (1951). 

314.  Dwyer  and  Hogarth,  J.  Am.  Chem.  Soc.,  75,  1008  (1953). 

315.  Anderson,  Briscol,  and  Spoor,  J.  Chem.  Soc.,  1943,  361. 

316.  Peters,  Zeit  anorg.  Chem.,  77,  137  (1912). 

317.  Ephraim,  Z.  phys.  Chem.,  81,  513,  539  (1913);  83, 196  (1913);  Ber.,  45, 1322  (1912); 

46,  3103,  3742  (1913);  47,  1828  (1914);  48,  41,  624,  629,  1638,  1770  (1915);  49,  2007 
(1916);  50,  529,  1069,  1088  (1917);  51,  130,  644,  706  (1918);  52,  236,  241,  940,  957 
(1919);  53,  548  (1920);  54,  973  (1921). 

318.  Biltz  and  co-workers,  Z.  phys.  Chem.,  82,  688  (1913);  Z.  anorg.  allgem.  Chem.,  83. 

L63,  177  (1913);  89,  97,  134,  141  (1914);  109,  89,  132  (1919);  114,  161,  174,  241 
1920);  119,  97,  115  (1921);  123,  31  (1922);  124,  235,  322  (1922);  125,  269  (1922); 
127,  1  (1923);  129,  1,  161  (1923);  130,  93  (1923);  Z.  Elektrochem.,  26,  374  (1920); 
Angew.  Chem.,  33,  313  (1920). 


GENERAL  SURVEY  61 

ammonia.  They  prepared  and  studied  hundreds  of  amminea  in  order  to  find 

out  what  factors  arc  important  in  determining  stability.  While  a  great  deal 
was  learned  about  the  stabilities  of  ammines,  little  light  was  thrown  on  the 
structures  of  such  compounds  as  A1C11:; ■  *)X  1 1:; ,  AlCl3-5NH.:  and 
AlClj-  INHj820  and  compounds  containing  very  large  amounts  of  ammonia, 
such  as  T1C1  r -OX  1 1/-1.  Doubtless  many  of  these  are  "lattice  compounds" 
only. 

The  ammines  which  are  of  chief  interest  are  those  of  the  transil  lob  metals 
and  the  metals  of  periodic  groups  IB  and  IIB.  Even  among  these,  there  are 
great  differences  in  stability.  For  example,  iron  ammines  cannot  be  ob- 
tained in  the  presence  of  water;  copper  ammines  and  cobalt(II)  ammines 
exist  in  water  solution,  and  can  be  crystallized  from  such  solutions,  but  they 
are  immediately  destroyed  by  acids.  Cobalt  (III)  and  platinum  ammines 
can  be  recrystallized  from  solutions  of  strong  acids,  and  the  hydroxides 
[Co(XH3)6](OH)3  and  [Pt(XH3)6](OH)4  are  sufficiently  stable  to  allow  their 
easy  preparation322.  This,  of  course,  may  be  a  measure  of  rate  of  decomposi- 
tion rather  than  of  intrinsic  stability,  but  it  is  of  tremendous  practical  im- 
portance. 

The  nature  of  the  anion  is  of  great  importance  in  determining  the  stability 
of  some  metal  ammines.  Weitz323  observed  that  the  ammines  of  gold  are 
stable  if  the  anion  is  an  oxy-anion  such  as  nitrate,  perchlorate,  phosphate  or 
oxalate,  and  the  ammonia  groups  cannot  be  removed  by  the  action  of  the 
oxyacids.  They  are  destroyed,  howrever,  by  halides,  presumably  because  the 
halide  ion  replaces  part  of  the  ammonia  in  the  coordination  sphere.  Tomlin- 
son,  Ottoson,  and  Audrieth324  have  called  attention  to  the  explosive  charac- 
ter of  cobalt  (III)  and  chromium(III)  ammines  in  which  oxidizing  groups  are 
present  in  the  coordination  sphere  or  as  anions. 

It  is  of  interest  that  the  ammines  which  are  easily  decomposed  by  acids 
(e.g.,  those  of  Cu,  Ag,  and  Zn)  are  easily  formed  by  the  addition  of  am- 
monia to  a  solution  of  the  metal  ion.  The  ammines  which  are  not  rapidly 
destroyed  by  acids  are  not  readily  formed.  Thus,  the  addition  of  an  excess 
of  ammonia  to  a  solution  of  a  chromium(III)  salt  ordinarily  precipitates  the 
hydroxide;  the  hexammine  is  formed  in  good  yield  only  by  the  action  of 
liquid  ammonia  on  anhydrous  chromium(III)  chloride  in  the  presence  of  a 
catalyst325.  The  hexammine  cobalt  (III)  ion  is  not  obtained  by  aerial  oxida- 

319.  Clark,  Quick,  and  Harkins,  ./.  Am.  ('hem.  Soc,  42,  2438  (1920);  Clark  and  Buck- 

ner,  J.  Am.  Chem.  Soc,  44,  230  (1922). 

320.  Clark,  An,.  ./.  Set.,  7,  1  (1924). 

321.  Young,/.  Am.  Chem.  Soc,  57,  997  (1935). 

322.  Hecht,  Z.  anorg.  aUgem.  Chem.,  270,  215  (1952). 

323.  Weitz.  Ann.,  410,  117    1915). 

324.  Tomlinson,  Ottoson,  and  Audrieth,  ./.  .1///.  Chi  m.  Soc,  71,  375  (1949). 

325.  Oppegard  and  Bailar,  Inorganic  Syntheses,  III,  153  (1950). 


62 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


tion  of  an  ammoniacal  eobalt(II)  solution  except  in  the  presence  of  a 
catalyst826.  Dwyer  and  Hogarth327  could  prepare  the  ion  [Os(NH3)6]+++ 
only  by  the  treatment  of  [Os(\]IH)5Br]++  with  ammonia  under  pressure. 

Ammonia  can,  of  course,  share  the  coordination  sphere  with  other  donor 
groups.  In  his  first  paper116,  Werner  pointed  out  that  ammonia  molecules  can 
be  displaced,  one  by  one,  from  the  coordination  sphere,  either  by  other 
neutral  groups  such  as  water,  or  by  negative  groups.  If  the  metal-ammonia 
bond  is  stable1,  the  groups  which  share  the  coordination  sphere  with  ammonia 
may  be  replaced  by  other  groups  to  form  a  great  variety  of  compounds. 
The  following  reactions  are  typical327: 

HCl 

[Os(NH3)5Br]++     Ag2° .    [Os(NH3)6OH]++  >  [Os(NH3)5Cl]++ 

H20 


The  amide  group,  like  the  hydroxide  group,  has  two  pairs  of  unshared 
electrons  and  coordinates  readily  with  certain  metals.  Mercury  amido  chlo- 
ride illustrates  this.  The  NH2~  group  can  also  act  as  a  bridge  between  two 
acceptor  atoms  (p.  23).  The  imino  group  frequently  acts  as  a  bridge  also, 
as  in 

NH 
III/     \IV 

en2  Co  Co  en2 


\  / 


and 


K2 


(NH8)8PtI 


<">: 


NH 


NH 


PtI(NH3)s 


Aliphatic  Amines 

The  aliphatic  monoamines  coordinate  less  readily  than  does  ammonia, 
and  the  compounds  so  formed  are  less  stable  than  the  ammines.  However, 
this  p<»int  is  often  overemphasized,  for  some  rather  stable  coordination 
compounds  of  the  aliphatic  amines  do  exist.  The  secondary  amines  co- 
ordinate less  readily  than  do  the  primary,  and  the  tertiary  amines  are  al- 
in<»t  devoid  of  ability  to  coordinate  with  metal  ions.  This  is  probably  due 

326    Bailar  and  Work, ./.  Am.  Chem.  Soc,  67,  176  (1945). 

327.  Dwyer  and  Hogarth, ./.  Proc.  Roy.  Soc.  N.S.  Wales,  84,  117  (1951). 

328    Werner,  Ann.,  375,  74  (1910). 

Ofven.  K.  Vet.  Akad.  Fork.,  27 ,  777  (1870);  28,  175  (1871). 


GENERAL  SURVEY  63 

tosteric  factors,  for  the  tertiary  amines  coordinate  (irmly  with  the  hydrogen 
ion;  that  is,  they  arc  strong  bases.  Straumanis  and  Circulis880  have  de- 
scribed compounds  of  the  mercury  and  copper  halides  with  ethylamine, 
propylamine,  butylamine,  dimethylamine,  and  diethylamine.  Jorgensen881 
prepared  platinum(II)  complexes  containing  methyl,  ethyl,  and  propyl- 
amines, and  Drew  and  Tress332  have  extended  his  study  to  include  the 
preparation  of  the  stereoisomers  forms  of  [Pt(CHsNHs)£)ls].  These  are 
Btable  enough  thai  they  can  be  oxidized  to  |Pt((1H3NH2)2Cl4].  Gil'denger- 
shel333  prepared  [PtCCHsNI^iClsJCli  by  the  action  of  methylamine  on 
potassium  chloroplatinate,  and  purified  it  by  recrystallization  from  hydro- 
chloric  acid.  Chernyaev334  has  prepared  three  of  the  four  possible  isomers  of 
[[Pt  en(CHsNH2)(NOi)CyClJ  and  has  resolved  one  of  them,  as  well  as 
[Pt  en(CHsNH2)(NOi)tCl]Cl.  Finally,  Meisenheimer  and  Kiderlen335  have 
introduced  various  primary  amines  into  the  coordination  sphere  of  cobalt 
by  the  reaction 

[Co  en2Clo]Cl  +  amine  — »  [Co  en2  amine  Cl]Cl2 

Even  aromatic  amines  form  fairly  stable  compounds  in  this  way.  Primary 
■nines  which  are  weaker  bases  than  aniline,  and  secondary  amines,  do  not 
enter  the  complex,  but  bring  about  more  complicated  reactions336337. 

If  chelation  can  take  place  to  form  five-membered  rings,  the  stability  of 
the  compounds  is  greatly  enhanced  (Chapter  5).  Ethylenediamine  is  the 
simplest  and  the  most  important  of  such  bases,  and  its  compounds  have 
played  an  important  part  in  the  development  of  the  coordination  theory. 
1,2-Diaminopropane  (propylenediamine)  also  forms  stable  compounds, 
which  are  similar  to  those  containing  ethylenediamine,  but  are  usually 
more  soluble.  Isobutylenediamine338,  2,3-diaminobutane339,  stilbenedi- 
amine33S-  34°,  and  several  other  1,2-diamines  have  been  shown  to  form 
stable  chelate  rings.  Pearson,  Boston,  and  Basolo341  have  prepared  com- 

130.  Straumanis  and  Circulis,  Z.  anorg.  allgem.  Chem.,  230,  65  (1936). 

531.  Jorgensen,  J.  prakt.  Chem.,  33,  530  (1886). 

B2.  Drew  and  Tress, ./.  Chem.  Soc.,  1935,  1212. 

m.  Gil'dengershel,  Zhur.  Priklad.  Khim.  (J.  Applied  Chem.),  23,  487  (1950). 

534.  Chernyaev,  Ann.  inst.  platine  No.  8,  37  (1931). 

535.  Meisenheimer  and  Kiderlen,  Ann.,  438,  238  (1924). 

136.  Ablov,  Bull.  soc.  chim.,  [5]  3,  2270  (1936);  4,  1783  (1937). 

537.  Bailar  and  Clapp,  ./.  Am.  Chun.  Soc.,  67,  171  (1945). 

B.  Mills  and  Quibell,  ./.  Chem.  Sue,  1935,  839;  Lidstone  and  Mills,  ./.  Chem.  Soc, 

1939,  1764. 
m.  Bailar  and  Balthie,  J.  Am.  Chem.  Soc.,  68,  L474  (19 
14i).  Williams,  thesis,  I'nivcrsit  y  of  Illinois,  1961. 
Ml.  Pearson,  Boston  and  Basolo.  J.  Am.  Chem.  Soc,  76,  3089  (1963 


64  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

pounds  of  the  type 


CI; 


CI 


in  which  the  R's  represent  hydrogen  or  methyl.  As  the  number  of  methyl 
groups  is  iii<  reased,  crowding  becomes  pronounced,  and,  in  water  solution, 
the  coordinated  chlorides  are  more  easily  replaced  by  water  molecules. 

Trimethylenediamine  forms  six-membered  rings,  which  compare  favor- 
ably in  stability  with  those  of  ethylenediamine297 • 342.  Mann343  has  prepared 
coordination  compounds  of  several  metals  with  bases  of  the  type 
(XIf2CH2)2CHX,  where  X  =  CH3 ,  Br,  SCN,  and  OH.  Tetramethylene- 
diamine  and  the  higher  homologs  in  the  series  apparently  cannot  form  rings 
at  all  in  aqueous  solution.  Diamines  having  four,  five,  ten  and  eighteen 
carbon  atoms  have  been  investigated344.  Pfeiffer345  has  shown,  however,  that 
tetramethylenediamine  and  hexamethylenediamine  will  form  chelates  from 
alcohol  solution. 

The  polyamines  NH2CH2CH2(NHCH2CH2)nNH2  (n  =  1,  2,  3,  or  4)  are 
strong  coordinators,  (even  though  part  of  the  nitrogen  atoms  are  second- 
ary), because  they  form  multiple  ring  systems.  Diethylenetriamine  acts  as 
a  tridentate  base  toward  copper(II)  and  nickel(II)  ions,  giving  complexes 
of  the  types  [Cu  dien  Cl]+  and  [Cu  dien2]++.  In  the  second  case,  because  of 
the  stereochemical  properties  of  the  base,  copper  assumes  a  coordination 
number  of  six346347.  Jonassen  and  his  students  prepared  platinum  and 
palladium-triethylenetetramine  complexes  [Pt  trien]++  and  [Pd  trien]++348, 
and  [Xij  trion3]4+.  The  [Ni2  trien3]4+  is  paramagnetic,  so  must  consist  of  two 
tetrahedra849.  Basolo350  prepared  a  series  of  cobalt  complexes  of  the  types 


342. 
343. 
344. 


346. 
346. 
347. 
348. 


Bailar  and  Work,  J.  Am.  Chem.  Soc,  68,  232  (1946). 

Mann,  ./.  Chem.  Soc,  1927,  2904;  1928,  1261. 

Pfeiffer  and  Haimann,  Ber.,  36,  1063  (1903);  Pfeiffer  and  Lubbe, ./.  prakt.  Chem! 

[2]  136,  321    (1933);  Tschugaeff,  Ber.,  39,  3190  (1906);  TschiiKaefT,  ./•  prakt. 

Chem.,  [2]  75,  159  (1907);  Werner,  Ber.,  40,  61  (1907);  McReynolds,  thesis, 

University  of  Illinois,  L938. 
Pfeiffer,  Naturwiss.,  36,  190  (1948). 
Mann,  ./.  Chem.  Soc,  1934,  466. 
Breckenridge,  <'<n,a</nin  ./.  Research,  26B,  11  (1948). 
Jonassen  and  Cull, ./.  Am.  Chem.  Soc.,  71,  1097  (191!)). 
.Joniisscii  and  Douglas,  J.  Am.  Chem.  Soc,  71,  1091  (1919). 
Basolo,  J\  .1///.  Chem.  Soc,  70,  2634  .1948). 


GENERAL  SURVEY  65 

[Co  trien  \..|  and  [Co  trien  Y],  where  X  is  CI,  NOs  and  Ml,  and  Y  is  (  '<  ». 
or  en.  He  also  obtained  [C02  triens]^",  an  ion  of  unusually  high  ionic  charge. 
Jonassen  and  Fry351  have  isolated  the  cobalt(II)  complex  of  tetraethylene- 
pentamine. 

(SjjS'^^-Triaminotriethylamine  behaves  as  a  quadridentate  amine  in 
spite  of  the  reluctance  of  tertiary  nitrogen  to  coordinate.  Mann  and  Pope852 
prepared  the  platinum(II)  and  platinum(IV)  complexes  |Pt  tren]Cla  and 
[Pt  tren  C12]C12 .  The  palladium(II)  and  nickel  ions  form  the  ion  |M  tren]++ 
and  nickel  forms  also  the  ion  [Nij  tren;1]1+,  in  which  the  coordination  number 
o\  nickel  is  evidently  six388.  Mann884  prepared  several  salts  of  the  ion  [Co 
tren(SCN)2]+.  By  treatment  of  [Co  enJ(,l..|+  with  the  same  base,  Jaeger  and 
Koets355  obtained  salts  of  an  ion  which  they  thought  to  be  [(Co  ei^trer^]94", 
but  at  tempts  to  repeat  this  work356  have  been  unsuccessful,  and  it  seems  that 
Jaeger  and  Koets  probably  had  [Co  tren  en]+++. 

Cases  are  known  in  which  the  polyamines  coordinate  without  using  all 
of  their  nitrogen  atoms357.  a,jS,7-Triaminopropane  can  act  either  as  a  bi- 
dentate  or  tridentate  group358  depending  upon  the  metal  ion  involved  and 
the  conditions  of  the  experiment.  If  only  two  amino  groups  coordinate,  they 
are  on  adjacent  carbon  atoms. 

Ethylenediamine,  and  presumably  other,  similar  bases,  sometimes  co- 
ordinate through  only  one  nitrogen.  Chernyaev  and  Fedorova359  prepared  a 
compound  whose  formula  they  write  [Pt(en-HCl)  'NHj-Clj].  Mild  alkalies 
close  the  ring  with  the  formation  of  [Pt  enXH3Cl]Cl,  and  chlorine  oxidizes 
the  compound  to  [Pt(enHCl)(XH3)Cl4].  This  platinum(IY)  compound  hy- 
drolyzes  to  [(XH:5)(H20)Cl3PtenPtCl3(H20)(XH3)]Cl2 ,  in  which  the  ethyl- 
enediamine acts  as  a  bridge  between  the  two  platinum  atoms.  Job360  has 
adduced  evidence  for  the  existence  of  [Ag  en2]+  and  [Tl  en]+  ions,  which  are 
kalagous  to  [Ag(NH3)2]+  and  [T1(XH3)]+,  and  hence  contain  monoco- 
Irdinated    ethylenediamine.    Di-n-propylgold(III)    bromide    reacts    with 

351.  Jonassen  and  Fry,  ./.  .1///.  Chem.  Soc.,  75,  1524  (1953). 

352.  Mann  and  Pope,  Proc.  Roy.  Soc.  London,  109A,  444  (1925). 
:-;:>:;.  Mann  and  Pope  ./ .  Chem.  Soc.,  1926,  482. 

354.  Mann.  ./.  Chem.  Sue,  1929,  40!). 

■5.  Jaeger  and  Koets,  Z.  anorg.  allgem.  Chem.,  170,  347  (1928). 

■6.  Middleton    1952)  and  Rebertus  (1964),  unpublished  work,  University  of  Illinois. 

357.  Mann. ./ .  Chem.  80c.,  1934,  466;  Job  and  Brigando,  Compt.  rend.,  210,  138    L940). 

■B.  Mann  and  Pope,  ./.  Chem.  Soc,  1926,  2675;  Nature,  119,  351   (1927);  Mann,  ./. 

Chem.  Soc.,  1926,  2681;  1927,  1224;  1928,  890;  1929,  651. 

hernyaev  and  Fedorova,  Ann.  secteur  platine,  Inst.  chim.  gen.   (U.S.S.R.), 

No.  14,  9  (1937). 
360.  Job,  (  nd.,  176,  4  12    1923) ;  184,  1066  (1927). 


.ill 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


ethylenediamine  to  form  a  compound  which  is  formulated 
Pr  Br  Br  Pr  361- 


Au 


Au 


Pr 


NH2 — CH2 — CH2-  Ml 


Pr 


On  heating,  one  of  the  gold  atoms  loses  its  two  propyl  groups,  retaining  its 
hold  on  the  bromine  and  the  nitrogen: 


Au 


Pr 


Br 


NH2CH2CH2NH2 


Au  . 


The  treatment  of  Zeise's  salt,  K[Pt(C2H4)Cl3],  with  ethylenediamine  re- 
sults in  the  formation  of  a  dinuclear  complex  in  which  the  ethylenediamine 
acts  as  a  bridging  group362 


C2H4 


CI 


CI 


Pt 

/  \ 

Cl  NH2CH2CH2NHS 


Pt 


C2H4 


Cl 


Gilman  and  Woods363  have  prepared  a  compound  which  they  believe  to  have 
the  structure  (CH3)3AuNH2CH2CH2NH2Au(CH3)3 .  In  this  case  ring  forma- 
tion is  impossible  because  only  one  coordination  position  is  open  on  each 
gold  atom. 

Pfeiffer  and  Glaser364  have  studied  the  donor  properties  of  X-substituted 
ethylenediamines.  With  copper(II)  perchlorate,  N-methyl  and  X,X'-di- 
ethylethylenediamine  give  blue-violet  compounds  analgous  to  [Cu  en2] 
(C104)2 .  The  corresponding  N-diethyl  compound  is  ruby  red  at  room 
temperature,  but  assumes  the  blue-violet  color  above  44°.  The  same  in- 
vestigators report  that  the  reaction  of  X-methyl-X'-diethyl  etlrylenedi- 
amine  and  N-triethyl  ethylenediamine  with  copper  perchlorate  do  not  give 
compounds  which  are  analagous  to  those  of  the  less  highly  substituted 
bases,  but  correspond  to  the  formula  [Cu  OH  diamine]C104 .  They  are 
probably  dimeric,  the  copper  atoms  being  linked  together  through  two  ol 
bridges.  These  compounds,  like  the  others,  are  thermochromic,  changing 
from  blue-violet  to  ruby-red  when  they  are  cooled  in  liquid  air. 

361.  Burawoy  and  Gibson,  ./.  Chem.  Soc.,  1935,  210;  Burawoy,  Gibson,  and  Holt. 

./.  Chem.  Soc,  1935,  1024. 

362.  Hel'man,  Compt.  rend.  acad.  set.  U.R.S.S.,  38,  243  (1043). 

363.  Gilman  and  Woods,  •/.  Am.  Chem.  Soc,  70,  550  (1948). 

364.  Pfeiffer  and  Glaser,/.  prakt.  Chem.,  12],  161,  134  (1938);  153,  300  (1939). 


GENERAL  SURVEY  67 

The  remarkably  stable  tris-(N-hydroxethylethylenediamine)cobal1  (IIIj 
complex  ""'  shows  none  of  the  characteristic  read  ions  of  aliphatic  hydroxy] 
■roups,  even  though  the  usual  formulation  would  indicate  that  the  hydroxy] 

groups  are  not  coordinated  to  the  metal. 

Aromatic   Vmines 

Aromatic  diamine-  form  quite  unstable  coordination  compounds.  Hieber 

and  his  co-workers  have  shown  that  ortho-phenylenediamine  usually  occu- 
pies only  one  coordination  position366  but  that  the  para  isomer  occupies 
two367.  They  give  the  latter  the  rather  improbable  formula 


MI.— R— XH2 

/  \ 

X*M  MX. 

"    \  / 

XHo— R— XH2 


Diamino-biphenyls  seem  to  have  somewhat  stronger  donor  properties. 
2,2'-Diamino-biphenyl  forms  cobalt  (III)  complexes  corresponding  to  those 
of  ethylenediamine368  and  several  stable  compounds  of  benzidine  and  tolidine 
have  been  reported369.  The  empirical  formulas  indicate  that  these  bases 
occupy  two  coordination  positions,  but  there  is  no  evidence  that  both  amino 
groups  attach  themselves  to  the  same  metal  atom. 

Heterocyclic  Amines 

The  heterocyclic  amines,  although  they  contain  tertiary  nitrogen,  co- 
ordinate readily,  and  a  large  number  of  pyridine  complexes  has  been  de- 
scribed. In  general,  these  resemble  the  corresponding  ammonia  compounds. 
Davis  and  his  students370  have  found  the  stability  of  certain  nickel  and  zinc 
pyridine  compounds  to  decrease  as  the  temperature  is  lowered.  For  example, 
Nipy4(SCX)2  is  stable  at  room  temperatures,  but  decomposes  at  —3°.  It 
may  be  that  the  coordinating  tendency  of  the  thiocyanate  group,  relative 
to  that  of  pyridine,  increases  with  falling  temperature  till,  at  —3°,  it  dis- 
places the  pyridine. 

In  this,  as  in  other  cases,  chelation  greatly  enhances  coordination,  and 
metals  which  ordinarily  do  not  coordinate  with  nitrogen  form  stable  corn- 
Keller  and  Edwards,  ./.  Am.  Chem.  Soc,  74,  215  (1952). 
■6.  Bieber,  Schlieezmann,  and  Ries,  Z.  anorg.  allgem.  Chem.,  180,  89  (1929);  Hieber 

and  Ries,  /  </.  allgem.  Chem.,  180,  225  (1929). 

;».:    Bieber  and  Ries,  Z.  anorg.  allgem.  Chem.,  180,  105  (1929). 

Middleton:  thesis,  University  of  Illinois,  1938. 
569.  Tettamanzi,  Atii  accad.  Torino,  Clause  sci.fis.,  mat.  not.,  69,  225  (1935);  Spacu 

and  Dima,  Bull.  Soc.  Stiinte  Cluj,  8,  549  (1937). 
70.  Davis  and  Batchelder,  /.  Am.  Chem.  8oc.,  52,  4069  (1930);  Davis  and  Ou,  J. 
Am.  Chem.  Soc.,  56,  1061,  1064  (1934). 


68  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

pounds  with  a-pyridyl  hydrazine371,  a-pyridyl  pyrrole372,  2,2/-dipyridy] 
and  1  .  10-phenanthroline.  Many  coordination  compounds  of  2,2/-dipyridy] 
have  been  prepared.  As  far  as  is  known,  dipyridyl  always  acts  as  a  bidentate 
coordinating  agent.  The  stability  of  some  of  its  coordination  compounds 
is  i  ruly  remarkable.  For  example,  [Ni  dipya]4"4"  is  destroyed  only  very  slowly 
by  sodium  hydroxide  or  ammonium  sulfide373.  Prussian  blue  is  completely 
destroyed  in  the  cold  by  the  addition  of  2,2/-dipyridyl374. 

Research  on  the  dipyridyl  complexes  has  centered  largely  on  their  stereo- 
chemistry, the  stabilization  of  unusual  valence  states  by  coordination  with 
dipyridyl,  and  the  usefulness  of  the  complexes  in  analytical  chemistry 
(Chapter  20). 

While  many  substituted  derivatives  of  2,2'-dipyridyl  form  complexes, 
substituents  in  the  6,6'  positions  may  prevent  coordination.  Thus,  2-pyri- 
dyl-2/-quinoline,  and  2,2/-diquinoline  fail  to  react  with  octahedral  meta] 
ions375,  as  does  G^'-dimethyl^^'-dipyridyl376. 

kAN^|s|Aj  ChJIn^-^n^-CH3 

The  stabilization  of  valence  states  by  coordination  with  dipyridyl  is 
illustrated  by  the  cases  of  silver  and  chromium.  If  present  as  the  dipyridyl 
complex,  Ag(I)  can  be  oxidized  to  the  Ag(II)  complex  and  isolated  as 
[Ag  dipy2]++ 377  • 378.  Hein  and  Herzog379  report  that  the  reduction  of 
[Cr  dipy3]+++  in  the  presence  of  perchlorate  ion  gives  [Cr  dipy3]C104 ,  a 
deep  blue  compound,  unstable  in  air,  insoluble  in  water,  but  soluble  in 
methanol,  ethanol  and  pyridine. 

2,2',2"-Terpyridyl  and  2,2,,2",2'"-tetrapyridyl  coordinate  through  all 
of  their  nitrogen  atoms.  The  iron (II)  ion  fills  its  coordination  sphere  by 
combination  with  two  molecules  of  terpyridyl380;  the  platinum (II)  ion, 
having  a  coordination  number  of  only  four,  forms  compounds  of  the  type 
[Pt  tripyCl]Cl381.  Tetrapyridyl  gives  compounds  such  as  [Ag  tetrapy]N03 

371.  Emmert  and  Schneider,  Ber.,  66,  1875  (1933). 

372.  Emmert  and  Brandl,  Ber.,  60,  2211  (1927). 

373.  Jaeger  and  Van  Dijk,  Proc.  Acad.  Sci.  Amsterdam,  37,  10  (1934);  37,  618  (1934) 

39,  164  (1936);  Z.  anorg.  allgem.  Chem.,  227,  273  (1936). 

374.  Barbieri,  Atti  X°  congr.  intern,  chim.,  2,  583  (1938). 

375.  Smirnoff,  Helv.  chim.  Acta,  4,  802  (1921). 

376.  Willink  and  Wibaut,  Rec.  Trav.  Chim.,  54,  275  (1935). 

377.  Barbieri  and  Malaguti,  Atti  acad.  nazl.  Lincei,  Rend,  classe  sci.fis.,  mat.  e  nat. 

8,619  (1950). 

378.  Malaguti,  Atti  acad.  nazl  Lincei,  Rend,  classe  sci.fis.,  mat.  e.  nat.,  9,  349  (1950) 

379.  Bein  and  Herzog,  Z.  anorg.  allgem.  Chem.,  267,  337  (1952). 

380.  Morgan  and  Burst  all,  J.  Chem.  Soc,  1932,  20. 

381.  Morgan  and  Burstall,  J.  Chem.  Soc,  1934,  1498. 


GENERAL  SURVEY  69 


Co  I!    tetrapy]Clj ,  and  [Pt(II)  tetrapy][Pt< 
1 .  LO-Phenanthroline 


resembles  2 , 2'-dipyridyl  in  its  coordinating  ability,  but  gives  somewhat 
more  stable  complexes.  Even  beryllium  and  magnesium,  which  seldom  co- 
ordinate with  nitrogen  compounds,  form  complex  ions  containing  three 
molecules  of  1 ,  10-phenanthroline 


The  complexes  of  1 , 10-phenanthroline  are  chiefly  of  interest  because  of 
their  stereochemistry  (Chapter  8),  their  usefulness  in  analytical  chemistry 
(Chapter  20),  and  the  ability  of  1 , 10-phenanthroline  to  stabilize  unusual 
valence  states  of  some  of  the  metals  (Chapter  11). 

Hydrazine  Coordination 

Hydrazine  forms  many  coordination  compounds,  though  their  number 
is  somewhat  limited  because  of  the  reducing  action  of  hydrazine.  Com- 
pounds of  the  noble  metals,  and  of  metals  in  their  higher  oxidation  states, 
are  thus  quite  unstable.  Efforts  to  prepare  compounds  of  cobalt(III),  for 
example,  have  been  unsuccessful.  Most  hydrazine  complexes  which  have 
been  isolated  as  solids  do  not  contain  enough  hydrazine  molecules  to  fill  the 
coordination  sphere,  so  it  has  been  suggested  that  hydrazine  serves  as  a 
■■dentate  ligand.  This,  however,  necessitates  the  formation  of  a  three- 
membered  ring.  Goremykin884  treated  potassium  chloroplatinate(II)  with 
■iHi-HCl  and  obtained  a  product  which  he  believes  to  be  [PtCl2(X2H5)2] 
Cb--_MI,<>.  nn  heating,  this  goes  to  [PtCUNVH4)(X,II5)JCl,  which  reacts 
with  pyridine  to  form  [PtCb(\oH4).,].  If  this  interpretation  is  correct ,  hydra- 
zine is  acting  a-  a  monodentate  donor. 

Schwarzenbach  and  Zobist*8*,  using  the  Bjerrum  technique,  have  shown 
that  in  Bolution,  zinc  ion  can  coordinate  with  four  molecules  of  hydrazine 
and  nickel  ion  with  six.  Etebertus,  Laitinen  and  Bailar888  have  shown, 
polarographieally,  thai  zinc  ion  forms  a  tetrahydrazine  complex. 


K2.  MorgaD  and  Buret  all, ./.  Chem.  8oc.t  1938,  1072.  1675. 

'feifferand  Werdelman,  Z.  anorg.  Chem.,  261,  197    1950  , 

mykin,  Compt.  n  »<1 .  acad.  set.  I  .R  8  8    33.  227    1941). 
Schwarzenbach  and  Zobist,  //</>■.  ckitn.  Acta,  35,  1291     1952 

Laitinen,  and  Bailar,  •/.  .1/".  Chem.  Soc.,  75,  3051     1053);  Rebertus, 
thc-is,  University  of  Illinois,  L954. 


70  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Biguanide  Coordination 
Biguanide, 

NH2— C— NH— C— NH2 , 

II  II 

NH  NH 

is  a  remarkable  coordinating  material  which  has  been  studied  extensively 
by  Ray  and  his  students.  Only  two  of  the  five  nitrogen  atoms  coordinate 
these  are  on  different  carbon  atoms.  When  coordination  takes  place,  i 
hydrogen  atom  is  lost  from  each  molecule  of  biguanide.  The  uncoordinatec 
nitrogen  atoms  still  have  basic  properties,  so  salts  may  be  formed.  Man} 
substituted  biguanides  have  powerful  coordinating  ability.  Among  thes< 
are  phenylbiguanide, 

C6H5NHC— NH— C— NH2, 

II  II 

NH  NH 

N,N'  diphenylbiguanide,  N,N'-diethylbiguanide,  N-phenyl-N'methyl  bi 
guanide,  ethylenedibiguanide 

NH2C— NH— C— NHCH2CH2NH— C— NH— C— NH2 

II  II  II  II 

NH  NH  NH  NH 

and  meta-phenylenedibiguanide 

NH— C— NH— C—  NH2 

II  II 

NH  NH 

NH— C— NH— C— NH2 

II  II 

NH  NH 

Bivalent  metal  ions  such  as  Cu++  and  Ni++  form  stable  complexes  witl 
the  biguanides.  The  copper  complex  is  stable  enough  that  the  metal  in  i 
is  not  reduced  by  iodide,  sulfite,  thiosulfate,  or  other  anions  that  commonh 
reduce  copper(II)  to  copper(I)387.  When  the  biguanide  is  unsymmetricalh 
substituted,  as  in  phenylbiguanide,  the  copper(II)  and  nickel  complexe 
exist  in  two  cis  and  trans  forms388.  Ghosh  and  Chatterjee,  however,  isolate* 
only  one  form  of  each  of  the  metal  bis(methylphen}dbiguanides)389. 

387.  Ray  and  Bagchi, ./.  Indian  Chem.  Soc,  16,  617  (1939). 

388.  Ray  and  Chakravarty,  ./.  Indian  Chem.  Soc,  18,  609  (1941). 
"  rhosh  and  Chatterjee,  ./.  Indian  Chem.  Soc,  30,  369  (1953). 


GENERAL  SURVEY 


The  tervalent  metal  ions  give  remarkably  stable  complexes  of  the  type 


H      1 


N  =  C 


NH  =  C 


NH       -3HX    or     M; 


NH*/ 
73 


NH 


JMH-C 


'HN=C 


NNH2'HX 


NH 


Kay  and  his  students  have  published  a  long  series  of  articles  on  these  in- 
teresting substances390.  The  chromium  complexes  undergo  slow  hydrolysis: 

[Cr(BigH)3]X3  +  2H,0  ->  [Cr(BigH)2(OH)H20]\,  . 

The  hydroxoaquobis(biguanides)  can  hydrolyze  further  to  monobiguanides, 
but  these  are  unstable.  The  cobalt  (III)  complexes  are  more  stable  than 
those  of  chromium,  and,  in  fact,  have  been  shown  to  be  more  stable  than 
the  cobalt  (III)  ammines391.  The  tris(phenylbiguanide)cobalt(III)  ion  has 
been  resolved  into  its  optical  antipodes392.  Bis(biguanide)  cobalt  (III)  com- 
plexes of  the  types  [Co(BigH)2X2]  and  [Co(BigH)2XY]  exist  in  cis  and  trans 
forms393.  The  dibiguanides  are  quadridentate394,  apparently  attaching  them- 
selves to  the  metal  through  the  a, a  , 7, y'  positions.  Ray  and  Das  Sarma395 
have  prepared  the  cobalt  (III)  meta-phenylenedibiguanide  complexes 
[Co  phenylene(BigH)2X2]+++,  where  X  =  XH3  or  H20.  These  apparently 
have  the  trans  configuration,  for  oxalate  ion  does  not  seem  to  be  able  to  re- 
place the  two  coordinated  X  groups. 

Among  the  most  remarkable  derivatives  of  biguanide  is  the  silver(III) 
compound  of  ethylenedibiguanide: 


CHs NH 


CH2 NH 


NH3 


NH- 


NH 


The  high  valence  state  of  silver  is  quite  stable,  but  is  reduced  to  silver  (I ) 

390.  Ray  and  co-workers  (Sana,  Ghosh,  Dutt,  Battacnarya,  Buddhanta,  Chakravarty, 

Majumdar,  Das  Sarma):  ./.  Indian  Chem.  Soc,  14,  670  (1937);  15,  347,  350, 
353,  633  (1938);  16,  621,  629  (1939);  18,  289,  298  (1941);  19,  1  (1942);  21,  47 
(1944);  23,  73  (1946);  25,  589  (1948);  26,  137  (1949),  and  other  articles  not  in 
the  "series". 

391.  De,  Ghosh,  and  Ray,  ./.  Indian  Chem.  Soc.,  27,  193  '1950). 

392.  Shiddhanta,  Dutt,  and  Ray,  •/.  Indian  Chem.  Snr.,  27,  641    1950 

393.  Ray  and  Majumdar.  ./.  Indian  Chi  m.  8oc.,  23.  73    1946). 

394.  Ray  and  Shiddhanta.  ./.  Indian  Chem.  Soc.,  20,  200  (1943). 

395.  Ray  and  Das  Sarma.  ./.  Indian  Cht  26,  137    1949). 


72  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

by  iodide  ion'16.  The  conductivity  and  magnetic  susceptibility  are  consistent 
with  the  assumption  that  the  compound  contains  trivalent  silver  with  dsp2 
bonds897.  Measurement  of  stability  constants  shows  this  to  be  a  very  stable 
Bubstance898. 

Quinoline  and  Its  Derivatives 

The  nitrogen  of  quinoline  has  very  weak  donor  properties,  but  properly 
substituted  quinolines  form  stable  coordination  compounds.  8-Hydroxy- 
quinoline  is  a  strong  complexing  agent,  and  has  found  wide  use  in  analytical 
chemistry  (Chapter  20).  The  important  compounds  are  inner  complexes 


which  are  insoluble  in  water,  but  soluble  in  organic  solvents;  this  property 
is  utilized  in  the  separation  of  metal  ions,  just  as  it  is  with  the  inner  com- 
plexes of  the  1,3  diketones  (page  44)399.  Substituted  8-hydroxyquinolines 
can  often  be  used  to  advantage400. 

Inner  complexes  can  often  be  given  water  solubility  by  the  introduction 
of  a  highly  polar  group  into  the  complexing  agent.  If  this  substituent  is 
distant  in  the  molecule  from  the  donor  atoms,  it  does  not  disturb  the 
stability  or  nature  of  the  coordinate  bonds.  Thus,  Liu  and  Bailar401  pre- 
pared the  soluble  zinc  compound 


H    SOj 


and  resolved  it  into  its  optical  antipodes. 

396.  Ray,  Nature,  151,  643  (1943). 

397.  Ray  and  Chakravarty, ./.  Indian  Chem.  Soc.,  21,  57  (1944). 

398.  Sen,  Ghosh,  and  Ray,  ./.  Indian  ("hem.  Soc.,  27,  619  (1950). 
Mueller:  Ind.  Eng.  Client.,  Anal.  Ed.,  15,  270,  346  (1943). 

inn    Moeller  and  Jackson,  Anal.  Chem.,  22,  1393  (1950);  Moeller  and  Ramaniah. 

./.  Am.  Chem.  Soc,  76,  2022  (1954). 
101.  I- in  and  Bailar,  J   .1///.  Chem. Soc, 78, 5432  (1951). 
402,  Ley  and  Ficken,  Ber.,  50,  1133  (1917). 


CENERAL  SIRVEY 


73 


Picolinic  acid,  like  other  alpha  amino  acids,  forms  stable  coordination 

compounds,   many  of  which  arc  inner  complexes.   The  cobalt  (III)   com- 
pound40- is  illustrative  of  this  group. 

Phthalocyanines  and  Porphins 

When  o-dicyanobenzene,  o-cyanobenzamide,  or  related  substances  are 
heated  with  metals  or  their  sails,  a  vigorous  exothermic  reaction  takes 
place  and  metal  derivatives  of  phthalocyanine 


*^ 


are  formed.  The  metal  occupies  a  position  in  the  center  of  the  molecule.  If 
it  be  a  divalent  metal,  it  displaces  the  two  hydrogen  atoms,  and  coordinates 
with  all  four  of  the  nitrogen  atoms.  Trivalent  ions  seem  to  form  compounds 
of  the  type  [Phthalocyanine  M]X.  The  metal  derivatives,  like  the  parent 
substance,  are  deep  blue.  Many  of  them  are  extremely  stable,  being  un- 
affected by  any  but  the  most  vigorous  chemical  agents;  some  of  them  can 
be  sublimed  in  vacuo  above  500°C.  This  combination  of  properties  makes 
them  valuable  as  pigments  (Chapter  22). 

Phthalocyanine  is  closely  related  to  porphin: 


HC 


J NH        HN— ^ 

>CH 


N 1 

Which  also  gives  highly  colored  metallic  derivatives**.  Porphin  is  the  parent 
403.  Fischer  and  Gleim,  Ann.,  521,  157  (1935). 


71 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


substance  of  chlorophyll  and  hemin 
CH=CH2  CH3 

ch3y^>^ch^yVc?H5 


r\  /A 


CH 


coo  phytyl 

Chlorophyll    a 


CH2CH2COOH      CH 
r^SY^CH"^Y^CH=CH2 


CI 


Hemin 


Haemocyanin,  the  blood  pigment  of  molluscs  and  crustaceans,  is  a  copper 
compound  of  the  porphin  family. 

The  Azo  Group 

The  donor  properties  of  the  azo  group  are  weak404,  but  azo  compounds 
which  contain  a  strong  donor  (e.g.,  carboxyl  or  hydroxy)  in  a  position  ortho 
to  the  azo  group  form  very  stable  chelate  rings.  The  complexes  so  formed 
are  usually  highly  colored  and  find  use  as  dyes  and  pigments  (Chapter  22). 

The  diazo  amino  compounds  have  been  the  subject  of  an  interesting 
study  by  Dwyer405.  The  imino  hydrogen  atom  is  replaced,  at  least  in  some 
cases,  and  the  nitrogen  chain  forms  a  chelate  ring  with  the  metal,  thus 
occupying  two  coordination  positions.  Examples  are 


><<3 


N=N— NH 


)( 


-N=N— N- 


■O), 


Cu(<3-N==N-N-<3) 


and 


KO 


-N=N— N< 


Nitriles 


The  nitrogen  atoms  in  organic  nitriles  have  fairly  strong  donor  proper- 
lies,  especially  toward  the  heavier  metals.  The  halides  of  platinum  add 

404.  Kharasch  and  Ashford,  J.  Am.  Chem.  Soc,  58,  1736  (1936). 

405.  Dwyer,  J.  Am.  Chem.  Soc,  63,  78  (1941). 


GENERAL  SURVEY  75 

nitrttes  directly*1  to  form  PtX8(RCN)2  and  PtX.dUNi,  (R  may  be  either 
aliphatic  or  aromatic).  Halogens  readily  convert  the  platinum(II)  com- 
pounds to  the  platinum(IV),  which  arc  not  readily  reduced  again,  even  l>.\ 
formaldehyde4,  sulfur  dioxide  or  aluminum  and  hydrochloric  acid. 

Lebedinskii  and  Golovnya407,408  have  carried  out  the  following  reactions: 

|1V('1I<\     Cl       -    C  II  Nil:    -  [Pt(C\.H5NII.>)4((,,H,(,\>,][PtC,l1| 

IK'l 

[Pt(C2H5XH2),(M,l 

[Pt(NH,),Cl,]  +  CII3CX  -»  lPt(XH,)>(CH3CX)Cl]Cl 

+  NH4OH  '-+  [Pt(XH3)4(CH3CX)]Cl2 

KlPt(XH3)Cl3]  +  CH3CX  -*  [Pt(XH3)(CH3CN)Cl2] 

+  XH4OH  -»  [Pt(XH3)4(CH3CN)]Cl2 

The  platinum  in  the  compound  [Pt(XH3)4(CH3CX)]Cl2  does  not  seem  to 
show  the  usual  coordination  number  for  platinum,  and  doubtless  needs 
further  study.  Upon  heating  with  hydrochloric  acid,  this  compound  is 
converted  to  [Pt(XH3)3Cl]Cl. 

In  the  presence  of  acetonitrile,  copper(I)  coordination  compounds  are 
readily  formed.  They  oxidize  slowly  in  the  air409. 

Pseudohalides 

The  cyanide  ion  has  unshared  pairs  of  electrons  on  both  the  carbon  and 
the  nitrogen  atoms,  and  theoretically,  it  might  coordinate  to  metals  through 
either  of  these  pairs.  Actually,  it  seems  to  combine  preferentially  through 
the  carbon  atom,  and  the  simple,  mononuclear  cyanides  are  characterized 
by  a  metal-carbon  link  (page  87).  The  formation  of  the  carbon-metal  bond, 
however,  does  not  preclude  the  formation  of  a  coordinate  bond  between 
the  nitrogen  and  another  metal  atom.  The  "super-complex"  heavy  metal 
cyanides,  such  as  Prussian  blue,  are  probably  built  up  in  this  way,  as  are 
the  organo  gold  cyanides. 

The  thiocyanate  group  also  has  pairs  of  electrons  on  two  atoms,  and  con- 
ceivably can  coordinate  through  either  nitrogen  or  sulfur   (p.  57).   The 

406.  Ashford,  thesis,  University  of  Chicago,  1936.  Ashford  gives  references  to  several 
earlier  articles  on  platinum-nitrile  addition  compounds,  the  more  important 
being  Hofman  and  Bugge,  Ber.,  40,  1772  (1907);  Ramberg,  Ber.t  40,  2578 
(1907);  TschugaofT  and  Lebedinskii,  Compt.  and.,  161,  563  (1915). 
Lebedinskii  and  Golovnya,  Izvest.  Sektora  Platiny  i  Drugikh  Blagorod.  Metal., 
Inst.  Obschei  i  Neorg.  Kkim.,  A  had.  Nauk.  S.S.S.R.,  No.  22,  168  (1948);  cf. 
Chem.  Abs.,  44,  10566a,  (1950). 

408.  Lebedinskii  and  Golovnya,  Ann.  seeteur  platine,  Inst.  Chim.  pen      I    S  8.R.), 

No.  16,  57  '1939). 

409.  Morgan,  ./.  Chem.  80c.,  123,  2901  (1923). 


7<i  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

easy  formation  of  the  highly  colored  iron  (III)  and  cobalt  (II)  complexes 
makes  them  suitable  for  the  qualitative  detection  of  these  ions  in  solution. 
Several  investigations  have  been  made  to  determine  the  nature  of  the  ferric 
bhiocyanate  complex  which  exists  in  such  solutions.  M0ller410  showed,  by 
conductivity  measurements,  that  there  are  not  more  than  three  thiocyanate 
groups  attached  to  the  iron.  Bent  and  French411  and  Edmonds  and  Birn- 
baum412,  from  a  study  of  the  absorption  of  light  by  solutions  containing 
iron  (III)  and  thiocyanate,  came  to  the  conclusion  that  the  formula  of  the 
complex  is  Fe(NCS)++,  neglecting  hydration.  Schlesinger  and  Van  Valken- 
burgh413  showed  that  in  ether  solution,  the  [Fe(NCS)e]~  ion  is  present.  The 
entire  subject  has  been  well  reviewed  by  Lewin  and  Wagner414. 

The  coordinating  ability  of  the  azide  ion  was  first  studied  by  Strecker 
and  Oxenius415,  who  prepared  a  series  of  cobalt  (III)  compounds.  They  ob- 
tained the  ions  as-[Co(NH3)4(N3)2]+  cis-  and  trans-[Co  en2(N3)2]+,  and 
[Co  py4ClN3]+.  Linhard  and  Flygare416  prepared  several  salts  of  the  ion 
[Co(NH3)6N3]++  which  they  report  to  be  similar  in  color  to  [Co(XH3)5Cl]++. 
The  action  of  sodium  azide  on  a  solution  of  [Co(NH3)4(H20)2]+++  gave  a 
mixture  of  the  cis  and  trans  forms  of  [Co(XH3)4(X3)2]+417.  That  the  azide 
group  has  strong  donor  properties  is  shown  by  the  fact  that  triazidotriam- 
minecobalt  can  be  prepared  by  treatment  of  [Co(NH3)4(H20)2]+++, 
[Co(NH3)4(N3)2]+,  or  [Co(NH3)bN3]++  with  azide  ion418.  Straumanis  and 
Circulis419  prepared  stable,  slightly  soluble  compounds  which  they  believed 
to  be  nonelectrolytes  of  the  type  [Cu  R2(N3)2]  in  which  R  is  ammonia  or 
any  one  of  a  number  of  aliphatic  or  aromatic  amines.  They  also  obtained 
the  anions  [Cu(N3)6]4_,  [Cu(N3)4]==,  and  [Cu(N3)3]~.  All  of  the  azido  com- 
plexes are  unstable  and  explosive. 

Oximes 

The  coordinating  tendency  of  the  oximes  is  well  known.  A  lone  oxime 
group  does  not  coordinate  firmly,  but  when  it  forms  part  of  a  chelate  ring, 
the  oxime  nitrogen  has  very  strong  donor  properties  and  oximes  are  fre- 
quently  used  in  inorganic  analysis  (Chapter  20).  Metallic  ions  having  a 
coordination   number  of  six  combine  with  only  two  dioxime  groups,  the 

HO.  M0ller,  Kern.  Maunedsblad,  18,  138  (1937). 

111.  Bent  and  French,  ./.  Am.  Chem.  Soc,  63,  568  (1941). 

U2.  Edmonds  and  Birnbaum,  ./.  .1///.  Chem.  Soc,  63,  1471  (1941). 

H3.  Schlesinger  and  Van  Valkenburgh,  ./.  Am.  Chew.  Soc,  53,  1212  (1931). 

II  I.  lewin  and  Wagner,  ./.  Chem.  Ed.,  30,  445  (1953). 

n:>    Strecker  and  Oxenius,  Z.  anorg.  allgem.  Chem.,  218,  151  (1934). 

416.  Linhard  and  Flygare,  Z.  anorg.  Chem.,  262,  328  (1950). 

417.  Linhard,  Weigel,  and  Flygare,  Z.  anorg.  allgem.  ('hem.,  263,  233  (1950). 
lis.  Linhard  and  Weigel,  Z.  anorg.  allgem.  Chem.,  263,  245  (1950). 

il'.i.  Straumanis  and  Circulis,  Z.  anorg.  allgem.  Chew..  251,  341  (1943);  252,  9,  121 
(1943). 


GENERAL  SCh'VKY  77 

remaining  coordination  positions  being  filled  by  other  donors,  as  the  follow- 
ing cobalt  compounds  illustrate: 

[Co(HD),(\II..V.]\,  [Co(HD)sNH,X]  and  [( \><H1)),.\,]-  12l>. 

Aniline  and  substituted  anilines  can  replace  the  ammonia  in  the  cobalt 
compounds1-1.  Compounds  of  the  type  M'|M"'(  II 1  > ) - X - 1  containing  rho- 
dium4'- and  iridium1-"  have  also  been  described.  Only  one  =X()II  group 
from  each  onlt-dioxime  molecule  liberates  a  hydrogen  ion,  as  is  shown  by 
the  fact  that  the  mono-ethers, 

R C C R  424.425,426 

II  II 

lio— X     N — OCH, 
and  the  imino  and  methylimino  compounds 

CtHs— C— C— CH,  C>Ho— C— C— CH3  426 

II       II  and  ||       || 

H— X     X— OH  CH8— N     X— OH 

give  entirely  analagous  compounds.  On  the  other  hand,  both  hydrogens  of 
an  am  phi -dioximv  are  replaceable.  Nickel,  for  example,  forms  rather  poorly 
defined  compounds  of  the  type 

R— C C— R 

II  II 

X  X 

/  \     / 

O  Xi— o 

in  which  the  metal  is  evidently  attached  to  one  nitrogen  atom  and  one 
oxygen  atom1-7  m  1'-"-'.  Acids  rearrange  these  to  the  more  stable  red  modifi- 
cation. The  anti  and  amphi  forms  of  benzil  dioxime  react  with  palladium(II) 
ion  just  as  they  do  with  the  nickel  ion.  Syn  dioximes  do  not  yield  nickel 
derivatives1-"'  but  syn  benzildioxime  readily  forms  a  crystalline  palladous 

420.  Tschugaeff,  Ber.t  39,  2692    1906);  41,  2226  (1908). 

421.  Ablov,  Bull,  toe.  ckim.,  7,  151  (1940 

422.  Lebedinskii  and   Fedorov,  Ann.  secteur  platine,  Inst.  ckim.  </<  n .   (U.S.S.R.  , 

No.  15,  19    1038). 

423.  Lebedinskii  and  Fedorov.  Ann.  sectew    platine,  Inst.  ckim.  gen.     I    SSI:    . 

No.  15,  27     )'• 
421.  Thilo  and  Friedrich,  Ber.,  62,  2990    L929 

125.  Bradj  and  Muere,  •/    Chem.  Soc.,  1930,  1599. 

126.  Pfeiffer,  Ber.,  63.  1811     1930 

B7,  .Mack.  ./.  Chem.  8oc.,  103,  1317    1913 

42v  Hieber  and  Leutert,  Ber.,  62,  L839    1929 

420.  Meisenheimer  and  Theilacker,  Ann..  469,  L28    L95 


78  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

compound  which  is  said  to  have  the  structure 

(h  —  c  i C  — <j)430 

II  H 

N-O  O-N 

Pd 

Bryson  and  Dwyer431  report  that  /3-furfuraldoxime  reacts  with  copper, 
silver,  nickel,  and  cobalt.  The  a  isomer  does  not,  but  on  standing  in  solu- 
tion with  the  metal  salt,  it  changes  to  the  /?  isomer. 

The  Donor  Properties  of  Phosphorus  and  Arsenic 
Phosphine  Coordination 

The  action  of  phosphine  on  metallic  salts  has  been  studied  by  several 
investigators.  Most  metallic  ions  are  reduced  to  metal  or  to  phosphides432, 
433 ,  434^  kut  some  form  phosphine  addition  compounds.  Riban435  found  that 
a  solution  of  copper(I)  chloride  in  hydrochloric  acid  absorbs  phosphine 
readily,  forming  the  rather  unstable  compounds  CuClPH3  and 
CuCl-2PH3  .  Upon  gentle  warming,  these  compounds  liberate  phosphine, 
while  stronger  heating  generates  copper  phosphide.  These  results  have 
been  confirmed  and  extended  by  Scholder  and  Pattock436. 

Holtje  and  his  co-workers437  have  made  a  systematic  study  of  the  donor 
properties  of  phosphine.  They  found  that  in  its  ability  to  form  coordination 
compounds,  phosphine  resembles  hydrogen  sulfide  more  closely  than  it 
does  ammonia.  The  phosphinates  are  more  stable  than  the  sulfhydrates  in 
every  case  investigated.  Among  the  more  stable  compounds  reported  by 
these  investigators  is  A1I3-PH3,  which  may  be  sublimed  in  vacuo. 

Tertiary  Phosphine  and  Arsine  Coordination 

The  tertiary  organic  phosphines  and  arsines  have  strong  donor  properties, 
in  which  regard  they  are  in  sharp  contrast  to  the  tertiary  amines,  but  are 
similar  to  the  thioethers.  Even  the  stibines  can  form  addition  compounds438. 

430.  Dwyer  and  Mellor, ./.  Proc.  Roy.  Soc.  N.S.  Wales,  68,  107  (1935). 

431.  Bryson  and  Dwyer,  ./.  Proc.  Roy.  Soc,  N.S.  Wales,  74,  107  (1940). 

432.  Winkler,  Ann.  chim.  phys.,  Ill,  443. 

133.  Keilisch,  Ann.,  231,  327  (1885);  "Ueber  die  Einwirkung  des  Phosphorwasser- 

stoffs  auf  Metallsalzlosungen,"  Berlin,  1885. 
i:u    Scholder,  Apel,  and  Haken,  Z.  anorg.  allgem.  Chem.,  232,  1  (1937). 
135    Riban,  Compt.  rend.,  88,  581  (1879);  Bull.  soc.  chim.,  [2]  31,  385  (1879). 
436.  Scholder  and  Pattock,  Z.  anorg.  allgem.  Chem.,  220,  250  (1934). 
137.  Holtje,  Z.  anorg.  allgem.  Chem.,  190,  241   (1930);  209,  241  (1932);  Holtje  and 

Meyer,  Z.  anorg.  allgem.  ('hem.,  197,  93  (1931);  Holtje  and  Schlegel,  Z.  anorg. 

allgem.  Chem.,  243,  246  (1940). 
438.  Jensen,  Z.  anorg.  allgem.  Chem.,  229,  225  (1936). 


GENERAL  SURVEY  79 

The  strong  trans  influence  of  tertiary  phosphines  is  emphasized  by  the 
failure  of  Kumakov's  rule  (Chapter  9)  in  the  reaction  of  thiourea  with 
[PtCPEtjJJBrJ489.  The  use  of  several  of  the  phosphine  compounds  as  anti- 
knocks has  been  patented440. 

Organic  phosphines111  and  arsines442  are  often  identified  through  their 
highly  crystalline  mercuric  halide  complexes.  These  are  true  coordination 
compounds,  and  are  soluble  in  organic  solvents. 

The  most  common  phosphines  and  arsines  of  copper  are  CuX-AsRa  and 
CuX  >2AsRs ,  where  X  is  a  halide  ion.  Those  containing  a  single  coordinated 
arsine  group  are  tetrameric  while  those  containing  two  arsine  groups  are 
presumably  monomelic.  Nyholm448  has  reported  that  four  molecules  of 
diphenylmethylarsine  may  be  associated  with  a  single  copper(I)  ion,  as  in 
the  compounds  [Cu(AsMePh2)4][CuX2]  and  [Cu(AsMePh2)4]X.  This  ter- 
tiary arsine  also  forms  the  nonelectrolytic  complex  [Cu(AsMePh2)3X]. 
Similar  behavior444  was  also  noted  among  the  o-phenylenebis(dimethyl- 
arsine)  complexes  of  copper(I).  Gold  complexes  of  the  form  AuX-MR3, 
where  X  is  Cl~,  Br~,  or  XCS~,  and  M  is  arsenic  or  phosphorus,  are  mono- 
meric,  and  some  of  them  can  be  distilled  under  reduced  pressure.  There  is 
evidence445  that  the  corresponding  cyanides  and  iodides  are  polymeric.  The 
extreme  stability  of  these  substances  is  shown  by  the  fact  that  tributyl- 
phosphinegold(I)  chloride  may  be  volatilized  at  atmospheric  pressure  and 
triethylphosphinegold(I)  chloride446  dissolves  in  concentrated  hydrochloric 
acid  and  in  potassium  hydroxide  without  decomposition,  and  is  only  slowly 
reduced  to  metallic  gold  by  sulfur  dioxide.  The  vapors  of  AuCl-PBu3  deposit 
a  fine  film  of  gold  when  passed  through  a  heated  tube447.  Both  gold  (I)  and 
gold  (I  II)  complexes448  with  o-phenylenebis(dimethylarsine)  have  been 
reported449. 

439.  Grirrberg  and  Razumova,  Zhur.  Obschei  Khim.,  18,  282  (1948). 

440.  Bataafsche  Petroleum  Maatschappij,  French  Patent  805  666  (1936);  Peski  and 

Melsen,  U.S.  Patent  2  150  349  (1938). 

441.  Davies  and  Jones,  J.  Chem.  Soc,  1929,  33;  Da  vies,  Pearce,  and  Jones,  J .  Chem. 

Soc,  1929,  1262;  Jackson,  Davies,  and  Jones,  ./.  Chem.  Soc,  1930,  2298;  Jack- 
son and  Jones,  J .  ('hem.  Soc,  1931,  575;  Jackson,  Davies,  and  Jones,  ./.  Chi  m . 
Sue,  1931,  2109. 

442.  Jones,  Dyke,  Davies,  Griffiths,  and  Webb,  J.  Chem.  Soc,  1932,  2284;  Challenger, 

Higginbot torn,  and  Ellis,  ./.  Chem.  Soc,  1933,  95;  Challenger  and  Ellis,  J. 
Chem.  Soc,  1935,  398;  Challenger  and  Rawlings,  ./.  Chem.  Soc,  1936,  264; 
Blicke  and  Cataline,  ./.  Am.  Chem.  Soc,  60,  419  (1938). 
44:i.  Nyholm,  ./.  Chem.  Soc,  1952,  1257. 

444.  Kabesh  and  Nyholm,  ./.  Chem.  Soc,  1951,  38. 

445.  Dwyer  and  Stewart, ./ .  Proc.  lion.  Soc,  .V.N.  Wales,  83,  177  (1949). 
146.  Levi  Malvano,  .1///  accad.  Lincei,  [5]  17,  i.  847  (1908). 

447.  Mann  and  Wells,  Natun  .  140,  502  (1937);  Mam..  Wdls.  and  Purdie,  •/.  Chem. 

Soc,  1937,  1828. 

448.  Mann  and  Purdie,  ./ .  Chem.  Soc,  1940,  1235. 

449.  Nyholm,  Nature,  168,  705  (1951). 


80  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Recently,  compounds  of  o-phenylenebis(dimethylarsine)  (PDA)  have 
been  prepared  with  four  or  six  arsenic  atoms  coordinated  to  one  metal 
atom.  Iron  forms  complexes  of  the  formulas  [FeIII(PDA)2Cl2]C104  and 
[Fcn(Pl)A)2X,l(X  -  Br",  I",  or  SCN-)450.  The  magnetic  moments  of  the 
complexes  indicate  that  the  iron  atom  is  covalently  bound. 

Rhodium  (III)  halides  react  with  o-phenylenebis(dimethylarsine)  forming 
analagous  compounds451.  However,  upon  reaction  with  a  monodentate 
tertiary  arsine,  rhodium  (III)  halides  form  two  isomeric  compounds  con- 
taining three  moles  of  arsine  per  mole  of  rhodium.  These  are,  presumably, 
|Rh(AsR3)G][RhX6]  and  [Rh(AsR3)3X3]452.  Rhodium(II)  forms  a  variety  of 
other  complexes453  with  tertiary  arsines,  such  as  [Rh(AsR3)6]3[RhX5(AsR3)]2 
and  [Rh(AsR3)6][RhX4(AsR3)2]. 

Iridium(II)  and  iridium(III)  also  form  complexes  with  tertiary  arsines454. 
Dwyer,  Humpholtz,  and  Nyholm455  have  investigated  the  complexes  of 
diphenylmethylarsine  with  ruthenium  (II)  and  ruthenium  (III).  Ruthe- 
nium(II)  forms  the  complex  [Ru(AsR3)4X2]  while  ruthenium(III)  forms 
[Ru(AsR3)3X3]. 

The  preparation  of  nickel  complexes  of  trialkyl  compounds  of  the  group 
V  elements  has  been  especially  fruitful,  as  higher  valence  states  of  nickel 
are  probably  best  characterized  among  these  derivatives.  Jensen  and  Ny- 
gaard456  prepared  a  rather  unstable  pentacoordinate  triethylphosphine 
complex  of  tripositive  nickel  [NiBr3(PEt3)2].  The  corresponding  cobalt(III) 
complex,  CoCl3-2PEt3,  has  been  studied457;  it  is  probably  of  the  same 
configuration  as  the  nickel  complex  (see  Chapter  10,  page  392).  Nyholm458 
has  reported  [Xi(PDA)2X2]X,  containing  nickel(III),  and  [Ni(PDA)2X2] 
(C104)2 ,  which  contains  nickel (IV). 

This  work459  on  the  o-phenylenebis(dimethylarsine)  complexes  of  the 
metals  of  the  first  transition  series  has  been  quite  significant  from  the  the- 
oretical standpoint.  It  has  been  found  that  this  ditertiary  arsine  reacts  with 
transition  metal  ions  with  the  formation  of  strongly  covalent  bonds  only 
when  the  metal  ion  contains  d-electrons  which  are  not  involved  in  the 

450.  Nyholm,  J.  Chem.  Soc,  1950,  851.     . 

451.  Nyholm,  J.  Chem.  Soc,  1950,  857. 

452.  Dwyer  and  Nyholm,  J.  Proc  Roy.  Soc,  N.S.  Wales,  75,  140  (1942). 

453.  Dwyer  and  Nyholm,  J.  Proc.  Roy.  Soc.,  N.S.  Wales,  76,  133  (1942). 

i:»  1 .  Dwyer  and  Nyholm , ./ .  Proc.  Roy.  Soc,  N.S.  Wales,  77,  116  (1943) ;  79,  121  (1946) . 

455.  Dwyer,  Humpholtz,  and  Nyholm,  J.  Proc.  Roy.  Soc,  N.S.  Wales,  80,  217  (1947). 

456.  Jensen  and  Nygaard,  Acta.  Chem.  Scand.,  3,  474  (1949). 
157.  Jensen,  Nature,  167,  434  (1951). 

458.  Nyholm,  ./.  Chem.  Soc.,  1950,  2061;  1951,  2602. 

150.   Hurst  all  and  Nyholm,  /.  Chem.  Soc,  1952,  3570;  Nyholm  and  Sharpe,  /.  Chem. 
Soc,  1952,  3579. 


GENERAL  SURVEY  81 

hybridized  group  (see  Chapter  1).  It  has  been  concluded  thai  the  stability 
of  arsine  and  phosphine  complexes  depends  on  the  formation  of  double- 
bonds  between  the  metal  and  the  donor  atom.  This  conclusion  is  not  incon- 
sistent with  the  observation  that  the  more  stable  complexes  containing 

phosphorus-metal  or  arsenic-metal  bonds  occur  among  the  group  VIII  and 
IB  metals. 
Complexes  of  the  tritertiaryarsine,  methylbis(3-dimethylarsinopropyl)- 

arsine(TAS),  have  been  prepared  by  Barclay  and  Nyholm460  The  iron(III) 
complexes,  [Fein(TAS)Xj],  are  nonelectrolytes  and  exhibit  magnetic  mo- 
ments corresponding  to  one  unpaired  electron.  Cobalt  (II)  iodide  forms  a 
similar  complex,  [Con(TAS)I]I,  which  contains  a  single  unpaired  electron; 
air  oxidation  produces  diamagnetic  [Com(TAS)L].  Copper(I)  and  nickel(II) 
form  the  diamagnetic,  nonelectrolytic  complexes  [Cu(TAS)I]  and 
[Ni(TAS)IJ.  The  possibility  of  pentacoordinate  nickel(II)  here  is  especially 
interesting  in  view  of  the  previously  mentioned  observations  of  Jensen  and 
Nygaard. 

By  far  the  best  known  compounds  in  this  group,  however,  are  those  of 
platinum  and  palladium.  Cahours  and  Gal,  in  1870,  isolated  isomeric  forms 
of  PtCl2-2P(CH3)3,  PtCV2P(C2H5)3,  and  PtCl2-2As(C2H5)3.  Their  work 
was  confirmed  by  Klason461  and  by  Jensen,  who  extended  it  to  the  stibines438. 
Chatt  and  Wilkins462  studied  the  isomerization  of  palladium  compounds  of 
this  general  type  by  following  the  variation  in  dielectric  constant  of  their 
solutions.  Xo  detectable  amount  of  the  cis  isomer  of  the  arsine  or  phosphine 
complexes  appears  to  exist  in  solution,  while  as  much  as  40  per  cent  of  the 
stibine  complex  may  be  cis. 

Complexes  of  platinum (IV)  with  tertiary  arsines463  and  phosphines464 
have  been  prepared  in  isomeric  forms  by  oxidation  of  the  appropriate  iso- 
mers of  PtX2-2MR3  . 

Upon  treatment  with  ammonium  tetrachloropalladate(II),  the  bis(phos- 
phine)palladium(II)  compounds,  [Pd(PR3)2Cl2],  are  converted  to  the  di- 
nuclear  complexes,  [Pd2(PR3)2Cl4].  Mann  and  his  co-workers465  have  studied 
these  bridged  compounds  in  some  detail.  They  were  at  first  of  the  opinion 
that  several  forms  could  exist 

460.  Barclay  and  Xyholm,  (hem.  and  Ind.,  1953,  378. 

461.  Cahours  and  Gal,  Compt.  rend.  70,  1380;  71,  208  (1870).  Klason  and  Wanselin, 

./.  prakt.  Chem.,  [2]  67,  41  (1903). 

462.  Chatt  and  Wilkins,  J.  Chem.  Soc,  1953s  70. 

463.  Xyholm,  J.  Chem.  Soc,  1950,  843. 

464.  Chatt,  J.  Chem.  Soc,  1950,  2301. 

465.  Mann  and  Purdie,  Chem.  and  Ind.,  54,  M4  (1935);  ./.  Chem.  Sue,  1935,  1549; 

1936,  873;  Chatt  and  Mann.   ./.  Cheni.  Soc.  1938,    1949;    Chatt,   Mann,    and 
Wells,  J.  Chem.  Soc.  1938,  2086. 


Xo  <\  7rs 


82 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


FV\     /CIN     /CI 

Pd  Pd 

r3p//    xcr      Nci 


R3P\    /CIN    /PR3 
Pd  Pd 

c/    Nc/    XCI 


AND 


R3pn   /Ck     /Ci 

Pd  Pd 


Evidence  for  the  first  formula  was  found  in  the  fact  that  dipyridyl  and 
nil  rites  read  with  these  substances,  and  with  the  corresponding  arsine 
derivatives,  to  give  mixtures  of  compounds: 

(R3P)2PdCl2PdCl2  +  dipy  -»  (R3P)2PdCl2  +  [(dipy)PdCl,] 

(R3As)2PdCl2PdCl2  +  6KNO2  ->  (R3As)2Pd(N02)2  +  K2[Pd(N02)4]  +  4KC1 

On  the  other  hand,  aniline,  toluene,  and  pyridine  give  good  yields  of  mono- 
phosphine  (or  arsine)  derivatives: 


\     /     \     S 

Pd  Pd 

C\y     NCIX    XPR3 


+  4. 


NH; 


R3PX    /CI 


CI 


Pd 


NH2((> 


Ethylenediamine  splits  the  butyl  phosphine  compound  unsymmetrically  in 
benzene,  but  symmetrically  in  alcohol. 

Later  evidence,  however,  showed  the  earlier  hypothesis  to  be  incor- 
rect466, 467;  the  dimeric  molecule  apparently  always  has  the  symmetrical 
structure.  It  was  shown  that  compounds  of  the  type  PtCl2-(PR3)2  are  not 
primary  products  of  the  splitting,  but  are  formed  by  secondary  reactions. 
The  unsymmetrical  formulas  for  the  bridged  complexes  would  indicate 
that  compounds  of  the  types 


(CH3)2 

Pd  Pd  AND 


■As'     ^cr 

(CH3)2 


kCI 


(C6H5)2 
^As  .CL  /CI 

HaC^      \     /     \     / 
Pd  Pd 

^As         Ncr         Nci 

(C6H5)2 


should  exist.  Chatt  and  Mann466  were  unable  to  prepare  any  such  com- 
pounds, but  obtained 


[PdCI4]     , 


ar 


PdCI, 


466.  Chatl  and  Mann, ./.  Chem.  Soc,  1939,  1622. 

167 .  Mann  and  Wells,  ./.  Chem.  Soc,  1938,  702;  Wells,  Proc.  Roy.  Soc.  London,  A167, 
169  (1938). 


GENERAL  SURVEY 


83 


and  several  other  interesting  substances.  Chatl  has  extended  this  work  to 
include  the  tripropylstibine complex488  of  platinum(II).  This  Bpecies  behaves 
essentially  as  the  arsine  and  phosphine  complexes. 

[nteresting  examples  of  phosphine  complexes  with  bridging  groups  other 

than  the  halide  ions  are  found  in  the  ethyl  mercaptan  and  oxalate  bridge* 


Et 

pc        Pt 
a'    xs^  xpr. 

Et 


R3PV       .0-0  =  0.  .CI 

xPt/  I  XPt 


CI 


/  \ 


o— c=o 


s    \ 


PR. 


The  ethyl  mercaptan  complex  exists  in  two  (cis-trans?)  forms.  A  related 
compound  with  thiocyanate  bridges  is  reported  to  exist  in  the  isomeric 
forms 


R3P 


\ 


,C\ 


Pt 


\ 


Pt 


SCN 


NCS/   ^c/    NPR3 


R3PX     y 


CN 

^Pt         'Pt 

c/    NS^    XPR3 
CN 


The  reported  isomerism  of  the  bridged  compound  trichlorotris(diphenyl- 
methylarsine)copper(I)copper(II)  is  interesting470.  Copper(I)  is  tetrahed- 
ral471  while  copper  (I  I)  is  planar  so  that  the  isomerides  were  thought  to  be 


CL             CL            AsMe(J)p 
Cu             Cu 

AND 

())2MeASx       Clx       CI 
Cu          Cu 

(J^MeAs^  XCI^    NAsMe<t>2 

^MeAs7      CI        XAsMe())2 

L          TETRAHEDRAL               PLANAR 

~~    TETRAHEDRAL               PLANAR 

However,  it  has  since  been  contended  that  these  substances  are  actually 
complexes  of  diphenylmethylarsine  oxide  and  that  the  reported  isomerism 
was  associated  with  an  impurity  in  one  of  the  form- 

The  existence  of  bridged  arsine  and  phosphine  complexes  containing  two 
different  metals  is  reported  by  Mann  and  his  co-workers47*.  A  series  of 
compounds  involving  palladium(II)  or  cadmium  bridged  to  mercury  is 
exemplified  by 

468.  Chart.  ./.  Chem.  So,-.,  1951,  652. 

hatl  and  Hart,  /.  Chem.  Soc.,  1953,  260;  Nature,  169,  673  (1952  ;  Chatt,  Mann, 
and  Wells,  ./    -  Joe.,  1938,  2086. 

470.  Mellor,  Burrows,  and  Morris,  Nature,  141,  114    1038). 

471.  Mellor  and  Craig,  ./.  Proc.  Roy.  Soc.,  X  S    Wales,  75.  27    1941 

472.  Nyholm,  J.  CI         -      .  1951,  L767. 

473.  Mann  and  Purdie,  ./ .  Chem.  Soc,  1940,  1230;  Allison  and  Mann,  ./.  Chen     - 

1949,  2915. 


SI  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Pr3Asx     ^Br\      ^Br 


M  Hq 

Br  NBr  ^AsPr3 

The  compounds  SnX4-2PR3(X   =    Cl~  or  Br~)  also  form  mixed-metal, 
bridged  complexes  with  mercury (II)  or  palladium(II). 

CI 

R3PX|   /C!N     /PR3 
Sn  M 

ci'|  NcK   xc. 

CI 

Tertiary  phosphines  react  with  the  carbonyls  of  iron,  cobalt  and  nickel 
to  produce  mixed  phosphine-carbonyl  complexes: 

Ni(CO)4  +  PR3  (or  2PR3)  ->  [Ni(PR3)(CO)3]  (or  [Ni(PR3)2(CO)2]) 

Their  catalytic  behavior  in  the  reactions  of  acetylene  has  been  discussed 
by  Reppe  and  Sweckendich474. 

Cacodyl  oxide,  (CH3)2As — O — As(CH3)2 ,  which  might  be  expected  to 
coordinate  through  both  arsenic  atoms,  does  so  with  difficulty,  and  it 
usually  occupies  only  one  coordination  position475. 

Phosphorus  (III)  Halide  Coordination 

The  "double  compounds"  formed  by  phosphorus(III)  chloride  and  bro- 
mide with  metal  halides  certainly  contain  true  coordinate  links,  the  phos- 
phorus acting  as  the  donor  atom.  Platinum(II)  chloride  and  phospho- 
rus(III)  chloride,  for  example,  give  the  highly  crystalline  compounds 
PtCl2  •  PC13  and  PtCl2  •  2PC13476.  These  react  with  water  to  give  PtCl2  •  P(OH)3 
and  PtCl2-2P(OH)3 ,  and  with  alcohols  to  form  the  corresponding  esters. 
Molecular  weight  determinations  have  shown  the  ethyl  ester  of  the  mono- 
phosphine  complex  to  be  dimeric,  and  hence  (presumably) 

Ck        ^Clx      /P(OR)3 
.Pt  Pt 

(ro)3p^    xcr      XCI 

474.  Reppe  and  Sweckendich,  Ann.,  660,  104  (1948). 

475.  Jensen  and  Frederiksen,  Z.  anorg.  allgem.  Chem.,  230,  34  (1936);  Baudrimont, 

Compt.  rend.,  55,  363  (1862);  Ann.  chim.  phijs.,  [4]  2,  5  (1864);  "Recherches  sur 
les  chlorures  et  les  bromures  de  phosphore,"  Paris,  1864. 

476.  Schutzenberger,  Compt.  rend.,  70,  1287,  1414  (1870);  Bull.  soc.  chim.,  [2]  14, 

97,   178   (1870);  Schutzenberger  and   Fontaine,  Bull.  soc.  chim.,  [2]  17,  386, 
82  (1872). 


GENERAL  SURVEY  85 

The  ester  PtCl2-2P(OCH8)i ,  however,  is  monomeric477: 

(CH30)3P  /CI 

Pi 

(CH30)3P^     XCI 

The  acids  and  esters  react  with  silver  salts,  with  replacement  of  the  chlo- 
ride groups,  the  acids  at  the  same  time  forming  silver  salts476: 

(AgO)v  N     /.w, 


kPN     /NO: 
Pt/ 
(AgO)3P/     XN03 


The  dimeric  esters  are  readily  split  by  substances  which  have  fairly  strong- 
donor  properties17'''  ,7s  179.  Aniline,  for  example,  gives  cis  and  trans 
[PtClo  P(OC2H5)3(C6H5XH2)]480.  [PtCl22P(OC2H5)3]  adds  two  molecules  of 
ammonia,  both  chlorides  becoming  ionic.  Platinum(II)  chloride  also  forms 
white,  crystalline  [PtCV  (PF3)2]  and  red  [PtCl2- (PF3)]2  when  treated  with 
phosphorus(III)  fluoride481.  Both  substances  are  sensitive  to  moisture;  how- 
ever, the  white  compound  is  thermally  stable  and  may  be  refluxed  in  a  dry 
atmosphere  without  substantial  decomposition.  It  is  interesting  that  phos- 
phorus(III)  fluoride,  which  has  no  appreciable  basic  character,  should  form 
such  stable  complex  compounds.  This  behavior  is  attributed  by  Chatt482 
to  the  formation  of  a  double-bond  between  the  phosphorus  and  the  plati- 
num. 

As  might  be  expected,  palladium(II)  chloride  forms  analagous  com- 
pounds483. The  corresponding  iridium  compounds,  which  have  been  studied 
by  Geisenheimer484  and  by  Strecker  and  Schurigin485,  are  reported  to  be 
much  more  stable  than  those  of  platinum  and  palladium.  IrCl3-3PCl3  does 
not  react  with  cold  alcohol,  with  cold  concentrated  sulfuric  acid,  or  with 
organic  bases. 

477.  Rosenheim  and  Loewenstamm,  Z.  anorg.  Chem.,  37,  394  (1903). 

178.  Schutzenberger  and  Fontaine,  Bull.  soc.  chim.,  [2]  18,  101,  148  (1872). 

479.  Rosenheim  and  Levy,  Z.  anorg.  Chem.,  43,  34  (1905). 

480.  Troitskaya,  Zhur.  Priklad.  Khim.s  26,  781  (1953). 

481.  Chatt  and  Williams.  ./.  Chun.  Soc,  1951,  3061. 

482.  Chatt :  Nature,  165,  637    I960). 

///.  rend.,  116,  176  (1892);  123,  603  (1896);  The  author's  name  is  spelled 
Pinck  in  the  second  reference,  but  it  evidently  refers  bo  the  same  man. 

484.  ( ieisenheimer,  Ann.  chim.  phye.,  [6123,231  (1891);  "Sur  lea  chlorures  el  bromures 
double  d'iridium  e1  de  phoephore,"  Paris,  1891. 

185.  Strecker  and  Schurigin,  Ber.,  42,  1 7 « . 7  1909  ;  Schurigin,  "Die  Einwirkung  von 
Phosphor-halogeniden  auf  die  Metalle  der  Platingruppe,"  Grieswald,  1909 


86  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

( rold(I)  halides  form  a  similar  series  of  compounds,  e.g.,  AuCl  •  PC13486 • 487. 
It  is  not  possible  to  obtain  AuCl  -P(OH)3  ,  for  the  phosphorous  acid  reduces 
the  gold  to  the  metallic  state,  but  AuCl-P(OC2H5)3  is  quite  stable,  and  is 
not  reduced  by  sulfur  dioxide.  It  is  soluble  in  ammonium  hydroxide  with 
the  formation  of  AuClP(OC2H5)3-2NH3 ,  from  which  acids  reprecipitate 
it  in  the  original  form.  The  methyl  ester,  AuCl-P(OCH3)3 ,  has  been  pre- 
pared by  the  action  of  methanol  on  AuCl-PCl3  and  by  the  union  of  tri- 
methyl  phosphite  and  gold  (I)  chloride.  The  phenyl  ester  was  prepared  by 
the  second  method488. 

A  series  of  nickel (0)  compounds  with  phosphorus(III)  halides  has  been 
prepared  from  nickel  tetracarbonyl489.  These  compounds  are  of  the  composi- 
tion Ni(PX3)4  .  Phosphorus(III)  fluoride  does  not  completely  replace  the 
carbonyl  groups  from  nickel  tetracarbonyl;  however,  the  compound 
[Xi(PF3)4]  can  be  prepared  by  the  following  reactions 

[Ni(PCl3)4]  or  [Ni(PBr3)4]  +  4PF3  ->  [Ni(PF3)4]  +  4PC13  or  4PBr3 

[Ni(PCl3)4]  +  4SbF3  ->  [Ni(PF,)4]  +  4SbCl3 

Antimony  (III)  chloride  reacts  with  nickel  and  iron  carbonyls  giving  the 
products  [Ni(CO)3SbCl3]  and  [Fe(CO),(SbCl3)2],  respectively490. 

Copper(I)  chloride  reacts  with  phosphorus(III)  chloride491,  but  the 
compound  so  formed  is  reactive  and  unstable.  With  methyl  alcohol  it  gives 
a  mixture  of  copper(I)  chloride  and  CuClP(OCH3)3  .  Iron  (III)  chloride 
gives  the  volatile  compound  FeCl3  ■  PC13492. 

The  Donor  Properties  of  Carbon 
There  are  three  great  classes  of  coordination  compounds  in  which  carbon 
apparently  shares  electrons  with  metals — the  ethylenic  compounds,  the 
metal  carbonyls,  and  the  complex  cyanides.  The  first  two  of  these  are  the 
subjects  of  special  chapters  in  this  book,  so  this  section  will  be  devoted  to 
the  cyanides  and  the  closely  related  complexes  of  metal  ions  with  iso- 
n  it  riles. 

Cyanide  Coordination 

The  cyanide  ion  has  unshared  electrons  both  on  the  carbon  atom  and  on 
the  nitrogen  atom,  and  one  might  expect  to  find  isomeric  series  of  complexes 

486.  Lindet,  Compt.  rend.,  98,  1382  (1884);  101,  164  (1885);  103,  1014  (1886);  Bull, 
soc.  chim.,  [2]  42,  70  (1884);  Ann.  chim.  phys.,  [6]  11,  177  (1887).  Most  of 
Lindet 's  conclusions  were  Later  confirmed  by  Levi-Malvano  (Ref.  446). 

187.  Arbuzov  and  Lovoastrova,  Doklady  akad.  Nauk.  S.S.S.R.,  84,  503  (1952). 

188.  Arl.uzov  and  Shavska,  Doklady  akad.  Nauk.  S.S.S.R.,  84,  507  (1952). 

189.  Irvine  and  Wilkinson,  Science,  113,  742  (1951);  Wilkinson,  J.  Am.  Ckem.  Soc, 

73,  559  (1951). 

190.  Wilkinson,  ./.  .1///.  Chem.  Soc,  73,  5502  (1951). 

191.  Davis  and  Ehrlich,  J.  Am.  Chem.  Soc,  58,  2151  (1936). 

192.  (Jrbain,  British  Patenl  312  685  (May  31,  1928). 


GENERAL  SURV1-)  87 

corresponding  to  the  nitriles  and  isonitriles  of  organic  chemistry.  Such,  how- 
ever,  have  not  been  observed,  so  it  is  concluded  thai  the  attachment  of  the 

cyanide  ion  to  any  given  metal  ion  always  takes  place  through  the  same 
atom.  It  is  conceivable  that  some  metals  share  electrons  with  the  carbon 
and  others  with  the  nitrogen,  but  there  is  no  experimental  support  for  such 
a  hypothesis.  The  preponderance  of  the  evidence  indicates  thai  in  complexes 
of  the  type  [M(CN)J*~,  union  is  always  through  the  carbon. 

Carbon  and  nitrogen  are  so  close  together  in  atomic  number  that  only 
the  most  accurate  x-ray  measurements  can  distinguish  between  them.  Such 
distinction  is  particularly  difficult  in  the  complex  metal  cyanides,  where  the 
heavy  metal  atom  masks  the  lighter  nonmetals.  A  few  such  accurate  meas- 
urements have  been  made,  and  all  of  them  support  the  hypothesis  that  the 
metal  is  attached  to  carbon493.  Holzl  and  his  co-workers  have  come  to  the 
same  conclusion  from  chemical  studies.  They  alkylated  a  number  of  metal 
cyanide  complexes,  and  obtained  compounds  which  upon  decomposition 
yielded  alkyl  isonitriles494.  In  some  cases,  alkyl  amines  were  also  obtained, 
but  in  no  case  were  ammonium  salts  formed  in  significant  amounts.  Infrared 
spectral  work  by  L.  H.  Jones495  indicates  the  existence  of  the  carbon-metal 
bond.  He  found  that  the  pattern  of  infrared  active  vibrational  frequencies 
for  the  compounds  KAu(CM-X14)2 ,  KAu(C12N14)(C13N14),  KAu(C13N14)2 , 
and  KAu(C12X14)(C12X15)  indicates  that  the  bonding  is  through  the  carbon. 
Jones  also  found  that  in  [Au(CX)2]-  theC=N  force  constant  is  greater  than 
in  CH3C=X.  In  CH3X=C  the  C=X  force  constant  is  considerably  less 
than  in  CH3C=X.  This  also  indicates,  but  does  not  prove,  that  the  CN  is 
bound  to  the  gold  through  the  carbon. 

The  cyanide  ion  is  a  powerful  coordinating  agent,  and  it  frequently  dis- 
places all  other  groups  from  the  coordination  sphere,  forming  ions  of  the 
type  [M(CX)X]V~.  Exceptions  to  this  are  found  among  the  carbonyl  and 
nitrosy]  cyanides  (Chapter  16),  and  in  such  complexes  as  [Co  en2(CX)2]+  496a, 
[Co(CX)5OH]=496b,  [Co(CX)4(OH)2]=  and  [Fe(CX)5H20]=497. 

Examples  of  unusual  and  variable  coordination  numbers  are  fairly  com- 
mon among  the  cyano  complexes.  Thus,  Adamson498  believes  that  the 
formula  for  potassium  cobalt (II)  cyanide,  which  has  long  been  written 

Hoard,  Z.  Krist.,  84,  231  (1933);  Hoard  and  Nordflieck,  ./.  .1///.  Chetn.  Soc,  61, 
2853    L939  ;  Powell  and  Bartindale,  •/.  Chem.  Soc.,  1945,  799. 

194.  Holzl,  Monats.,  48,  71    (1927);  51,  1,  397  (1929);  Holzl  and  Xenakis,  Monats., 

48,689    L927  ;  Holzl  and  Viditz,  Monats.,  49, 241  (1928);  Holzl  and  Krichmayr, 
Monats.,  51,  397    1929);  Holzl,  Meier-Mohar,  and  Viditz.  Monats.,  52,  73;  53  54, 
237    L929  . 

195.  I..  II.  Jones,  private  communication. 

496a.  Ray  and  Sauna,  ./.  Indian  clem.  Soc.  28,  59    1951 
496b.  Smith..  Kleinberg,  and  Griswold,  /.  Am.  Chem.  Soc.,  75,  149  (1953). 
197.  Hieber,  Nast,  and  Bartenstein,  Z.  anorg.  allgem.  Chem.,  272,  32    1953). 
A-damson,  /.  Am,  Chem.  Nor.  73,  5710    1951). 


88  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

K4Co(CN)6 ,  is  actually  K3Co(CN)5 .  Even  in  aqueous  solution,  the  co- 
ordination number  of  five  is  maintained.  The  familiar  copper  cyanide  plating 
hath  contains  both  [Cu(CN)2]~  and  [Cu(CN)3]=,  the  latter  predominating. 
The  infrared  spectral  studies  of  L.  H.  Jones  and  Penneman499  have  shown 
that  in  aqueous  solutions  containing  silver  ion,  increasing  concentration  of 
cyanide  ion  brings  about  the  successive  formation  of  [Ag(CN)2]~, 
[Ag(CN)3]=,  and  [Ag(CN)4]-.  The  tricyano  complex  exists  over  a  wide  range 
of  concentrations,  the  equilibrium  constants  between  the  successive  com- 
plexes being  K3l2  =  0.20  d=  0.05  and  K4,3  =  13.4  ±  4.  Under  the  same 
conditions,  gold  (I)  forms  only  [Au(CN)2]~.  The  gold  and  silver  complexes 
are  both  adsorbed  on  anion  exchange  resins,  but  the  gold  complex  is  held 
much  more  firmly  than  is  that  of  silver. 

Adamson,  Welker,  and  Volpe500  have  studied  the  exchange  of  radiocy- 
anide  with  some  heavy  metal  cyanides.  The  rate  of  exchange  for  complexes 
in  which  the  metal  shows  a  coordination  number  of  two  or  four  was  found 
to  be  immeasurably  rapid.  With  hexacyano  manganate(III)  it  is  rapid  but 
measurable,  and  with  the  other  hexacyano  complexes  it  is  negligible.  Thus, 
the  rate  seems  to  be  a  function  of  coordination  number  rather  than  thermo- 
dynamic stability.  A  more  detailed  study  of  the  exchange  between 
[Mn(CN)6]-  and  CN~  showed  that  the  rate  of  this  reaction  is  proportional 
to  the  concentration  of  cyanomanganate(III),  but  independent  of  the  con- 
centration of  cyanide  ion.  The  authors  postulate  the  existence  of  an  unstable 
intermediate,  [Mn(CN)6H20]-,  in  which  manganese  shows  a  coordination 
number  of  seven.  This  is  possible  for  manganese  (III),  but  not  for  chrom- 
ium(III),  iron(III),  or  cobalt(III). 

The  cyanide  group  acts  as  a  bridging  group  in  polynuclear  complexes, 
both  the  carbon  and  the  nitrogen  atoms  sharing  electrons  with  the  metals. 
An  interesting  example  of  this  is  found  in  dipropyl  gold  cyanide,  which  has 
been  shown  to  be  tetrameric  and  to  which  the  structure 

R  R  501 

I  I 

R— Au— C=N— Au— R 

I  I 

N  C 

III  III 

C  N 

I  I 

R— Au— N=C— Au— R 

I  I 

R  R 

Jones  and  Penneman,  J.  ('Item.  Phys.,  22,  965  (1954). 

500.  Adamson,  Welker,  and  Yolpo,  ./.  Am.  Chem.  Soc,  72,  4030  (1950);  Adamson, 

Welker,  and  Wrighl ,  ./ .  Am.  Chem.  Soc,  73,  4786  (1951). 

501.  Phillips  and  Powell,  I'roc.  Roy.  Soc.  London,  A173,  147  (1939). 


GENERAL  SURVB1  89 

has  been  assigned.  The  polymeric  structure  is  dictated  by  the  necessity 
of  coordinating  tour  donor  groups  to  each  gold  atom. 

Upon  heating,  a  compound  of  the  type  [R»Au(CN  )]i  decomposes  to  form 
a  substance  of  the  empirical  formula  R  An  CN,  which  Gibson802  believes  is  a 
linear  polymer 

R  R 

I  I 

— Au— C  X— Au— C  X— Au— 

I  I 

R  R 

In  spite  of  the  stability  of  the  gold-carbon  bond,  the  tetramer  [R2Au(CN)]4 
is  destroyed  by  ethylenediamine,  giving  [R2Au  en][R2Au(CN)2]503. 

In  the  "simple"  cyanides  of  the  heavy  metals,  linking  between  the  metal 
atoms  takes  place,  the  complexity  of  the  resulting  structure  depending  upon 
the  relative  numbers  of  cyanide  ions  and  metal  atoms,  and  the  coordination 
number  of  the  latter.  Silver504  and  gold505  cyanides  have  been  shown  to  con- 
tain infinite  chains  of  metal  atoms  held  together  by  cyanide  bridges.  Mer- 
cury! Hi  cyanide  is  also  said  to  have  a  linear  structure506,  while  the  closely 
related  zinc507  and  cadmium503  compounds  are  three-dimensional  super 
complexes.  Tetracovalent  metals  which  form  planar  bonds  form  layer  struc- 
tures. Thus,  palladium  cyanide  is 

i  i 

— Pd — CseN— Pd— 

I  I 

N  C 

III  III 

C  N 

I  I 

— Pd— N=C— Pd— 

I  I 

Long509  has  studied  the  rate  of  exchange  between  [Ni(CN)4]=  and  CN""  and 
between  [Ni(CN)4]"  and  Xi++.  The  first  of  these  is  fast  but  the  second  is 
Blow  compared  with  the  rate  of  precipitation  of  these  ions  when  they  are 

502.  Gibson,  Proc.  Roy.  Soc.  London,  A173,  160  (1939). 

503.  Brain  and  Gibson,  ./.  Chem.  >SW.,  1939,  762. 

504.  Braekken,  K<jl.  Norske  Yidensk.  Sehkohs.  Forh.t  II  1929,  123;  West,  Z.  Krist., 

90,  555  (1 
105.  Zhandov  and  Shugam,  Acta  Physicochim.  V  R  S  S    20,  253  (1945). 

506.  Hassel,  Z. Krist.,  64, 218  (1926);  Zhandov  and  Shugam,  C.R.  Acad.  Sci.  U.R.S.S., 

45,  295    1944). 

507.  Zhandov,  C.  /.'.  Acad.  Set.  \    R  8.S    31.  360    I'M  I   . 

508.  Shugam  and  Zhandov,  Acta  Physicochim.  ('/:.<  >'..  20,  _'!7     1946). 

509.  Long,  ./.  .1-.  Chem.  Soc,  73,  537  (1951). 


90  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

mixed  with  each  other.  Long  concludes  that  the  nickel  is  bound  in  two 
different  ways,  and  that  nickel  cyanide  may  be  formulated  as  nickel  tetra- 
cyanonickelate(II),  Ni[Ni(CN)J.  Hume  and  Kolthoff510  have  come  to  the 
same  conclusion  from  polarographic  studies. 

Heavy  metal  salts  of  the  hexacyano  complexes  have  been  studied  ex- 
tensively, especially  the  ferro-  and  ferricyanides.  It  has  long  been  known 
that  the  heavy  metal  ferrocyanides  are  not  simple  salts  of  H4[Fe(CN)6]. 
For  example,  Reihlen  and  Zimmermann511  showed  that  ammonia  will  ex- 
tract only  part  of  the  cadmium  from  cadmium  ferrocyanide.  The  com- 
plexity of  these  materials  is  indicated  also  by  their  great  insolubility  and 
their  colloidal  nature.  Turnbull's  blue,  made  from  an  iron(II)  salt  and  a 
hexacyanoferrate(III),  and  Prussian  blue,  made  from  an  iron(III)  salt  and 
a  hexacyanoferrate(II),  were  long  thought  to  be  different  materials,  but 
both  chemical  and  physical  studies  have  shown  them  to  be  identical.  This 
comes  about  because  the  ions  involved  react  with  each  other  readily : 

Fe+++  +  [Fe(CN)6]4-  ^±  Fe++  +  [Fe(CN)6]s  512. 

When  union  between  the  simple  cation  and  the  complex  anion  takes  place, 
the  nitrogen  of  each  cyanide  group  shares  electrons  with  an  iron  atom,  which 
in  turn  shares  electrons  with  nitrogen  atoms  from  other  complex  anions. 
Thus,  a  super  complex  is  built  up. 

The  x-ray  studies  of  Keggin  and  Miles513  have  revealed  the  structure  of 
the  ferro-  and  ferricyanide  pigments.  In  Berlin  green,  Fe[Fe(CN)6],  which 
is  made  by  the  reaction  of  Fe+++  and  [Fe(CN)6]-,  the  iron  atoms  form  a 
cubic,  face-centered  lattice.  (Fig.  1.2).  This  arrangement  is  retained  in 
"soluble"  Prussian  blue  (Fig.  1.3),  in  which  half  of  the  iron  atoms  are  in 
the  3+  state  and  half  in  the  2+  state.  It  is  impossible  to  distinguish  between 
these,  and  it  is  probable  that  they  are  identical,  the  charge  distribution 
being  levelled  out  by  resonance.  One  potassium  ion  (or  another  univalent 
ion)  must  be  present  for  each  iron (II)  ion  to  maintain  electroneutrality. 
These  univalent  cations  are  located  in  the  centers  of  alternate  small  cubes. 
If  all  of  the  iron  atoms  are  in  the  dipositive  state,  there  is  an  alkali  ion  at 
the  center  of  each  small  cube;  the  arrangement  of  the  iron  atoms  is  not 
changed  (Fig.  1.4). 

510.  Hume  and  Kolthoff,  J.  Am.  Chem.  Soc.,  72,  4423  (1950). 

511.  Reihlen  and  Zimmermann,  Ann.,  475,  101  (1929). 

512.  Bhattacharya,  ./.  Indian  Chem.  Soc,  11,  325  (1934);  Davidson,  J.  Chem.  Ed., 

14,  238,  277  (1937). 
513    Keggin  and  Miles,  Nature,  137,  577  (1936). 


GKXKKAL  SI  L'\  /■:) 


91 


0 


V 

I 


Hi-TH 


m\ 


Fig.  1.2.  Structure  of 

Feiii[Feiii(CX)6]. 


Fig.  1.3.  Structure  of*         Fig.  1.4.  Structure  of 
KFe111  [Fe"(CN)6]  K2FeII[FeII(CN)6] 


A  structure  similar  to  this  is  probably  common  to  all  of  the  heavy  metal 
ferrocyanides,  variations  being  introduced  as  the  nature  of  the  second  metal 
ion  is  changed.  For  example,  assuming  that  the  coordination  number  of 
silver  is  two,  Ag4Fe(CN)6  should  be  formulated  as  Ag[Ag3Fe(CN)6].  Since 
the  covalences  of  silver  are  linear,  each  of  the  coordinated  silver  atoms  must 
share  electrons  with  the  nitrogen  atoms  of  two  different  Fe(CN)c  units, 
thus  forming  a  giant  polymer. 

Examples  are  known  in  which  coordination  with  carbon  tends  to  stabilize 
high  oxidation  states  of  the  metal  ions  (i.e.,  the  hexacyanocobaltate(TII) 
Ion),  but  in  most  cases,  metals  coordinated  to  carbon  show  very  low  oxida- 
tion states.  In  the  metal  carbonyls,  for  example,  the  metals  are  in  the  zero 
oxidation  state  and  in  the  salt-like  carbonyls  and  the  coordination  com- 
pounds containing  ethylenic  substances,  the  metals  are  always  in  their 
lower  oxidation  states.  The  same  tendency  appears  in  the  complex  cyanides, 
as  is  exemplified  by  the  compounds  K2Ni(I)(CN)3  and  K4Ni<°>(CN)4 ,  The 
compound  of  monovalent  nickel  was  first  prepared  by  Bellucci  and  Corelli514 
by  reducing  Kj[Ni(CN)4]  with  potassium  amalgam.  Hydrazine  can  also  be 
used  as  the  reducing  agent515,  but  the  best  method  of  preparation  involves 
reduction  with  metallic  potassium,  using  liquid  ammonia  as  the  solvent516. 
Bellucci  and  Corelli  supposed  that  it  is  similar  in  structure  to  K2[CuI(CN)3], 
but  the  fact  that  it  is  diamagnetic  shows  that  it  must  be  a  polymer.  Mellor 
and  Craig617  proposed  that  it  may  be  a  dimer  containing  a  metal-metal  bond, 
but  the  x-ray  work  of  Xast  and  Pfab515  indicates  the  presence  of  a  double 
bridge  and  they  write  the  structure 

514.  Bellucci  and  Corelli,  Z.  anorg.  Chem.,  86,  88  (1914). 

515.  Xast  and  Pfab,  Naiurwissenschaften,  39,  300  (1952). 

516.  Eastes  and  Burgess,  J.  Am.  Chem.  Soc.t  64,  1187  (1942). 

517.  Mellor  and  Craig,  J.  Proc.  Roy.  Soc.  X.  S.  Wales,  76,  281  (1943). 


92 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


\ 


NC 


NC 


Ni 


C 


\ 


CN 


CN 


The  complex  ion  readily  adds  nitric  oxide  and  carbon  monoxide : 

K2[Ni(CN)3]  +  NO  ->  K2[Ni(CN)3NO]518 

K2[Ni(CN)3]  +  CO  -»  K2[Ni(CN)3CO]519 

The  products  are  actually  more  complex  than  these  equations  indicate;  the 
carbonyl  compound,  at  least,  is  evidently  a  polymer,  for  it  is  diamagnetic. 

The  action  of  excess  potassium  on  K2[Ni(CN)4]  in  liquid  ammonia  gives 
K4[Ni(CN)4]516  as  a  copper-colored  solid  of  very  strong  reducing  powers. 
Compounds  of  the  formula  K4[M(CN)4]  containing  palladium520  and  co- 
balt521 have  been  prepared  in  analogous  fashion. 

Kleinberg  and  Davidson522  have  reduced  the  hexacyanomanganate(III) 
ion  in  liquid  ammonia,  obtaining  a  product  of  the  formula  K5Mn(CN)6- 
K6Mn(CN)6-2NH3 .  They  also  have  evidence  for  the  existence  of  a  cyano 
complex  of  chromium  (I)523. 

Isonitrile  Coordination 

Metal  complexes  of  the  isonitriles  have  been  known  for  a  long  time,  but 
have  received  little  attention  until  recent  years.  Hartley524  prepared  two 
isomers  of  "methyl  ferrocyanide,"  and  showed  that  upon  treatment  with  a 
mixture  of  alkyl  iodide  and  mercury  (II)  iodide  both  isomers  were  con- 
verted to  [Fe(CH3NC)4(RNC)2]l2-2Hgl2.  Both  isomers  gave  the  same 
product  when  methyl  iodide  was  used,  but  ethyl  iodide  gave  two  isomers525. 
These  have  been  subjected  to  x-ray  analysis526,  and  have  been  shown  to  be 
cis-  and  trans-  isomers. 

There  is  a  close  relationship  between  the  isonitrile-metal  complexes  and 
the  metal  carbonyls.  In  both,  the  metal-carbon  bond  possesses  a  consider- 

518.  Hieber,  Nast,  and  Proeschel,  Z.  anorg.  Chem.,  256,  145  (1948). 

519.  Nast  and  Krakkay,  Z.  anorg.  Chem.,  272,  233  (1953). 

520.  Burbage  and  Fernelius,  J.  Am.  Chem.  Soc,  65,  1484  (1943). 

521.  Hieber  and  Bartenstein,  Naturwissenschaften,  39,  300  (1952). 

522.  Kleinberg  and  Davidson,  J.  Am.  Chem.  Soc,  75,  2495  (1953). 

523.  Davidson  and  Kleinberg,  J.  Phys.  Chem.,  57,  571  (1953). 

524.  Hartley,  J.  Chem.  Soc.,  103,  1196  (1913). 

525.  Hartley,  J.  Chem.  Soc.,  1933,  101. 

526.  Powell  and  Stanger,  J.  Chem.  Soc,  1939,  1105. 


GENERAL  SURVEY  93 

able  degree  of  double  bond  character4980.  The  isonitrile  complexes  can  be 
made  by  displacement  of  carbon  monoxide  from  metallic  carbonyls.  Thus, 
phenyl  isonitrile  reacts  with  nickel  carbonyl  to  give  |\i(0\C)4]  as  long, 

canary-yellow    needli  They    arc   stable,   soluble    in    many    organic 

Bolvents,  1  >ut  insoluble  in  water.  Iron  and  chromium  carbonyls  also  read, 
though  more  slowly.  The  resulting  compounds  have  not  been  fully  charac- 
terized. 

Methyl  isonitrile  reacts  incompletely  with  nickel  carbonyl,  giving 
[Xi(CO)(CH3XC1)3].  However,  the  same  read  ants  in  the  presence  of  iodine 
and  pyridine  give  [Ni(CH»NC  )<].  A  cobalt  complex  of  the  empirical  formula 
[Co2(CO)3(<£XC)o]  is  obtained  by  the  action  of  phenylisonitrile  on 
Hg[Co(CO).i]2  in  the  presence  of  iodine  and  pyridine.  Klages,  Monkemeyer, 
and  Heinle-9  have  prepared  a  series  of  copper(I)  complexes,  CuCl-x</>XC, 
(x  =  1  -  4),  the  silver  compounds  AgXOaCp-CI^Cel^XC)*  (x  =  2  and  4), 
and  the  mercury(II)  and  zinc  compounds  MCl2(p-CH3CeH4X'C)2  . 

The  Nomenclature  of  Coordination  Compounds 

Werner's  system  of  nomenclature  is  the  basis  for  the  S3'stem  which  has 
been  adopted  by  the  International  Union  of  Pure  and  Applied  Chemistry, 
and  which  is  now  almost  universally  used530.  These  rules  may  be  summarized 
p&  follows531-532: 

(1)  If  the  substance  is  an  electrolyte,  the  cation  is  named  first,  then  the 
anion. 

(2)  The  names  of  all  negative  coordinating  groups  end  in  -o,  but  those  of 
'utral  groups  have  no  characteristic  ending.  In  deference  to  long  established 

practice,  the  coordinated  water  molecule  is  called  aquo. 

(3)  The  numbers  of  coordinating  groups  of  each  kind  are  indicated  by  the 
Greek  prefixes  mono-,  di-,  tri-,  tetra-,  etc.,  unless  these  groups  are  complex. 
In  that  case,  the  prefixes  bis-,  tris-,  tetrakis-,  etc.,  are  used. 

(4)  Xegative  coordinated  groups  are  listed  first,  then  neutral  coordinated 
groups,  then  the  metal.  (The  reverse  order  is  followed  in  writing  formulas 
of  complexes.) 

(5)  The  oxidation  state  of  the  metallic  element  is  indicated  by  a  paren- 
thetical Roman  numeral.  With  cations  and  neutral  molecules,  this  numeral 

527.  Hieber,  Z.  Xaturforsch.,  5b,  129  (1950);  Hieber  and  Bockly,  Z.  anorg.  Chem., 
262,  344  (1950). 
Klagea  and  Monkemeyer,  Ber.,  83,  501  (1950). 
B9.  Klages,  Monkemeyer,  and  Heinle,  Ber.,  85,  109,  126  (1952). 

530.  Jorissen,  Bassett,  Damiens,  Fichter,  and  Remv,  J.  Am.  Chem.  Soc,  63,  889 

(1941). 

531.  Fernelius,  I  Veto*,  26,  161   (1948):  Advances  in  Chemistry  Series, 

[81,  9  (1953).  American  Chemical  Society. 

532.  Fernelius,  Larsen,  Marchi,  and  Rollinson,  Chem.  Eng.  News,  26,  520  (1948). 


04  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

follows  the  name  of  the  metal  directly.  With  anions,  the  Roman  numeral  is 
placed  after  the  name  of  the  complex,  which  always  bears  the  suffix  -ate. 

Werner's  system  differed  from  this  chiefly  in  the  mode  of  designation  of 
oxidation  state  of  the  metal.  Werner  indicated  the  oxidation  state  of  the 
metal  in  cations  by  the  suffixes  -a,  -o,  -i,  and  -e,  indicating  1  +  ,  2+,  3+, 
and  4  +  ,  respectively.  In  anions,  the  same  suffixes  were  used,  followed  by 
the  ending  -ate.  In  neutral  molecules,  no  suffixes  were  used. 

Fernelius  and  his  co-workers531  - 532  have  suggested  some  useful  additions 
to  the  system  adopted  by  the  International  Union.  The  more  important  of 
these  have  been  summarized  by  Moeller  as  follows533 : 

(1)  The  names  of  coordinated  positive  groups  end  in  -ium. 

(2)  Positive  groups  are  listed  last,  after  negative  and  neutral  groups. 

(3)  Groups  of  the  same  general  nature  (i.e.,  all  negative,  all  neutral,  all 
positive)  are  listed  in  alphabetical  order  without  regard  to  any  prefixes 
designating  the  numbers  of  such  groups  present. 

(4)  Zero  oxidation  state  for  the  central  element  is  designated  by  the 
Arabic  character  0  placed  in  parentheses. 

(5)  Coordinated  hydrogen  salts  are  named  as  acids  by  dropping  the 
word  hydrogen  and  replacing  the  suffix  -ate  by  4c. 

(6)  Oxidation  state  of  the  central  element  is  designated  in  the  usual 
manner  even  though  the  complex  is  a  neutral  molecule. 

(7)  Use  of  prefixes  such  as  bis-,  tris-,  and  tetrakis-,  followed  by  the  name 
of  the  coordinated  group  set  off  by  parentheses  is  preferred  to  that  of  the 
old  designations  di-,  tri-,  and  tetra-  to  indicate  numbers  of  coordinated 
groups  if  the  names  of  those  groups  are  complex. 

In  both  the  Werner  and  the  I.U.C.  systems,  the  names  of  bridging  co- 
ordinated groups  (i.e.,  those  which  are  coordinated  to  two  metal  atoms 
simultaneously)  are  given  after  the  names  of  all  the  other  coordinating 
groups,  and  are  preceded  by  the  Greek  letter  ju.  Bridging  groups  have  their 
usual  names,  except  the  OH  group,  which  is  designated  as  ol. 

Geometrical  isomers  of  planar  ions  may  be  distinguished  either  by  the 
terms  cis-  and  trans-  or  by  the  numbers  1 ,2-  and  1,3-.  For  octahedral  com- 
plexes, these  become  cis-  and  trans-  or  1,2-  and  1,6-.  Where  there  are  more 
than  two  kinds  of  coordinating  groups,  or  more  than  two  of  any  one  kind, 
the  number  system  is  much  to  be  preferred. 

The  sign  of  rotation  of  optical  isomers  is  indicated  by  d-,  I-  (or  meso-). 
If  the  complex  contains  optically  active  coordinating  groups,  the  small 
letter  may  be  used  to  designate  that  fact,  and  the  capital  letters  d-  and  m 
to  indicate  the  sign  of  rotation  of  the  complex  as  a  whole. 

These  rules  are  exemplified  in  Table  1.2. 

It  is  customary,  in  writing  formulas  of  metal  coordination  compounds,  to 

Moeller,  "Inorganic  Chemistry,"  New  York,  John  Wiley  &  Sons,  Inc.,  1952. 


GENERAL  SURVEY 


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Table  1.3.  Symbols 

for  Names  of 

Some  Ligands 

Name  of  ligand 

Symbol 

acetate  ion 

ac1 

acetylacetonate  ion 

acac 

alanine  anion 

alan 

amino  acid  anion 

amac 

ammonia 

a2 

benzidine 

bzd 

benzoylacetate  ion 

benzac 

benzylamine 

bzl 

biguanide 

BigH 

2,3-butanediamine 

bn3 

isobutanediamine 

ibn 

citrate  anion  (monobasic) 

ci 

cj'anide  ion 

cy2 

1,2,  frans-cyclohexanediamine 

chxn 

1,2,  irans-cyclopentanediamine 

cptn 

diallylamine 

dim 

2, 4-diaminopentane 

ptn 

dibenzoylmethane 

dibenz 

diethylenetriamine 

dien 

dimeth}-lglyoxime  monobasic  anion 

DMG  or  HD 

2,2'-dipyridyl 

dipy 

ethylamine 

etn 

ethylenebiguanide 

enBigH 

ethylenediamine 

en 

ethylenediamine-acetylacetone 

enac 

ethylenediaminetetraacetic  acid 

H4Y  or  EDTA 

ethylenethiourea  or  ethylenethiocarbamide 

etu 

glycine  anion 

gly 

halide 

X 

hydroxyl  amine 

hx 

methyl  bis  (3-dimethylarsinopropyl) 

arsine 

TAS 

oxalate  dibasic  anion 

ox 

1, 10-phenanthroline 

o-phen4 

phenylalanine  anion 

4>  ala 

phenylbiguanide 

0  BigH 

ortho-phenylenediamine 

ph 

o-phenylenebis  (dimethyl  arsine) 

PDA 

phthalocyanine  (dinegative  group) 

pc 

propylenediamine  (1 , 2-diaminopropane) 

pn 

pyridine 

py 

st ilbenediamine  (1 , 2-diphenylethyle] 

lediamine) 

stien 

2,2',2",2'"-tetrapyridyl 

tetrpy 

thenoyltrifluoroacetone 

TTA 

t  hiourea 

tu 

i  hiosemicarbazide 

thio 

1,2, 3-t  riaminopropane 

tn 

2,2',2"-triaminotricthylamine 

tren  or  trin 

t  net hylenetet  ramine 

trien 

i  timet  hylenediamine 

trim 

2, 2', 2"  -tripyridyl 

tripy 

1  Not    to  be  contused  with  "Ac"  used  in  organic 

chemistry  to  denote  the  ac 

etyl 


group. 

I  Ibsolet  e. 

2  May  be  preceded  by  <l ,  I,  or  m  (dextro,  levo,  or  meso). 
4  Other  symbols  commonly  used  are  "phenan"  and  "ph". 


96 


Table  1.4.  Soin  Complex  Compoi  nds  Named  Aih.k  Thbib  Discovebsbs 


Name 

Strut  tare 

Cleve'a  Salt 

ci«-[Pi  \n  ica4] 

1 

Cleve'a  Triammine 

[Pt(NHi),Cl]Cl 

-  l'a  Firs!   Salt 

K[Pt(NH,)ClJ 

"s  Second  Salt 

K[Pt(NH,)ClJ 

Drechsel's  Chloride 

[Pt(XH3)6]Cl4 

OH 

Durrani  'a  Salt 

K4(C204)2Co<^      ^>Co(C204)2] 
OH 

Krdmann's  Salt 

/■•«/is-K[Co(NH3)2(N02)4] 

Fischer'a  Salt 

K,[Co(NO,),] 

( rerard'a  Salt 

trans-[Pt(XH3)2Cl4] 

2 

Gibba'  Salt 

[Co(NH3)3(N02)3] 

Gro's  Salt 

erona-[Pt(NHa)4ClilClj 

Litton'a  Salt 

Na.[Pt(SO,)4] 

3 

Magnus'  Green  Salt 

[Pt(NH,)4][PtCl4] 

4 

Magnus'  Pink  Salt 

Two  substances  of  this  name  are  known. 
The  common  one  is  [Pt(NH3)3Cl]2[PtCl4] 

Melano  chloride 

A  mixture,  chief!}* 

XH2 

/       \ 

(XH3)3Co— OH— Co(NH3)3 

Oli 

\       / 

OH 

Morland's  Salt 

CN»H.[Cr(NH,)»(SCN)4] 

5 

Peyrone's  Salt 

c*s-[Pt(XH3)2Cl2] 

6 

Recoura's  Sulfate 

[Cr(H20)5Cl]S04 

Reinecke's  Salt 

NH4[Cr(NH,),(SCN)4] 

7 

Rieset's  First  Chloride 

\P\   XH3)4]C12 

Rieset's  Second  Chloride 

//•ans-[Pt(XH3)2Cl2] 

8 

Roussin's  Red  Salts 

M  Fe(NO)«S  (M  =  Xa,  K,  XH4) 

Roussin's  Black  Salts 

M  Fe4(XO)7S3  (M  =  Xa,  K,  Rb,  Cs,  XH4  , 
or  Tl 

Vaquelin'a  Salt 

[Pd(XH3)4][PdCl4] 

9 

Vortmann's  Sulfate 

A  mixture,  chiefly 

OH 

III/       \III 

CNHi)4Co                Co(XH3)4 

(S04)2 

\       / 

Ml. 

containing  some 

02 

III/     \IV 

CNH                       Co  NH,)4 

[804)1 

\     / 

Ml 

and  other  materials 

WolfTram's  Red  Salt 

PI   (    II  \H2)4C13-2H,0— contains     Pt(II) 
and  Pt    IV 

10 

se'a  Salt 

K[PtCl3C  II 

''7 


98 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Table  1.4 — Continued 
1  cf.  Gerard's  Salt. 
>  cf.  Cleve's  Salt. 
Sails  of  this  type,  in  which  the  platinum  is  replaced  by  other  divalent  metals, 
ho  ammonia  is  replaced  by  other  nitrogen  bases,  or  the  chloride  is  replaced  by  other 
talides,  are  often  referred  to  as  Magnus  salts. 

4  Cox,  Pinkard,  Wardlaw,  and  Preston,  J.  Chem.  Soc.  1932,  2527. 

5  The  guanidinium  analog  of  Reinicke's  salt. 

6  cf.  Rieset's  Second  Chloride. 

7  Discovered  by  Morland  in  1861;  investigated  by  Reinicke  in  1863. 

8  cf.  Pey rone's  salt. 

9  cf.  Magnus'  Green  Salt. 

10  Jensen,  Z.  anorg.  allgem.  Chem.,  229,  252  (1936). 

Table  1.5.  Some  Names  of  Complexes  Based  on  Color 


Name 

Color 

Structure 

Note 

Croceo 

Yellow 

*mns-[Co(NH3)4(N02)2]+ 

Flavo 

Brown 

czs-[Co(NH3)4(N02)2]+ 

Luteo 

Yellow 

[Co(NH3)6]+++ 

Praseo 

Green 

/mns-[Co(NH3)4Cl2]+ 

l 

Purpureo 

Purplish-red 

[Co(NH3)5Cl]++ 

l 

Roseo 

Rose-red 

[Co(NH3)5H20]+++ 

Violeo 

Violet 

czs-[Co(NH3)4Cl2]+ 

l 

1  Often  used  to  denote  other  halopentammines,  sometimes  with  a  designation  as 
to  the  halogen  present;  thus,  [Co(NH3)5Br]++  is  referred  to  as  the  bromopurpureo 


use  symbols  for  the  names  of  many  organic  ligands.  Table  1.3  lists  the 
symbols  used  in  this  book,  as  well  as  some  others  which  may  be  encountered 
in  other  reading.  Unfortunately,  there  is  not  complete  uniformity  in  the 
use  of  these  abbreviations,  which  may  lead  to  some  confusion.  Because  of 
this,  some  notes  and  recommendations  are  included  in  Table  1.3. 

The  early  workers  in  the  field  of  complex  inorganic  compounds  did  not 
understand  the  nature  of  these  substances,  so  were  not  able  to  give  them 
names  based  upon  structure.  It  was  customary,  therefore,  to  name  each 
compound  after  its  discoverer.  A  few  of  these  early  names  persist  in  the 
current  literature,  and  are  listed  in  Table  1.4. 

In  1840,  Fremy634  suggested  that  the  ammines  of  cobalt  be  given  names 
descriptive  of  their  colors.  He  derived  such  names  from  the  Latin.  The  sys- 
tem was  easily  extended  to  the  cobalt  compounds  containing  ethylenedi- 
amine  and  other  amines,  and  to  the  chromium(III)  salts,  the  colors  of  which 
are  similar  to  those  of  their  cobalt(III)  analogs.  These  names  are  now 
frequently  used  to  describe  classes  of  compounds.  For  example,  the  term 

534     Fremy,  Ann.  chim.  phys.  [3],  35,  257  (1852);  J.  prakt.  Chem.,  57,  95  (1852). 


GENERAL  SURVEY  99 

"luteo,"  originally  used  to  describe  the  ion  [Co(NH3)6]:i+,  has  been  extended 
to  include  also  [Co  en3]3+,  [Co  dipy3]3+,  [Co  trien2]:H~,  and  other  cobalt  (III) 
complexes  in  which  six  amine  nitrogen  atoms  are  coordinated  to  the  cobalt. 
The  terms  are  sometimes  used  to  describe  ammines  of  metals  other  than 
cobalt  and  chromium,  even  though  the  colors  are  quite  at  variance  with 
the  names  suggested  by  Fremy.  For  example,  Gleu  and  Etehm686  use  the 
term  "luteo"  in  reference  to  the  hexammine  ruthenium(III)  ion,  which  Is 
colorless.  When  any  metal  other  than  cobalt  is  meant,  it  is  usual  to  in- 
clude the  name  of  the  metal.  Thus,  luteo  chromium(III)  chloride  is 
[Cr(XH3)6]Cl3 .  The  more  important  of  Fremy's  "color  names"  are  as- 
sembled in  Table  1.5. 

535.  Gleu  and  Rehm,  Z.  anorg.  allgem.  Chem.,  227,  237  (1936). 


A.   The  Early  Development  of  the 
Coordination  Theory 

John  C.  Bailor,  Jr. 

University  of  Illinois,  Urbana,  Illinois 

The  history  of  chemistry  in  the  nineteenth  century  is  largely  an  account 
of  the  growth  of  our  knowledge  of  molecular  structure.  When  the  doctrine 
of  constant  valence  proved  so  successful  in  explaining  the  structures  of  or- 
ganic compounds,  it  was  natural  that  every  effort  should  be  made  to  apply 
it  also  to  the  structures  of  inorganic  substances.  Thus  it  happened  that  the 
growth  of  inorganic  chemistry  was  retarded  for  over  twenty  years  by  the 
same  factor  which  contributed  most  to  the  phenomenal  development  of  our 
knowledge  of  the  compounds  of  carbon.  Inorganic  chemistry  is  the  older  of 
the  two  fields,  and  the  study  of  inorganic  "complex  compounds"  antedated 
the  rise  of  organic  chemistry  by  over  fifty  years.  The  structures  of  hydrates, 
double  salts,  and  metal  ammonia  compounds  were  widely  discussed  even 
before  the  beginning  of  the  nineteenth  century.  Of  these,  the  ammonia  com- 
pounds attracted  the  most  attention,  for  they  lent  themselves  to  study  by 
classical  methods.  The  early  history  of  the  theory  of  complex  compounds  is 
therefore  the  history  of  the  ammonates.*  The  discovery  of  these  substances 
is  usually  attributed  to  Tassaert1,  who  observed  in  1798  that  cobalt  salts 
combine  with  ammonia. 

Early  Theories  of  the  Structure  of  Ammines 

Berzelius'  Conjugate  Theory 

The  first  logical  attempt  to  explain  the  metal  ammonia  compounds  was 
made  by  Berzelius2,  who  observed  that  a  metal  in  "conjugation"  with  am- 
monia did  not  lose  its  capacity  for  combination  with  other  substances.  He 

*  The  term  "ammonate"  was  displaced  by  the  simpler  term  "ammine"  at  the  sug- 
gestion of  Werner. 

1.  Tassaert,  Ann.  chim.  phys.,  [1]  28,  92  (1798). 

2.  Berzelius,  "Essai  sur  la  theorie  des  proportions  chimique  et  sur  Pinfluence  chimi- 

que  de  l'electricite,"  Paris,  1819. 

100 


EARLY  DEVELOPMENT  OF  THE  COORDINATION  THEORY        101 

attempted  to  extend  this  theory,  but  without  great  success,  to  the  double 
salts  and  complex  cyanidi 

Graham's  Ammonium  Theorx 

According  to  Graham's  "ammonium"  theory*,  metal  ammonates  are 
considered  to  be  substituted  ammonium  compounds.  This  view,  in  one  form 

or  another,  was  generally  accepted  until  the  time  of  Werner.  (iraham  made 
this  suggestion  in  an  attempt  to  explain  the  structure  of  diammoniuin 
copper  II  salts,  in  which  he  supposed  one  hydrogen  atom  from  each  of 
two  ammonium  groups  had  been  displaced  by  copper.  Obviously,  such  a 
formula  can  apply  only  when  the  number  of  ammonia  groups  in  the  mole- 
cule is  the  same  as  the  electrovalence  of  the  metal — a  condition  which 
usually  does  not  hold.  Gerhardt4,  Wurtz6,  Rieset6,  A.  W.  Hofmann7  and 
Boedecker8  suggested  modifications  of  the  theory  to  take  care  of  other  cases. 
According  to  Rieset  and  Hofmann,  the  hydrogen  atoms  of  an  ammonium 
group  are  replaceable,  not  only  by  metals,  but  also  by  other  ammonium 
groups.  Hofmann  represented  the  compound  of  cobalt  (III)  chloride  with 
six  molecules  of  ammonia,  for  example,  as 

Co/XH,-  XH4 


0/XH2-  XHA 

J.     ). 


Some  years  later  the  experiments  of  Jorgensen  showed  this  argument  to  be 
fallacious.  It  does  not  allow  for  the  existence  of  similar  compounds  of  ter- 
tiary amines9,  and  it  does  not  explain  why  the  removal  of  one  molecule  of 
ammonia  completely  alters  the  function  of  one  of  the  chlorine  atoms. 
Boedecker  avoided  the  branching  of  the  chain  by  assuming  that  the  metal 
substitutes  in  an  ammonium  group  which  is  itself  a  substituent  group: 
Co(XH3 — NH3 — Cl)3 .  The  diammonate  and  tetrammonate  of  platinum(II) 
chloride  were  represented  as  Pt(XH3— Cl)2  and  Pt(XH3— XH3— Cl)2 .  The 
question  "What  prevents  further  lengthening  of  the  ammonia  chain?"  was 
never  answered,  and  was  an  insurmountable  objection  to  this  type  of  theory. 

3.  Graham,  "Elements  of  Chemistry,"  London,  1837.  This  book  is  rare,  and  is  best 

known  in  Otto's  German  translation  "Lehrbuch  der  Chemie"  Braunschwieg, 
1840.  Graham's  suggestion  of  the  ammonium  theon-  appears  in  Vol.  2,  page  741 
of  the  German  edition. 

4.  Gerhardt,  Jahresber.  Fortshr.  pharm.,  tech.  chem.  physik  <  Living),  3,  335  (1850). 

5.  Wurtz.  .!/.//.  rhim.  phys.,  [3]  30,  488  (1850). 

6.  Rieset,  Ann.  chim.  phys.,  [3]  11,  417  (1844 

7.  Hofmann,  Ann., 78, 253  (1851). 

8.  Boedecker,  Ann.,  123,  56  (1862). 

rgensen,  J.  praht.  Chem.,  [2]  33,  489  (1886). 


102  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Glaus'  Theory 

The  theory  of  Claus10  met  with  vigorous  opposition,  but  is  all  the  more 
interesting  on  that  account,  for  the  parts  of  it  which  were  most  vigorously 
attacked  appeared  in  only  slightly  modified  form  in  Werner's  theory.  Claus 
believed  that,  when  combined  with  metallic  oxides,  ammonia  not  only  does 
not  affect  the  saturation  capacity  of  the  metal,  but  becomes  "passive"  as 
regards  its  own  basicity.  His  views  may  be  summarized  as  follows:11 

(1)  The  union  of  several  equivalents  of  ammonia  with  one  equivalent  of 
a  metal  chloride  leads  to  the  formation  of  a  neutral  substance,  in  which  the 
basic  property  of  ammonia  is  lost,  so  that  the  ammonia  can  no  longer  be 
determined  by  the  usual  means  nor  eliminated  by  double  decomposition. 
Thus,  the  ammonia  is  in  a  different  condition  than  in  ordinary  ammonium 
salts.  This  hypothesis  met  with  a  storm  of  protest,  just  as  Werner's  similar 
suggestion  did  forty  years  later.  The  attack  was  led  by  Weltzein12,  who  held 
the  term  "passive  molecule"  to  be  indefinite  and  confusing,  and  who  be- 
lieved that  every  part  of  a  molecule  influences  every  other  part,  so  that 
no  part  can  be  said  to  be  "passive". 

(2)  If  these  chlorides  are  converted  to  oxides,  strong  bases  are  formed. 
The  saturation  capacity  of  these  is  the  same  as  that  of  the  metal  oxides 
themselves,  and  cannot  be  calculated  from  the  number  of  ammonia  mole- 
cules combined  with  the  oxide.  Schiff13  criticized  this  conclusion  by  pointing 
out  that  the  oxides  of  the  "ammonia  bases"  of  the  metals  are  much  stronger 
bases  than  the  metal  oxides  themselves.  This  criticism  seems  to  rest  on  a 
confusion  between  the  "strength"  of  a  base  and  its  "saturation  capacity" 
(i.e.,  equivalence).  It  is  true  that  the  hydroxides  of  the  metal  ammines  are 
strong  bases,  but  the  ammonia  present  in  them  does  not  readily  combine 
with  the  hydrogen  ion. 

(3)  The  number  of  ammonia  molecules  combined  with  a  molecule  of 
metallic  salt  is  determined  by  the  same  factors  as  the  number  of  molecules 
of  water  in  the  hydrate  and  the  two  will  be  the  same.  This  point  of  Claus' 
theory  was  easy  to  attack,  for  many  hydrates  were  known  for  which  analo- 
gous ammonia  compounds  did  not  seem  to  exist.  The  conclusion  which  Claus 
drew,  however,  was  restated  as  an  integral  part  of  Werner's  theory  and  has 
been  amply  verified. 

Blomstrand's  Chain  Theory 

Odling14  suggested  that  metallic  atoms  can  substitute  for  the  hydrogen 
atoms  in  ammonia  just  as  organic  radicals  do.  The  diammonate  of  plati- 

10.  Claus,  "Beitrage  zur  Chemie  der  Platinmetalle,"  Dorpat,  1854;  Zentralblatt,  25, 

789  (1854);,4nn.,98,317  (1856). 

1 1 .  Reitzenetein,  Z.  anorg.  Chem.,  18,  152  (1898). 

12.  Weltzein,  Ann.,  97,  19  (1856). 

13.  Schiff,  Ann.,  123,  1  (1862). 

14.  Odling,  Chem.  News,  21,  289  (1870). 


EARLY  DEVELOPMENT  OF  THE  COORDINATIOh    THEORY        103 

num(II)  chloride  was  construed  as  being  analogous  to  ethylenediamine 
hydrochloride:  Pt(NH,),-2HCl  and  C,,lI.J(MI2),-i-)Il('l.  The  chaining  of 
ammonia  molecules  was  compared  to  the  chaining  of  methylene  groups  in 

the  hydrocarbons. 

Blomstrand1'  made  this  the  basis  of  his  famous  theoiy.  Ammonium 
chloride  was  represented  as  11  Ml,  CI,  XII,X(),XH,  as  II  ML 
Nil,  N( >,  ,  and  MI,I  r>MI3  as  H(XH3)7L  The  terminal  hydrogen  atom 
can  be  replaced  by  other  positive  atoms,  such  as  metals.  The  metal,  in  fact, 
stabilizes  the  chain,  and  its  nature4  determines  the  length  and  stability  of 
the  chain.  Chains  of  three  ammonia  molecules  are  often  found  in  union 
with  nickel,  cobalt,  iridium  and  rhodium,  but  platinum  and  copper  seem 
unable  to  stabilize  chains  of  more  than  two  nitrogen  atoms.  On  the  basis  of 
these  postulates,  Blomstrand  wrote  the  formulas  for  the  tetrammonate  of 
platinum(II)  chloride  and  the  hexammonate  of  cobalt(II)  chloride  as 

\  II  —  XH3— CI  XH3— XH3— XH3— CI 

/  / 

Pt  and  Co 

\  \ 

X 1 1 ,— XH.  —  CI  NH3— XH3— XH3— CI 

According  to  Blomstrand,  the  stability  of  the  ammonia  chain  is  not 
dependent  on  its  length.  Although  platinum  is  unable  to  stabilize  chains  of 
any  great  length,  platinum(II)  chloride  ammonate  is  not  attacked  by  hy- 
drogen sulfide  or  by  sodium  hydroxide.  Chlorine  oxidizes  the  platinum 
without  attacking  the  ammonia,  converting  the  compound  to: 

CI     XH3— NH3— CI 

1/ 
Pt 

l\ 
CI     XH3— XH3— CI 

in  which  chlorine  is  attached  to  the  molecule  in  two  different  ways.  The 
validity  of  this  postulate  is  borne  out  by  experiment,  for  only  half  the 
chlorine  is  replaced  by  the  action  of  sodium  carbonate,  and  the  second  half 
i-  only  slowly  precipitated  by  silver  nitrate.  Blomstrand  referred  to  the 
two  types  of  chloride  as  the  "farther"  and  "nearer".  This  expression  may 
have  inspired  Werner's  postulate  of  "first"  and  "second"  spheres11. 

JorgenserTs  Theories 

Blomstrand's  formulas  for  the  cobalt  ammonia  compounds  became  the 
center  of  a  Long  controversy  between  Jorgensen  and  Werner,  and  are  there- 
fore  of   considerable    interest.    Blomstrand    believed     and    the    belief   wa> 

15.  Blomstrand,  "Chemie  der  Jetztzeit,"  Heidelberg,  1869;  Ber.,  4,  40  (1871). 


104  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

universal  until  189016 — that  cobalt(III)  chloride  and  its  ammonia  com- 
pounds  were  dimolecular.  In  that  year,  Jorgensen  adduced  evidence  for  the 
simpler  molecular  weights,  and  halved  Blomstrand's  formulas.  This  did  not 
affeel  the  postulates  of  Blomstrand's  theory,  but  without  this  change, 
Werner's  theory  might  not  have  been  conceived.  Blomstrand  first  supposed 
the  Luteo  cobalt  salts  (e.g.,  Co2Cl6-12NH3)  to  have  the  completely  sym- 
metrical structure: 

NH3— NH3— CI 
NH3— NH3— CI 
NH3— NH3— CI 
NH3— NH3— CI 
NH3— NH3— CI 
[NH3— NH3— CI 


Co2 


and  the  purpureo  salts  (Co2Cl6  •  10NH3)  the  structure : 

NH3— CI 
NH3— NH3— CI 
NH3— NH3— CI 
NH3— NH3— CI 
NH3— NH3— CI 
NH3— CI 


Co; 


But  this  was  soon  seen  to  be  incorrect,  for  the  purpureo  salt  contains  chlo- 
rine in  two  very  different  modes  of  combination11, 17.*  In  a  cold  solution, 
silver  nitrate  precipitates  two-thirds  of  the  chlorine  at  once,  and  the  other 
third  only  after  long  standing.  The  slight  functional  difference  shown  in 
the  formula  above  can  hardly  explain  such  a  difference  in  behavior.  Jorgen- 
sen18 prepared  a  whole  series  of  salts  in  which  the  more  readily  precipitated 
chlorine  is  replaced  by  other  groups.  He  concluded  that  the  chlorine  in  these 
salts  is  combined  directly  with  the  metal,  while  the  other  negative  groups 
are  united  with  the  ammonia.  Similar  relationships  were  shown  to  hold  for 
the  chromium19  and  rhodium20  pentammonate  salts.  Jorgenson  also  dem- 
onstrated that  the  "masked"  chloride  can  be  replaced  by  bromine21,  sul- 
fate22, and  other  negative  groups.  These  groups,  like  the  chloride  in  the 
original  purpureo  salt,  have  lost  their  ionic  properties. 

*  For  explanation  of  nomenclature,  see  Chapter  1. 

16.  Jorgensen,  J.  prakt.  Chcm.,  [2]  41,  429  (1890);  Petersen,  Z.  phys.  Chem.,  10,  580 

(1892). 

17.  Gibbs  and  Genth,  "Researches  on  the  Ammonia  Cobalt  Bases,"  Washington, 

1856. 

18.  Jorgensen,  J.  prakt.  Chcm.,  [2]  18,  209  (1878). 

19.  Jdrgenaen,  ./.  prakt.  ('hem.,  [2]  20,  105  (1879);  25,  83  (1882). 

20.  Jorgensen,  J.  prakt.  Chem.,  [2]  25,  346  (1882);  27,  433  (1883);  40,  309  (1886). 

21.  Jorgensen,  J.  prakt.  Chcm.,  [2]  19,  49  (1879). 

22.  Jorgensen,  J.  prakt.  Chem.,  [2]  31,  262  (1885). 


EARLY  DEVELOPMEXT  OF  THE  COOUD1 X ATIOX   THEORY         10.") 


When  Jorgensen  found11  thai  two-thirds  of  the  chlorine  in  the  tetram- 
monates  of  the  trivalent  metals  is  "masked",  he  concluded  that  this  should 
be  represented  as  in  direcl  union  with  the  metal.  He  formulated  these  salts 
as: 


Co: 


CI 

CI 

Ml 

Ml 

CI 

CI 


N 1 1 

Ml 


-XH3— NH3— CI 
-NH3— NH3— CI 


and  the  purpureo  and  the  luteo  salts  as: 


Co, 


CI 

MI. CI 

MI    -XH3— XH3— XH3— CI 
)  XH3— NH3— NH3— NH3— CI 

MI.C1 

la 


and 


Co2< 


(NH3CI 
NH3C1 

NH3— NH3— NH3— NH3— CI 
NH3— NH3— NH3— NH3— CI 
NH3— CI 
NH3-C1 


Jorgensen  showed  that  the  "roseo"  salts,  which  had  been  thought  to  be 
isomeric  with  the  purpureo  salts,  contain  two  molecules  of  water24.  This 
water  is  lost  at  elevated  temperatures,  leaving  a  residue  of  the  purpureo 
salt.  The  roseo  salts  resemble  the  luteo  salts  in  that  all  of  the  negative 
groups  are  ionic  as  well  as  in  solubility,  crystalline  form,  and  appearance. 
Jorgensen  concluded  that  they  are  luteo  salts  in  which  one-sixth  of  the 
ammonia  molecules  are  replaced  by  water. 

The  roseo  tetrammonate  salts  were  also  shown  to  be  analogous  to  the 
luteo  salts,  but  they  contain  water  in  place  of  one-third  of  the  ammonia 
molecules.  Xo  compounds  were  known  in  which  more  than  a  third  of  the 
ammonia  was  replaced  by  water,  so  it  was  assumed  that  the  "unchained" 
ammonia  molecules  were  the  ones  replaced.  The  roseo  tetrammonia  salts 
were  therefore  represented  as: 


Co2s 


H20— CI 
H20-C1 
NH3— XH3-XH3 
NH3— NH3— NH3 
H20— CI 
H20— CI 


-NH3— CI 
-NH3-CI 


or,  using  the  simplified  formula,  as: 

H20— CI 


/ 
Co— H20— CI 

\ 

MI  — XH3— XII  — XII 


-CI 


23.  Jorgensen,  J.  prakt.  Chem.,  [2]  27,  433  (1883). 

2-4.  Jorgensen,  J.  prakt.  Chem.,  [2]  29,  409  (1884);  31,  49  (1885), 


10()  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

These  postulates  suggest  many  questions,  some  of  which  Jorgensen  at- 
tempted to  answer  by  modifications  or  elaborations  of  the  theory: 

Why  can  cobalt  hold  only  six  ammonia  molecules?  If  one  of  the  valences 
holds  a  chain  of  four,  why  cannot  the  others  also?  Can  chains  of  more  than 
four  ammonia  groups  exist?  How  shall  we  explain  the  existence  of  isomeric 
compounds? 

Jorgensen  felt  that  the  chains  contain  a  maximum  of  four  — NH3-groups25, 
because  of  the  many  examples  of  tetrammonated  compounds,  and  because 
the  penta-  and  hexaammonated  salts  seemed  to  contain  one  and  two  am- 
monia molecules,  respectively,  which  are  different  from  the  other  four.  He 
answered  the  other  questions  by  developing  Blomstrand's  hypothesis  that 
the  three  valences  of  cobalt  are  different.  An  example  or  two26  will  illustrate 
the  argument:  The  luteo  chloride, 

7NH3CI 

M  <*(NH3)4C1        (M  represents  a  trivalent  metal) 
0NH3C1 

readily  loses  one  molecule  of  ammonia  to  form 

7C1 
M«(NH3)4C1, 
/3NH3CI 

which  in  water  is  converted  to  the  aquo  (roseo)  salt,  which  must  therefore 
be 

7H20— CI 
M«(NH3)4C1. 
/3NH3C1 

The  diaquo  roseo  salt, 

7H20— CI 

M<*(NH3)4C1 

/3H20— CI 

readily  loses  one  molecule  of  water  to  form  a  compound  which  must  contain 
the  groups  — CI  and  — H20 — CI.  But  the  — H20 — CI  group  in  this  com- 
pound is  not  like  the  — H20 — CI  group  in  the  roseo  pentammine.  The  former 
loses  a  molecule  of  water  when  heated  to  100°C  or  lower,  while  no  water  is 
lost  from  the  latter  until  the  temperature  is  well  above  100°C.  According 
to  Jorgensen,  this  difference  indicates  that  the  0  and  7  valences  are  not  the 
same.  He  cited  the  fact  that  the  tetrammonates  take  up  one  molecule  of 


25.  Jorgensen,  Z.  anorg.  Chem.,  5,  147  (1894). 

26.  Jorgensen,  Z.  anorg.  Chem.,  7,  289  (1894). 


EARLY  DEVELOPMENT  OF  THE  COORDINATION*    THEORY        107 

ammonia  or  water  easily,  and  a  second  with  difficulty,  as  further  evidence 

for  this  view.  The  isomerism  of  the  "flavo"  and   "croceo"  chlorides  was 
explained  by  i lie  formulas: 

yNO  7NOj 

Coo  Ml      CI        and        CoaNOj 
UNO  0(NH,)4C1 

Early  Theories  of  the  Structure  of  Hydrates 

While  these  theories  of  the  metal  ammonia  compounds  were  being  dis- 
cussed, attempts  were  also  being  made  to  elucidate  the  structures  of  the 
hydrates.  The  best  known  of  the  hydrate  theories  was  that  of  Wurtz27,  who 
postulated  that   the  water  molecules  link  themselves  to  the  metal  and  to 

each  other  in  rings: 

H20— H20  H20— H20— H20 

/  \  /  \ 

S04  Cu  H20  and  S04  Mg  H20 

\  /  \  / 

H,0— H20  H20— H20— HoO 

The  assumptions  underlying  the  theory  were  unsupported  by  experimental 
evidence,  and  it  met  with  little  favor. 

Early  Theories  of  the  Structure  of  Double  Salts 

The  double  salts,  especially  the  double  halides,  were  of  great  interest, 
and  numerous  theories  of  their  constitution  were  advanced.  Bonsdorff28  and 
Boullay29  compared  the  chlorides  to  oxides,  some  of  which  are  acidic  and 
others  basic,  and  they  supposed  double  salts  were  formed  by  a  sort  of 
neutralization  reaction.  Others3031  took  exception  to  this  theory,  but  it 
found  wide  acceptance.  Xaquet32  expressed  the  view  that  two  chlorine  atoms 
are  equivalent  to  one  oxygen,  and  Blomstrand15  went  so  far  as  to  suppose 
these  two  chlorine  atoms  to  be  linked  together  through  a  double  bond.  On 
this  basis  3KClFeClj  and  2KClPtCl4  become 

C1=C1— K  CI  C1=*C1 — K 

/  \    / 

Fe— C1=C1— K  and  Pt 

\  /     \ 

C1=C1— K  CI  C1=C1— K 

27.  Wurtz,  "La  Theorie  Atomique,"  Paris,  1879. 

28.  Bonsdorff,  Ann.  ckim.  phys.,  34,  142  (1827). 

29.  Boullay,  Ann.  ckim.  phys.,  34,  337  (1827). 

30.  Liebig,  Ann.  ckim.  phys.,  35,  68  (1827). 

31.  Bcrzelius,  Jahresbcr.  Forfsch.  chem.  mineral.  (Berzelius),  8,  138  (1829). 

82.  Naquet,  "Principea  de  Chemie  fondee  sur  les  Theories  Modernes,"  Paris,  1867. 


108 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


There  was  little  experimental  evidence  to  support  Blomstrand's  suggestion, 
and  it  was  not  widely  accepted33.  Such  formulas  do  not  indicate  why  the 
potassium  should  be  ionic  and  the  iron  and  platinum  nonionic,  nor  do  they 
allow  for  the  formation  of  double  chlorides  such  as  CdCl2-4KCl,  in  which 
the  number  of  molecules  of  alkali  metal  chloride  exceeds  the  number  of 
chlorine  atoms  in  the  heavy  metal  chloride.  Remsen34  "solved"  the  latter 
difficulty  by  assuming  the  formation  of  halogen  rings: 

K— CI  CI— K 


CI— Cd— CI 

/  N 

K— CI  CI— K 

In  1885  Horstmann35  wrote  the  reaction: 

CI     CI  CI     CI     CI 

\l  \\^ 

Pt  +  2KC1  -»  Pt— K 

/I  /l\ 

CI     CI  CI     CI     K 

in  analogy  to 

H  H  H 


CI 


H— N  +  HC1  -*  H— N 


II 


B 


CI 


which  was  the  generally  accepted  mechanism  for  the  reaction  of  ammonia 
with  hydrochloric  acid.  By  assuming  large  enough  valences  for  the  metals, 
we  can  apply  this  theory  to  complexes  of  all  sorts.  It  is,  of  course,  mislead- 
ing in  its  implication  that  all  of  the  groups  are  attached  to  the  central  atom 
in  the  same  way  (the  chlorine  and  the  potassium,  in  the  example  given). 
With  this  feature  modified,  Horstmann's  formulas  become  almost  identical 
with  those  of  Werner. 

Werner's  Coordination  Theory 

This,  then,  is  the  background  on  which  Werner  built.  In  his  paper  "Con- 
tribution to  the  Theory  of  Affinity  and  Valence"36  published  in  1891,  he 
suggested  thai  an  atom  does  not  have  a  certain  number  of  valence  bonds, 
but  that  the  valence  force  is  exerted  oxer  the  whole  surface  of  the  atom,  and 

33.  Remsen,  .1///.  Chem.  J.,  11,  291  (1889). 

34.  Remsen,  Am.  Chi  »,.J.,  14,  81  (1892). 

35.  Horstmann,  "Lehrbuch  der  Physikalischen  und  Theoretischen  Chemie,"  Braun- 

Bchweig,  1885. 

36.  Werner,  "Beitrage  sue  Theorie  der  Affinital  und  Valenz,"  1891. 


EARLY  DEVELOPMENT  OF  THE  COORDINATION  THEORY        L09 

can  l>o  divided  into  several  units  of  varying  strength,  depending  on  the 
demands  of  the  atoms  which  unite  with  it.  Sonic  of  its  valence  force  may  be 
left  unexpended.  This  thoughl  is  differenl  from  the  postulate  of  "primary" 
and  "secondary"  valences,  but  is  certainly  a  forerunner  of  it.  The  wide- 
spread belief  that  the  coordination  theory  had  no  roots  in  earlier  theories  or 
in  the  experience  of  its  author  is  a  mistaken  one.  It  is  true,  however,  thai 
the  theory  was  different  from  anything  which  had  previously  been  proposed 
and  that  it  came  in  a  spectacular  way.  Pfeiffer"7  lias  writ  ten  :  "According  to 
his  own  statement,  the  inspiration  came  to  him  like  a  flash.  One  morning 
at  two  o'clock  he  awoke  with  a  start  ;  the  long-sought  solution  of  this  prob- 
lem had  lodged  in  his  brain.  He  arose  from  his  bed  and  by  five  o'clock  in 
the  afternoon  the  essential  points  of  the  coordination  theory  were  achieved." 
Werner  was  then  twenty-six  years  old.* 

Fundamental  Postulates 

The  fundamental  postulate  in  Werner's  coordination  theory  is  stated  in 
the  following  way88  "Even  when,  to  judge  by  the  valence  number,  the  com- 
bining power  of  certain  atoms  is  exhausted,  they  still  possess  in  most  cases 
the  power  of  participating  further  in  the  construction  of  complex  molecules 
with  the  formation  of  very  definite  atomic  linkages.  The  possibility  of  this 
action  is  to  be  traced  back  to  the  fact  that,  besides  the  affinity  bonds  desig- 
nated as  principal  valencies,  still  other  bonds  on  the  atoms,  called  auxiliary 
valences,  may  be  called  into  action."  The  rest  of  the  theory  is  an  elucidation 
of  the  nature,  the  number,  and  the  spatial  distribution  of  these  "auxiliary" 
valences. t  The  auxiliary  valences  were  originally  conceived  as  being  very 
different  from  principal  valences,  since  they  do  not  allow  ionization  while 
the  principal  valences  do.  Yet  according  to  Werner,  there  is  a  connection 
between  them,  for  if  an  atom  forms  strong  primary  bonds  with  certain  other 
atoms,  ii  usually  forms  strong  secondary  bonds  with  them  too.  Thus  the 
alkaline  earth  oxides  are  extremely  stable,  and  they  combine  with  water 
(by  secondary  valence)  with  great  avidity.  Similarly,  the  very  stable  sul- 
fides of  copper,  mercury  and  arsenic  readily  form  thio  complexes.  It  is  pos- 

For  biographical  sketches  of  Werner,  see  G.  T.  Morgan: ./.  Chem.  Soc.,  117,  1639 

1920);  J.  Lifschitz,  Z.  Elektrochem.,  26,  514  (1920);  and  I'.  Karrer,  Helv.  ckim.  Acta, 

3,  l'.»6  (1920).  These  give  brief  accounts  of  his  theory.  The  art  iele  by  Karrer  contains  a 

portrait  and  B  list  of  Werner's  publications.  P.  Pfeiffer,  /.  Chem. Ed.,  5, 1090  (1928) 

gives  a  description  of  Werner's  personal  life  and  a  portrait  of  him. 

t  The  terms  "primary"  and  "secondary"  were  often  used  instead  of  "principal" 
and  "auxiliary." 

Pfeiffer, ./.  Chem.  A''/.,  5,  L096    L928);  Ostwald's  "Klassiker  der  Exakten  Wissen 

schaften,"  No.  212,  p.  ">,  Leipzig,  Akademiache  Verlagsgesellschaft,  1924. 
Werner,  "Neuere  Anschauungen,"  1th  Ed.  p.  44,  \  ifwi^,  Braunschweig,  1920. 
Quoted  from  Bass'  translation  of  Schwarz,  "The  Chemistry  of  Inorganic  Com 
plex  Compounds,"  p.  '.»,  \ew  York,  John  Wiley  a-  Bona,  [nc,  i'»23. 


110  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

sible,  too,  for  a  primary  valence  to  be  converted  into  a  secondary  one.  In 
solutions  of  hexammine  chromic  chloride,  [Cr(NH3)6]Cl3  ,*  all  of  the  chlorine 
is  at  once  precipitated  by  solutions  of  silver  nitrate.  If  the  dry  hexammine 
be  heated  somewhat  above  100°C,  a  molecule  of  ammonia  escapes,  and 
simultaneously  one-third  of  the  chlorine  loses  its  ionic  properties.  Werner 
argued  that  this  means  it  has  become  attached  by  a  secondary  valence, 
though  of  course  this  does  not  release  a  primary  valence,  and  the  new  com- 
pound contains  only  two  chloride  ions39.  Jorgensen  and  Werner  both  be- 
lieved the  nonionic  chlorine  to  be  attached  directly  to  the  metal,  in  place 
of  the  ammonia  which  had  been  lost.  On  standing  in  water  solution,  the 
pentammine  undergoes  a  slow  change  by  which  the  third  chlorine  again 
becomes  ionic.  Upon  evaporation  at  room  temperature,  the  resulting  solu- 
tion yields  crystals  of  a  rose-red  pentammine,  containing  a  molecule  of  water. 
Jorgensen40  had  shown  that  this  "roseo"  compound  is  closely  analogous  to 
the  hexammine,  and  he  recognized  it  as  a  hexammine  in  which  one  ammonia 
molecule  is  replaced  by  water.  In  this,  he  and  Werner  agreed.  They  dis- 
agreed, however,  on  the  fate  of  the  chlorine  atom  which  the  water  molecule 
had  displaced.  Jorgensen  believed  it  to  be  attached  to  this  water  molecule 
through  the  quadri valence  of  oxygen  while  Werner  felt  that  it  was  not  at- 
tached to  any  particular  atom  in  the  complex,  but  was  attracted  by  the 
complex  ion  as  a  whole.  Werner's  postulate  clearly  foreshadows  the  theory 
of  ionization  of  salts  in  the  crystalline  state,  and  has  been  amply  confirmed 
by  x-ray  measurements  and  by  other  means.  At  the  time  of  its  proposal, 
however,  it  was  a  most  revolutionary  doctrine,  and  for  many  years  it  met 
with  widespread  criticism41. 

The  relationship  between  primary  and  secondary  valence  became  closer 
and  closer  in  Werner's  mind,  and  he  was  finally  led  to  the  conclusion  that 
there  is  no  essential  difference  between  the  two.  This  came  about  through 
his  study  of  the  tetrakis(ethylenediamine)-ju-amino-nitro-dicobalt(III)  ion, 

*  The  term  "amrain"  proposed  by  Werner  to  designate  the  metal  ammonia  com- 
pounds, is  translated  into  English  as  "ammine".  Its  use  in  this  place  is  somewhat 
anachronic,  as  it  was  not  used  in  Werner's  earlier  papers,  but  we  shall  use  it  through- 
out. The  term  "ammonate"  is  still  used  by  some  authors  to  designate  simple  addition 
compounds  of  ammonia  with  metallic  salts.  Such  compounds  can  be  called  "ammines" 
equally  well,  however.  In  the  earlier  papers,  Werner  indicated  the  constituents  of  the 
complex  ion  by  enclosing  them  in  parentheses,  but  he  later  adopted  the  use  of  square 
brackets. 

39.  JSrgensen,  ./.  prakt.  Chem.,  [2]  20,  105  (1879). 

40.  Jorgensen,  ./.  prakt.  Chem.,  [2]  29,  409  (1884). 

41.  See  for  example,  Friend,  ./.  Chem.  Soc,  109,  715  (1916);  119,  1040  (1921). 


EARLY  DEVELOPMENT  OF  THE  COORDINATION  THEORY 

Ml 

ens  Co  Co  enj 

\       / 
NO 

This  ion  contains  two  asymmetric  cobalt  atoms  (See  Chapter  8)  which  ap- 
parently arc  not  identical.  One  of  them  is  attached  to  the  amino  group  by  a 
primary  valence  and  to  the  nitro  group  by  a  secondary  valence,  while  for 
the  other  one,  these  relationships  are  reversed.  Resolution,  then,  should  give 
a  dextro,  a  levo,  and  two  meso  forms.  Careful  experimentation,  however, 
yielded  only  one  meso  form.  This  compound  is  completely  inactive,  indicat- 
ing the  identity  of  the  two  asymmetric  atoms.  Werner  may  not  have  been 
surprised  at  this  discovery,  for  his  first  paper43  draws  an  analogy  between 
the  metal  ammine  ions  and  the  ammonium  ion,  in  which  the  hydrogen  which 
is  held  by  "secondary"  valence  is  indistinguishable  from  the  rest. 

It  has  long  been  known  that  many  of  the  metal  ions  form  hexammonates 
and  hexahydrates,  and  that  tetraammonates  are  common.  The  tetra-  and 
hexacyanides  have  also  long  been  known  as  stable,  well-defined  compounds. 
From  such  facts,  Werner  deduced  that  each  element  has  only  a  certain 
number  of  secondary  valences.  Groups  attached  to  the  central  element  by 
these  valences  are  said  to  be  "coordinated"  to  it.  The  "coordination  number" 
of  an  atom  or  ion  is  the  number  of  groups  which  can  be  coordinated  to  it.* 
While  four  and  six  are  the  most  common  coordination  numbers,  coordina- 
tion numbers  of  two,  three,  five,  seven  and  eight  are  known. 

In  terms  of  Werner's  theory,  the  secondary  valences  of  an  atom  must  be 
satisfied.  In  the  case  of  hexamminechromium(III)  chloride,  if  a  molecule  of 
ammonia  is  driven  out,  one  of  the  chloride  ions  will  take  its  place  to  main- 
tain the  coordination  number  six.  A  wide  variety  of  neutral  groups  or  nega- 
tive ions  can  enter  the  coordination  sphere.  When  these  latter  become  co- 
ordinated, they  cease  to  be  ions,  of  course,  and  this  is  indicated  by  the 
suffix  -o  on  their  names  or  abbreviated  names;  thus,  "cyano,"  "chloro," 
"nitro,"  and  "hydroxo". 

If  a  trivalent  metal  hexammine  chloride  loses  one  molecule  of  ammonia, 
one  of  the  three  chlorides  loses  its  ionic  properties,  as  has  been  pointed  out. 
If  a  second  molecule  of  ammonia  is  lost,  a  second  chloride  becomes  non- 
ionic41.  What  will  happen  if  a  third  ammonia  molecule  is  lost?  According  to 

*  When  applied  to  the  structure  of  crystals  the  term  "coordination  Dumber"  is 
given  a  somewhat  different  meaning;  it  refers  to  the  number  of  atoms  (or  ions)  which 
surround  the  atom  or  ion  i  in  question,  and  arc  ;it  equal  distances  from  it,  no  matter 
what  the  natin<-  of  i he  bond  between  them. 

VI    Werner,  Ber.,  46,  3674  (1913);  47,  L964,  1978    I'd  I 

13.  Werner.  Z.  anorg.  Chem.,  3,  267    18 

1  1 .  Jdrgensen .  Z   anortj.  ('firm .,  5,  117  (1894). 


112  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Jorgensen 's  own  statement44,  he  had  never  considered  this  point,  but  it 
became  very  important,  for  his  theory  and  Werner's  predicted  different 
behaviors.  According  to  the  coordination  theory,  the  third  chloride  should 
become  nonionic,  and  a  nonelectrolytic  molecule  should  result.  Jorgensen 
had  to  assume  that  his  postulated  ammonia  chain  would  simply  be  shortened 
by  one  nitrogen  atom,  which  would  still  leave  the  chloride  in  the  ionic  state. 
Very  few  triammines  of  trivalent  metals  were  known  at  that  time,  and  when 
Werner  pointed  out43  that  their  properties  supported  his  own  theory, 
Jorgensen  objected44  that  the  compounds  were  not  sufficiently  understood 
to  justify  the  conclusion. 

One  of  these  compounds,  Ir(NH3)3Cl3 ,  had  been  described  by  Palmaer45, 
who  found  that  it  did  not  liberate  hydrochloric  acid  when  heated  with  con- 
centrated sulfuric  acid.  He  suggested  that  it  had  twice  the  simplest  formula, 
and  was  a  double  salt,  Ir(NH3)6Cl3-IrCl3  .  Jorgensen  showed  that  a  cor- 
responding rhodium  double  salt  could  be  prepared  from  the  components, 
and  that  it  did  not  liberate  hydrogen  chloride  when  warmed  with  sulfuric 
acid.  He  pointed  out  also  that  Magnus'  salt  Pt(NH3)4Cl2-PtCl2*  is  resistant 
to  concentrated  sulfuric  acid,  and  concluded  that  this  reagent  cannot  be 
relied  upon  to  indicate  the  presence  pf  ionic  chlorine. 

The  other  example  cited  by  Werner  was  Erdmann's  Co(NH3)3(N02)346, 
which  was  admittedly  not  a  well  characterized  compound47.  Several  sub- 
stances of  the  same  composition  had  been  discovered,  and  Erdmann's 
description  of  his  compound  was  incomplete.  Investigation  of  the  com- 
pound convinced  Jorgensen  that  it  has  the  structure 

N02 

/ 
Co— NH3— N02 

\ 

NH3— NH3— NO  2 

He  converted  it  to  the  chloride,  which  however,  contains  one  molecule  of 
firmly  held  water;  to  this  compound  he  assigned  the  structure 

H20— CI 

/ 
Co— NH3— CI 

\ 

NH3— NH3— CI 

*  We  would  now  give  these  "double  salts"  the  formulas  [Ir(NHs)e]  [IrCl6]  and 
[Pt(XH3),]  [PtCl*],  which  indicate  that  they  do  not  contain  chloride  ions. 
15.   Palmaer,  Oefvers,  af  k.  Vet.  Acad.  Fdrh,  No.  6,  373  (1889);  Ber.,  22,  15  (1889). 
L6.  Erdmann,  ./.  prakt.  Chew.,  97,  412  (1866). 
47.  Gibbs,  Proc.  Amer.  Acad.,  10,  16  (1875). 


EARLY  DEVELOPMENT  OF  THE  COORDINATION  THEORY        113 

because  all  of  the  chlorine  is  precipitated  at  once  by  silver  nitrate.  This 
compound  is  readily  converted  to  Erdmann's  "trinitrite,"  which  must  then 

have  the  structure  shown. 

The  tWO  theories  differ  also  in  their  predictions  a>  to  the  roult  of  the  loss 
oi  another  molecule  of  ammonia,  with  the  production  of  a  diammine.  No 
such  compounds  were  known  and  this  was  in  accord  with  Werner'.-  theory. 
To  him,  an  ammonia  molecule  cannot  he  "lost";  it  must  he  replaced  by 
another  group.  Thus  far  in  the  process,  the  halide  ions  which  accompany 
the  complex  have  been  able  to  carry  out  this  replacement,  hut  now  ;i  new 
group  must  he  supplied.  If  this  be  a  negative  ion,  it  will  give  the  complex 
a  negative  charge.  Keinecke's  salt,  NH4[Cr(XH3)2(SCX)4]48  and  Erdmann's 
salt.  XII; [Col  XH3)2(X02)4]49  *  are  examples  of  this  type  of  compound. 
There  were  no  examples  of  the  monoammonates,  Ms'tM'^NHaXs],  pre- 
dicted by  Werner,  but  numerous  examples  of  the  final  step  in  the  replace- 
ment were  known;  e.g.,  the  heavy  metal  cyanides,  the  cobaltinitrites,  and 
the  double  chlorides. 

The  tetravalent  elements  furnish  a  similar  series.  Platinum(IV)  chloride 
yields  ammines  containing  six,  five,  four,  three,  two,  and  one  molecules  of 
ammonia.  All  the  chloride  is  readily  removed  from  the  first  of  these.  Blom- 
strand15  had  observed  that  two  of  the  four  chlorine  atoms  in  the  tetram- 
monate  are  much  less  reactive  than  the  other  two.  There  are  two  isomeric 
forms  of  the  diammonate,  which  therefore  elicited  great  interest.  In  accord- 
ance with  the  demands  of  Werner's  theory,  both  of  these  are  nonionic.  The 
end  member  of  the  series  is  potassium  hexachloroplatinate(IV),  which  does 
not  react  with  silver  nitrate  to  give  silver  chloride,  but  gives  silver  chloro- 
platinate,  Ag*[PtCU]. 

Conductivity  Studies 

To  give  further  support  to  these  views,  Werner  and  Miolati  measured 
the  conductivities  of  a  large  number  of  metal  ammines50.  Again,  the  results 
med  to  substantiate  the  coordination  theory,  but  Emil  Petersen51  raised 
objections  to  this  conclusion.  The  number  of  ions  found  was  in  some  cases 
greater  than  predicted  by  the  theory.  A  case  in  point  is  Co(XH3)3(X02)j(,l. 
which  the  theory  demands  musl  be  a  nonelectrolyte,  but  which  showed  the 
conductivity  of  a  uni-univalent  electrolyte.  Werner  explained  this  by  assum- 
ing the  reaction  [Co(NH,),(N02),Cl]  +  H20  ->  [C0(XH3)3(X02)2H20)C1, 

*  Erdmann's  salt  is  not  to  be  confused  with  Erdmann's  trinitrotriamminecobalt 
(III),  mentioned  above. 
48.  Reinecke,  Ann.,  126,  113  (1863). 

Erdmann, ./.  prakt.  char.,  97,  406  (1866). 

50.  Werner  and  Miolati,  Z.  pkysik.  Ch  -.12,  35  (1893);  14,  506  (1804  ;  21,  225  (1896). 

51.  Petersen,  Z   pi  22,  410  (1897). 


114  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

T  \  ble  2.1 .  Effect  of  Aging  on  the  Molar  Conductivity  of  an  Aqueous  Solution 

of  [Co(NH3)4Br2]Br 
(Molar  Concentration,  0.2%) 

m  = 
Freshljr  prepared  solution  190.6 

5  minutes  after  the  first  measurement  288.0 

10  minutes  after  the  first  measurement  325.5 

15  minutes  after  the  first  measurement  340.7 

20  minutes  after  the  first  measurement  347.8 

40  minutes  after  the  first  measurement  363.5 

and  supported  this  by  the  fact  that  at  0°C,  where  the  hydration  reaction 
cannot  proceed  readily,  the  conductivity  is  indeed  very  low.  Petersen 
countered  by  pointing  out  that  all  salts  show  much  lower  conductivities  at 
0°  than  at  25°C. 

Werner  and  Miolati  reported  several  instances  of  this  kind,  and  in  some 
of  them,  had  good  evidence  that  reaction  with  the  water  does  take  place. 
The  dark  green  Co(NH3)4Br3  dissolves  to  give  a  deep  green  solution,  which 
rapidly  becomes  red.  At  the  same  time,  the  conductivity  rises,  as  shown  in 
Table  2.1.  It  seems  to  approach  that  of  the  diaquotetrammine  salt  (see 
Table  2.2),  which  is  bright  red.  Werner  and  Miolati  wrote  the  equation: 

[Co(NH3)4Br2]Br  +  2H20  ->  [Co(NH3)4(H20)2]Br3 

The  "dichro"  salt,  Co(NH3)3(H20)Cl3  gave  similar  results,  the  solution 
turning  from  green  through  blue  to  violet; 

[Co(NH3)3(H20)Cl2]Cl  +  2H20  ->  [Co(NH3)3(H20)3]Cl3  . 

This  reaction  proceeds  so  rapidly  at  room  temperature  that  Werner  and 
Miolati  made  their  conductivity  studies  at  1°C.  The  molecular  conductivity 
was  compared  with  those  of  potassium  chloride,  barium  chloride,  and 
hexamminecobalt(III)  chloride  at  the  same  temperature,  and  found  to 
correspond  to  that  of  the  first;  in  other  words,  the  salt  is  composed  of  two 
ions. 

With  those  compounds  which  do  not  contain  readily  displaced  groups  in 
the  coordination  sphere,  Werner  and  Miolati  obtained  results  entirely  in 
accord  with  their  expectations.  Many  of  their  results  are  elegantly  shown  in 
graphical  form  in  the  second  paper  of  their  series,  and  two  are  reproduced 
in  Figs.  2.1  and  2.2.  The  conductivities  of  aquoammine  salts  are  significant 
in  that  they  support  Werner's  contention  that  water  molecules  and  am- 
monia molecules  occupy  equivalent  positions  in  the  coordination  sphere. 
Some  of  these  are  shown  in  Table  2.2.  Petersen51  repeated  some  of  this  work, 


J 


EARLY  DEVELOPMENT  OF  THE  COORDINATION*    THEORY 


115 


522  9 


256 


[Pt(NH3)6]ci4 
[Pt(NH3)5Cl]cis 
[Pt(NH3)4CI2]ci2 
[Rt(NH3)3CI3]ci 


>^NH3)2CI4] 

K[Pt'(NH3)C,5] 
K2[PtCI6] 


Fig.  2.1.  The  molar  conductivities  of  0.1  molar  per  cent  aqueous  solutions  of  some 
platinum  (IV)  ammines.* 


99.29 


A-  [C0(NH3)6]CI3 

B-  [C0(NH3)5(N02)]CI2 

c-  i,6[co(nh3)4(no2)2]ci 

D-  [C0(NH3)3(N02)3] 
E-       K[C0(NH3)2(N02)4] 


Fig.  2.2.  The  molar  conductivities  of  0.1  molar  per  cent  aqueous  solutions  of  some 
cobalt  (III)  ammines. 


Table  2.2.  Molab  Conductivities  of  Some  Cobalt(III)  Ammines  at 

Various  Dilutions 

(25°C) 

V    liters      [Co{NHi)t]Bn    [Co(XH,)&(H,0)]Br3    [Co(NH2)4(H20)2]Br3   [Co(NHi)iNOt](NO»)j  [Co(\H3)4C03]Br 

125           343.8               333.6                       325.5  98.58 

250                 .1)               365.4                       354.8                     206.1  101.3 

500           401.6               390.3                      379.8                    225.1  103.5 

1000           426.9               412.9                      399.5                    234.4  106.0 

2000           442.2               436.4                        117.1                     242.8  111.8 


*  The  value  for  [Pt(NHi)iCl]Clj  w&e  doI  given  in  the  original  paper,  but    has 
since  been  determined  by  Tschugaeff  and  Wladimiroff:  Compt.  ri  //'/.,  160,  840    1915 


I  16  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

and  his  results  agree  with  those  of  Werner  and  Miolati.  Particularly  inter- 
esting is  his  value  for  the  molecular  conductivity  of  Co(NH3)3(N02)3  (8.4 
a1  a  dilnt  ion  of  800  liters  at  25°C)  which  fully  confirms  that  of  Werner  and 
Miolati,  and  clearly  shows  the  compound  to  be  nonionic.  Petersen  also 
at  tempted  to  determine  the  number  of  ions  formed  from  many  of  the  metal 
ammonia  compounds  by  measuring  the  freezing  points  of  their  solutions. 
The  results  did  not  agree  in  all  cases  with  those  obtained  from  the  con- 
ductivity studies.  They  did  not  support  Jorgensen's  beliefs  any  better  than 
they  did  Werner's,  but  they  were  used52  to  discredit  the  conductivity 
method,  upon  which  Werner's  crucial  experiments  rested. 

The  coordination  theory  handles  metals  of  coordination  number  four  just 
as  it  does  those  of  coordination  number  six,  and  one  example  will  suffice: 
Platinum(II)  chloride  forms  ammines  with  two,  three,  and  four  molecules 
of  ammonia.  The  first  of  these  is  especially  interesting,  because  two  isomeric 
forms  exist.  The  Blomstrand-Jorgensen  theory  supposed  these  to  be 


Pt 


whereas,  according  to  the  coordination  theory  they  are  stereoisomeric  forms 
of  [Pt(NH3)2Cl2].  The  older  theory  would  postulate  that  form  (I)  can  lib- 
erate two  chloride  ions  whereas  form  (II)  can  liberate  only  one,  but  the 
coordination  theory  allows  no  ionization  in  either  case.  As  far  as  form  (II) 
is  concerned,  the  data  of  Table  2.3  clearly  support  the  latter  contention. 

Table  2.3.  Effect  of  Aging  on  the  Molar  Conductivity  of  an 

Aqueous  Solution  of  "Platosemidiamminchlorid" 

(Molar  Concentration,  0.1%) 

m  = 
Freshly  prepared  solution  1.17 

2  minutes  after  first  measurement  1.81 

4  minutes  after  first  measurement  2.41 

10  minutes  after  first  measurement  2.61 

15  minutes  after  first  measurement  4.33 

30  minutes  after  first  measurement  11.03 

180  minutes  after  first  measuremenl  21.87 

Form  (I),  (the  "platosamminchlorid"),  goes  into  solution  very  slowly,  and 
then  only  with  warming,  so  it  was  possible  to  measure  the  conductivity 
only  after  some  rend  ion  with  the  water  had  taken  place.  The  molar  con- 

52.  Jdrgensen,  Z.  anorg.  Chem.,  14,  404  (1897);  19,  132  (1899). 


NH3— CI 

NH3— NH3— CI 
/ 

and 

Pt 

\ 

NH3— CI 

CI 

(I) 

(ID 

BARL1    DEVELOPMENT  OF  THE  COORDINATION  THEORY        117 

ductivity,  at  25°C  and  for  a  0.1  molar  per  cent  solution,  was  found  to  be 
22.42.  Platinum (II)  chloride  docs  not  form  a  monainmine,  bu1  the  com- 
pound K[PtClj-NHj]  takes  its  place  in  the  scries.  Potassium  tetrachloro- 

platinate(II)  represents  the  complete  replacement  of  ammonia  by  the 
chloride  ion. 

His  views  on  the  ion  forming  properties  of  the  metal  annuities  thus  over- 
thrown, Jorgensen  turned  his  attack  on  the  coordination  theory  to  Werner's 
postulate4  thai  all  of  the  coordinated  groups  occupy  equivalent  positions  in 
the  complex-',  lie  cited  several  reactions  of  the  hexammines  to  indicate  that 
four  of  the  ammonia  molecules  are  attached  to  the  metal  ion  more  firmly 
than  the  other  two.  Thus,  the  aquopentamminecobalt(III)  salts,  on  heating 
with  ammonium  carbonate,  give  carbonatotetrammine  salts,  and  the  nitro- 
pentammines  give  dinitrotetrammines  when  treated  with  sodium  nitrite. 
In  neither  case  is  more  ammonia  readily  removed. 

Jorgensen  felt  also  that  the  reactions  of  Co(XH3)j(X02)3  and  "croceo" 
dinitrotetrammine  salts  indicate  that  all  of  the  nitro  groups  are  not  held  to 
the  cobalt  in  the  same  way.  In  each  case,  the  action  of  hydrochloric  acid 
eliminates  one  nitro  group  more  readily  than  the  others.  Werner  had  as- 
sumed the  existence  of  nitro  ( — N02)  and  nitrito  ( — OXO)  groups  (in  agree- 
ment with  Jorgensen)  to  explain  the  existence  of  isomeric  salts  of  the  com- 
position  Co(NH3)5NOjXs  .  Why,  then,  argued  Jorgensen,  does  he  assume 
that  the  "flavo"  and  "croceo"  salts  must  be  stereoisomers  rather  than 
structural  isomers?  If  the  "croceo"  compounds  are  frans-dinitro  salts  as 
Werner  suggested,  the  two  nitro  groups  will  show  identical  chemical  reac- 
tions. In  reality,  they  do  not.  One  of  them  resembles  the  nitrous  group  of 
the  "isoxantho"  (nitritopentammine)  compounds,  and  is  readily  liberated 
by  dilute  acids;  the  other  is  not  attacked. 

Jorgensen  also  found  fault  with  Werner's  theory  because  it  predicted  the 
existence  of  many  compounds  which  were  then  unknown.  Most  important 
among  these  were  the  "violeo"  (cis)  dichlorotetramminecobalt(III)  salts, 
which  might  be  expected  to  be  formed  upon  replacement  of  the  nitro  groups 
of  "flavo"  (cis)  dinitrotetrammine  compounds  by  chloride.  Such  replace- 
ment can  be  effected  by  the  action  of  dilute  hydrochloric  acid,  but  "praseo" 
Baits,  rather  than  "violeo", are  formed.  Jorgensen  called  upon  Werner,not 

only  to  explain  the  nonexistence  of  the  "violeo"  salts,  hut  also  the  rear- 
rangements which  the  coordination  theory  implied  in  this  and  similar  re- 
actions. Jorgensen  also  pointed  out  that  many  compounds  exist  which 
Werner*-  theory  does  not  satisfactorily  explain.  Commonest  of  these  are 

the  hydrate-,  many  of  which  contain  more  than  six  molecule-  of  water. 
Werner's  assumption  of  double  water  molecule-,   II.u...  was  without  ex- 

53.  Jorgensen,  7. .  anorg.  Ckem.,  19,  109  (188 


118  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

perimental  support,  and  could  explain  only  a  small  fraction  of  the  examples 
known. 

Finally,  Jorgensen53  criticized  the  suggestion  that  the  entrance  of  a  nega- 
tive  group  into  the  complex  ion  should  lower  the  valence  of  the  complex. 
In  support  of  his  criticism,  he  quoted  Werner  to  the  effect  that  "the  co- 
ordinated groups  do  not  change  the  valence  of  the  metal  atom."  He  argued 
that  if  this  negative  group  still  saturates  one  of  the  primary  valences  of  the 
metal,  it  cannot  be  coordinated. 

While  some  of  these  criticisms  were  obviously  not  well  founded,  others 
were  thoroughly  sound,  and  challenged  Werner's  ingenuity  and  experi- 
mental skill  to  the  utmost.  Many  of  the  missing  compounds  were  dis- 
covered, among  them  the  crucial  "violeo"  cobalt  salts54;  a  theory  of  re7 
arrangements  was  devised55;  the  relationship  between  the  primary  and 
secondary  valences  was  clarified42;  and  the  octahedral  structure  of  the 
hexacoordinated  complexes  was  firmly  established  by  the  resolution  of 
many  compounds  into  their  optical  antipodes.  The  coordination  theory,  as 
originally  devised,  was  supported  in  almost  every  particular. 

54.  Werner,  Ber.,  40,  4817  (1907). 

55.  Werner,  Ann.,  386,  1  (1912). 


vj.    Modern  Developments — The  Electro- 
static Theory  of  Coordination 
Compounds 

Robert  W.  Parry 

University  of  Michigan,  Ann  Arbor,  Michigan 

and 

Raymond  N.  Keller 

University  of  Colorado,  Boulder,  Colorado 

Although  Werner's  ideas  regarding  the  stereochemistry  of  complex  com- 
pounds were  well  substantiated  by  experiment,  widespread  dissatisfaction 
with  his  postulates  of  primary  and  secondary  valences  served  as  a  strong 
deterrent  to  the  general  acceptance  of  his  entire  theory  even  as  late  as  19161. 
Since  data  available  to  Werner  did  not  always  permit  a  sound  differentiation 
between  the  assumed  valence  types,  the  coordination  theory  led  to  the 
prediction  of  a  variety  of  unusual  valence  states  for  many  common  metals. 
It  was  justly  held  that  such  a  theory  led  to  confusion,  and  Werner's  postu- 
lates concerning  primary  and  secondary  valence  bonds  were  called  vague 
and  unfounded1  •  2. 

It  was  not  until  the  development  of  the  electronic  theory  of  valence  by 
Lewis.  Kossel,  Langmuir,  Sidgwick,  Fajans,  Pauling  and  others  that  a  self- 
consistent  explanation  of  valence  types  evolved.  The  models  which  were 
developed  for  the  electronic  theory  were  so  successful  in  resolving  the  con- 
fusion surrounding  the  ideas  of  primary  and  secondary  valence  that  almost 
general  acceptance  of  Werner's  views  soon  followed  the  work  of  Lewis  and 
In-  contemporari 

.Modem  x-ray  diffraction  data  have  now  provided  unequivocal  experi- 
mental support  for  Werner'.-  ideas  on  stereochemistry.  In  addition,  quantum 

1.  Friend,  /.  Chem.  Soc.,  93,  260,  1006    L908);  109,  715  (1916);  110,  1040  (1921). 

2.  Briggs,/.  '  8oe.,  93,  1564    1908  ;  Proc.  '  hem  Sot.,  24,  94     L908);  Jorgen- 

Ben,  Z    pi  ,  144,   ]s7     L929  ;  Pfeiffer,  Z.  anorg.  allgem.  Chem.  112, 

81     1920  ;   Povamin,  ./.   Ri        PI        Chem.  Soc.  47,  217,  501,  980   (1915  ; 
cf    '  be.  10,  138  (1916). 

Hit 


120  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

mechanics  now  provides  the  framework  for  a  more  detailed  solution  of 
valence  problems.  Unfortunately,  the  quantum  mechanical  approach  is 
extremely  complex  unless  many  simplifying  assumptions  are  made;  as  a 
result,  the  simple  molecular  models  suggested  by  Lewis,  Kossel,  and  others 
are  still  of  fundamental  importance  in  correlating  fact  and  theory. 

The  Electrostatic  Model 

The  Charge -size  Ratio 

According  to  the  viewpoint  first  clearly  developed  by  Kossel3,  complexes 
are  held  together  by  the  electrostatic*  attraction  between  oppositely 
charged  ions  or  between  ions  and  dipolar  molecules.  For  example,  the  fluoro- 
borate  ion,  (BF4)~,  can  be  pictured  as  a  triply  charged  central  boron  ion 
to  which  four  fluoride  ions  are  symmetrically  bound  by  electrostatic  forces. 
The  hydrated  calcium  ion,  [Ca(H20)6]++,  may  be  pictured  as  a  central  cal- 
cium cation  to  which  six  water  dipoles  are  electrostatically  bound  with 
octahedral  symmetry.  Complex  ammines,  halides,  hydrates,  and  many 
other  compounds  may  be  represented  in  a  similar  manner.  From  considera- 
tions of  elementary  electrostatics,  Kossel  suggested  that  those  metal  ions 
with  high  ionic  charget  and  small  ionic  radius  would  form  coordination  com- 
pounds of  greatest  stability.  De5  pointed  out,  apparently  independently, 
that  the  metals  whose  ions  have  the  highest  coordinating  ability  are  those 
of  small  atomic  volume  (and  thus  of  small  ionic  radius),  such  as  Cr,  Fe,  Co, 
Ni,  Cu,  Ru,  Rh,  Pd,  Os,  Ir,  Pt,  and  Au.  Since  ionic  charge  and  ionic  size 
have  opposite  effects  in  determining  the  electrostatic  field  of  an  ion,  Cart- 
ledge6  suggested  a  single  arbitrary  parameter  called  the  ionic  potential, 
which  is  denned  as  the  charge  of  the  ion  divided  by  its  crystal  radius  in 
Angstrom  units.  In  general,  coordinating  ability  increases  with  an  increase 
in  the  ionic  potential  of  the  central  ion,  although  a  number  of  qualitative 
exceptions,  such  as  the  high  relative  stability  of  the  complexes  of  Hg++  and 

*  Electrostatic  interaction  was  implied  by  earlier  workers2d  •  4  but  never  developed. 

f  In  general,  the  stability  of  ammines  frequently  does  increase  with  increasing 
charge  on  the  central  ion,  but  this  is  not  always  so  as  is  shown  by  the  fact 
that  FeCl  2 -6NH3  is  more  stable  than  FeCl3-6NH3  . 

3.  Kossel,  Z.  Elektrochem.,  26,  314  (1920);  Z.  Phys.,  1,  395  (1920);  Naturwissen- 

schaflen,  7,  339,  360  (1919);  11,  598  (1923);  Ann.  Phys.,  49,  229  (1916). 

4.  Nelson  and  Falk,  J.  Am.  Chem.  Soc,  37,  274  (1915). 

5.  r><\  ./.  Chem.  Soc,  115,  127  (1919). 

6.  Cartledge,  ./.  Am.  Chem.  Soc,  50,  2855,  2863  (1928);  52,  3076  (1930);  J.  Phys. 

Colloid  Clnm.  55,  248  (1951). 

7.  Bjerrum,   "Metal  Ammine  Formation  in  Aqueous  Solution,"   pp.  75,  87.   P. 

Hasse  and   Son,   Copenhagen,   1941;    Irving    and    Williams,  J.   Chem.   Soc, 
1953,  3202;  Bjerrum,  Chem.  Revs.,  46,  381  (1950). 


ELECTROSTATIC  THEORY  OF  COORDINATIOh   COMPOUNDS      121 

Cu+,  are  known.  As  early  as  L928  Fajans8  pointed  ou1  thai  the  concepts  of 
ion  deformation  and  interpenetrat ion  must  be  u>rd  along  with  any  ionic 
model  in  order  to  obtain  reasonable  agreemenl  between  fad  and  theory. 
The  problem  is  considered  under  polarization  (see  page  12.")).  More  recently 

Irving  and  Williams71'  have  demonstrated  in  a  most  convincing  manner 

that  the  ionic  potential  alone  is  not  adequate  as  a  parameter  for  the  estima- 
tion of  complex  stability  constants. 

Aeid-base  Phenomena  in  Coordination  Compounds 

An  extension  of  the  charge-size  ratio  principle  to  the  hydrolysis  of  the 
ions  of  tin1  first  two  periods  of  the  periodic  table  permitted  Kossel  to  treat 
aqueous  acid-base  phenomena  as  a  natural  consequence  of  the  coordination 
theory.  (See  references  3c,  3d,  6,  9  and  Chapter  12  for  a  more  thorough 
treatment  of  this  topic.)  This  viewpoint  readily  justifies  the  acid  character 
of  the  complex  ion,  [Pt(XH3)6]4+  and  is  effective  in  explaining  acid-base 
behavior  in  nonaqueous  solvents.* 

Polarization  as  a  Factor  in  the  Ionic  Model 
Nature  of  Polarization 

Many  of  the  early  energy  calculations  based  on  the  electrostatic  model 
had  two  rather  serious  limitations.  No  provision  was  made  for  energy 
changes  involved  in  lattice  expansion  or  in  solution  processes;  only  inter- 
action energy  between  ion  and  ligand  was  considered.  Secondly,  the  exist- 
ence of  rigid,  spherically  symmetrical  ions  or  molecules  was  assumed  (i.e., 
the  ionic  potential  was  considered  as  a  suitable  differentiating  parameter). 
Actually,  the  electronic  clouds  of  each  atom  or  ion  are  deformed  by  the 
fields  which  are  set  up  by  neighboring  ions  or  dipolar  molecules. f 

This  deformation  of  ions  is  related  to  their  polarization.  The  amount  of 
distortion  is  determined  by  the  strength  of  the  distorting  field  and  by  the 

*  The  ideas  expressed  by  Kossel  were  anticipated  to  some  extent  in  1899  by  Abegg 
and  Bodlander10  who  discussed  the  factors  influencing  coordination.  They  noted  that 
certain  weak  liases,  such  as  Co203-H20  become  strong  bases  when  coordinated  to 
form  complexes  such  as  [Co(NH3)e](OH)3,u  and  that  weak  acids  such  as  HCN  form 
Btrong  acids  when  coordinated  to  metal  ions,  as  is  illustrated  by  H3[Fe(CN)6].12 

t  The  inaccuracy  of  the  approximation  of  rigid  ions  was  mentioned  by  Kossel,3'1  ■  13 
but  not  considered  as  a  major  factor  in  compound  stability. 

8.  Fajans,  Z.  Krist.,  A66,  321  (1928). 

9.  Foster,  J.  Chi  m.  Ed.,  17,  509  (1940). 

10.  Abegg  and  Bodlander,  Z.  anorg.  Chem.,  20,  453  (1899). 

11.  ham!)  and  Yngve,  ./.  .1//'.  Chem.  Soc,  43,  2352  (1921). 

12.  Brigando,  Compt.  rend.,  208,  197  (1939);  Ray  and  Dutt,  Z.  anorg.  allgem.  Ch 

234,  65  (1937). 

13.  Kossel,  Naiurwissenschaften,  12,  703  (1924). 


L22  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

magnitude  of  the  force  binding  the  electron  cloud  to  the  atomic  nucleus. 
It  the  electrons  are  tightly  bound  (low  polarizability),  little  distortion 
occurs.  If  they  are  loosely  bound  (large  polarizability),  the  ion  may  be 
seriously  deformed  from  its  spherical  symmetry. 

Polarization  as  a  factor  in  binding  forces  was  first  suggested  by  Haber14 
in  L919  and  independently  by  Debye15  in  1920.  The  development  of  the 
concept  and  its  applications  to  chemical  theory  were  due  largely  to  Fajans. 
Some  attempt  was  also  made  to  apply  the  idea  to  structural  problems. 
Hund16  and  Heisenberg17  used  the  ideas  of  polarization  to  account  for  the 
fact  that  the  water  molecule  is  angular  instead  of  linear,  as  the  concept  of 
rigid  spherical  ions  would  suggest18.  The  effects  of  polarization  have  been 
reviewed  by  Fajans19,  Clark20,  and  Debye18.  Quantitative  data  on  the 
polarizability  (deformability)  of  various  ions  as  measured  by  their  molar 
refraction  were  reported  by  Fajans  and  Joos22  and  others21, 23, 24, 25.  These 
data  in  the  hands  of  Fajans  permitted  the  modification  of  the  original  ionic 
model  to  correct  for  deformation  effects.  The  modified  ionic  model  has  been 
used  to  correlate  both  the  chemical  and  physical  properties  of  complexes. 

Chemical  Properties  and  the  Polarization  Model 

Stability  of  Ammines  and  Hydrates.  It  is  a  well  known  fact  that  cations 
such  as  those  of  the  alkalies  and  the  alkaline  earths  do  not  form  stable 
ammonia  complexes  in  water  solution.  In  aqueous  solution  the  hydrate  is 
far  more  stable  than  the  ammine.  For  these  cations,  the  metal  ion-ammonia 
bond  in  solution  is  weaker  than  the  metal  ion- water  bond.  On  the  other 
hand,  cations  such  as  copper(II),  silver  (I),  cadmium(II),  and  zinc(II), 
which  are  found  in  Periodic  Groups  IB  and  IIB,  form  ammine  complexes 
which  are  much  more  stable  in  aqueous  solution  than  are  the  hydrated  ions. 
For  these  metals,  the  metal-ammonia  bond  is  significantly  stronger  than  the 
metal-water  bond.  It  is  also  interesting  that  the  coordinating  ability  of 

14.  Haber,  Verhandl.  deut.  physik.  Ges.,  21,  750  (1919). 

15.  Debye,  Z.  Phys.,  21,  178  (1920);  22,  30  (1921). 

16.  Hund,  Z.  Phys.,  31,  81  (1925);  32,  1  (1925). 

17.  Heisenberg,  Z.  Phys.,  26,  196  (1924). 

18.  Debye,  "Polar  Molecules,"  p.  63,  New  York,  The  Chemical  Catalog  Co.,  Inc. 

(Reinhold  Publishing  Corp.),  1929. 

19.  Fajans,  "Radioelements  and  Isotopes — Chemical  Forces,"  pp.  63  and  76,  New- 

York,  McGraw-Hill  Book  Company,  1931. 

20.  Clark,  "The  Fine  Structure  of  Matter,"  Vol.  II,  Part  II,  p.  405,  "Molecular 

Polarization,"  New  York,  John  Wiley  &  Sons,  Inc.,  1938. 

21.  Wasastjerna,  Z.  Phys.  Chem.,  101,  193  (1922). 

22.  Fajans  and  Joos,  Z.  Phys.,  23,  1  (1924). 

23.  Horn  and  Heisenberg,  Z.  Phys.,  23,  388  (1924). 

24.  Mayer  and  Mayer,  Phys.  Rev.,  43,  610  (1933). 

25.  Bauer  and  Fajans,  /.  Am.  Chem.  Soc.,  64,  3023  (1942). 


ELECTROSTATIC  THEORY  OF  COORDINATIOh   COMPOUNDS      123 

many  metal  cations  with  amines  varies  in  the  order  Nib  equal  to  or  greater 
than  a  primary  amine  >  secondary  >  tertiary  amine,4  while  the  coordinat- 
ing ability  of  the  phosphines  appears  to  increase  in  the  order  phosphine  to 
trisubstituted  phosphine2*. 

The  elements  oxygen  and  sulfur  in  (Iroup  VI  show  relations  similar  to 
those  observed  for  the  Group  V  elements.  Coordinating  ability  decreases  in 
the  series  water,  alcohol,  ether  in  a  manner  analogous  to  the  decrease  on 
going  from  ammonia  to  the  tertiary  amines.  On  the  other  hand,  coordinating 
ability  increases  in  the  series  hydrogen  sulfide,  mercaptans,  thioethers,  just 
as  in  the  case  of  the  phosphines  and  substituted  phosphines.  In  short,  alky] 
substitution  on  the  first  short  period  elements,  oxygen  and  nitrogen,  de- 
creases their  coordinating  ability,  while  alkyl  substitution  on  the  second 
short  period  elements,  sulfur  and  phosphorus,  increases  their  coordinating 
ability.  While  one  is  probably  not  justified  in  claiming  that  such  generaliza- 
tions are  completely  explained  by  the  electrostatic-polarization  treatment, 
it  is  significant  that  the  treatment  permits  a  good  correlation  between  the 
stability  of  some  of  the  complexes  and  certain  fundamental  properties  of  the 
coordinated  groups  and  metal  ions. 

The  fact  that  some  ions  coordinate  with  ammonia  more  strongly  than  with 
water  while  others  coordinate  with  water  in  preference  to  ammonia  has  been 
treated  by  a  number  of  different  investigators,27- 28  using  the  electrostatic 
model.  Verwey25d  first  recognized  that  the  attraction  between  an  ion  and  a 
molecule  will  depend  upon  the  strength  of  the  electrostatic  field  around  the 
central  cation  and  upon  the  total  dipole  moment  of  the  coordinated  mole- 
cule. In  turn,  the  total  dipole  moment  of  the  coordinated  group  depends 
upon  its  permanent  dipole  moment,  P,  and  upon  the  induced  moment,  p'.f 
(Total  Moment  =  P  +  p').  The  moment  induced  in  a  given  molecule  (pf) 
is  determined  by  the  strength  of  the  inducing  electrostatic  field,  E,  and  the 
electronic  polarizability,  a,  of  the  molecule  (Total  Moment  =  P  +  p'  = 

*  Sidg\vick26a  pointed  out  that  in  general  the  ability  to  coordinate  decreases  in  the 
order  XH3  ,  XH2R,  XHR2  ,  NR3  ,  but  the  rule  is  not  inviolate.  In  the  case  of  SnCb  , 
all  amines  coordinate  almost  equally  well.  For  the  iron  (III)  ion,  data  are  uncertain, 
hut  the  trend  seems  to  be  reversed.  Useful  data  are  limited  in  number. 


f  The  energy  for  such  a  system  is  approximated  by  the  expression  x 


-*(-?) 


where  the  factor  1£  in  the  second  term  compensates  for  energy  expended  in  inducing 
the  dipole. 

26.  Sidgwiek,  J.  Chem.  Soc,  1941,  433;  Hertel,  Z.  anorg.  Chem.,  178,  200  (1929); 
Carlson.  McReynoldfl,  ami  Verhoek,  J.  Am.  Chem.  Soc. ,67,  1336  (1945);  Spike 
and  Parry, ./  Joe.,  75,  2726  (1953). 

■27.  Van  Arkr-1  arid  de  Boer,  Rec.  trnv.  chim.,  47,  593  (1928). 

28.  Garrick,  Phil.  .Mag.,  [7]  9,  131  (1930);  [7],  10,  76  (1930);  (b)  [7]  11,  741  (1931); 
(c)  Magnus,  Z.  Phys.,  23,  241  (1922);  (d)  Verwey,  Chem.  Wcekblad.,  25,  250 
(1928). 


124  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

P  +  aE).  While  water  has  a  higher  permanent  dipole  than  ammonia,  am- 
monia has  a  much  higher  polarizability  which  gives  a  higher  induced  dipole 
under  the  same  conditions.  Thus  the  total  dipole  of  the  ammonia,  (P  +  aE), 
ina  strong  field  may  easily  exceed  the  total  dipole  moment  of  the  water 
molecule  in  the  same  field.  This  line  of  reasoning  then  suggests  that  for 
inert  gas  type  ions  of  low  charge  and  large  size  (small  external  field,  E) 
water  will  coordinate  more  strongly  because  the  induced  dipole  contribution 
is  small,  while  for  smaller  central  ions  with  greater  external  fields  (i.e., 
greater  polarizing  power),  ammonia  will  coordinate  more  easily. 

A  semiquantitative  electrostatic  treatment  of  hydrate  and  ammine  for- 
mation by  Van  Arkel  and  De  Boer27  suggested  that  for  univalent,  noble  gas 
type  ions,  which  are  larger  than  the  lithium  ion,  the  hydrate  should  be  more 
stable  than  the  ammine;  for  the  lithium  ion,  they  should  be  about  equally 
stable;  and  for  smaller  ions  of  higher  field  strength  than  lithium,  the  am- 
mine should  be  the  more  stable.  These  predictions  are  in  agreement  with 
fact.  Bjerrum7a  was  unable  to  detect  any  potassium  ammine  formation  in 
aqueous  solution,  but  the  lithium  ion  forms  detectable  amounts  of  ammine 
complexes  in  solutions  containing  ammonia  at  concentrations  above  one 
normal7\*  In  addition,  the  heat  of  reaction  between  lithium  bromide  and 
two  moles  of  gaseous  ammonia  is  12.7  kcal,  while  that  for  lithium  bromide 
with  two  moles  of  gaseous  water  is  15.3  kcal.  The  difference  of  2.6  kcal  is 
small  and  in  favor  of  greater  hydrate  stability.  On  the  other  hand,  the  small 
doubly  charged  beryllium  ion  forms  a  much  more  stable  ammine,  as  is  sug- 
gested by  comparing  the  heats  of  reaction  for  the  processes : 

BeCl2(s)  +  4NH3(ff)  ->  Be(NH3)4Cl2(s)  +  34.1  kcal 

BeCl2(s)  +  4H20((7)  ->  Be(H20)4Cl2(s)  +  20.8  kcal 

The  behavior  of  the  very  small  hydrogen  ion  is  in  accord  with  this  principle, 
since  it  forms  an  ammine  which  is  much  more  stable,  NH4+,  than  the  cor- 
responding hydrate,  H30+. 

The  importance  of  ion  type  (i.e.,  inert  gas,  palladium,  or  transition  types) 
in  determining  field  strength  around  the  metal  ion  must  not  be  overlooked 
in  the  electrostatic  treatment.  Although  copper(I)  and  sodium  ions  have 
approximately  the  same  charge-size  ratio,  the  palladium-type  copper  (I)  ion 
has  a  much  stronger  field  than  the  inert  gas-type  sodium  ion.  (The  ionization 
potential  of  sodium  is  5.14  ev,  that  of  copper  is  7.72.)  Failure  to  recognize 
this  fact  has  led  to  unwarranted  criticism  of  the  electrostatic  approach. 
The  existence  of   stable  ammines  of  silver(I),  copper(I),  zinc(II),  cad- 

*  It  should  be  noted  that  this  relationship  may  be  obscured  if  the  field  is  strong 
enough  to  force  a  proton  from  the  water  to  form  a  hydroxide  ion,  [i.e.,  B(OH)3  forms 
instead  of  a  complex  B(OH2)3+++]. 


ELECTROSTATIC  THEORY  OF  couHMX AT/u.\   (OMl'OUXDS      125 

miunu  II ),  copper (II),  and  other  related  ions  iii  water  solution  seems  reason- 
able, if  ion  type  is  considered,  since  the  dipole  moment  induced  in  the  D0- 
laiizable  ammonia  molecule  by  the  Strong  field  of  the  metal  ions  more  than 
compensates  for  the  difference  between  the  permanenl  dipoles  of  water  and 
ammonia. 

Representation  of  the  greater  field  strength  around  palladium-  and 
transition-type  ions  in  terms  of  any  physical  model  is  difficult;  however,  a 
rather  crude  illustration  may  be  obtained  if  the  18  electron  shell  of  the 
palladium  and  transition  types  of  ions  is  regarded  as  being  softer  and  hence 
more  easily  deformed  and  penetrated  than  the  inert  gas  type  shell.  The  ease 
of  such  deformation  is  related  to  the  polarizability  of  the  central  ion.  The 
silver  ion  is  much  more  easily  polarized  than  the  potassium  ion  of  sup- 
posedly equal  size19, 30, 31.  The  role  of  polarization  and  interpenetration  in 
complex  formation  may  be  illustrated  by  the  following  drawings  which  were 
first  suggested  by  Fajans  (Fig.  3.1).  In  Fig.  3.1A  no  deformation  of  either 


A-   NO    POLARIZATION 


B-  POLARIZATION    OF 
COORDINATED    DIPOLAR 
MOLECULE 


C- POLARIZATION    OF 

BOTH      CATION    AND 

COORDINATED    DIPOLAR 

MOLECULE 


Fig.  3.1.  The  role  of  deformation  in  coordination 


the  cation  or  dipolar  molecule  has  occurred  and  the  charges  are  separated 
by  the  distance  rA  ;  in  Fig.  3. IB  the  coordinated  groups  have  been  de- 
formed and  the  negative  pole  of  the  groups  is  pulled  in  toward  the  positive 
cation.  In  this  case,  the  distance  between  the  positive  and  negative  charges, 
rB  ,  is  shorter  than  the  distance  rA  (Fig.  3.1A)  and  the  resulting  potential 
energy  of  the  system  is  reduced,  giving  a  greater  stability.  In  Fig.  3.1C  both 
ntral  ion  and  the  coordinated  groups  have  been  deformed,  producing  a 


30.  Pauling,  "Nature  of  the  Chemical  Bond, 

University  Press,  \\)Y2. 

31.  Fajans,  Ceramic  Age,  64,  288  (1949). 


p,  376,   Ethaca,  New    York,  Cornell 


L26  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

still  smaller  distance  of  separation,  rc  ;  case  C  represents  the  most  stable 
bond.* 

As  the  positive  charge  on  the  ('(Mitral  cation  increases,  its  polarizability 
decreases.  As  a  result,  cation  polarizability  and  deformability  are  of  greatest 
importance  in  ions  of  low  charge.  Cation  deformability  and  ion  size 
are  of  major  importance  in  differentiating  the  A  and  B  subgroups  of 
the  periodic  table.  The  A  group  ions,  with  8  outer  electrons,  are  not  de- 
formed easily,  while  the  B  type  ions,  with  18  outer  electrons,  are  more 
easily  deformed  and  penetrated.  Since  deformation  differences  are  most 
pronounced  with  cations  of  low  valence,  subgroups  I A  and  IB  of  the  periodic 
table  exhibit  the  most  startling  contrasts  in  behavior.  The  differences  (di- 
minish as  the  charges  on  the  ions  increase.  As  a  result,  tetravalent  ions  of 
both  Groups  IV  A  and  IV  B  are  of  low  deformability  and  are  very  similar 
in  their  complexing  properties. 

The  above  discussion  suggests  at  least  five  major  factors  which  must  be 
considered  in  estimating  the  amount  of  energy  released  when  a  free  gaseous 
metal  ion  unites  with  a  gaseous  dipolar  molecule  to  form  &free  gaseous  complex 
ion  (i.e.,  Ag+(6)  +  2NH3(ff)  -*  [Ag(NH3)2]+(,)  .  These  factors  include:  (1) 
the  charge  and  size  of  the  central  ion  (ionic  potential) ;  (2)  the  deformability 
of  the  central  ion,  which  is  in  turn  determined  by  the  electronic  structure  of 

*  Van  Arkel  and  de  Boer27  used  the  following  equation  to  represent  the  phenome- 
non in  C .  Situation  A  is  represented  by  omission  of  terms  2,  3,  4,  and  5,  while  B  is 
represented  by  omission  of  terms  3  and  5. 

^  =  zeP  _  ep^  _  2(P  +  p')PA        (pO2       P\ 
r2  r2  r3  2a  2aA 

where      t  =  the  potential  energy  of  the  gaseous  complex  ion. 

e  =  the  charge  on  the  electron. 

P  =  permanent  dipole  moment  of  the  coordinated  molecule. 

p'  =  the  additional  dipole  moment  induced  in  the  coordinated  molecule. 

r  =  the  distance  between  the  center  of  the  central  ion  and  the  center  of  the 
dipole  of  the  coordinated  molecule. 

a  =  polarizability  of  the  coordinated  molecule. 

P A  =  the  dipole  or  quadripole  moment  induced  in  the  central  metal  ion. 

aA  =  the  polarizability  or  ease  of  deformation  of  the  central  metal  ion. 

The  first  term  in  the  expression  represents  the  energy  change  due  to  interaction 
of  the  permanent  dipole  and  the  cent  tal  ion;  the  second  term,  the  energy  change  due 
to  interacl  ion  of  1  lie  induced  dipole  and  t  he  cent  ral  cation;  the  third  term,  interaction 
between  the  induced  dipole  of  the  cation  and  the  total  dipole  of  the  coordinated 
group;  while  the  fourth  and  fifth  terms  represent  the  energy  required  to  polarize  the 
coordinated  molecule  and  the  central  cation,  respectively. 


ELECTROSTATIC  THEOR]    OF  COORDINATION*   COMPOUNDS      L27 


Table  3.1.  Some  Phthcal  1 

BOPBB1  [B8  ml 

AlMMONIA,   PHOSPHINE,    \\n  Aiwm. 

Molecule 

Dipole  Moment  (< 

H  X   Distance 

1  [  X  I  [  Angle 

Ht.  of 

Pyramid  (A) 

Polariza- 

bility 
X   10"  (a) 

MI 

I'll 
\>H3 

1.46  X  10-18ab 
0.55  X  10~18  b 
0.16  X  10-18b 

1.016Aa 
L.46A* 

1.523Ad 

108°  a 
99°  c 
91°  34' •' 

3.60" 
0.67° 
0.93" 

.22'' 

.48b 
.58b 

•  Martin.  ./.  Phys.  Colloid  Chan.,  51,  14(H)  (1947). 

h  Maryott  and  Buckley,  '■Table  of  Dielectric  Constants  and  Electric  Dipole  Mo- 
ments." Natl.  BUT.  Stats.  Circular  537  (1953). 

'  Pauling, ./.  Chem.  Soc.t  1948,  1461 ;  "Valence  Commemeratiff  Victor  Henri,  Liege, 
47. 
d  Nielsen,  ./.  Chem.  Phys.,  20,  1955  (1952). 

•  Meisenheimer,  Z.  Phys.  Chan.,  97,  304  (1921). 


194: 


the  ion  (i.e.,  inert  gas,  palladium,  or  transition  type);  (3)  the  magnitude  of 
the  permanent  dipole  in  the  coordinated  molecule;  (4)  the  polarizability  of 
the  group  to  be  coordinated  (this  is  important  in  determining  the  size  of  the 
induced  dipole);  and  (5)  the  size  of  the  group  being  coordinated  (this  influ- 
ences the  distance  between  the  central  ion  and  the  center  of  negative  charge 
in  the  coordinated  group).  If  a  charged  ion  is  being  coordinated  instead  of  a 
dipolar  molecule,  the  charge  on  the  ion  will  also  be  important. 

Coordination  Compounds  of  Phosphine  and  Hydrogen  Sulfide.  Experi- 
mentally, it  is  found  that  phosphine  coordinates  much  less  strongly  than 
ammonia  with  all  of  the  metal  ions  which  have  been  studied.  This20  is  not 
unexpected  since  phosphine  has  a  much  smaller  permanent  dipole  moment 
and  a  larger  central  atom  than  ammonia.  Comparative  data  for  ammonia, 
phosphine,  and  arsine  are  cited  in  Table  3.1.  Although  the  polarizability  of 
the  phosphine  molecule  is  twice  as  large  as  that  of  ammonia,  the  magnitude 
of  the  induced  dipole  is  not  large  enough  to  overcome  the  adverse  effects  of 
low  permanent  moment  and  large  molecular  size.  Holtje  and  Schlegel32 
prepared  the  following  phosphine  complexes: 

CuCl-2PH3  CuClPH,  AgI0.5PH3 

CuBr-2PH3  CuBrlMl  AgIPH3 

CuI2PH,  CuIPH,  AuIPH, 

The-'-  were  unstable  as  compared  to  the  ammines.  One  would  expect  the 
mosl  -table  coordination  compounds  of  phosphine  with  cations  of  high  po- 
larizing power  such  a>  Ag  .  <>r  I  ly.  " .  In  such  a  case  the  induced  dipole  con- 
tribution would  be  relatively  large. 

Arsine,  of  -mailer  permanent  moment  (0.15  X  10  w  e.s.u.)  than  phos- 
phine, coordinate-  with  even  greater  difficulty,  despite  the  fact  that  arsine 
is  more  polarizable. 

Holtje  and  Schlegel,  '/.   anorg.  Allot  m.  Chi  m..  243,  246    1940] 


128  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Hydrogen  sulfide  bears  the  same  relationship  to  water  that  phosphine 
bears  to  ammonia.  Though  hydrogen  sulfide  is  more  polarizable  than  water 
(refractivity:  H>0  =  3.7  cc;  H2S  =  9.5  cc19),  the  larger  size  and  smaller 
permanent  moment  of  the  H2S  molecule  (H20  =  1.89  X  10-18  e.s.u.;  H2S 
=  about  1.1  X  10-18  e.s.u.20)  reduce  its  coordinating  ability  to  a  point 
below  that  of  water  for  ions  of  low  field  strength.  For  ions  of  high  field 
strength  (Hg++,  Ag+  etc.)  the  hydrogen  sulfide  coordinates  and  the  pro- 
tons are  forced  off  to  give  insoluble  metal  sulfides. 

Coordinating  Ability  of  Alkyl  Substituted  Hydrides  of  Group  V  and  Group 
VI  Elements.  The  coordinating  abilities  of  the  alkyl  and  aromatic  deriva- 
tives of  ammonia,  phosphine,  water,  and  hydrogen  sulfide  also  show  a  fairly 
good  correlation  with  the  permanent  dipole  moments  of  the  molecules.  The 
decrease  in  coordinating  ability  from  water  to  alcohol  to  ether  and  from 
ammonia  to  primary  amine  to  secondary  amine  to  tertiary  amine  runs 
parallel  to  a  decrease  in  the  permanent  dipole  moment  of  the  molecules. 
This  is  shown  in  Table  3.2.  Polarizabilities,  where  available,  are  also  in- 
cluded. The  increase  in  the  coordinating  ability  in  the  series  H2S,  RHS,  R2S 
runs  parallel  to  an  increase  in  the  dipole  moment  of  the  compounds.  A 
similar  relationship  is  noted  for  the  phosphines.  Very  stable  tertiary  phos- 
phine complexes  have  been  described  by  many  investigators37  (see  Chap- 
ter l,p.  78). 

In  a  similar  manner,  the  fact  that  the  cyclic  tertiary  amine,  pyridine, 
coordinates  more  strongly  than  most  other  tertiary  amines  can  be  correlated 
with  its  higher  dipole  moment,  which  is  even  higher  than  that  of  ammonia 
(Table  3.2). 

It  will  also  be  observed  that  the  polarizability  of  the  bonding  electrons33 
(i.e.,  the  electrons  on  the  nitrogen  or  phosphorus  atom)  is  decreased  in  all 
cases  by  alkyl  substitution,  but  the  per  cent  decrease  in  going  from  H20  to 
R20  is  much  greater  (about  24  per  cent)  than  the  decrease  in  going  from 
H2S  to  R2S  (about  5  per  cent).  The  per  cent  decrease  in  going  from  NH3 
to  R3N  (about  12  per  cent)  is  likewise  greater  than  the  per  cent  decrease  in 
going  from  PH8  to  R3P  (about  5  per  cent).  From  this  it  appears  that  the 
polarizability  factor  also  favors  the  differences  in  relative  stabilities  out- 
lined above. 

33.  Reference  34,  p.  152. 

'M.  Smyth,   "Dielectric  Constant    and  Molecular  Structure,"  p.   192,  New  York, 
Chemical  Catalog  Co.,  Inc.,  (Reinhold  Publishing  Corp.),  1931. 

35.  Kodama  and  Parry,  unpublished  results. 

36.  Sidgwick,  "The  Electronic  Theory  of  Valency,"  p.  152,  London,  Oxford  Uni- 

versity Press,  1927. 
37     Mann  and  Purdie,  Chem.  and  Intl.,  1935,  814;  Mann,  Wells,  and  Purdie,  J.  Chem. 
Boo.,  1937,  1828. 


Table  3.2.  Moi  \n  EIefbactivitibs  \m>  Dipole  Moments  oj   A.lkyl  Si  bbtiti  rso 

Hydrides 


Molecule 

Refractivit) 

X  in   R  X    R 

Permanent  Dipole 
Moment  X  10'»  < 

Coordinating  Ability 

11  o 

3.7  cc 

L.89     20 

1 
■2 

('II  oil 
C  lUOH 
n-CM.OU 

Aboul  3.2  cc 

1.68 
1.69 
1.66 

0 
- 

■~ 

P 

(CH3)20 

rii,).:0 

(m-C3H7)20 

About  2.8  cc 

1.29 
1.15 
1.16 

3 

H> 

9.6  cc 

1.1   (20) 

3 

CH3SH 

C  11   SH 
/<-C3H7SH 

About  9.4  cc 

1.39 
1.33 

2 

0 

Q 

(CH3)2S 
(C2H5)2S 
t//-C3H7)2S 

About  9.1  cc 

1.40 
1.58 
1.55 

1 

XH3 

5.6 

1.49   (34) 

1 

CH3NII. 
C2H5NH2 

About  5.1  cc 

1.23 
1.3 

2 

0J 

Q 

(CH,)»NH 

(C2H5)A'II 

About  4.8 

0.96 
1.20 

3 

(CH,)iN 

C  IU)3X 
(C,H6),X 

About  4.7  cc 

0.6 

0.90 

0.26 

4 

Pyridine 

— 

2.1 

1 

Unusual    ability    for 
tert.  amine. 

PH3 

About  11.9  cc 

0.55 

1 

CB  I'll 
( MliPHo 
»-CH7PB 

— 

1.17   (35) 

3 

c 

a 

e 

u 

0> 

Q 

(CH3)2PH 
(CH^aPH 

— 

1.4  (35) 

2 

(CH.),P 

(C6H5)3P 
(C,H.),P 

Aliout   11.3  cc 

L.45  (36) 
1.1.-)  (35) 

1 

L29 


130  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Instability  Constants  for  Complexes  and  the  Polarized  Ionic  Model 

In  1953  Irving  and  Williams7b  completed  a  most  thorough  analysis  of 
essentially  all  the  data  available  on  the  instability  constants  of  complexes 
of  dipositive  ions  of  the  transition  metals  of  the  first  period.  The  order 
Mn  <  Fe  <  Co  <  Ni  <  Cu  >  Zn  was  found  to  hold  for  the  stability  of 
nearly  all  such  complexes  irrespective  of  the  nature  of  the  coordinated 
ligand  or  the  number  of  ligand  molecules  involved.  They  demonstrated  the 
failure  of  an  electrostatic  model  which  neglects  polarization  terms  and 
showed  that  Pauling's  theory39  (Chapter  4)  fails  to  account  even  qualita- 
tively for  the  order  of  stability  of  metal  complexes.  On  the  other  hand,  they 
showed  in  a  most  convincing  manner  that  the  above  Irving- Williams  order 
of  the  transition  metal  (II)  cations  follows  logically  from  considerations  of 
the  reciprocal  of  the  ionic  radii  and  the  second  ionization  potentials  of  the 
metals  concerned.  It  is  apparent  that  these  are  the  very  parameters  which 
are  indicative  of  the  electrostatic  field  strength  of  the  cations  of  the  transi- 
tion metals  involved.  They  point  out  that  if  attempts  are  made  to  introduce 
other  cations  such  as  Cd^-1-  into  the  sequence,  difficulties  arise.  This  is 
readily  understood  as  they  describe,  and  can  also  be  correlated  with  the 
fact  that  the  cation  polarizabilities  (deformabilities)  of  the  transition  metal 
and  palladium  type  ions  differ;  thus  the  order  of  stability  would  be  de- 
pendent upon  the  ligand  selected  [i.e.,  compare  treatment  of  ammines  and 
hydrates  of  Na+  and  Ag+  in  which  cation  polarizabilities  differ.]  As  noted 
by  these  authors,  other  factors  such  as  steric  hindrance  and  entropy  terms 
must  also  be  considered  for  a  thorough  analysis  of  complex  stability. 

Physical  Properties  of  Complex  Compounds  and  the  Ionic  Model 

Color  and  Structure.  The  remarkable  colors  commonly  associated  with 
coordination  compounds  were  attributed  by  Fajans41  to  a  strong  deforma- 
tion of  the  electron  clouds  of  the  coordinated  groups.  This  concept  was 
amplified  by  Pitzer  and  Hildebrand42.  Orgel43  has  recently  considered  the 
similarity  in  the  spectra  of  Cr+++  and  Co+++  as  a  consequence  of  the 
Stark  splitting  of  the  d  levels  by  the  strong  crystal  field.  The  crystal  field 
theory  is  discussed  in  connection  with  magnetism  and  may  yet  provide  a 
sound  interpretation  of  the  color  of  complex  ions.* 

39.  Pauling,  J.  Chem.  Soc,  1948,   1461;   "Valence  Commemoratiff  Victor  Henri, 
Liege,  1917 

41.  Fajans,  Naturwissenschaften,  11,  165  (1923);  Remarks  to  this  paper,  circulated 

privately,  1946. 

42.  Pitzer  and  Hildebrand,  J.  Am.  Chem.  Soc,  63,  2472  (1941). 

43.  Orgel,  ./.  Chem.  Soc,  1952,  4756. 

*  Note  added  in  proof:  In  a  recent  series  of  papers  from  J.  Bjerrum's  laboratory, 
Bjerrum,  Jdrgensen  and  others  have  treated  the  color  of  complexes  of  Cu"1"1:,  etc., 
using  .ni  electrostatic  model.  Acta.  Chem.  Scand.,  8,  1289  (1954);  9,  116,  1362  (1955). 


ELECTROSTATIC  THEORY  OF  COORDINATION  COMPOUNDS      131 

Stereochemistry  and  fh<  Polarized  Ionic  Model.  The  rigid  ionic  model  of 
Kossel  leads  to  a  linear  molecule  for  coordination  number  two,  a  planar 
Structure  for  coordination  number  three,  a  tetrahedral  molecule  for  coor- 
dination number  four,  and  a  regular  Octahedron  for  coordination  number 
six.  Deviations  from  these  forms  have4  been  attributed  to  polarization16 •  18. 
Because  of  the  success  of  the  polarization  treatment  in  justifying  the 
stereochemistry  of  the  water  molecule,  several  attempts  have  been  made  to 
justify  the  planar  structure  of  platinum(Il)  complexes  on  the  basis  of  the 
large  polarizability  of  the  central  platinum(II)  ion27.*  Xekrasov44  used  polari- 
zation and  the  radius  ratio  to  justify  the  planar  structure.  Values  of  the 
radius  ratio  below  0.41  supposedly  favor  a  tetrahedral  arrangement,  while 
high  polarizability  of  the  coordinated  ligand  and  values  of  the  radius  ratio 
greater  than  0.41  presumably  favor  a  planar  arrangement^ 

Tsuchida"  and  co-workers  developed  a  stereochemical  theory  which  might 
be  considered  as  a  compromise  between  the  ionic  model  and  the  electron 
pair  bond  model.  They  considered  that  all  coordination  compounds  are 
built  up  from  ions,  polar  molecules,  and  stereochemical^  active  electron 
pairs  (or  odd  electrons  in  some  cases).  The  shape  of  a  molecule  would  then 
be  determined  by  the  most  symmetrical  grouping  of  these  ligands  around  a 
cation.  Walsh46  has  recently  given  a  molecular  orbital  treatment  to  simpler 
molecules  which  leads  essentially  to  the  rules  of  Tsuchida,  but  without  the 
ionic  implications.  According  to  Tsuchida,  the  charge  of  the  cation  would  be 
equal  to  its  position  in  the  periodic  table  except  for  the  transition  elements, 
whose  charge  would  be  equal  to  the  accepted  oxidation  state  of  the  ion  under 
consideration  (i.e.,  Fe4-1-1-).  In  such  a  scheme  molecular  shape  wrould  be 
determined  by  the  number  of  coordinating  groups  (including  stereochem- 
ically  active  electron  pairs).  The  shapes  proposed  for  different  numbers  of 
groups  are:  linear  for  2;  planar  for  3,  tetrahedral  for  4;  octahedral  for  6,  and 
cubic  for  8. 

Special  attention  was  given  to  transition  elements  with  a  coordination 
number  of  four  in  planar  arrangement.  It  was  noted  that  such  metals  con- 

*  Cases  of  planar  coordination  have  been  experimentally  established  only  for 
complexes  in  the  solid  states  or  in  solution.  Fajans  has  raised  the  interesting  possi- 
bility that  the  planar  arrangement  may  be  due  in  part  to  electric  field  effects  in  the 
crystal  or  in  solution.  If  so,  a  planar  structure  might  not  appear  in  the  vapor  state. 
f  The  conclusions  regarding  radius  ratio  are  the  same  as  those  advanced  by  Strau- 
bel  and  Huttig  in  1925.  (p.  143,  ref.  75  and  76). 

44.  Xekrasov,  J.Gen.Chem.  U.S.S.R.,  16,  341  (1946);  cf.  Chem.  Abs.,  41,  633  (1947). 
15.  Tsuchida,  Bull.  Chem.  Soc.  Japan,  14,  101  (1939);  J.  Chem.  Soc.  Japan,  60,  245 
(1939);  Rev.  Phys.  Chem.  Japan,  13,  31  (1939);  Tsuchida  and  Kobayaahi    /.'<  i 

('hem.  Japan,  13,  61   (1939);  Tsuchida,  Kobaya&hi,  and  Kuroya,  Rev. 
Phys.  Ok  in.  Japan .  13,  151   (1939);  Tsuchida,  Collected  Papers  Faculty  Sci., 
Osaka  Imp.  Univ.  [C]  6,  No.  35  (1938). 
46.  Walsh, ./.  Chem.  Soc,  1953,  2260,  2266,  2288,  2296,  2306. 


132  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

tain  nearly  full  d  levels  (i.e.,  8  electrons);  hence,  two  pairs  of  electrons  could 
become  stereochemically  active,  one  above  and  one  below  the  plane  to  give 
an  octahedral  configuration  instead  of  the  apparent  planar  structure.*  If  the 
d-level  contains  less  than  four  electrons,  such  coordination  would  be  im- 
possible and  a  tetrahedral  structure  would  be  mandatory.  The  basis  for 
determining  which  electron  pairs  would  be  stereochemically  active  in  planar 
complexes  was  never  clearly  defined  although  one  could  now  make  reason- 
able decisions  on  the  basis  of  the  crystal  field  splitting  of  the  d  levels44. 

Tsuchida's  theoryf  is  interesting  in  that  it  provides  a  simple  empirical 
scheme  for  many  stereochemical  predictions,  but  it  is  unrealistic  in  its 
chemical  implications.  For  example,  attributing  hydridic  character  to  the 
hydrogens  of  water  and  ammonia  is  obviously  unreasonable  in  view  of  the 
latent  acid  character  of  these  two  solvents. 

The  fundamental  stereochemical  ideas  of  Tsuchida  without  the  accom- 
panying chemical  objections  are  embodied  in  the  modern  quanticule  theory 
of  Fajans49.  The  electron  pair  is  retained  as  a  coordination  group  in  certain 
formulations  but  chemical  contradictions  are  avoided.  For  example,  water 
is  considered  as  a  polarized  oxide  ion  with  two  imbedded  protons.  Ammonia 
is  considered  as  a  nitride  ion  with  three  imbedded  protons.  In  both  cases  the 
correct  geometry  can  be  obtained,  if  polarizability  of  the  anion  is  considered 
in  a  quantitative  fashion18.  Fajans  also  differentiates  certain  chemically 
recognizable  groups  as  a  single  "quanticule"  or  group  of  atoms  with  common 
quantization.  For  example,  the  peroxide  ion  would  represent  a  quanticule 
composed  of  two  oxygen  atoms  with  essentially  molecular  quantization  of 
the  electrons  between  them.  In  this  respect  and  others,  it  has  much  in 
common  with  the  qualitative  aspects  of  the  molecular  orbital  theory.  The 
CH3~"  quanticule  (ion)  would  be  considered  as  a  starting  point  for  a  polari- 
zation treatment  of  [Pt(CH3)4]4  in  order  to  avoid  the  problem  of  hexaco- 
valent  carbon  (see  p.  165).  More  detailed  examples  are  given  by  Fajans. 

Magnetism  and  the  Polarized  Ionic  Model.  It  is  a  well  known  fact,  widely 
used  in  spectroscopy,  that  the  energy  levels  in  an  atom  or  ion  will  be  altered 
by  the  presence  of  a  magnetic  or  electrostatic  field  [Zeeman  effect  and 
Stark  effect].  If  the  magnetic  field  is  very  strong,  the  spin  and  orbital  vec- 
tors of  angular  momentum  can  no  longer  be  combined  to  give  the  quantum 
number  J,  but  each  vector  is  space  quantized  independently  to  give  inde- 

*  Others47  have  also  raised  this  possibility. 

f  A  set  of  empirical  structural  rules  which  utilize  a  stereochemically  active  elec- 
tron pair  was  also  proposed  by  Helferich.48 

47.  Sidgwick,  J.  Chem.  Soc,  123,  730  (1923);  Fowler,  Trans.  Faraday  Soc,  19,  468 

(1923);  Sidgwick  and  Powell,  Proc.  Roy.  Soc.  London,  176A,  159  (1940). 

48.  Helferich,  Z.  Naturforsch,  1,  666  (1946);  cf.  Chem.  Abs.,  41,  6086  (1947). 

49.  Fajans,  Chem.  Eng.  News,  27,  900  (1949). 


ELECTROSTATIC  THEORY  OF  COORDINATION  COMPOUNDS      L33 

pendent  orbital  and  spin  interactions  with  the  field.  This  is  known  as  the 
Paschen-Back  effect  and  indicates  thai  the  field  is  stronger  than  the  spin- 
orbit  coupling.  The  uncoupling  of  the  L  and  S  vectors  by  a  strong  electro- 
static held  [i.e.,  an  electrostatic  Paschen-Back  effect  |  is  also  possible  though 
not  as  widely  recognized.  The  electrostatic  field  in  crystals  is  strong  and  it 

is,  indeed,  this  resulting  "electrostatic  Paschen-Back  effect"  which  makes 

the  magnetic  properties  of  the  first  transition  elements  differ  from  those 

of  the  rare  earth-. 

If  an  even  stronger  field  is  imposed  upon  the  d  electrons  of  a  cation,  their 

interaction  with  the  field  becomes  so  strong  that  the  ground  state  of  the 
ion  can  no  longer  be  obtained  by  using  Himd's  rules  for  electron  distribu- 
tion (i.e.,  rule  of  maximum  multiplicity)  and  then  combining  individuals 
values  by  means  of  Russell-Saunders  coupling.  New  formulas  are  then 
arv  to  calculate  the  magnetic  moment  of  the  ion;  the  value  is  no 
longer  determined  by  the  procedures  used  for  the  simple  ion.  This  situation 
is  applicable  to  many  complex  compounds. 

In  recent  years  the  powerful  new  tool  of  paramagnetic  resonance  absorp- 
tion has  been  developed,  permitting  a  much  more  detailed  knowledge  of  the 
magnetic  properties  of  complexes  than  has  been  possible  heretofore.* 
Crystal  field  theory  has  frequently  been  applied  to  treat  the  detailed  data. 

The  details  of  the  crystal  held  theory  may  be  outlined  as  follows.  A  cen- 
tral metal  cation  is  surrounded  by  anions  or  dipoles,  i.e.  [Ir4+  Cl6_]=  or 
[Fe"l~H"(CX~)6]-,  which  set  up  a  strong  electrostatic  or  crystalline  field.  In 
this  electrical  field  the  normally  degenerate  d  levels  are  split  as  in  the 
familiar  spectroscopic  Stark  effect,  the  extent  of  the  splitting  depending 
upon  the  central  cation  and  upon  the  symmetry  and  strength  of  the  ap- 
plied field.!  The  behavior  of  the  ion  in  this  field  is  approximated  by  the 
methods  of  wave  mechanics.  The  three  cases  of:  (1)  weak  field  as  in  the  rare 
earths,  (2)  moderate  field  as  in  the  so-called  "ionic"  complexes  of  transi- 

*  The  paramagnetic  resonance  absorption  phenomenon  is  a  phase  of  microwave 
spectroscopy.  It  has  been  reviewed  in  masterful  fashion  by  Bleaney50,  and  Bleaney 
and  Stevens51. 

f  For  example,  changing  the  field  by  changing  the  ligand  in  a  complex  has  a  sig- 
nificant effect  upon  the  moment,  even  when  the  same  orhitals  are  ostensibly  used. 
For  example,  in  [CoX4]~  complexes,  the  moment  along  the  sequence  mci  >  MBr  >  m  > 
kern  falls'2.  Xvholm53  has  recently  utilized  the  results  of  the  crystal  field  treatment 
iicv  and  8chlaapMand  by  Van  Vleck**  as  a  basis  for  suggesting  thai  in  "ionic" 
Co++  complexes  a  larger  orhital  contribution  indicates  octahedral  coordination  while 
the  smaller  orbital  value  indicates  tetrahedral.  A  particularly  large  orbital  cent  ribu- 
tion  was  reported  empirically  for  planar  Co'T  comple 
Bleaney,  ./.  Phys.  Chem.}  57,  508  (1953). 

51.  Bleaney  and  Stev<  /<  Physics.,  16,  ins  (1953). 

52.  Nyholm,  Quart.  Revs.,  7,  104  (1953). 

53.  Xvholm,  ./.  Chun.  Soc,  1954,  12. 


134  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

t  ion  metals,  and  (3)  strong  field  as  in  the  so-called  covalent  complexes  can 
be  di  fferentiated.  Because  of  the  importance  of  case  three  in  the  electrostatic 
theory  of  complexes,  it  will  be  considered  more  carefully. 

In  the  presence  of  a  strong  field,  the  degenerate  d  levels  are  split  into 
sublevels.  If,  then,  the  distribution  of  electrons  in  orbits  is  based  on  these 
sublevels  rather  than  the  original  five  degenerate  d  levels,  the  magnetic 
properties  must  follow.  The  manner  in  which  the  d  levels  are  split  is  deter- 
mined by  the  field  geometry  as  shown  in  Fig.  3.2.  For  the  case  of  K2PtCl6 
(Fig.  3. 2 A)  the  normally  degenerate  d  levels  are  split  into  three  lower  and 
two  upper  levels.  Filling  the  lower  triplet  with  six  electrons  as  indicated 
gives  the  expected  diamagnetic  result.  The  cases  of  tetrahedral  Ni11,  planar 
Ni11,  and  duodecahedral  MoIV  are  also  worked  out.  In  every  case  the  quali- 
tative agreement  between  predictions  of  the  atomic  orbital,  molecular 
orbital,  and  crystal  field  theories  is  gratifying. 

These  ideas,  which  are  an  extension  of  generally  applicable  magnetic 
theory,  were  first  used  to  explain  the  magnetism  of  complex  compounds 
by  Penney  and  Schlaap54  by  Van  Vleck55  and  Van  Vleck  and  Penney57.  How- 
ard58 accounted  for  not  only  the  gross  magnetic  moment  of  K3[Fe(CN)6] 
by  this  method  but  accounted  for  the  magnetic  anisotropy  and  temperature 
dependence  of  the  moment  in  the  solid.  Kotani59  gave  a  more  rigorous 
treatment  of  the  temperature  dependence  for  several  transition  complexes. 
The  method  has  been  applied  extensively  in  recent  years  to  the  interpreta- 
tion of  paramagnetic  resonance  absorption  data50, 51-  60'  61  for  complex  ions, 
and  appears  to  be  more  tractable  than  the  orbital  theories  in  the  quantita- 
tive interpretation  of  modern  detailed  data. 

The  essential  physical  ideas  of  electron  distribution  according  to  the 
crystal  field  theory  and  their  applications  to  magnetism,  color,  planar  con- 
figuration, and  heat  of  hydration  of  the  transition  metal  cations  have  been 
considered  in  an  outstanding  paper  by  Orgel43.  The  electrons  tend  to  avoid 
those  regions  where  the  field  due  to  the  attached  negative  ions  and  dipoles 
is  largest,  a  fact  which  accounts  for  the  field  splitting  of  d  levels.  The  two 
high  energy  orbitals  correspond  to  a  high  electron  density  along  the  lines 
joining  the  central  metal  cation  with  the  attached  ligands,  whereas  the  three 
low  energy  orbitals  correspond  to  a  high  electron  density  between  these 

54.  Penney  and  Schlapp,  Phys.  Rev.,  41,  194  (1932). 

55.  Van  Vleck,  /.  Chem.  Phys.,  3,  812  (1935). 

56.  Kimball,  ./.  Chem.  Phys.,  8,  198  (1940). 

57.  Van  Vleck  and  Penney,  Phil.  Mag.,  17,  961  (1934). 

58.  Howard,  ./.  Chem.  Phys.,  3,  813  (1935). 

59.  Kotani,  ./.  Phys.  Soc.  Japan,  4,  293  (1949). 

60.  Abragam  and  Pryce,  Proc.  Roy.  Soc.  London,  206A,  164,  173  (1951). 

61.  Stevens,  Proc.  Roy.  Soc.  London,  219A,  542  (1953);  Griffiths,  Owen,  and  Ward, 

Proc.  Roy.  Soc.  London,  219A,  526  (1953). 


ELECTROSTATIC  THEORY  OF  COORDINATION  COMPOUNDS      L35 

lines.  In  this  sense  tin1  former  doublet  would  be  bonding  for  the  Ligands  and 
the  latter  triplet  would  be  Donbonding,  as  is  also  suggested  by  both  atomic 
and  molecular  orbital  theories.  The  separation  between  these  levels  can  be 
found  in  some  cases  from  the  optical  spectrum  of  the  complex,  a  fact  which 
indicate.-  that  it  may  be  possible  to  correlate  color  as  well  as  magnetism  iii 
more  definite  theoretical  terms48.  The  relationship  between  these  ideas  and 
cation  deformability  (Fig.  3.1)  is  obvious. 
Another  way  of  viewing  the  transition  from  the  paramagnetic  to  the  di- 


DEGENERATE     ORBITALS 
WITH    6d       ELECTRONS 


WEAK     OR 
MODERATE    FIELD 


UPPER     DOUBLET 


LOWER     TRIPLET 


RESULT     EQUIVALENT     TO 


SEE  P.  170 


A) 


STRONG 
OCTAHEDRAL 
FIELD    AS    IN 
K2[PtCla] 


DEGENERATE      ORBITALS 
WITH      8^.         ELECTRONS 


UPPER     TRIPLET 


•  •       *       .       . 

•  •       • 


LOWER     DOUBLET 


"STRONG  TETRAHEDRAL  FIELD;  Le.,  .NifNHj)^] 
MAGNETIC  SUSCEPTIBILITY  IS  IDENTICAL  TO 
THAT   OF   ORIGINAL    ION.  HENCE   "lONIC" 


+  + 


DCGENERATE     ORBITALS 
WITH    S  d      ELECTRONS 


:     :  •  • 


D 

m 


STRONG    PLANAR    FIELD     AS    IN 
Efl(CN)4] 

RESULT    EQUIVALENT     TO 
dsp2  HYBRIDIZATION    (56)     SEE   P.  170 


DEGENERATE      ORBITALS 
WITH    2d      ELECTRONS 


•       • 


DUODECAHEDRAL     FIELD     AS     IN     K4Mo(CN)e 
DIAMAGNETIC      (ft>) 

EQUIVALENT     TO    d4jp3   HYBRIDIZATION 
SUGGESTED    BY     KIMBALL     [j    CHEM    PHYS     fl. , 
196      (19  4  0)] 


Fig.  3.2.  Crystal  field  theory  of  magnetism 


130 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


amagnetic  state  can  be  seen  in  Fig.  3.3.  The  case  of  cobalt(III)  is  taken  as 
an  illustrative  example,  although  any  other  ion  with  paramagnetic  and 
diamagnetic  configurations  could  be  used  equally  well.  The  ground  state 
for  the  cobalt(III)  ion  is  obtained  by  Hund's  rules  as  the  lower  representa- 
tion [5d]  on  the  left-hand  side  of  Fig.  3.3.  An  excited  state  of  this  ion  [I] 
is  shown  at  a  higher  energy  on  the  left-hand  side.  If  now  a  crystalline  field 
is  applied  to  both  states,  the  relative  energies  of  each  will  undergo  change 
dependent  upon  field  direction  and  geometry.  [Each  state  will  be  split  into 


ENERGY  OF 
A  GIVEN 

EXCITED   STATE 

FOR    ISOLATED  Co"  +  10N               ^ 

ELECTRONIC 
CONFIGURATION 

I 

• 
• 

• 
• 

• 
• 

GROUND    STATE    FOP 
5D     ISOLATED    CO+++ION 

• 
• 

• 

• 

• 

• 

<^A 

• 
• 

• 

• 

• 

• 

STATES 
REVERSED 
IN    STRONG 
FIELD. 

• 
• 

• 
• 

• 
• 

INCREASING 

FIELD    STRENGTH 

^ 

Fig.  3.3.  Crystal  field  effects  on  cobalt  (III) 

different  levels  by  the  field].  If  the  excited  state  changes  in  energy  more 
rapidly  than  does  the  ground  state  [i.e.,  slope  of  line  X  greater  than  line  Y], 
the  two  configurations  will  reverse  at  the  intersection  of  lines  X  and  Y 
("A"  on  the  diagram).  The  point  "A"  then  indicates  the  strength  of  the 
crystal  field  required  to  bring  about  the  transition  from  the  "ionic"  to  the 
"covalent"  configuration.  It  is  now  immediately  apparent  that  the  location 
of  A  is  dependent  upon  the  original  energy  separation  of  the  two  levels  and 
upon  the  slopes  of  lines  X  and  Y,  (i.e.,  upon  electronic  structure  of  cation). 
11  is  interesting  to  oote  that  no  discontinuous  energy  change  is  involved 
in  the  transition  from  "ionic"  to  "covalent"  configuration  although  the 
rate  of  change  of  energy  with  field  strength  is  altered  at  this  point.  This 
fact  justifies  the  observation  of  Orgel  that  "covalent"  bonds  in  one  system 
are  ao1  necessarily  stronger  than  "ionic"  bonds  in  another  system  (see  also 
Taube82). 

62.  Taube,  Chem.  Revs.,  50,  69  (1951). 


ELECTROSTATIC  THEORY  OF  COORDINATION  COMPOUNDS      137 

Filially,  a  word  should  be  said  concerning  the  argument  involving  termi- 
nology which  arose  when  the  crystal  field  theory  was  first  introduced61.  Ob- 
jections were  raised  to  the  crystal  field  treatment  on  the  ground  thai 
[FeF«]™,  which  is  •"ionic"  according  to  magnetic  measurements, should  have 
a  stronger  crystal  field  than  [Fe  I  "\  .j  which  is  "covalent."  Such  an  argu- 
ment involves  a  matter  of  definition  of  the  terms  "ionic"  and  "covalent "  in 
relation  to  field  strength*4.  If  polarization  is  included,  the  cyanide  crystal 
field  is  stronger  than  the  fluoride  (see,  for  example,  Fig.  3.1)  and  the  observed 
moments  are  in  line  with  this  expectation.  One  might  argue  that  the  polari- 
zation of  the  cyanide  «>;roup  is  in  itself  indicative  of  covalent  character  in 
the  bond.  Such  an  argument  is  valid,  however,  solely  because  of  the  way 
chosen  to  define  the  term  "covalent"  and  in  no  way  alters  the  fundamental 
validity  of  the  crystal  field  theory.  In  short,  an  approach  involving  polariza- 
tion of  ions  leads  to  the  same  gross  qualitative  result  as  a  model  involving 
the  perturbation  of  atoms  by  mutual  interaction.  The  former  approach 
is  currently  most  useful  for  quantitative  interpretation  of  detailed  data 
on  the  magnetism  of  complexes. 

The  Thermochemical  Cycle  ix  Complex  Formation 

The  relationship  between  dipole  moment  and  coordinating  ability  is  not 
always  as  simple  as  the  section  on  chemical  properties  would  suggest. 
Hertel26b  compared  the  stability  of  complexes  formed  between  nickel(II) 
cyanide  and  methyl  amine,  ethyl  amine,  propyl  amine,  and  butyl  amine. 
Stability  was  determined  by  measuring  and  comparing  the  vapor  pressures 
of  the  amines  above  the  complexes.  The  complexes  identified  were 
Ni(CN)2-R  and  Xi(CX)2-2R  (R  =  the  original  amine).  Though  the  size 
of  the  dipole  increases  slightly  in  the  series  MeNH2,  EtXH2 ,  PrNHj  , 
BuXHo ,  the  stability  of  the  coordination  compounds  decreases  markedly 
from  methyl  amine  to  butyl  amine.  Data  are  summarized  in  Table  3.3. 

Table  3.3.  Dependence  of  Dipole  Moment  on  Size  of  Alkyl  Group  in  Primary 

Amines 


Amine 

Permanent  Dipole  Moment14 
X  10»8  e.s.u. 

Relative  Complex  Stability 

MI 

1.46 

1  Most  stable 

MeNH, 

1.23 

2 

EtNH, 

1.3 

3 

PrNHi 

about  1.3  to  1 

4 

4 

BuNH2 

about  1.3 

5  Least  stable 

63.  Paulinn.  /.  An  .  ('hem.  Soc,  54,  988  (1932);  Pauling  and  Huggins,  Z.  Krist.,  87, 

205  (1934);  Van  Vleck,  J.  Chem.  Phys.,  3,  807  (1935). 
04.  Moeller,  "Inorganic  Chemistry,"  p.  205,  New  York,  John  Wile}  and  Sons,  Inc., 

1952. 


138  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Obviously,  some  factor  which  was  neglected  in  the  simplified  treatment  is 
now  of  importance.  The  factors  previously  discussed  (page  126)  were  re- 
stricted to  the  formation  of  a  free  gaseous  complex  ion  from  a  gaseous  metal 
ion  and  the  gaseous  amine.  The  energy  released  in  this  reaction  is  the 
energy  of  coordination.  The  actual  process  which  is  usually  considered  in  the 
laboratory  involves  reaction  between  a  solid  metal  salt  and  the  amine  to 
form  the  solid  complex  compound.  In  this  process  other  energy  terms  may 
overshadow  small  differences  in  the  coordination  energy.  The  relative  im- 
portance of  each  energy  term  may  be  illustrated  by  describing  the  laboratory 
process  with  a  thermochemical  cycle. 

The  simple  crystalline  salt  is  vaporized  and  ionized  (if  it  is  not  already 
ionized) ;  then  the  gaseous  metal  ions  combine  with  the  amine  to  give  the 
complex  cation,  and  finally  the  complex  cation  and  the  salt  anion  combine 
to  give  the  solid  complex  compound.  The  process  is  represented  in  Fig.  3.4. 
All  values  are  exothermic  and  positive  in  the  direction  of  the  arrows;  then 
Q  =  E  +  Vi  —  XJ\ .  Since  accurate  entropy  data  are  not  available,  the 
heat  of  formation,  Q,  (or  —  AHiOTm),  may  be  considered  as  an  approximate 
measure  of  the  relative  stability  of  comparable  complexes.  Differences  in 
the  energies  of  coordination,  E,  are  frequently  sufficiently  large  to  over- 
shadow the  effects  of  differences  in  the  lattice  energies,  U\  and  Ui ;  i.e., 
A(Ui  —  U2),  is  small  in  comparison  to  AE  (the  differences  in  energies  of 
coordination).  In  such  a  case  the  stability  of  the  complex  can  be  correlated 
with  factors  influencing  only  the  energy  of  coordination,  E.  Such  a  situation 
is  illustrated  by  the  water,  alcohol,  ether,  and  hydrogen  sulfide,  mercaptan, 
thioether  series  discussed  earlier.  However,  in  the  cases  of  the  different  pri- 
mary alkyl  amines,  the  differences  in  the  lattice  energy  terms  A(U2  —  Ui) 

MX(soud)      +       nNH*R(9)    —       [M(NH*R)n> 


U 


(SOLID) 

U2 


M_(9)    +     X(9)        +       nNH2R(9)  -^[M(NH*R)n]  (gj      +      X    (g) 

Fig.  3.4.  Ammine  formation  as  represented  by  a  thermochemical  cycle. 

Ui  =  lattice  energy  of  solid  "simple"  salt. 

U2  =  lattice  energy  of  solid  "complex"  salt. 

Q     =  heat  evolved  in  formation  of  solid  complex  from  solid  salt  and  gaseous  amine. 
E    =  energy  of  coordination  =  heat  evolved  in  reaction  between  gaseous  metal  ion 
and  caseous  nmine  to  cive  a  caseous  eomnlex  ion. 


and  gaseous  amine  to  give  a  gaseous  complex  ion 


*  If  stabilities  are  compared  in  solution,  solvation  energies  for  the  simple  cation, 
the  complex  cation,  and  the  ligand  replace  the  lattice  energy  terms  U\  and  Ui . 


ELECTROSTATIC  THEORY  OF  COORDINATIOh   COMPOX  NDS      L39 
Table  3.4,  Expansion  oi  nn   Cbtstal  Lattk  b  oj    \  Complex  B  llt  a.8  ras  Size 

01    THE    (  lOORDIN  \  PED   <  rBOl   l'    \s<  RE  LSES 


Ige  of  Cube 

Metal  iodide 

Complex 

of  Unit  Cell     \ 

Dista 

[Ni  NH,),]I, 

10.88 

1.71 

[Ni(MeNHi).JIi 

L2.03 

5.  Ill 

Co  l*H,).]I, 

10.91 

4.73 

[Co  MeNH«),]Ii 

12.05 

5.20 

become  of  greater  significance  than  the  small  differences  in  the  coordination 

energy,  A  A'.  Differences  in  coordination  energy,  A',  tor  the  series  methyl, 
ethyl,  propyl,  and  butyl  amine  arc  not  large  because  the  dipole  moment- 
and  polarizabilities  do  not  change  appreciably  throughout  the  series.  On 

the  other  hand,  appreciable  differences  are  observed  in  the  lattice  energy 
terms  for  the  series.  Going  from  ammonia  successively  to  methyl  amine, 
ethyl  amine,  propyl  amine,  and  butyl  amine  brings  about  a  progressive 
expansion  in  the  size  of  the  lattice.  The  larger  distance  between  the  complex 
cation  and  the  salt  anion  reduces  the  electrostatic  lattice  energy,  Us  .  Since 
c'i  is  the  same  as  long  as  only  a  single  simple  salt  is  being  considered  and 
since  differences  in  the  energy  of  coordination  are  not  particularly  large  for 
the  primary  amines,  the  differences  in  the  values  of  Q  and  thus  the  differ- 
ences in  stability  of  the  amine  complexes  can  be  attributed  largely  to  differ- 
ences in  the  lattice  energy  of  the  complex,  Us  •  As  the  size  of  the  R  group 
on  the  amine  increases,  the  lattice  energy,  Us ,  usually  decreases.  Since 
Q  =  E  +  Us  —  L\  ,  a  decrease  in  lattice  energy  will  bring  about  a  decrease 
in  Q  and  a  lesser  stability  of  the  solid  complex.  This  deduction  is  in  agree- 
ment with  the  observations  of  Hertel. 

The  expansion  of  the  complex  lattice  as  the  size  of  the  R-group  increases 
is  indicated  by  x-ray  data  on  hexammine-nickel(II)  iodide  and  hexammine- 
cobalt(II)  iodide  and  the  corresponding  methyl  amine  complexes50.  All 
crystallize  in  the  fluorite  type  lattice.  The  length  of  the  unit  cell,  and  the 
metal-halogen  distances  are  as  indicated  in  Table  3.4. 

The  Influence  of  Anions  on  the  Stability  of  Solid  Complex  Com- 
pounds 

The  preceding  discussion  suggests  that  any  factor  which  might  influence 
the  lattice  energy  of  the  simple  salt  or  of  the  complex  might  influence  the 
stability  of  the  entire  complex  compound.  1  )ata  of  Ephraim,  Biltz,  and  their 
co-workers  on  anion  effects  in  complexes  provide  adequate  support  for 
such  a  conclusion.  Biltz  and  Messerknecht65  measured  the  heat  evolved  in 
the  formation  of  a  number  of  ammines  of  zinc  chloride,  zinc  bromide,  and 
zinc  iodide  (Fig.  3.5).  Similar  data*  showing  the  heal  evolved  in  the  forma- 

65.  Biltz  and  Mes.scrkncchT ;  /    anc       a  .  129,  ltil     1923). 

66.  Biltz  and  Hansen,  Z  <m<>nj.  cdlgem.  Chem.,  127,  1  (1923). 


140  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

30  r 


UJ 

Z 

I* 

<  o 

o3 

_l  * 
O    • 

CO    Q 

*■   Q 
< 

o    « 

!«i 

£"• 

o° 

Lu    UJ 

SB 


28 


26  - 


24 


22 


20 


18 


16 


ZnlNHjXg 


ZnCI2  ZnBr2  Znl2 

Fig.  3.5.  Heats  of  formation  of  zinc  ammine  halides 

tion  of  ammines  of  lithium  chloride,  lithium  bromide,  and  lithium  iodide 
are  shown  in  Fig.  3.6.  If  one  considers  a  simple  salt  such  as  zinc  chloride, 
the  energy  of  coordination  per  ammonia  molecule  falls  sharply  as  the  number 
of  ammonia  molecules  increases.  Such  behavior  is  in  agreement  with  quali- 
tative predictions  based  on  electrostatics.  In  this  case,  the  only  variables 
considered  are  the  energy  of  coordination  and  the  lattice  energy  of  the  solid 
complex  crystal. 

If  one  considers  variations  in  any  given  set  of  ammines  such  as 
[Zn(NH8)4]Br2  and  [Zn(NH8)4]l2 ,  the  energy  of  coordination,  E,  will  be  the 
same  in  each  case  (e.g.,  Zn++((/)  +  4NH3(ff)  ->  [Zn(NH3)4]++(a)).  The  differ- 
ence between  lattice  energies  of  the  simple  salt  and  the  complex  salt  of 
each  halide  will  account  for  the  observed  differences. 

A  similar  t  reatmenl  is  useful  in  correlating  other  generalizations  on  anion 
effects  in  complex  ammines.  Ephraim67  found  that  the  nickel  salts  of  strong 


67    Ephraim,  Ber.t  46,  3103  (1913), 


ELECTROSTATIC  THEORY  OF  COORDINATION  COMPOUNDS      14L 

22 


20 


18 


16 


I    14 

z 


o  ^ 

£° 

x 


12 


10 


LiX-NH. 


LiX  -2NH. 


LiX-3NH3 
UX-4NH, 


-O    LiX-5NH3 


1 1 1 

LiCI  LiBr  Lil 

Fig.  3.6.  Heats  of  formation  of  lithium  ammine  halides 


acids  have  greater  affinity  for  ammonia  than  nickel  salts  of  weak  acids, 
affinity  being  almost  parallel  to  acid  strength.  Spacu  and  Voichescu68  found 
that  the  stability  of  the  solid  ammines  of  copper  salts  of  organic  acids  rims 
almost  parallel  to  the  strength  of  the  organic  parent  acid.  Shuttleworth69 
reports  similar  behavior  for  complexes  of  the  chromium  salts.  If  one  makes 
the  plausible  assumption41  that  those  anions  which  bind  the  proton  strongly 

*  The  correlation  between  the  binding  of  a  proton  and  the  binding  of  ;i  metal  ion 
has  received  considerable  experimental  support.  Calvin  and  Wilson70,  Bruehlman 
and  Verhoeck71,  and  others  have  noted  an  almost  linear  relationship  between  the 
ability  of  a  coordinating  group  to  hind  a  metal  ion  and  its  ability  to  bind  an  II+ 
ion.  Groups  of  comparable  type  must  he  considered. 

68.  Spacu  and  Voichescu,  Z  anorg.  <ili</<  m.  Ch  m.s  226,  27:5  d<)36). 

69.  Shuttleworth,  ./.  .w.   Leaihei    Trade*  Chem.,  30,  342    1946);  cf.  Chetn.  A 

41,   1572     10  17 

70.  Calvin  and  Wilson,  /.  Am.  Chem.  8oc.t  07,  2003  (1946 

71.  Bruehlman  and  Yerhoek,  ./ .  Nor..  70,  1  101    (1948 


142  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

will  also  bind  the  nickel,  copper,  or  chromium  ion  strongly,  one  can  draw  a 
parallel  between  low  acid  strength  of  the  parent  acid  and  high  lattice  energy 
for  the  simple  salt,  Ui .  Since  Q  =  E  +  U2  —  U\ ,  a  high  value  for  U\ , 
the  lattice  energy  of  the  simple  metal  salt  of  the  organic  acid,  will  reduce  Q 
and  lower  the  stability  of  the  ammine. 

Quantitative  Treatment  of  the  Thermochemical  Cycle 

Biltz  and  Grimm72  were  the  first  to  recognize  and  outline  the  importance 
of  the  various  energy  terms  in  complex  formation.  They  attempted  a  quanti- 
tative  treatment  of  the  factors  involved.  From  the  expression  E  =  Q  + 
(Ui  —  c72)  they  estimated  E  for  the  coordination  of  six  ammonia  molecules 
to  calcium  ion.  Q  was  measured  directly  and  (Ui  —  t/2)  was  estimated  from 
electrostatics.  Using  an  E  value  of  30  kcal  per  mole  of  ammonia  as  the 
average  energy  for  the  coordination  of  each  of  six  ammonia  molecules  around 
a  calcium  ion,  they  predicted  that  the  reaction  between  calcium  fluoride 
and  gaseous  ammonia  would  be  endothermic  because  of  the  very  large 
amount  of  energy  required  to  expand  the  calcium  fluoride  lattice.  Subse- 
quent attempts  by  Biltz  and  Rahlfs73  to  prepare  ammoniates  of  the  alkali 
and  alkaline  earth  fluorides  were  unsuccessful,  thus  offering  experimental 
support  for  the  earlier  theoretical  predictions.  Fluoride  salts  of  more 
strongly  polarizing  metal  cations  such  as  silver  (I),  copper  (II),  man- 
ganese(II),  iron (II),  cobalt(II),  and  nickel(II)  add  ammonia  to  form 
complexes73.  This  fact  may  be  correlated  with  the  much  larger  amount  of 
energy  released  in  coordinating  the  polarizable  ammonia  molecules  around 
the  strongly  polarizing  cation.  The  large  coordination  energy  overcomes 
the  high  fluoride  lattice  energy. 

One  of  the  most  thorough  and  generally  satisfactory  electrostatic  treat- 
ments of  the  coordination  process  was  carried  out  by  Garrick28b.  He  evalu- 
ated the  energy  of  coordination,  E,  by  two  more  or  less  independent  meth- 
ods. First,  the  coordination  energ}^  was  estimated  from  a  thermochemical 
cycle  by  the  methods  of  Biltz  and  Grimm72  and  of  Grimm  and  Herzfeld74. 
Then  the  coordination  energy,  E,  was  estimated  directly  from  the  electro- 
static interaction  between  the  cation  and  the  coordinated  dipoles  in  a 
manner  similar  to  that  of  Van  Arkel  and  de  Boer27.  Three  ammines  were 
considered:  [Zn(NH,)J++  [Fe(NH3)6]++,  and  [Mn(NH3)6]++.  The  results 
appear  in  Table  3.5. 

The  agreement  between  values  for  E  obtained  by  the  two  methods  is 
fairly  good,  and  suggests  that  for  the  so-called  ionic  or  normal  ammines  the 
pure  electrostatic  model  (E}i ,  Table  3.5)  may  be  fairly  reliable.  It  is  signifi- 


7_\  Biltz  mid  Grimm,  Z.  anorg.  allgem.  Chem.}  145,  63  (1925). 
73.  Biltz  :.nd  Rahlfs,  Z.  anorg.  allgem.  Chan.,  166,  351  (1927) 
7  1    Grimm  and  Berzfeld,  Z.  Phys..  19,  141  (1923). 


ELECTROSTATIC  THEORY  OF  COORD/ \  AT/oX  COMPOUNDS      143 


Table  3.5.  Enbbgt  of  Coordination 


Complex  Compounds 


[Zn  NH,),]C1, 
[Fe  NH,)e]Cl, 

;.\I:    XHj)6]C12 


Lattice 

Simple  Salt 
kcal  mole 


634 
615 


Lattice  1 

of  Complex 

Salt    /      . 

mole 

Heat  of 

Reaction 

Salt    -     \1I 

mole 

:V27 
327 
323 

88 
82 

Ba 

hneiv 

ordination 

from  Thermo- 

chem.  Cycle 

Real  mole 


i:;s 
395 
374 


1  H 

Coordination 
from  Eli  i  tro 

Real  'mole 


139 

423 
391 


Table  3.6.  The  Coordination  Number  as  Determined  by  the  Radii  -  Ratio 

Radius  Metal  Ion 


Number  of 

Coordinated 

Spheres 

3 

4 


Radius  Coordination  Group 
(Radius  Ratio) 

.1548  to  .2164 
.2165  to  .4142 
.4143  to  .5912 
.4143  to  .5912 
.4143  to  .5912 
.6455  to  .7323 


Spatial  Distribution  of  Coordinated  Ions 

Equilateral  triangle 
Tetrahedron 
Plane 

Trigonal  bipyramid 
Octahedron 
Cube  or  regular 
square  prism 


cant,  however,  that  the  two  methods  give  values  differing  by  as  much  as 
28  kcal.  This  difference  emphasizes  the  difficulty  in  quantitative  correlation 
of  chemical  properties  and  electrostatic  energy  terms,  since  even  one  or 
two  kcal.  may  be  of  great  chemical  significance. 

The   Coordination   Number   in   Relation   to   the  Thermochemical 
Cycle 

Straubel75  and  Hiittig76  considered  the  problem  of  predicting  the  coordina- 
tion number*  from  the  geometry  of  the  packing  of  rigid  spherical  ions  or 
molecules  around  a  central  spherical  ion.  Since  the  relative  sizes  of  the  ions 
will  be  of  major  importance  in  determining  the  packing,  it  is  convenient  to 
consider  the  radius  ratio  as  a  differentiating  parameter77.  The  coordination 
numbers  and  the  configurations  are  summarized  in  Table  3.6. 

In  many  cases  the  radius  ratio  is  not  an  adequate  criterion  for  determin- 
ing the  coordination  number  of  complex  compounds.  For  example,   the 

*  Bidgwick  pointed  out  in  1928  that  the  maximum  coordination  number  for  ele- 
ments of  the  first  short  period  is  usually  4;  for  elements  of  the  second  short  period 
and  first  long  period  it  is  usually  6;  while  the  maximum  coordination  number  for  the 
remaining  elements  is  usually  8. 

75.  Straubel,  Z.  anorg.  oUp  ,  142,  133    L926). 

76.  Efittig,  /.  anorg.  dUgem.  Ckern.,  142,  135  (1 

77.  Rice,  ''Electronic  81  met  ure  and  Chemical  Binding," p. 317,  Ne*  York.  M<-<  ira*  - 

Hill  Book  Co.,  1940. 


ic  Radius,  A. 

Coordination  No. 

0.93A. 
0.69A. 

4 
6 

144  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

smaller  ions  of  higher  valence  state  almost  invariably  have  a  greater  coor- 
dination number  than  the  larger  ions  of  lower  valence  state.  Penney  and 
Anderson78  illustrated  this  point  with  the  complexes  of  platinum. 

Ion 

Pt++ 
Pt4+ 

An  alternative  method  for  evaluating  the  number  of  coordinated  groups 
was  suggested  by  Kossel3d.  It  is  possible,  at  least  in  principle,  to  estimate 
from  electrostatics  and  polarization  the  amount  of  energy  released  by  the 
grouping  of  negative  ions  or  dipolar  molecules  around  a  central  positive 
ion.  It  may  then  be  assumed  that  the  arrangement  of  coordinated  groups 
which  releases  the  most  energy  will  give  the  most  stable  coordination  com- 
pound. If  two  arrangements  release  about  the  same  amount  of  energy,  two 
forms  may  exist  in  equilibrium. 

These  ideas  were  used  by  a  number  of  investigators280  • 79  to  calculate  the 
most  probable  formulas  for  many  compounds.  The  early  investigators 
assumed  rigid  spherical  ions  and  made  no  provision  for  lattice  energy  or 
hydration  energy  terms;  however,  Garrick28a' 80  refined  the  methods  by 
considering  polarization  of  ions  and  by  using  a  thermochemical  cycle.  Using 
the  refined  technique,  he  calculated  the  coordination  number  to  be  expected 
when  water  or  ammonia  is  coordinated  around  a  free  gaseous  metal  ion. 
His  calculated  values*  are  in  fair  agreement  with  experimental  results.  A 
more  complete  treatment  involving  a  thermochemical  solution  cycle  was 
used  to  calculate  coordination  numbers  and  formulas  of  metal  chloride  and 
fluoride  complexes  in  solution.  In  general  his  theoretical  results  were  in 
striking  agreement  with  experiment,  indicating  stable  ions  such  as  [A1F6]- 
and  [BF4]~.  Similar  calculations  were  carried  out  for  the  solid  complexes. 
In  view  of  the  uncertainties  of  the  calculations,  the  agreement  must  be 
regarded  as  rather  fortuitous. 

*  Values  obtained  by  Garrick  for  coordination  of  water  molecules  are: 
Coordination  No.  4:  Li+,  Be++ 
Coordination  No.  6:  Na+,  K+,  Mg++ 
Coordination  No.  8:  Cs+,  Ba++ 
For  ammonia: 

Coordination  No.  4:  Mg++ 
Coordination  No.  6:  Na+,  K+,  Ca++,  Sr++ 
Coordination  No.  8:  Rb+,  Cs+,  Ba++ 
No  coordination  number  of  2  was  reported. 

78.  Penney  and  Anderson,  Trans.  Faradmj  Soc,  33,  1364  (1937). 
7!).  Remy  and  Laves,  Ber.,  66,  401,  571   (1933);  Remy  and  Pellens,  Ber.,  61,  862 
(1928);  Remy  and  Rothc,  Ber.,  58,  1565  (1925);  Remy  and  Busch,  Ber.,  66,  961 
(1933). 
Garrick,  Phil.  Mag.,  [7]  14,  914  (1932). 


ELECTROSTATIC  THEORY  OF  COORDINATION  COMPOUNDS      145 

Ablov"1  attempted  to  relate  the  coordination  number  to  the  nature  of  the 
anion  in  the  simple  salt.  When  the  coordination  of  pyridine  with  nickel  and 
copper  salts  of  organic  acids  was  investigated,  he  found  an  increase  in  co- 
ordination number  as  the  acid  strength  of  the  parent  organic  acid  was  in- 
creased; these  facts  are  readily  understandable  in  view  of  the  close  rela- 
tionship between  the  coordination  number  and  the  energy  of  the  complete 
t  hermochemical  cycle.  Changing  the  anion  of  the  nickel  salt  alters  the  hit  t  ice 
energy  of  both  the  simple  and  complex  salts  and  thus  brings  about  a  change 
in  the  total  energy  released  in  the  formation  process.  Also,  the  weak  acid 
radical  may  fill  a  position  in  the  coordination  sphere. 

Thermochemical  considerations  also  suggest  that  the  nature  of  the  co- 
ordinated amine  may  be  important  and  that  different  results  may  be  found 
if  different  amines  are  used.  In  a  separate  stud}',  Ablov82  considered  com- 
plexes between  nickel  trichloroacetate  and  a  series  of  organic  amines,  mostly 
substituted  anilines.  He  observed  a  rather  indistinct  relationship  between 
the  dipole  moment  of  the  amine  and  the  coordination  number  of  the  nickel. 
A  relatively  large  increase  in  dipole  moment  frequently  increased  the  number 
of  amine  molecules  bound  to  the  nickel.  Again,  such  factors  are  intelligible 
if  the  entire  thermochemical  cycle  is  considered,  but  consideration  of  a 
single  factor  such  as  the  dipole  moment  is  inadequate. 

Much  of  the  data  in  the  literature  on  coordination  number,  such  as  that 
of  Ablov  and  of  Remy79' 81>  82,  assumes  that  the  coordination  number  can  be 
obtained  from  the  empirical  formula  of  the  complex  compound.  Such  evi- 
dence, however,  is  subject  to  the  criticism  that  water  molecules  may  co- 
ordinate in  solution  to  give  a  coordination  number  of  six  for  ions  such  as 
[FeF5]=  and  that  comers  of  the  individual  octahedra  may  be  shared  in  the 
solid  state  to  give  coordination  numbers  which  are  larger  than  those  indi- 
cated by  the  empirical  formula.  A  coordination  number  which  is  smaller 
than  that  indicated  by  the  empirical  formula  may  also  exist  if  extra  mole- 
cules of  the  coordinated  ligand  can  be  packed  into  the  lattice  interstices. 
Ephraim  and  his  co-workers83  and  Clark84  observed  that,  when  cations  such 
as  N1++  Co4-4",  and  Fe4^,  which  normally  show  a  coordination  number  of 
six.  are  associated  with  very  large  anions,  such  as  the  benzoate  ion  or 
[Co(NH3)2(X02)4]_,  eight  or  ten  ammonia  molecules  may  appear  to  be  co- 
ordinated to  the  central  metal  cation  at  room  temperature.  They  suggested 
that  the  last  two  or  four  molecules  of  ammonia  are  probably  trapped  in  the 
lattice  interstices,  since  they  differ  appreciably  from  the  first  six  ammonias 

81.  Ablov,  Bull.  soc.  chim.,  [51  1,  731,  1489  (1934). 

82.  Ablov,  Bull.  soc.  chim.,  [5]  2,  1724  (1935);  3,  1673  (1936). 

Ephraim  and  Moser,  Ber.,  53,  548  (1920);  Ephraim  and  Rosenberg,  Ber.,  51,  644 
(1918). 
84.  Clark,  Quick,  and  Harkins,  ./.  -4m.  Chem.  Soc,  43.  2496,  2488  (1920). 


146  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

in  their  heats  of  coordination  and  in  their  effect  on  complex  color.  Somewhat 
more  recently,  Lamb  and  Mysels85  have  reported  that  the  water  in 
[Co(NH3)5C03]N03-H20  has  no  structural  significance  but  may  be  con- 
sidered as  lattice  water. 

It  may  be  concluded  that  the  coordination  number  has  been  successfully 
estimated  by  the  electrostatic  treatment  for  the  simplest  cases  involving 
normal  or  ionic  type  complexes.  For  the  more  polarized  covalent  or  pene- 
1  rat  ion*  type  compounds  the  electrostatic  treatment  is  completely  inade- 
quate in  its  present  state  of  development.  In  cases  where  the  electrostatic- 
treatment  can  be  successfully  applied,  all  the  terms  in  the  thermochemical 
cycle  must  be  considered.  In  general,  the  interactions  of  such  factors  as 
appear  in  these  cycles  are  too  numerous  and  involved  to  permit  close  general 
correlation  with  any  single  molecular  or  ionic  property,  such  as  dipole 
moment,  ion  charge,  or  ionic  potential. f 

The  Application  of  the  Complete  Ionic  Model  to  the  Properties 
and  Structures  of  Selected  Complex  Compounds 

The  Trans  Effect 

One  of  the  useful  concepts  suggested  by  Werner  in  the  development  of 
his  theory  was  the  idea  of  "trans  elimination"  in  substitution  reactions.  In 
brief,  the  rule  of  "trans  elimination"  suggests  that  the  "reactivity"  of  a 
given  group,  A,  in  a  coordination  compound  is  dependent,  in  large  measure, 
upon  the  nature  of  the  group  coordinated  in  the  position  trans  to  group  A. 
(By  "reactivity"  we  mean  the  ease  with  which  the  group  A  may  be  replaced 
in  the  coordination  sphere  by  other  donor  molecules.)  In  general,  acid  anions 
and  neutral  groups  which  are  easily  polarized  show  a  much  greater  trans 
effect  than  groups  such  as  water  or  ammonia.  Thus,  a  group  which  is  trans 
to  chloride  or  bromide  is  much  more  labile  than  a  group  trans  to  a  neutral 
molecule  such  as  water.  The  idea  of  trans  elimination  has  been  applied  to 
compounds  of  platinum,  cobalt,  chromium,  osmium,  palladium,  rhodium, 
and  iridium.  The  principle  has  been  widely  used  and  developed  by  Tscher- 
niaev87,  Grinberg88  and  their  co-workers.  A  comprehensive  review  on  the 
trans  effect  has  been  published  by  Quagliano  and  Schubert89. 

*  For  description  of  penetration  complexes,  see  page  151. 

t  A  most  interesting  treatment  of  the  heats  of  formation  in  oxyacid  salts  in  terms 
of  an  ionic  model  and  lattice  energies  has  been  given  by  Ramberg86. 

85.  Lamb  and  Mysels, ./.  Am.  Chem.  Soc.,  67,  468  (1945). 

86.  Ramberg,  J.  Chem.  Phys.,  20,  1532  (1952). 

87.  Tscherniaev,  Ann.  Inst.  Platine,  U.S.S.R.,  4,  261  (1926);  5,  118,  134  (1927). 

88.  Grinberg,  Shulman,  and  Khorunzhenkov,  Ann.  Inst.  Platine,  U.S.S.R.,  12,  69, 

119  (1935);  cf.  Chem.  Abs.,  29,  3253  (1935);  Ann.  Inst.  Platine,  U.S.S.R.,  11, 
17  (1933);  Ann.  Inst.  Platine,  U.S.S.R.,  10,  58  (1932);  cf.  Chem.  Abs.,  28,  1447 
(1934). 

89.  Quagliano  and  Schubert,  Chem.  Revs.,  50,  201  (1952). 


ELECTROSTATIC  THEORY  OF  COORDINATION  COMPOl  NDS      147 

Grinberg91  suggested  the  following  explanation  of  the  trans  efifeci  based 
on  the  ideas  of  electrostatics  and  polarization,  [f  a  central  metal  ton  is 
surrounded  by  four  identical  groups,  the  cation  is  in  a  symmetrical  field  and 
all  dipoles  induced  in  the  central  ion  arc  compensated  by  one  another.  Now, 
if  one  of  the  coordinated  groups  is  replaced  by  a  relatively  more  negative 
or  more  easily  polarized  group  ("Y"  in  Fig.  3.7),  the  symmetry  of  the  field 
around  the  central  ion  is  destroyed  and  a  noncompensated  dipole  is  induced 
in  the  central  metal  ion.  The  group  X2  which  is  adjacent  to  the  negative  end 
(A  the  induced  dipole  is  labilized,  and  trans  elimination  can  easily  occur. 

On  the  basis  of  this  explanation,  the  trans  effect  will  be  exhibited  by  any 
group  which  possesses  mobile  electrons  that  can  be  dislocated  in  the  direc- 


NEGATIVELY  CHARGED 
OR  RELATIVELY  EASILY 
POLARIZED   GROUP  "Y" 


LABILIZED    BY     TRANS 
EFFECT    OF    "Y" 


SYMMETRICAL  UNSYMMETRICAL 

Fig.  3.7.  The  trans  effect  according:  to  the  electrostatic  concept 

tion  of  the  central  ion91.  Tronev  and  Chulkov92  report  the  decreasing  efficacy 
of  a  substituent  in  labilizing  the  group  trans  to  it,  as: 

CX-,  C,H,  ,  NOr,  I"  Bi-,  C1-,  XH3  ,  OH",  H,0 
Decreasing  Trans  Influence 

Chatt  and  Williams93  give  the  order:  CX~  >   C2H4   >   CO   >   NO»~  > 
-      Ml,  ,  >  RS~PR3  ~  r  >  Br"  >  CT  >  F"  ~  NH3  >  OBT  >  II  <  ». 
This  order  is  roughly  that  expected  on  the  basis  of  the  above  treatment.  Sub- 
stituted phosphines  have  been  reported  to  have  a  high  trans  influeni 
as  would  be  predicted. 

The  above  mechanism  for  the  trans  effect  suggests  that  the  effect  will 

90.  Grinberg  and  Ryabchikov,  Acta  Physicockim.  U.R.S.S.,  3,  555,  573  (1935);  cf. 

.  30,  ln:t    ]<)36). 

91.  Grinberg,  Bull.  acad.  get.,  U.R.SJ5.,  Clasu   sci.,  ckim.,  350    1943);  cf.  Chem. 

39.  -"J     L945  . 
2    Tronev  and  Chulkov.  Doklady  Akad.  Nauk.  S.S.S.R.,  63,  545    1948  ;  cf.  Chem. 
Abs.,  48,2854    1949). 

93.  Chatt  ami  Williams,  ./.  Chem.  Soc.,  1951,  3061. 

94.  Grinberg,  Razumova,  and  Troitskaya,  Hull.  acad.  sci.,  (hiss,  sci.,  ckim.,  3. 

(1946);  cf.  ..,43,417.'    1949);  Grinberg  and  Razumova,  Zkur.  Priklad. 

Khim.,  27,  105  (1954)  ;cf.  (  ,  48,  6308  (1954). 


148  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

be  promoted  by: 

(1)  A  central  cation  of  high  field  strength  which  is  itself  easily  deformed; 
both  Pd"1^  and  Pt++  meet  this  specification.  Chatt  and  Hart95  find  some 
evidence  to  indicate  that  palladium(II)  compounds  are  less  influenced  by 
trans  directing  groups  than  the  corresponding  platinum (II)  compounds. 

(2)  A  coordinated  group  which  can  release  electrons  toward  the  central 
cation;  thus  anions  and  easily  polarized  groups  would  be  more  effective 
than  neutral  molecules  of  low  polarizability  such  as  H20. 

The  two  most  serious*  objections  raised  to  the  treatment  of  Grinberg 
are:  (a)  that  the  diamagnetism  of  the  platinum(II)  compounds  indicates 
that  the  platinum  cannot  be  present  as  the  dipositive  ion  since  platinum (II) 
should  have  two  unpaired  electrons  and,  (b)  a  high  trans  effect  has  been 
attributed  to  PF3  by  Chatt  and  Williams93,  though  they  assume  that  the 
polarizability  of  the  attached  phosphorus  would  be  so  reduced  by  attached 
fluorine  atoms  that  its  trans  effect  would  be  reduced  rather  strongly. 

The  first  of  these  objections  has  been  answered  in  the  section  on  mag- 
netism where  it  has  been  shown  that  the  diamagnetism  in  the  platinum  (I I) 
is  a  direct  result  of  the  Stark  splitting  of  normally  degenerate  d  levels  in  the 
crystal  field.  This  cannot  be  considered  as  a  valid  objection.  The  second 
point  raised  by  Chatt93  cannot  be  accepted  as  unequivocal  and  must  be 
regarded  as  an  open  question  for  the  following  reasons: 

(1)  The  assumption  that  the  attached  fluorines  on  PF3  reduce  the  po- 
larizability of  the  free  electron  pair  on  phosphorus  to  a  point  where  it  would 
not  be  expected  to  be  trans  directing  has  no  direct  experimental  support. 

(2)  A  strong  trans  effect  for  PF3  has  never  been  established.  Coordination 
compounds  of  PF3  have  been  prepared  such  as  PtCl2(PF3)2  which  are  analo- 
gous to  the  corresponding  carbonyl  halidesof  platinum.  Hel'man  attributed 
a  strong  trans  effect  to  CO  since  it  directs  pyridine  trans  when  the  pyridine 
replaces  a  chloride  ion  in  [COPtCl3]~  and  since  it  is  analogous  to  C2H4  in 
many  of  its  coordination  compounds;  C2H4  is  reported  to  have  a  high  trans 
effect  (p.  490). 

On  the  other  hand,  the  only  direct  evidence  available  on  the  reactions  of 
PF3  which  is  comparable  in  nature  to  that  used  in  establishing  the  trans 
series,  would  suggest  that  PF3  is  not  highly  trans  directing.  The  complex 
solid  (a)  reacts  with  PF3  to  give  the  cis  isomer  (b)  as  indicated  by  dipole 


CI 

CI 

PF8 

PF3 

Cl 

\ 

/     \ 

/ 

\ 

/ 

Pt          Pt 

+ 

2PF3   - 

->  2               Pt 

/ 

\    / 

\ 

/ 

\ 

PF3 

CI 
(a) 

CI 

PF3 

cis- 

CI 
(b) 

*  Other  objections  cited89  are  trivial. 
95.  Chatt  and  Hart,  J.  Chem.  Soc,  1953,  2367. 


ELECTROSTATIC  THEORY  OF  COORDINATION  COMPOUNDS      149 

measurements93'  w,  yet  on  the  basis  of  a  high  trans  directing  influence  for 
PF3  a  trans  isomer  was  predicted  for  the  compound  by  Quagliano  and 
Schubert".  Chatl  and  Wilkins98  also  suggested  a  trans  structure  for  the 

product  obtained  by  the  analogous  reaction  between  the  slrongly  trans 
directing  CVH4  and  its  comparable  dimeric  complex.  Despite  such  predic- 
tions, the  PF3  product  is  cis.  It  has  also  been  shown  that  CO  gives  the  cifl 
product,  contrary  to  expectations  for  strong  trans  directing  properties.  Even 

the  trans  case  for  CO  is  established  on  very  meager  evidence  as  compared 
to  that  used  by  Werner  in  first  elucidating  the  concept. 

(3)  Further,  the  nature  and  operational  meaning  of  the  trans  effect  are 
very  uncertain.  No  definite  quantitative  method,  free  from  objections,  can 
be  applied  to  place  groups  in  the  series.  If  the  effect  is  considered  to  be  one 
of  thermodynamics  involving  bond  stabilities,  difficulties  are  legion.  Chatt" 
tried  to  evaluate  the  relative  coordinating  affinity  of  a  series  of  tertiary 
alkyls  in  Group  V.  He  reported  the  order:  PR3  >  AsR3  >  SbR3  >  NR3  > 
BiR3  .  Attempts  to  place  ethylene  in  this  series  led  to  conflicting  positions 
depending  upon  the  experimental  criterion  selected,  indicating  that  the 
relative  coordinating  ability  is  affected  by  many  other  variables,  such  as  the 
groups  already  attached  to  the  metal.  Chatt  and  Wilkins100  estimated  certain 
of  the  thermodynamic  constants  for  the  metal-tertiary  phosphine  linkage 
and  concluded:  "This  study  also  serves  to  emphasize  the  importance  of  the 
entropy  term  in  determining  the  position  of  equilibrium  in  reactions  in- 
volving the  formation  or  destruction  of  highly  polar  molecules,  and  how 
completely  erroneous  conclusions  regarding  relative  stability  can  be  arrived 
at  by  consideration  of  only  equilibrium  positions  or  decomposition  tempera- 
tures in  coordination  chemistry."  Much  of  the  trans  effect  series  is  based 
on  relative  yields  obtained  under  different  sets  of  conditions. 

On  the  basis  that  such  yields  are  determined  by  relative  rates  of  reaction 
rather  than  complex  stability,  mechanisms  have  been  suggested  for  various 
processes  which  are  designed  to  show  that  even  when  using  the  trans  effect, 
a  result  directly  contrary  to  that  normally  expected  can  be  obtained101.  In 
view  of  this  situation  the  trans  effect  must  be  considered,  at  the  present 
time,  as  only  a  broad  qualitative  generalization  covering  a  very  complex 
proc» 

Application  of  the  Polarization  Theory  to  a  Number  of  Unusual 
Compound^ 

The  strength  of  any  theory  lies  in  its  ability  to  adequately  describe  the 
unusual  as  well  as  the  commonplace.  In  the  following  section,  types  of  com- 

97.  button  and  Puny,  ./.  Am.  Chem.  Sue,  76,   1271    [1064). 
98. -Chatt  and  Wilkins,  ./.  Chem.  Sor.,  1952,  2822. 

chatt,  ./.  Chem.  8oe.,  1951,  652. 
100.  Chatt  and  Wilkins,  J.  Chem.  Soc,  1952,  276. 


150  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

plexes  which  differ  from  the  classical  Werner  coordination  compounds  are 
discussed.  The  relationship  of  each  complex  to  the  polarization  theory  is 
noted.  The  customary  successes  and  failures  are  observed. 

"Super  Complexes"  If  one  considers  the  positively  or  negatively  charged 
complex  ion  as  a  unit,  it  becomes  apparent  that  an  electrostatic  field  exists 
around  the  complex  ion  just  as  a  field  exists  around  a  simple  ion.  Because 
the  complex  is  in  general  much  larger  than  the  simple  ion102,  the  attraction 
of  the  complex  for  the  solvent  or  for  ions  of  opposite  charge  in  solution  is 
significantly  less  than  that  exerted  by  the  simple  ion.  Still,  the  fact  that  a 
complex  ion  may  enter  into  an  ionic  crystal  as  a  structural  unit  offers  con- 
clusive proof  that  the  residual  field  is  not  negligible.  The  existence  of  this 
field  would  lead  one  to  suspect  that  additional  ions  or  dipolar  molecules 
might  be  attracted  to  the  complex  ion  to  produce  a  second,  a  third,  or  per- 
haps even  a  fourth  coordination  sphere  in  solution.  Obviously,  groups  held 
in  these  outer  spheres  will  be  held  less  tightly  as  their  distance  from  the 
central  ion  increases.  Such  super  complexes  have  been  described  by  Brint- 
zinger103.  Definite  formulas  such  as  [Fe(H20)i8]+++  and  [Co(NH3)6(S04)4]5- 
have  been  reported  from  diffusion  studies.  Such  formulations  are  completely 
arbitrary  and  of  little  significance,  since  the  formula  is  dependent  upon  the 
nature  and  the  reliability  of  the  method  used  to  define  the  compound.* 
Laitinen,  Bailar,  Holtzclaw,  and  Quagliano104  obtained  polarographic  evi- 
dence for  such  complexes  and  suggested  that  the  super  complexes  formed 
between  the  hexamminecobalt(III)  ion  and  acetate  or  sulfate  ion  may  be 
strong  enough  to  cause  a  measurable  shift  of  the  reduction  potential  for 
the  hexammine  ion  and  a  lowering  of  the  polarographic  diffusion  current. 
The  existence  of  such  super  complexes  can  best  be  considered  as  an  electro- 
static phenomenon,  probably  more  comparable  to  the  Debye-Hiickel  ionic 
atmosphere  than  to  true  coordination  compounds. 

Ammoniates  of  the  Alkaline  Earth  Metals.  An  interesting  series  of  com- 
pounds is  the  alkaline  earth  metal  ammines:  Ca(NH3)6 ,  Sr(NH3)6 ,  and 
Ba(NH3)6  •  These  compounds  are  formed  by  simple  addition  of  ammonia 
to  the  solid  metal.  The  stability  decreases  from  calcium  to  barium.  Meas- 
urements by  Biltz107  indicate  that  the  metal-ammonia  complex  is  almost 

*  Brintzinger's  methods  have  been  criticized  by  a  number  of  investigators.  See 
particularly  J.  Bjerrum'05. 

loi .  Jonassen  and  Cull, ./.  .1///.  Chem.  Soc,  73,  274  (1951);  Jonassen,  Sistrunk,  Oliver, 
and  Helfrich,  ./.  .1///.  Chem.  Soc,  75,  5216  (1953). 

102.  B0dtker-Naess  and  Hassel,  Z.  anorg.  Chem.,  211,  21  (1933):  Z.  phys.  Chem.,  22B, 

171   (1933). 

103.  Brintzinger and OsBwald,Z.  anorg.  allgem.  Chem. ,223, 263  (1935); 225, 221  (1935). 
101.  Laitinen,  Bailar,  Holtzclaw,  and  Quagliano, ./.  .1///.  Chem.  Soc,  70,  2999  (1948). 
105.  Reference  7a.,  p.  77. 

107.  Biltz,  Z    Elektrochem.,  26,  374  (1920);  Z.  anorg.  Chem.,  114,  241  (1920). 


ELECTROSTATIC  THEORY  OF  COORDINATION  COMPOUNDS      151 

as  stable  as  the  ion-ammonia  complex,  ICmXIIJe]^  or  [Ba(NH3)6]++. 
Since  in  the  metal-ammonia  complex  there  lb  do  charged  ion  to  attract  the 
dipoles  of  the  ammonia,  any  explanation  based  on  electrostatics  musl 
assume  an  arbitrary  reassignment  of  charge  among  components  of  the 

molecule,  or  it  must  assume  that  dipoles  are  induced  in  the  central  metal 
atom  by  the  dipoles  of  the  ammonia. 

Watt108  and  his  students  have  reported  the  analogous  Pt(NH3)4  and 
Ir(XH3)5 .  Explanations  of  why  dipoles  or  multipoles  would  arise  in  such 
compounds  are  inadequate  at  present. 

M(  la!  ( 'arbonyls.  The  interesting  coordination  of  compounds  formed  by  the 
reaction  between  carbon  monoxide  and  many  metals,  particularly  those  of 
Group  VIII,  are  known  as  the  metal  carbonyls  (Chapter  16).  These  com- 
pounds, of  which  [Xi(CO)4]  and  [Fe(CO)5]  are  typical,  are  particularly  diffi- 
cult to  fit  into  the  electrostatic  polarization  scheme  since  the  central  metal 
atom  apparently  bears  no  charge  and  the  carbon  monoxide  has  such  a  low 
dipole  moment  that  bonding  based  on  dipole-induced  dipole  interaction  is 
completely  unrealistic.  The  effective  atomic  number  concept  of  Sidgwick 
(page  159)  has  been  particularly  fruitful  in  a  consideration  of  the  formulas 
and  chemistry  of  these  substances. 

The  interesting  [Xi(PF3)4]  and  [Ni(PCl3)4]  complexes109,  the  compound 
[XiH(CO)3]2110,  as  well  as  Ni(N4S4)m  and  the  metal  cyclopentadienes 
("Chapter  15)  provide  other  examples  of  the  same  type  of  substance. 

Types  of  Complexes:  Normal  (or  Ionic)  and  Penetration 
(or  Covalent)  Complexes 

The  ammoniates  of  the  alkali  halides  and  of  compounds  such  as 
[Fe(X"H3)6]Cl2  can  be  rather  accurately  described  with  a  polarized  electro- 
static model.  On  the  other  hand,  the  carbonyls  and  alkaline  earth  and 
platinum  metal*  ammoniates  are  not  particularly  wrell  adapted  to  treatment 
by  electrostatics.  With  many  compounds,  such  as  those  just  mentioned, 
the  electron-pair  bond  or  molecular  orbital  theory  is  more  useful  in  correlat- 
ing experimental  facts.  In  between  the  typical  electrostatic  or  ionic  alkali 
halides  on  one  hand  and  the  strongly  covalent  metal  carbonyls  on  the  other, 
lie  most  of  the  common  coordination  compounds.  Biltz112,  recognizing  this 

*  One   must    differentiate    metal-ammoniatos,    Ca(NH«)«  ,    from  ion-ammoniates 
MI  )J++. 
lev  Watt,  Walling,  and  May  field,  ./.  .1///.  Chem.  Soc,  75,  6175    196 
109.  Irvine  and  Wilkinson,  Science,  113,  7l_'    1951);  Wilkinson,  ./.  .1///.  Chem.  Soc, 

73,  5501     1951). 
no.  Brehrena  and  Lohofer,  /.  Naturforsch,  8b,  091   (1951 

111.  Goehring  and  Debo,  Z.  anorg.  cUlgem.  Chem.,  273,  319  (1953), 

112.  Biltz,  Z.  anorg.  Chem.,  164,  245  (1027). 


152  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

fact,  attempted  to  set  up  a  method  of  classifying  compounds  based  on  four 
experimental  criteria.  The  properties  selected  and  applied  to  the  cobalt 
ammines  were:  (1)  thermochemical  and  chemical  data  indicating  the  sta- 
bility of  the  complex  unit,  (2)  the  molecular  volume  of  the  coordinated 
groups,  (3)  molecular  distances  as  obtained  from  x-ray  data,  (4)  magnetic 
susceptibility  measurements.  On  the  basis  of  the  above  factors  it  is  possible 
to  divide  coordination  compounds  roughly  into  two  general  types,  though, 
as  Taube62  and  Orgel43  have  shown,  the  classification  is  not  unequivocal. 

The  first  group  is  characterized  by  a  comparatively  weak  bond  between 
the  central  group  and  the  coordinating  ligands.  Members  of  this  group 
can  be  readily  and  reversibly  dissociated  into  their  component  parts,  either 
in  the  solid  phase  or  in  solution;  they  show  a  comparatively  large  bond  dis- 
tance between  the  coordinated  ligand  and  the  central  atom ;  and  they  show 
no  deep-seated  electronic  rearrangement  as  measured  by  changes  in  the 
magnetic  susceptibility  of  the  central  ion.  These  compounds  were  named 
normal  complexes  by  Biltz112.  The  ammoniates  of  the  alkali  halides  and  of 
certain  divalent  metal  halides  such  as  cobalt  (II)  chloride  represent  typical 
examples  of  the  normal  complex.  The  term,  normal  complex,  is  often  used 
synonymously  with  the  term  ionic  complex,  although  the  terms  "ionic"  and 
"covalent"  as  applied  to  complexes  indicate  different  things  to  different 
workers. 

Members  of  the  second  group  are  not  in  facile  equilibrium  with  their 
components  in  either  the  solid  state  or  solution.  An  unusually  short  bond 
distance  between  the  coordinating  group  and  the  central  ion  is  usually 
characteristic  of  this  class  of  compounds,  and  a  deep-seated  electronic 
change  is  frequently  indicated  by  a  change  in  the  magnetic  susceptibility 
of  the  central  ion.  Such  compounds  were  called  Werner  complexes  by  Biltz. 
Since  the  so-called  normal  complexes  may  also  be  called  Werner  complexes, 
Ray113  introduced  the  term  "Durchdringungskomplexe"  or  penetration  com- 
plex for  the  second  group  because  of  the  apparent  penetration  of  the  co- 
ordinating ligand  into  the  central  ion.  The  term,  penetration  complex,  is 
frequently  considered  to  be  synonymous  with  the  term,  covalent  complex. 

The  two  types  of  compounds  are  illustrated  by  the  hexamminecobalt(II) 
ion  and  the  hexamminecobalt(III)  ion.  In  the  subsequent  discussion  experi- 
mental evidence  for  the  classification  will  be  reviewed. 

Chemical  Properties  as  a  Basis  for  Classification 

Thermal  decomposition  of  [Co(NH3)6]Cl2  is  characterized  by  the  re- 
versible evolution  of  ammonia  from  the  solid114. 

113.  Hay,  Z.  anorg.  ('hem.,  174,  189  (1928);  J.  Indian  ('hem.  Soc,  5,  73  (1928). 

114.  Biltz,  Z.  anorg.  Chcm.,  89,  97  (1914). 


ELECTROSTATIC  THEORY  OF  COORDINATION  COMPOUNDS      153 

150° 

[Co(NH3)6]Cl2  ^=  =±  Co(NH3)2Cl2  +  4NII 

below  200° 

>  CoCl2  +  2NH3 


The  hexammine  can  be  easily  reformed  by  exposing  the  anhydrous  cobalt  (I  I) 
chloride  to  ammonia  vapors.  The  compound  CoCl2-6NH3  exists  in  aqueous 
solution  in  labile  equilibrium  with  its  components: 

6H20  +  [Co(NH3)6]++  <=>  [Co(H20)6]++  +  6NH3. 

The  solid  hexammoniate  may  be  crystallized  from  a  concentrated  solution  as 
red  octahedra.  The  moist  complex  is  readily  oxidized  by  air  and  is  destroyed 
by  acids.  The  dry  ammoniate  is  fairly  stable  in  air;  in  fact,  ammonia  re- 
places water  from  cobalt (II)  chloride  6-hydrate  when  a  stream  of  ammonia 
is  passed  over  the  solid  compound84.  These  chemical  properties  are  typical 
of  normal  complexes. 

In  sharp  contrast  to  the  ammoniates  of  the  cobalt(II)  salts,  the  hexam- 
minecobalt(III)  salts  do  not  undergo  reversible  thermal  decomposition. 
When  [Co(XH3)6]Cl3  is  carefully  heated,  one  molecule  of  ammonia  is  given 
off  to  produce  chloropentamminecobalt(III)  chloride84. 

[Co(NH3)6]Cl3 : >  [Co(NH3)5Cl]Cl2  +  NH3 

The  reaction  is  slow  and  not  readily  reversible.  Further  heating  brings  about 
complete  decomposition  of  the  chloro  complex  with  reduction  of  the  co- 
balt (III)  ion  by  the  ammonia84-115. 

180°  to  220° 
6  [Co(XH3)5Cl]Cl2 >  6CoCl2  +  6NH4CI  -f  22NH3  +  N2 

The  hexamminecobalt(III)  phosphate  undergoes  immediate  and  complete 
decomposition  on  heating: 

6  [Co(NH3)6]PO<  -»  3Co2P207  +  34NH3  +  3H20  +  N2 

In  solution,  the  hexamminecobalt(III)  ion  does  not  undergo  dissociation 
into  its  component  parts,  as  is  demonstrated  by  the  fact  that  exchange 
studies  on  this  and  related  complex  ions  have  revealed  no  exchange  between 
the  central  metal  ion  and  radioactive  metal  ions  in  solution116,  and  by  the 

115.  Biltz,  Z.  anorg.  Chem.,  83,  190  (1913). 

116.  Lewis  and  Coryell,  Brookhaven  Conf.  Rept.  BNL-C-8,  Isotopic  Exchange  Re- 

actions and  Chem.  Kinetics,  Chem.  Conf.,  No.  2,  131  (1948);  Lewis,  Coryell 
and  Irvine,  J.  Chem.  Soc,  1949,  S386;  McCallom,  Brookhaven  Conf.  Rept. 
BNL-C-8,  Isotopic  Exchange  Reactions  and  Chem.  Kinetics,  Chem.  Conf., 
Xo.  2,  120  (1948);  McCallom  and  Hoshowsky,  ./.  Chem.  Phys.,  16,  254 
(1948);  Hoshowsky,  Holmes,  and  McCallom,  Can.  J.  Research,  27B,  258  (1949); 
Flagg,  J.  Am.  Chem.  Soc,  63,  557  (1941). 


154  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

fact  that  the  complex  ion  is  stable  even  in  strongly  acid  solutions  where  the 
coball  (II)  complex  is  rapidly  decomposed.  In  many  chemical  reactions  the 
complex  hexamminecobalt(III)  ion  participates  as  a  unit  in  a  manner 
analogous  to  that  of  sulfate,  phosphate,  and  other  stable  radicals: 

2[Co(NH3)6]Cl3  +  3H2S04  ->  [Co(NH8)6]2(SO<)3  +  6HC1 

These  chemical  properties  are  characteristic  of  penetration  or  covalent  com- 
plexes. Chromium  is  similar  to  cobalt.  Dipositive  chromium  forms  normal 
complexes  and  tripositive  chromium  forms  penetration  complexes.  It  might 
appear  that  the  greater  charge  and  polarizing  power  of  the  tripositive  ion 
could  account  for  the  differences  in  stability ;  however,  as  Klemm117  points 
out,  the  higher  charge  on  the  central  atom  cannot  explain  the  phenomenon 
by  itself  since  ammines  of  iron  (III)  are  apparently  less  stable  than  those  of 
iron  (II)118. 

Molecular  Volume  as  a  Criterion  for  Classification  of  Complexes 

The  chemical  properties  of  the  hexammine  of  tripositive  cobalt  suggest  a 
much  stronger  bond  between  cobalt  and  nitrogen  than  is  found  in  the 
hexammoniates  of  the  cobalt (II)  salts.  One  might  logically  expect  the  for- 
mation of  the  stronger  cobalt-nitrogen  bond  to  be  accompanied  by  a  de- 
crease in  the  distance  between  the  cobalt  and  nitrogen  nuclei.  Many  of  the 
early  German  workers  reasoned  that  the  decrease  in  the  bond  distances 
might  become  apparent  if  the  molecular  volumes  of  di-  and  tripositive 
metal  ammine  salts  were  compared.  For  this  reason  molecular  volume  was 
introduced  as  a  criterion  of  bond  type. 

Biltz  and  his  co-workers119  applied  Kopp's  rule  of  additive  volumes  to 
coordination  compounds.  They  were  able  to  show  that  the  molecular  vol- 
umes of  a  number  of  hexammines  of  the  divalent  metal  chlorides  are  roughly 
equal  to  the  sums  of  the  zero  point  volumes  of  the  components.  If  the 
additivity  relationship  were  applicable  to  the  hexammines  of  the  tripositive 
metal  chlorides,  one  would  expect  the  volumes  of  the  compounds  containing 
tripositive  metal  ions  to  exceed  the  volumes  of  the  complexes  containing 
divalent  metal  ions  by  an  amount  equal  to  the  volume  of  the  extra  chloride 
ion  (about  16  cc). 

It  is  then  somewhat  surprising  to  find  that  the  molecular  volumes  of  the 
hexammines  of  di-  and  tripositive  metal  ions  with  any  given  anion  are  prac- 
tically identical  in  a  very  large  number  of  cases.  The  extra  anion,  in  most 

117.  Klemm,  Jacobi,  and  Tilk,  Z.  anorg.  Chem.,  201,  1  (1931). 

118.  Thoinr  and  Roberts,  "Fritz  Ephraim's  Inorganic  Chemistry,"  pp.  252,  271. 

and  310,  New  York,  Interscience  Publishers,  Inc.,  (1946). 

119.  Biltz  and  Birk,  Z.  anorg.  Chem.,  134,  125  (1924);  Biltz,  Z.  anorg.  Chew.,  130,  116 

(1923). 


ELECTROSTATIC  THEORY  OF  COORDINATION  COMPOUNDS      155 


Table  3.7.  A  Comparison  of  Molecular  Volumes  for  Selected  Norm  a  i    a\i> 
Penetration  Complexes  Showing  the  Neab  [dentitt  of  Volume  i\ 

Com  PARABLE    Dl-    AND    TRIPOSITIVE    AmMINES 


Normal  Complexes 

Penetration    Complexes 

Ammine 

Ap- 

parent 
Mol. 
Vol. 
Ml. 
(cc) 

Mol. 

Vol. 

Ammine 

Ammiiu' 

Ap- 

parent 
Mol. 
Vol. 
MI, 
(cc) 

Mol. 

Vol. 

Ammine 

[Co(NH,),]Cli 

[Co(NH,).](NO,)a 

[Co(NH,),](C10«), 

[Co(NH,).](CNS), 

[Co(NH,),]Br, 

[Co(NH,).]I, 

[Cr(NH,).]Br, 

[Cr(NH,),]I2 

20 
22 

21 
21 
24 
22 

27 

156.9 
193.2 
225.4 
217.3 
171.6 
198.0 
182.8 
220.3 

[Co(NH3)6]Cl3 
[Co(NH,),](NO,), 

[Co(NH3)6](C104)3 

[Co(NH3)6](CNS)3 

[Co(NH3)6]Br3 

[Co(NH3)6]l3 

[Cr(NH3)6]Br3 

[Cr(NH3)6]I3 

17 

17 

14.5 

18 

19 

19 

22 

156.4 
192.5 

218.2 
171.3 
197.3 
183.2 
220.6 

The  Approximate  Additivity  Relationship  in  Certain  Di-  and  Tripositive 
Hexamminecobalt  Salts 


[Co(NH3)6]S04 

19.1 

155.5 

[Co(NH3)6]2(S04)3 

18.7 

339.8 
(169.9) 

[Co(NH,),]CO« 

19.5 

165.6 

[Co(NH3)6]2(C204)3 

19.1 

368.1 
(184.0) 

[Co(NH3)6](C10H7SO3)2 

18.3 

408.7 

[Co(NH3)6](C10H7SO3)3 

18.0 

553.4 

cases,  does  not  bring  about  a  significant  increase  in  the  volume  of  the 
crystalline  salt.  Data  illustrating  this  point  are  summarized  in  Table  3.7. 

Biltz  and  other  German  workers  of  the  early  1920's  attributed  this  un- 
usual situation  to  a  compression  of  the  coordinated  ammonias  during  the 
formation  of  penetration  complexes.  In  fact,  it  was  from  this  apparent 
compression  of  the  coordinated  ammonias  that  the  name  "penetration 
complex"  arose. 

It  has  been  shown,  however120,  that  the  equal  volume  relationship  is  not 
due  to  the  compression  of  the  coordinated  ammonias,  but  to  the  fact  that 
many  of  the  normal  complexes  such  as  [Co(XH3)6]X2  crystallize  in  a  lattice 
of  the  calcium  fluoride  type.  This  lattice  contains  holes  into  which  four 
extra  anion-  per  unit  cell  may  be  packed  without  destroying  the  basic 
crystal  pattern. 

Magnetic  Susceptibility   Measurements  and  Other  Data  as  Criteria 
for  the  Classification  of  Complexes 

In  his  original  discussion  of  penetration  complexes,  Biltz  noted  thai  i In- 
formation of  such  complexes  is  accompanied  by  profound  changes  in  elec- 

120.  Parry,  Chem.  Revs.,  46,  507  (1950). 


156  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

tronic  arrangement.  The  interpretation  of  these  changes  in  terms  of  a  highly 
polarized  ionic  model  has  been  given  in  the  section  on  magnetism.  A 
comprehensive  review  of  magnetic  data  in  coordination  compounds  was 
published  by  Selwood121  in  1943.  This  work  and  other  work  on  bond  type  is 
most  conveniently  considered  after  a  discussion  of  the  electron-pair  bond. 

121.  Selwood,  "Magnetochemistry,"  New  York,  Interscience  Publishers,  Inc.,  1943. 


4.   Modern  Developments:  The  Electron 

Pair  Bond  and  Structure  of 

Coordination  Compounds 

Raymond  N.  Keller 

University  of  Colorado,  Boulder,  Colorado 

and 

Robert  W.  Parry 

University  of  Michigan,  Ann  Arbor,  Michigan 

Early  Treatments  of  the  Covalent  Bond  en  Coordination 

Theory 

Werner's  Primary  and  Secondary  Valences 

The  advent  of  electronic  theories  of  valence  made  it  possible  to  reconcile 
the  coordination  theory  with  the  structural  theory  of  organic  chemistry. 
The  key  to  the  problem  was  found  by  G.  X.  Lewis1  in  a  postulate  to  the 
effect  that  the  covalent  bond  consists  of  a  shared  pair  of  electrons,  this 
pair  originating  in  one  of  two  ways:  each  of  the  two  atoms  forming  the 
bond  can  furnish  one  electron,  or  one  atom  can  furnish  both.  In  either  case, 
the  outer  shells  of  both  atoms  will  tend  to  be  filled  and  covalent  links  will 
be  formed.  Because  of  its  simplicity,  this  concept  has  served  as  the  founda- 
tion upon  which  much  of  our  present  valence  theory  has  been  built. 

An  electronic  picture  of  a  chemical  bond  did  much  to  make  Werner's 
postulates  of  primary  and  secondary  valences  more  acceptable.  For  ex- 
ample, in  the  ammonia  molecule  the  nitrogen  contributes  one  electron  to 
each  of  the  three  hydrogen  atoms  to  form  three  normal  covalent  bonds. 
These  were  Werner's  "primary  valences."  In  forming  the  ammonium  ion 
the  unshared  electron  pair  on  the  nitrogen  of  the  ammonia  molecule  binds 
a  fourth  proton  to  form  a  coordinate  covalent  bond  or,  a  "secondary  valence." 
Although  the  mode  of  forming  the  two  types  of  bonds  is  different,  the  bonds 
to  all  hydrogens  become  identical  once  they  arc  formed.  (  ha  the  other  hand, 
when  ammonia  is  coordinated  to  a  metal  ion,  the  metal-nitrogen  bond  will 

1.  Lewis,  /.  Am.  Chem.  Soc,  38,  778  (1916). 

157 


158  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

differ  from  the  hydrogen-nitrogen  bond,  not  because  one  bond  is  a  normal 
covalent  bond  and  one  a  coordinate  covalent  bond,  but  because  the  proton 
and  the  metal  ion  differ  in  their  abilities  to  interact  with  the  electrons  of 
the  nitrogen. 

If  the  electronic  interaction  between  two  atoms,  A  and  B,  results  in  a 
complete  transfer  of  an  electron  from  A  to  B,  the  ions  A+  and  B~  are  pro- 
duced to  give  the  conventional  electrovalent  or  ionic  bond. 

Recognition  of  these  different  modes  of  electron  interaction  did  much  to 
dispel  one  of  the  great  objections  to  Werner's  early  theory — that  some 
compounds  of  "first  order"  are  ionic  (e.g.,  NaCl)  and  others  are  not  (e.g., 
CC14).  It  soon  became  apparent  that  Werner's  compounds  of  the  "first 
order"  could  be  divided  into  two  extreme  groups,  ionic  and  covalent,  ac- 
cording to  the  extent  of  electron  transfer  and  that  the  covalent  group  dis- 
played many  properties  which  were  almost  identical  to  those  of  Werner's 
compounds  of  the  "second  order."  For  instance,  in  the  compound 
[Co(NH3)5Cl]Cl2  the  normal  covalent  cobalt-chlorine  bond  in 

[Co(NH8)6Cl]++ 

is  quite  similar  to  the  coordinate  covalent  cobalt-nitrogen  bond  in  terms 
of  chemical  behavior  (i.e.,  slow  reaction  of  CI-  with  Ag+,  etc.).  On  the  other 
hand,  the  ionic  bonds  binding  the  two  remaining  chlorides  to  the  cation 
are  very  different  chemically  from  their  covalent  counterpart.  In  short, 
one  form  of  Werner's  primary  valence  appears  to  be  quite  similar  to  his 
secondary  valence. 

Early  Theories  of  Electron  Quantization 

One  of  the  important  problems  which  followed  the  simple  electronic 
interpretation  of  Werner's  postulates  involved  the  quantization  of  the 
electrons  in  a  complex  molecule  in  a  manner  comparable  to  that  proposed 
by  Bohr2  for  a  simple  atom.  The  problem  is  still  an  active  one  and  many 
methods  of  approach  are  still  being  explored.  Many  early  proposals  as- 
sociated with  such  names  as  Huggins3,  Sidgwick4- 5,  Lowry6,  Main-Smith7, 
Pauling8,  Fowler9,  Butler10-  n,  and  Bose12  are  of  current  interest  in  that  they 

2.  Bohr,  Phil.  Mag.  [6],  26,  1,  476,  857  (1913). 

3.  Huggins,  Phys.  Chem.,  26,  601  (1922);  Science,  55,  459  (1922). 

4.  Sidgwick,  /.  Chem.  Soc,  123,  725  (1923);  Trans.  Faraday  Soc,  19,  469  (1923); 

Chemistry  &  Industry,  42,  901  (1923). 

5.  Sidgwick,  Chem.  and  Ind.,  42,  1203   (1923);    "The  Electronic  Theory  of  Va- 

lency," pp.  100,  172,  124,  Oxford,  Clarendon  Press,  1927. 

6.  Lowry,  Chemistry  &  Industry,  42,  316  (1923). 

7.  Main  Smith,  Chemistry  &  Industry,  42, 1073  (1923);  44,  944  (1925);  Trans.  Faraday 

Soc,  21,  356  (1925-26). 

8.  Pauling,  J.  Am.  Chem.  Soc,  53,  1367  (1931);  54,  988  (1932). 

9.  Fowler,  Trans.  Faraday  Soc,  19,  459  (1923). 


ELECTRON  PAIR  BOS  1)  A  SI)  STRUCTl  RE  L59 

suggest  much  of  our  modern  theory.  For  example,  the  modern  idea  of 
double  bonds  between  metal  and  ligand  was  implied  in  one  of  Sugden's 
early  papers1*.  A  number  of  early  proposals  involving  single  electros  bonds" 
were  severely  criticized8*'  u  and  are  of  little  present  day  value. 

Sidgwick's  Effective  Atomic  Number  Concept 

The  apparent  tendency  of  simple  atoms  to  achieve  an  inert  gas  configura- 
tion in  compound  formation  has  been  a  helpful  and  much  used  concept. 
Sidgwick1  extrapolated  this  idea  in  a  somewhat  modified  form  to  the  heavy 
metal  atoms.  He  postulated  that  the  central  metal  atom  or  cation  of  a 
complex  will  share  electron  pairs  with  coordinating  groups  (or  triplets  in 
some  cases,  as  in  coordination  with  XO)  until  the  "effective  atomic  num- 
ber'' (EAN  )16  of  the  next  higher  inert  gas  is  achieved. 

In  the  case  of  [PtCl6]=,  for  example,  the  effective  atomic  number  of  the 
platinum  atom  is  obtained  by  adding  74  electrons  from  the  Pt4+  ion  and  2 
electrons  from  each  of  the  six  coordinated  chloride  ions  to  obtain  a  total 
of  86.  This  is  the  atomic  number  of  the  inert  gas  radon. 

The  scheme  is  applicable  to  such  a  large  group  of  compounds  that  its 
validity  can  hardly  be  fortuitous.  The  metal  carbonyls  and  nitrosyls  are 
particularly  susceptible  to  treatment  by  this  scheme.  For  example,  the 
formulas  of  the  carbonyls  and  nitrosyls,  and  in  some  cases  their  substitu- 
tion products,  can  usually  be  predicted  by  an  application  of  the  following 
relatively  simple  EAN  rules: 

(1)  Carbon  monoxide  and  electron  pair  donors  such  as  pyridine  etc.,  are 
assumed  to  donate  an  electron  pair  to  the  metal  atom. 

(2)  Nitric  oxide  (XO)  is  assumed  to  donate  three  electrons  to  the  metal 
atom.  since  the  ion  XO+  is  isoelectronic  with  CO. 

(3)  Hydrogen  atoms,  halogen  atoms,  and  pseudo  halogens  such  as  CN 
are  assumed  to  donate  a  single  electron  to  the  metal  atom.  (One  can  also 
look  at  this  in  an  equivalent  manner  as  the  halide  ion  donating  an  electron 
pair  to  the  metal  ion.) 

10.  Butler,  Trans.  Faraday  Soc.,  21,  349  (1925-26). 

11.  Butler.  T  ant.  Faraday  Soc.,  21,  359  (1925-26). 

12.  Bose,  Phil.  Mag.,  [7],  5,  1048  (1928). 

13.  Sugden,  ./.  Chem.  So,-.,  1927,  117:;. 

14.  Main  Smith,  Chemistry  &  Industry,  43,  323   (1924);  Sugden,   "Parachor  and 

Valency,"  Chapts.  6  and  7.  Geo.  Routledge  and  Sons,  Ltd  .  1930;  Sugden,  ./. 
.  125,  1177    L924). 

15.  Samuel,  •/.  Ch*  ..  12,  167     L944  ;  Pauling,  ./.  .1///.  Chem.  Soc.,  53,  3229 

L931);  Emeleua  and  Anderson,  "Modern  Aspects  of  Inorganic  Chemie 
2nd  Ed.,  p.  173,  p.  169,  New  York,  I).  Van  Nostrand  Co.,  Inc.,  1952;  Lessheim 
and  Samuel.  Natv  e,  135,  230  (1935). 

16.  Sidgwick,  7   an*.  Fa  ada  ■  Soc.,  19,  172  (192 


160  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Table  4.1.  Effective  Atomic  Number  Concept  as  Applied  to  Metal 
Carbonyls,  Nitrosyls,  and  Their  Derivatives 


Compound 

Electrons 

From  Central 

Atom 

Electrons 

From 
Ligands 

Effective 
Atomic 
Number 

Deviation 
of  E.A.N. 

from  Inert 
Gas 

Normal  Degree  of 
Association  of  Molecule 

Ni(CO)4 

28 

8 

36 

0 

Monomer 

Fe(CO)4I2 

26 

10 

36 

0 

Monomer 

HCo(CO)4 

28 

9 

36 

0 

Monomer 

Fe(NO)2(CO)2 

26 

10 

36 

0 

Monomer 

Co(CO)4 

27 

8 

35 

1 

Dimer  [Co(CO)4]2 

HNi(CO)3 

28 

7 

35 

1 

Dimer 

Fe(CO)4 

26 

8 

34 

2 

Trimer 

(4)  If  the  effective  atomic  number  of  the  metal  in  the  compound  is  that 
of  an  inert  gas,  the  compound  will  be  a  monomer. 

(5)  If  the  effective  atomic  number  of  the  metal  in  the  compound  is  one 
short  of  that  of  an  inert  gas,  the  compound  will  be  a  dimer.  This  statement 
is  equivalent  to  the  hypothesis  that  the  two  metal  atoms  share  their  odd 
electrons  to  achieve  the  inert  gas  configuration.  More  sophisticated  treat- 
ments of  this  problem  in  terms  of  molecular  orbital  theory  have  been  given17, 
and  the  postulated  metal-metal  bond  seems  reasonable.  A  short  Fe-Fe 
distance  in  Fe2(CO)9  offers  experimental  support  for  the  metal-metal 
bond18. 

(6)  If  the  effective  atomic  number  of  the  metal  in  the  compound  is  two 
short  of  that  of  an  inert  gas,  the  compound  will  be  a  trimer. 

(7)  If  the  effective  atomic  number  of  the  metal  in  the  compound  is  three 
short  of  that  of  an  inert  gas,  the  compound  will  be  a  tetramer. 

The  formulas  in  Table  4.1  illustrate  the  application  of  these  rules. 

On  the  other  hand,  for  non-carbonyl  or  nitrosyl  compounds  there  are  a 
number  of  exceptions  to  the  rare  gas  generalization.  These  were  clearly 
recognized  by  Sidgwick.  The  stable  hexacoordinate  chromium(III)  and 
nickel (II)  complexes  (EAN  =  33  and  38,  respectively)  and  the  stable 
tetracoordinate  nickel(II),  palladium(II),  platinum (II),  and  gold(III) 
complexes  (EAN  =  34,  52,  84,  and  84,  respectively)  are  particularly  strik- 
ing examples. 

Several  well-known  complexes  of  the  alkali  and  alkaline  earth  metals  are 
particularly  damaging  deviations  from  the  rare  gas  rule.  The  rules  on 
polymerization  also  seem  to  be  violated  in  a  number  of  cases  involving 
normal  eovalent  bonds,  particularly  where  the  halo  carbonyls  or  thio 
carbonyls  are  concerned.  In  some  cases  these  exceptions  can  be  rationalized 
by  assuming  a  bridged  type  of  configuration  and  by  using  more  than  one 
pair  of  electrons  per  donor  group,  but  some  discrepancies  still  remain  un- 

17.  Dunitz  and  Orgel,  /.  Chem.  Soc,  1953,  2594. 

18.  Powell  and  Evans.  ./.  Chem.  Soc,  1939,  286. 


ELECTRON  PAIR  BOND  AND  STRUCTURE  161 

explained.  For  example,  the  compounds  Fe(CO)3SEt  and  Fe(NO)2SEt  mv 
dimeric  in  organic  Bolventa  whereas  a  strict  application  of  the  preceding 
rules  would  give  an  effective  atomic  number  of  33  in  each  case  with  a  re- 
sulting tetrameric  structure.  A  rationalization  of  this  apparent  exception  is 
obtained  it*  a  bridged  structure  is  assumed  involving  the  sulfides,  the  formu- 
las becoming 

Et  Et, 

(CO\Fe  .Fe(CO)3  AND  (NO)2Fe  /Fe(NO)2 

Et,  Et 

Each  iron  atom  in  the  above  structures  has  thus  achieved  an  EAN  of  36 
which  is  consistent  with  the  dimeric  formulation.  It  should  be  noted,  how- 
ever, that  such  an  explanation  begs  the  question  since  three  of  the  CO 
groups  in  Feo(CO)9  also  serve  to  bridge  the  two  iron  atoms, 

CO 
CO^  |        CO 


co;6c  ; 


CO^|     ^co 
CO 

yet,  in  this  case  each  bridge  CO  is  still  assumed  to  donate  only  two  electrons 
to  the  metal  atoms.  The  Fe-Fe  distance  and  the  possibility  of  forming  a 
metal-metal  bond  is  probably  important  in  differentiating  the  two  cases. 
Even  more  disturbing  is  the  compound  Fe(CO)3SC6H5  which  has  an  EAX 
of  33  yet  is  a  monomer  in  organic  solvents20.  A  similar  type  of  problem 
arises  in  the  case  of  the  compounds  Pt(CO)2Cl2  and  [Pt(CO)Cl2]2  which 
give  EAX  values  of  84  and  82,  respect  ively.  These  compounds,  which  in- 
volve normal  Pt — CI  bonds,  are  suggestive  of  the  well-known  compounds 
[PtfXH^Clo]  and  [Pt(XH3)4]Cl2 ,  which  are  well-recognized  exceptions  to 
the  EAX  generalization.  R3PAuCl  (EAX  =  82)  is  also  monomeric19,  as  is 
Zi,  en),++  (EAX  =  40). 

In  short,  the  rules  appeal-  to  be  strictly  applicable  to  the  pure  nitrosyls, 
carbonyls,  and  carbonyl  hydrides,  but  their  application  becomes  less  re- 
liable as  other  groups  forming  norma]  covalenl  bonds  are  attached  to  the 
metal  atom.  The  compound  Fe(XO)4  might  appear  to  be  an  exception  to 

19.  Maim,  Wella  and  Purdie,  ./.  Chem.  Soc.,  1937,  1828. 

2o    Hieber  and  Bcharfenberg,  Ber.,  73,  1012    1940  ;  Bieber  and  Spacu,  Z.  anorg. 
233,  363  (1937). 


162  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

this  statement,  but  available  information  on  its  chemical  properties  indi- 
cates (he  structure  can  be  regarded  as  NO+,  [Fe(NO)3]~  in  which  the  EAN 
rules  are  strictly  obeyed21. 

1 1 )  some  cases  magnetic  susceptibility  measurements  can  be  interpreted 
satisfactorily  by  the  EAN  concept.  For  example,  the  carbonyls, 
the  nitrosyls,  and  compounds  such  as  K3[Co(CN)6],  [Co(NH3)6]Cl3 , 
[Co(NH3)8(N02)2Cl],  and  K4[Fe(CN)6]  in  which  the  metal  has  an  EAN  of 
36  are  diamagnetic.  Moreover,  many  compounds  in  which  the  EAN  of  the 
central  atom  is  not  that  of  an  inert  gas  are  paramagnetic,  and  show  sus- 
ceptibilities which  correspond  closely  to  the  deficiency  or  excess  of  elec- 
trons22. 

Quantum  Mechanical  Theories  of  Directed  Valence 

Inherent  in  the  early  Lewis  concept  of  the  shared  electron  pair  and  all 
other  static  models  which  arose  as  variants  of  Lewis'  early  picture  of  the 
chemical  bond  was  the  implication  of  stationary  electrons  and  charges. 
Since  Earnshaw's  theorem  of  electrostatics  states  that  no  system  of  charges 
can  be  in  stable  equilibrium  while  at  rest,  such  models  did  violence  to 
established  rules  of  electrical  behavior  and  failed  to  describe  obvious  physi- 
cal phenomenon  such  as  absorption  and  radiation  of  energy  by  atomic 
systems. 

Bohr's  postulate  of  the  planetary  atom  in  which  electrons  rotate  about  a 
central  positively  charged  nucleus  obviated  some  of  these  difficulties,  but 
the  recognition  of  the  Uncertainty  Principle  by  Heisenberg  in  1927  in- 
dicated that  the  idea  of  definite  electron  orbitals  was  likewise  untenable. 
As  Heisenberg  showed,  there  is  no  way  of  measuring  exactly  the  velocity 
of  an  electron  at  any  given  point;  hence,  a  model  describing  the  electron 
in  such  exact  terms  is  unacceptable.* 

From  this  background  the  modern  discipline  of  wave  mechanics  de- 
veloped. The  theory  introduced  by  Shrodinger  rests  upon  two  concepts: 
(1)  the  wave  nature  of  the  electron  and  (2)  the  statistical  character  of  our 
knowledge  concerning  the  position  of  the  electron.  The  application  of  these 
ideas  to  general  questions  of  valence  is  admirably  done  by  Coulson23  and 
his  book  should  be  consulted  for  any  further  background   information. 

The  probability  of  finding  the  electron  in  any  given  direction  from  the 
nucleus  can  be  obtained  for  different  orbitals  of  the  hydrogen  atom  by  a 

21.  Sidgwick,  "The  Chemical  Elements  and  Their  Compounds,"  Vol.  II,  p.  1373. 

Oxford,  Clarendon  Press,  1950. 

22.  Selwood,  "Magnetochemistry,"  (a)  p.  174,  (b)  p.  161,  New  York,  Interscience 

Publishers,  Inc.,  1943. 
*  See  References  23,  24,  25. 

23.  Coulson,  "Valence,"  (a)  p.  201,  (b)  p.  216,  Oxford,  Clarendon  Press,  1952. 


ELECTRON  PAIR  BOND  AND  STRl'CTl  RE 


1C3 


proper  solution  of  the  wave  equation.  'The  electron  distribution  associated 
with  s,  />,  or  (/*  electrons  is  indicated  in  Fig.  4.1.  The  electron  can  be  found 


C.    d-ORBlTALS  * 

Fig.  4.1.  Shapes  of  atomic  orbitals 


inside  the  appropriate  boundary  surface  any  given  percentage  of  the  time, 
depending- upon  the  absolute  scale  chosen  for  the  drawing23, 26. 

*  See:  Ref.  26  for  more  general  representation  of  d  orbitals. 

24.  Lipscomb,  "Atomic  and  Molecular  Structure"  in  "Comprehensive  Inorganic 
Chemistry,"  edited  by  Sneed,  Maynard  and  Brasted.  New  York,  D.  Van 
Nostrand  Co.,  Inc.,  1953;  Pitzer,  "Quantum  Chemistry,"  New  York,  Prenl  ice- 
Ball,  Inc.,  1053. 

2.5.  Pauling   and  Wilson,    "Introduction   to   Quantum  Mechanics,"      New     York, 
McGraw-Hill  Book  Co.,  1035;  Eyring,  Walter  and  Kimball,  "Quantum  Chem 
istry,"  Now  York,  John  Wiley  &  Sons,  Inc.,  194  1. 

26.  White,  "Introduction  to  Atomic  Spectra,"  p.  63,  Now  York,  McGraw-Hill  Book 
Co.,  Inc.,  1931. 


164  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Two  well-established  approximate  methods  for  treating  molecular  struc- 
tures are  currently  in  use:  (1)  the  atomic  orbital  approximation,  and  (2) 
the  molecular  orbital  approximation.  These  two  approaches  to  the  structure 
of  molecules  differ  in  their  basic  philosophies  and  consequently  in  the 
mathematical  apparatus  used.  It  is  gratifying,  therefore,  that  the  same 
ideas  of  stereochemistry  and  magnetism  of  coordination  compounds  can 
usually  be  obtained  by  the  use  of  either  method. 

The  Atomic  Orbital  Approximation 

In  principle,  the  atomic  orbital  approximation  pictures  the  electron  pair 
bond  as  arising  wThen  two  atoms  are  brought  together  in  a  manner  such 
that  their  appropriate  electronic  orbitals  interact.  As  a  first  approximation, 
such  interaction  will  lead  to  bonding  if  (a)  the  electrons  in  the  two  orbitals 
have  opposite  spin  so  that  electron  pairing  may  result,  and  (b)  the  orbitals 
of  the  two  bonding  electrons  overlap.  In  fact,  it  is  frequently  assumed  that 
the  extent  of  the  overlap  will  determine  the  covalent  bond  strength.  Since 
the  electron  clouds  are  directed  in  space,  the  concept  of  directed  valence 
follows. 

Hybridization.  A  modification  of  the  above  theory  of  directed  valence, 
based  on  the  method  of  localized  electron  pairs8b- 27,  has  been  widely  applied 
in  the  correlation  and  interpretation  of  the  properties  of  coordination  com- 
pounds. It  recognizes  the  experimental  facts  that  all  coordinating  groups 
in  a  complex  ion  such  as  [PtCl6]=  are  bound  to  the  central  metal  ion  in 
exactly  the  same  manner  and  occupy  positions  about  the  metal  ion  which 
are  geometrically  equivalent.  It  follows  that  the  atomic  orbitals  involved 
in  forming  a  number  of  equivalent  covalent  bonds  must  differ  from  each 
other  only  in  direction. 

In  the  formation  of  complex  compounds  there  is  usually  an  insufficient 
number  of  equivalent  bonding  orbitals  available.  It  is  postulated  that  with 
atoms  or  ions  in  which  several  of  the  outer  electronic  levels  differ  little  in 
energy  the  normal  quantization  can  be  changed  or  broken  down  and  new 
equivalent  bonding  orbitals  can  be  formed.  This  is  usually  referred  to  as  a 
"hybridization"  process  and  the  resultant  equivalent  bonding  orbitals,  as 
"hybridized"  orbitals.  In  this  manner,  it  is  possible,  for  example,  to  get 
four  equivalent  orbitals  directed  toward  the  corners  of  a  tetrahedron  or 
square,  or  six  toward  the  corners  of  an  octahedron.  The  energy  for  this 
change  in  quantization  comes  from  the  interaction  energy  accompanying 

27.  Pauling,  "The  Nature  of  the  Chemical  Bond,"  2nd  Ed.,  Ithaca,  New  York,  Cor- 
nell University  Press,  1940;  Heitler  and  London,  Z.  Physik.,  44,  455  (1927); 
Heitler,  Z.  Physik.,  46,  47  (1928);  47,  836  (1928);  51,  805  (1928);  London,  Z. 
Physik.,  46,  455  (1928);  50,  24  (1928);  Naturwissenschaften,  16,  58  (1928); 
Physik.  Z.,  29,  558  (1928) ;  Eisenschatz  and  London,  Z.  Physik.,  60,  491  (1930); 
Slater,  Phys.  Rev.,  37,  481  (1931);  38,  1109  (1931). 


ELECTRON  PAIR  BOND  AND  STRUCTURE  L65 

the  formation  of  the  electron-pair  bonds.  Calculations  by  Pauling  Indicate 
that  this  orbital  hybridization  process  results  in  the  formation  of  stronger 
bonds*  than  would  result  from  bonding  with  pure  unhybridized  orbitals. 
In  general,  the  bonds  formed  between  atoms  will  be  those  with  the  greatest 

bond  strength,  i.e.,  the  condition  of  minimum  potential  energy. 

tiinalion  Number  VI.  Further  insight  into  this  theory  can  perhaps 
best  be  gained  by  considering  the  compounds  and  complex  ions  of  coordi- 
nation Dumber  six.  As  can  he  seen  by  reference  to  a  table  showing  the  elec- 
tronic structure  of  the  elements,  there  are  no  atoms  or  common  ions  which 
have  as  many  as  six  equivalent  peripheral  orbitals. 

For  elements  of  the  first  short  period  of  the  periodic  classification,  the 
single  2s  and  the  three  2p  orbitals  are  available  for  bonding  purpn- 
To  obtain  six  equivalent  orbitals  for  these  elements,  all  or  some  of  the  four 
L  (n  =  2)  orbitals  must  be  combined  or  hybridized  with  orbitals  of  higher 
energy.  Inasmuch  as  the  L  shell  contains  no  d  orbitals,  use  would  have  to 
be  made  of  orbitals  of  the  M  (n  =  3)  shell.  The  large  energy  separation 
between  the  n  =  2  and  n  =  3  levels  evidently  precludes  this  possibility, 
and  no  hexacoordinate  derivatives  of  these  elements  are  known. f 

In  the  case  of  the  elements  of  the  second  short  period,  the  situation  is 
somewhat  different.  The  M  shell  contains  five  3c?  orbitals  along  with  one  3s 
and  three  3p  orbitals,  but  various  lines  of  evidence  indicate  that  the  3d 
orbitals  lie  considerably  above  the  3p  orbitals  in  energy.  Evidently  for  this 
reason  s-p-d  hybridization  is  not  common  among  these  elements,  but  it  is 
not  excluded  and  may  be  operative  in  hexacoordinate  derivatives  such  as 
SF|  ,  [PC16]~,[  SiF6]=,  and  [A1F6]S.  Pauling  has  suggested  that  these  mole- 
cules may  exist  as  partial  ionic  structures  stabilized  by  considerable  reso- 
nance energy  or  may  involve  essentially  ionic  rather  than  covalent  bonds30. 

The  electronic  constitution  of  the  elements  of  the  first  long  period  is 
different  from  that  of  either  of  the  two  short  periods.  In  this  period  the 
elements  of  the  first  transition  series  occur.  These  elements  are  characterized 
by  the  building  up  of  the  3c/  sublevel.  Both  spectroscopic  and  chemical 
evidence  lead  to  the  conclusion  that  the  3d  electrons  in  these  elements  differ 
very  little  in  energy  from  the  4s  electrons.  As  Pauling  pointed  out,  it  is  in 

ding's  interpretation  of  bond  strength  as  the  product  of  the  "strengths"  of 
two  separate  orbitals,  ^A  and  i£B  ,  has  been  extensively  criticized23"-  28-  29.  Mulliken29 
suggests  the  overlap  integral  computed  at  the  experimental  bond  distance,  r,  as  a 
mon  bory  index  of  bond  energ 

f  The   compound     !'•   (  H3)4]<  represei.  tceptioo  to  this  statement,  since 

one  carlxm  in  each  methyl  group  actually  appe  coordinate  *  Rundle  and 

31    rdivant,  ./  goe.,  69,  1661     1947    ;  the  higher  hydrides  of  boron  can 

requiring  special  treatment . 
28.  i        _    M arcoll,  Xyholm,  Orgel  and  Sutton,  ./.  CJu  m.  Sac,  1954,  332. 
-      Mulliken, ./.  An,.  Cfu  m.  8oc.,  72,  440.3  (1950;; ./.  /  .,  56,  295  (19.52;. 

30.  Reference  27a,  pp.  92  and  228 


166  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

elements  of  this  very  type — elements  in  which  the  energy  of  the  inner  d 
orbitals  is  quite  similar  to  that  of  the  s  or  p  orbitals  of  the  valence  shell — 
1  hat  the  d  orbitals  are  most  prone  to  play  an  important  part  in  bond  forma- 
tion, provided  they  are  not  fully  occupied  by  electron  pairs  in  the  uncom- 
I lined  species.  It  has  been  shown31  using  the  atomic  orbital  approximation 
that  a  set  of  six  equivalent  bonding  orbitals  can  be  obtained  by  d2sp*  hy- 
bridization, and  that  these  hybridized  orbitals  are  directed  toward  the 
corners  of  a  regular  octahedron. 

These  concepts  can  be  illustrated  by  applying  them  to  the  cobalt(III) 
ion.  The  outer  electron  configuration  for  the  ground  state  of  the  neutral 
cobalt  atom  is  3d74s2,  and  for  the  cobalt(III)  ion,  3d6.  Application  of  Hund's 
rule  of  maximum  multiplicity  to  obtain  the  ground  state  of  the  ion  gives: 

-     ro+++ M. li         _i?  _ 

1M    11    1 

When  this  ion  combines  with  six  ammonia  molecules,  for  example,  six  pairs 
of  electrons  are  supplied  by  the  ammonias  to  make  the  six  bonds.  To  make 
six  equivalent  orbitals  available  for  these  electrons,  a  rearrangement  of 
electrons  and  levels  must  occur.  The  electrons  occupying  orbitals  singly 
pair  up,  thereby  freeing  two  of  the  3d  levels  for  the  hybridization  process. 
The  combination  of  two  d,  one  s,  and  three  p  orbitals  gives  the  six  equiva- 
lent hybrid  bonds;  the  resulting  configuration  is  abbreviated  as  d2sp3.  The 
final  electron  distribution  is  shown  below.  Since  all  the  electrons  are  paired, 
the  complex  ion  is  diamagnetic. 

— , 

[co(nh3)61++4"    _£1L     ff  _^__         i 

^5pJHYBRIDIZATI0Nj 


In  some  instances  the  total  number  of  electrons  involved  is  not  sufficient 
to  fill  all  the  d  orbitals  after  the  hybridization  process,  and  unpaired  elec- 
trons are  present  in  the  complex.  The  electronic  configurations  for  the 
iron(III)  ion  and  the  cyanide  complex  of  this  ion  are  given  as: 


*  \  \   \  \ 


[Fe(CN)g] 


_3d_\_  4s    _4p_ 

nnunn    \>    nnu 

!  d2sp3 

!__  I 


31.  Pauling,  ./.  Am.  Chem.  Soc,  53,  1386  (1931);  Mills,  ./.  Chem.  Soc,  1942,  465; 
Hultgren,  Phijs.  Rev.,  40,  891  (1932). 


ELECTRON  PAIR  BOND  AND  STRUCTURE  L67 

The  presence  of  one  unpaired  election  in  the  cyanide  complex  is  confirmed 
by  magnetic  susceptibility  measurements88. 

In  the  case  of  CoF|  ,  in  which  the  magnetic  moment  is  the  same  as  that 

of  the  To+++  ion  before  hybridization,  it  is  generally  assumed  that  the  six 
1  ions  are  bound  to  the  centra]  (,o+++  by  electrostatic  forces  such  that 
d-sp''  hybridization  is  not  required.  An  alternative  explanation  would  use 
\d  orbitals  SO  that  the  3d  pattern  would  not  be  disturbed  (see  page  214). 
In  contrast  to  the  cases  just  cited,  in  which  the  total  number  of  electrons 
is  insufficient  to  (ill  all  of  the  orbitals  remaining  after  hybridization,  is  the 
case  in  which  the  d  sublevel  in  the  simple  ion  already  contains  the  maxi- 
mum number  of  elections  allowable  by  the  Pauli  principle.  Copper(I)  serves 
as  an  illustration: 


Vi  \\  n\  u  n 


For  this  ion  to  form  six  covalent  bonds  involving  d2sp3  hybridization,  four 
electrons  would  have  to  be  forced  out  of  the  3d  level  and  promoted  to  a 
higher  state  such  as  the  4d;  alternatively,  4s4p34d2  hybridization  might 
occur.  However,  with  a  nuclear  charge  of  the  order  of  that  of  copper,  the 
energy  difference  between  the  4p  and  4d  levels  is  considerable,  and  either 
of  the  possibilities  for  providing  six  equivalent  bonding  orbitals  would  re- 
quire considerable  energy.  It  is  not  surprising,  therefore,  that  copper(I) 
shows  a  common  coordination  number  of  four  rather  than  six.* 

It  is  apparent  that  these  principles  apply  equally  well  (with  appropriate 
change  in  quantum  numbers)  to  the  4d  transition  elements  in  the  second 
long  period  and  to  the  5c?  transition  elements  in  the  third  long  period.  The 
existence  of  complexes  of  some  of  the  heavy  metals  in  which  the  underlying 
d  shell  is  already  filled,  as  for  example,  SnCl6=  and  SnBr6=,  suggests  that 
the  d  orbitals  of  the  valence  shell  of  the  central  atoms  are  utilized  in  these 
complexes,  or  that  the  complexes  are  essentially  ionic  in  character. 

Nearly  every  theoretical  treatment  in  coordination  chemistry  has  ap- 
parent exceptions  which  require  alteration  of  the  simple  picture.  In  this 
respect,  the  atomic  orbital  approximation  runs  true  to  form  since  disturb- 
ing exceptions  to  the  above  treatment  are  known.  The  ion  [Ru2ClioO]4_  has 
the  hexacoordinate  atomic  arrangement  shown  in  Fig.  4.232.  In  this  com- 

*  The  quest  ion  of  electron  promotion  is  discussed  in  more  detail  on  pages  160  and 

1st. 

32.  Mathieson,  Nfellor  arid  Stephenson,  Acta  Cryst.t  5,  185  (1952). 


168  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

0 


•  =Ru 
0=CI 
©=0 


Fig.  4.2.  The  structure  of  [Ru2Cl10O]4- 

pound  ruthenium  has  a  formal  oxidation  state  of  +4  and  the  complex 
should  be  analogous  to  the  well-known  ion  RuCl6=.  The  magnetic  moment 
of  the  latter  complex  indicates  two  unpaired  electrons  per  ruthenium  atom : 


[RuCi6r i£ ££   __££_ 

d*Sp3  HYBRIDIZATION 


i>  i  4  >n  u    n    u  u  n 


(/x  for  K^RuCle  is  3.07  Bohr  magnetons).  However,  the  oxo-complex  is 
diamagnetic.  The  obvious  conclusion  to  be  drawn  from  this  is  that  seven 
orbitals  of  each  ruthenium  atom  are  being  used  for  bond  formation  instead 
of  six.  Pauling8a-32  suggested  that  two  of  the  seven  bonding  orbitals  are  in- 
volved in  double  bond  formation  to  the  oxygen  (page  202),  Acceptance  of 
such  an  explanation  reduces  the  orbital  treatment  to  a  much  less  certain 
means  of  correlating  structure  and  magnetism,  since  a  decision  cannot  be 
made  in  advance  as  to  when  the  d2spd  hybridization  will  not  correlate  the 
facts  associated  with  the  octahedral  configuration.  Any  explanation  must 
do  violence  to  the  generally  accepted  d2sp3  hybridization  for  the  octahedral 
structure.  The  problem  has  been  treated  by  molecular  orbital  theory  (page 
201). 

Tetrahedral  Configuration.  A  tetrahedral  arrangement  of  orbitals  around 
a  central  ion  may  be  obtained  by  sp*  hybridization.  Elements  of  the  first 
short  period  exhibiting  this  type  of  symmetry  are  found  in  Be(NH3)4++, 
BF4-  CC14 ,  and  NH4+. 

Representative  species  of  the  first  long  period  of  elements  presumably 
showing  sp3  hybridization  are:  [Cu(CN)4]=  [Zn(CN)4]=  and  [Ni(CO)4]. 
Both  Cu(I)  and  Zn(II)  have  completely  filled  3d  sublevels;  hence,  utiliza- 
tion of  the  d  electrons  in  single  bond  formation  is  unlikely.  The  sp3  hybridi- 
zation appears  possible  for  all  the  elements  beyond  and  including  zinc  in 
the  first  long  period.  The  tetrahedral  configuration  seems  to  be  generally 
favored  except  in  the  cases  of  a  relatively  few  hexacoordinate  derivatives 
such  as  [Zn  ena]"*"1",  SeF6=  and  AsF6~  which  may  involve  predominantly  ionic 
bonding  or  utilization  of  4d  levels. 


ELECTRON  PAIR  BOND  AND  STRUCTURE  L69 

Tetracoordinate  derivatives  of  the  transition  elements  may  also  at  lain 
a  tetrahedraJ  arrangement  by  hybridization  of  three  of  the  penultimate  d 
orbitals  with  the  8  orbitals  of  the  valence  shell.  Such  behavior  is,  of  course, 

usually  limited  to  the  higher  oxidation  states  of  these  elements  as  in  Cr04~, 
MnOr,  MoOr,  and  WOr. 

Planar  Configuration.  When  only  one  d  orbital  of  the  penultimate  major 
quantum  shell  is  available,  dsp2  hybridization  occurs,  and  the  resulting 
equivalent  hybridized  orbitals  are  directed  in  space  toward  the  corners  of  a 
square.  It  is  remarkable  that  most  of  the  planar  molecules  and  ions  so  far 
discovered  are  compounds  of  nickel(II),  palladium(II),  platinum(II),  and 
gold  (III).  It  will  be  noticed  that  each  of  these  ions  has  only  eight  d  elec- 
trons, leaving  one  d  orbital  available  for  hybridization  with  s  and  p  orbi- 
tals. It  seems  quite  likely  that  all  tetracovalent  compounds  of  copper(II) 
are  planar33.  Since  the  copper(II)  ion  contains  9d  electrons,  dsp2  hybridiza- 
tion can  take  place  only  if  one  d  electron  is  promoted  to  a  4p  or  4d  level, 
a  process  requiring  energy.  However,  if  sufficient  energy  can  be  gained  by 
the  formation  of  dsp2  hybrid  bonds,,  the  combination  procedure  of  d-elec- 
tron  promotion  plus  dsp2  hybridization  is  favored  over  the  alternative  of 
sp3  hybridization.*  On  the  basis  of  a  limited  amount  of  experimental 
evidence,  silver(II)  and  silver(III)  as  well  as  copper(III)  show  a  square 
configuration  in  covalent  structures33, 34.f 

The  original  theory,  as  stated  by  Pauling,  predicted  a  planar  configura- 
tion for  ions  having  one  and  only  one  d  orbital  available  for  bond  formation, 
those  with  more  than  one  d  orbital  forming  either  tetrahedral  or  octahedral 
compounds.  However,  there  is  some  evidence  for  the  planar  configura- 
tion  of  cobalt(II)   and  manganese(II)33a' 34, 35. 

*  Xyholm34  has  pointed  out  that  there  are  serious  objections  to  this  hypothesis  of 
electron  promotion  in  copper(II)  complexes.  First,  promotion  of  the  electron  to  a  4p 
level  should  result  in  facile  oxidation  of  square  copper(II)  complexes  to  the  cop- 
per(III)  state.  This  is  not  observed.  Also,  theoretical  work28  leads  to  the  conclusion 
that  fairly  electronegative  groups  like  H20  and  Cl~  (which  do  give  square  copper(II) 
complexes)  are  more  likely  to  use  4d  rather  than  3d  bond  orbitals.  In  the  case  of 
Xi(II),  groups  of  low  electronegativity  are  required  to  form  SdisAp2  bonds.  Nyholm 
favors  a  4s4p24d  configuration  for  square  copper(II)  complexes. 

t  The  compound  K3CuF6  containing  copper  (III)  has  a  moment  of  2.9  Bohr  mag- 
netons; hence  the  structure  is  probably  ionic  and  octahedral  (p.  172). 

33.  Mellor,  Chem.  Rev.,  33,  137  (1943) ;  Helmholz,  J.  Am.  Chem.  Soc,  69,  886  (1947). 

34.  Xyholm,  Quart.  Revs.,  7,  392  (1953). 

35.  Calvin  and  Melchior,  J.  Am.  Chem.  Soc,  70,  3273  (1948);  Biltz  and  Fetkenheur, 

Z.  anorg.  Chem.,  89,  97  (1914);  Cambi  and  Malatesta,  Gazz.  chim.  ital.,  69,  647 
(1939) ;  Mellor  and  Craig,  J.  Proc.  Roy.  Soc,  N.  S.  Wales,  74,  495  (1940) ;  Bark- 
worth  and  Sugden,  Nature,  139,  374  (1937);  Mellor  and  Coryell,  ./.  Am.  Chem. 
Soc,  60,  1786  (1938);  Cox,  Shorter,  Wardlaw  and  Way,  ./.  Cfu  m.  Soc,  1937, 
1556;  Figgis  and  Xyholm,  ./.  Chem.  Soc,  1964,  12. 


170  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

-J 
Table  4.2.  Orbital  and  Spatial  Configurations  for  Coordination  Numbers 
Two  Through  Eight  Including  Bond  Strengths  and 
Representative  Compounds* 


Coordin- 
ation No. 

Orbital 
Configuration 

Spatial  Configuration 

Relative 
Bond 

Strengths* 

Examples 

2 

P2 

angular 

1.732 

H20,H2S,0F2,SC12 

sp 

linear 

1.932 

Ag(CN)2-,Hg2X2 

3 

p3 

trigonal  pyramid 

1.732 

NH3  ,  PH3  ,  AsCl3 

sp2 

trigonal  plane 

1.991 

B(CH3)3  ,  N03- 

4 

spz 

tetrahedron 

2.000 

[B(CH3)3NH3],Ni(CO)4, 
[Cu(CN)4]- 

dsp2 

tetragonal  plane 

2.694 

[Pt(NH,)4]++,  IA11CI4]- 
[Ni(CN)4]= 

d3s 

tetrahedron 

2.950 

Cr04=,  Mo04= 

d2p2 

tetragonal  plane 

— 

IClr  t 

5 

dsp3  or  d3sp 

trigonal  bipyramid 

— 

PCI5  ,  M0CI5  ,  TaF5 

d2sp2,   d4s, 

tetragonal  pyramid 

— 

IF6,[Ni(PEt3)2Br3] 

d2p3,  or  dAp 

6 

d2sp3 

octahedron 

2.923 

[PdCle]",  [Co(NH3)6]++*- 

d4sp 

trigonal  prism 

2.983 

MoS2 ,  WS2 

7 

d5sp  or 
d3sp3 

octahedron  with  an  atom 
at  the  center  of  one  face 

— 

[ZrF7]"3 

d4sp2  or 

trigonal    prism    with    an 

— 

[TaF7]=,  [NbF7]- 

d5p2 

atom  at  center  of  one  of 
the  square  faces. 

81 

d4sp3 

dodecahedron 

— 

[Mo(CN)8]*- 

d5p3 

antiprism 

— 

[TaF8]- 

d5sp2 

face-centered  prism 

— 

[OsF8] 

*  For  the  special  meaning  of  "bond  strength"  as  used  here,  see  references28-  29>  37. 

f  The  iodine  atom  in  this  compound  is  also  considered  to  possess  two  stereochemi- 
cally  active  unshared  electron  pairs  in  octahedral  positions,  a  structure  which  at 
the  present  time  appears  to  be  unique33*-  38. 

t  Van  Vleck39  has  expressed  the  opinion  that  a  complex  with  eight  attached  groups 
is  unlikely  to  be  stable  unless/  orbitals  are  available  on  the  central  atom.  This  may 
be  one  reason  why  relatively  few  atoms  exhibit  a  coordination  number  of  eight40. 

Other  Coordination  Numbers.  A  comprehensive  treatment  of  coordination 
involving  different  modes  of  hybridization  was  carried  out  by  Kimball36 
using  both  the  atomic  orbital  and  molecular  orbital  approximations.  A 
summary  of  the  stereochemical  implications  of  his  results  for  coordination 
numbers  two  through  eight  appears  in  Table  4.2. 

36.  Kimball,  J.  Chem.  Phye.,  8,  188  (1940). 

37.  Reference  27a,  Chap.  Ill;  Pauling  and  Sherman,  J.  Am.  Chem.  Soc,  59,  1450 

(1937)  ;Ref.  23,  p.  197. 

38.  Sidgwick  and  Powell,  Proc.  Roy.  Soc.  (London),  A176,  153  (1940);  Mooney,  Z. 

Krist.,98,  377  (1938). 

39.  Van  Vleck,  J.  Chem.  Phys.,  3,  805  (1935). 

40.  Penney  and  Anderson,  Trans.  Faraday  Soc.,  33,  1363  (1937). 


ELECTRON  PAIR  BOND  AND  STRUCTURE  171 

Stereochemistry  and  the  Nature  of  the  Central  Atom.*  As  has 
been  indicated  previously,  the  nickel(II)  ion  has  an  electronic  structure 
which  permits  formation  of  diamagnetic  square  planar  dqp*  bonds,  yet 
paramagnetic  tetrahedral  sp{  nickel(II)  complexes  are  also  known.  The 
nickel  gly oximes  and  Ni(CN)4",  for  example,  have  been  shown  to  be  dia- 
magnetic and  planar11,  whereas  [Nil  X 1 1.:')»] {  f  is  paramagnetic  and  presum- 
ably tetrahedral. | 

In  a  comprehensive  review,  Mellor33a  considered  which  electronic  configu- 
rations of  a  metal  will  favor  octahedral,  planar,  or  tetrahedral  structures. 
After  a  very  careful  review  of  the  data,  he  concluded  that,  "when  a  metal 
atom  of  the  transition  series  forms  a  covalent  complex,  it  tends  to  assume 
that  configuration  (tetrahedral,  square,  octahedral,  etc.)  which  involves 
the  least  possible  number  of  unpaired  electrons. "{  This  generalization 
appears  to  follow  from  an  inspection  of  Table  4.3,  which  is  reproduced 
from  Mellor's  paper.  The  relatively  few  ions  for  which  a  planar  configura- 
tion has  been  reported  are  underlined.  It  is  significant  that  the  planar 
configuration  is  most  common  among  the  elements  in  those  oxidation  states 
for  which  the  resulting  complex  contains  no  unpaired  electrons  (Ni++,  Pd++, 
Pt4"1",  Au+++)  or  one  unpaired  electron  (CU++  Ag++,  CO++);  the  planar 
configuration  is  much  less  common  or  even  doubtful  among  those  ions 
giving  <lsp2  bonded  complexes  with  two  or  three  unpaired  electrons  (Fe++, 
Mirf),  and  is  probably  not  existent  among  those  containing  the  maximum 
of  four  unpaired  electrons.  The  octahedral  configuration  is  invariably  as- 
sociated with  complexes  of  Co+++  Rh+++  Pd4+,  Ir+++  and  Pt4+;  and,  with 
few  exceptions,  these  complexes  are  diamagnetic. 

According  to  the  original  criteria  used  to  predict  planar  and  tetrahedral 
configurations,  a  change  in  the  oxidation  state  of  a  central  metal  ion  can 
lead  to  a  complete  change  in  bond  orientation  (Table  4.3).  This  is  confirmed 
by  the  existence  of  tetrahedral  Ni(CO)4  and  planar  [Ni(CN)4]=  which  are 
derivatives  of  nickel (0)  and  nickel(II),  respectively,  and  by  diamagnetic 

*  See  also  Chapter  9. 

f  It  is  interesting  that  unequivocal  experimental  proof  for  the  tetrahedral  con- 
figuration for  this  ion  is  not  yet  available— more  than  twenty  years  after  Pauling's 
suggestion — but  Xyholm42a  has  summarized  existing  evidence  for  the  tetrahedral 
form  in  a  rather  convincing  fashion.  The  complexes  assumed  to  be  tetrahedral  are 
generally  green  or  blue  in  color  as  compared  to  the  diamagnetic  complexes  which  are 
usually  red,  brown,  or  yellow12.  Mellor13  and  his  co-workers  have  reported,  however, 
that  the  correlation  between  configuration  and  color  is  not  always  clear-cut.  Xy- 
holnr'Ji  reports  thai  a  more  reliable  though  not  Infallible  criterion  of  diamagnetism 
is  a  sharp  absorption  band  in  the  vicinity  of  4,000  A. 

X  Van  Vleck*'  expressed  about  the  same  idea  in  calling  attention  to  the  fact  that 
while  a  large  spin  (due  to  unpaired  electrons)  might  be  an  advantage  as  far  as  a  free 
atom  is  concerned,  in  an  atomic  Bystem  the  interatomic  energy  may  be  decreased  by 
a  lowering  of  the  total  spin. 


L72 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


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ELECTROS  PAIR  BOND  AND  STRUCTURE  173 

tetrahedral  complexes  of  copper(I)  and  Bilver(I)  as  opposed  to  the  para- 
magnetic planar  derivatives  of  copper(II)  and  silver(II). 

The  question  as  to  which  of  two  possible  structures,  square  or  tetrahedral, 
will  be  assumed  by  t ho  nickel(II)  compounds  is  more  complex.  One  of  the 
major  factors  determining  the  geometry  appears  to  be  the  relative  differ- 
ences in  the  electronegativities  of  the  nickel  and  of  the  atoms  linked  to  it. 
Large  differences  appear  to  favor  predominantly  ionic  bonds  and  the 
tetrahedral  configuration,  although  the  nature  of  the  functional  group  in 
which  the  atom  bonded  to  the  nickel  occurs  may  also  be  significant.  Stone 
factors  are  sometimes  of  major  importance44. 

In  some  instances  the  type  of  crystal  lattice,  the  solvent,  and  the  tem- 
perature appear  to  be  important  in  determining  which  configuration  will  be 
assumed42*'  45, 46.  For  example,  [Ni  en2]  [AgIBr]2  is  diamagnetic  in  the  solid 
state,  whereas  compounds  of  [Ni  en2]++  with  anions  like  C104~  are  para- 
magnetic in  the  solid  state.  Similarly,  dipole  moment  measurements  and 
magnetic  data  indicate  that  [NiCl2- {(C2H6)3P}2]  and  [NiBiv  { (C2H5)3Pj2] 
are  trans-planar,  but  when  the  halogens  are  replaced  by  nitrate,  both  the 
dipole  data  and  magnetic  moments  indicate  a  tetrahedral  structure.  Lattice 
factors  are  of  importance  in  determining  the  reorientation  of  orbitals. 
The  compound  bis(salicylaldoxime)  nickel (II)  is  diamagnetic  both  in  the 
solid  state  and  in  benzene  solution,  but  has  a  magnetic  moment  indicating 
two  unpaired  electrons  in  pyridine  solution.  This  has  been  ascribed  to  octa- 
hedral coordination  in  pyridine  solution.  On  the  other  hand,  bis(N-methyl- 
salicylaldimine)  nickel(II)  is  diamagnetic  in  the  solid  state  but  paramag- 
netic in  benzene.  Since  benzene  does  not  usually  coordinate  with  nickel, 
one  might  assume  that  the  paramagnetic  form  represents  a  tetrahedral 
-tincture  in  benzene.  Actually,  Klemm  and  Raddatz47  have  reported  the 
isolation  of  paramagnetic  and  diamagnetic  forms  of  the  solid  salt;  the 
paramagnetic  form  changes  spontaneously  to  the  diamagnetic  form  on 
standing.  Recently  Basolo  and  Matoush46  reported  that  no  direct  correla- 

41.  Sugden,  J.  Chem.  Soc,  1932,  246;  Brasseur,  de  Rassenfosse  and  Pierard,  Compt. 

rend.,  198,  1048  (1934);  Cambi  and  Szego,  Ber.,  64,  2591  (1931). 

42.  Nyholm,  Chem.  Rev.,  53,  267   (1953);  Lifschitz,  Bos,  and  Dijkema,  Z.  anorg. 

aUgem.  Chem.,  942,91  (1930);  Lifschitz  and  Bos,  Rec.  trav.  cftim.,89,  107  (1940); 
Lifschitz  and  Dijkema,  Rec.  trav.  chim.,  60,  5S1  0941);  Ref.  27a,  p.  122. 

43.  Mills  and  Mellor,  •/.  .1///.  Chem.  Soc.,  64,  181  (1942);  Mellor,  Mills  and  Short, 

./.  Proc.  Roy.  Soc.,  N.  8.  Wales,  78,  70  (1911 

44.  Reference  22,  p.  180. 

45.  Willis  and  Mellor, ./.  Am.  Chem.  Soc.,  69,  1237  (1947);  French,  Magee,  and  Shef- 

field,./. .1///.  Chem.  Nor.,  64,  1924  (1942  ;  Johnson  and  Hall../.  .1///.  Chem.  Soc, 
70,23  17    litis  ;  Lifschitz,  Rec.  trav.  chim.,  66,  401  (1947). 

46.  Basolo  and  Matoush,  ./.  .1///.  Chem.  Soc.  75,  5663    1963  , 

17.  Klemm  and  Raddatz,  Z.  anorg.  allgem.  Chem.,  250,  207  (1942) 


174  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

tion  exists  between  the  magnetic  susceptibility  of  solutions  of  bis(formyl- 
camphor)ethylenediamine  nickel (II)  in  methylbenzenes  and  the  base 
strength  of  the  solvent.  If  the  paramagnetic  susceptibility  were  due  to  for- 
mation of  octahedral  complexes  by  expansion  of  the  coordination  shell  of 
nickel,  one  might  expect  such  a  correlation.  The  data  lead  to  the  conclusion 
that  tetrahedral  nickel (II)  compounds  are  formed  in  the  solvent.  Data 
delineating  the  effects  of  temperature  on  the  conversion  are  sparse. 

Sidgwick  and  Powell38a  studied  the  empirical  relationship  between  stereo- 
chemical types,  the  nature  of  the  valence  group  of  the  central  atom,  and 
the  number  of  shared  electrons.  Their  scheme  bears  considerable  resem- 
blance to  that  of  Tsuchida  (page  131)  in  application,  although  the  assumed 
charge  distribution  is  quite  different  in  the  two  cases.  The  results  are  em- 
pirically useful,  although  of  doubtful  theoretical  interest  at  present.* 

Stability  of  Complexes  and  the  Atomic  Orbital  Theory.  The  Role 
of  the  Metal.  The  stability  of  complexes  has  been  considered  in  terms  of  a 
thermochemical  cycle  on  page  137.  It  is  apparent  that  the  ultimate  stability 
of  any  given  compound  is  dependent  upon  small  differences  between  large 
energy  terms  (page  143) ;  thus,  the  degree  of  precision  required  in  making 
energy  estimates  for  any  given  step  in  the  cycle  must  be  very  high;  other- 
wise the  final  energy  of  formation  of  a  compound  may  even  be  reversed  in 
sign  as  a  result  of  relatively  small  errors  in  any  one  term.  Fortunately,  in 
many  cases  of  complex  formation,  particularly  in  aqueous  solution,  the 
stabilities  of  compounds  of  similar  type  can  be  compared  under  such  condi- 
tions that  differences  in!  the  energy  of  coordination,  E,  for  different  metals 
will  be  relatively  large  compared  to  differences  in  other  energy  terms  such 
as  heats  of  hydration  of  the  gaseous  ions  and  the  ligands  involved.  Under 
such  conditions  the  stabilities  of  the  complexes  may  be  correlated  with  those 
factors  influencing  the  energy  of  coordination: 

M(g)++  +  Ligand(g)  ->  M  Ligand(g)++ 

Since  nitrogen,  oxygen  and  sulfur  serve  as  the  actual  bonding  atoms  in 
a  large  majority  of  complex  compounds,  Sidgwick49  divided  the  metals  into 
three  categories  on  the  basis  of  their  relative  abilities  to  combine  with  oxy- 
gen (usually  through  a  normal  covalent  bond)  or  nitrogen  (usually  through 

*  Several  general  rules  applying  to  molecular  configurations  and  electronic  con- 
stitution of  simple  molecules,  which  are  almost  identical  to  portions  of  the  scheme  of 
Sidgwick  and  Powell,  were  advanced  more  recently  by  Helferich48. 

48.  Helferich,  Z.  Naturforsch.,  1,  666  (1946). 

49.  Sidgwick:  J.   Chem.  Soc,  433   (1941);  "The  Electronic  Theory  of  Valency," 

Oxford  University  Press,  1927. 


ELECTRON  PAIR  BOND  AND  STRUCTl  RE  175 

a  coordinate  covalent  bond).  These  categories  are: 

(1)  Bond  to  oxygen  Btronger  than  to  nitrogen: 

Mg,  Ca,  Sr,  Ba,  Ga,  In,  Tl,  Ti,  Zr,  Th,  Si,  Ge,  Sn,  Vv,  \  ,v,  \l>\ 
Tav,  Mov,  QVI,  Fem,  Co11. 

(2)  Bond  to  oxygen  and  nitrogen  with  about  equal  strength: 

Be,  Crm,  Fe11,  platinum  metals 

(3)  Bond  to  nitrogen  stronger  than  to  oxygen : 

Cu1,  Ag\  Au1,  Cu11,  Cd,  Hg,  V111,  Co111,  Ni11. 

It  will  be  noticed  that  nearly  all  of  the  ions  of  group  (1)  are  of  the  inert 
iia>  type;  those  of  group  (3)  are  of  the  palladium  type  or  are  small  and  have 
a  nearly  full  d  level  (i.e.,  Nr*"*),  whereas  the  intermediate  ions  are  the  very 
small  beryllium  ion  and  the  larger  transition  ions.  Some  justification  for 
this  grouping  has  been  given  in  Chapter  3. 

It  must  be  recognized  that  broad  generalizations  such  as  the  above  will 
have  many  exceptions,  particularly  in  certain  intermediate  regions,  but  it 
is  significant  that  in  a  recent  survey  of  the  coordinating  ability  of  a  number 
of  different  ligands  YanUitert  and  Fernelius50  reported  that  "compounds 
formed  by  chelating  agents  bonding  through  nitrogen  show  a  greater  de- 
pendency upon  metal  ion  electronegativity  than  those  bonding  through 
oxygen,"  an  observation  which  supports  admirably  the  foregoing  generali- 
zation. In  particular  it  was  found  that  Ca++  and  Mg"^  coordinate  more 
effectively  through  oxygen  whereas  CU++  and  Xi++  coordinate  best  through 
nitrogen. 

A  number  of  investigators  in  recent  years  have  attempted  to  list  the 
metal  cations  on  the  basis  of  their  ability  to  coordinate  with  one  or  two 
specific  ligands.  Using  a  chelating  agent  involving  oxygen  and  nitrogen 
bonds,  Pfeiffer,  Thielert,  and  Glaser51  obtained  the  following  order  of  de- 
creasing stability  of  complex:  Cu++,  Ni++,  Fe^,  Zn++,  Mg++.  Mellor  and 
Maley52  studied  the  stability  of  salicylaldehyde  complexes  in  50  per  cent 
water-dioxane  solution  using  the  method  developed  by  Bjerrum53.  Their 
order  of  decreasing  stability  was:  Pd++,  Cu++  Ni++,  Co++  Zn++,  Cd++ 
Fe++,  Mn44,  Mg^+.  With  minor  exceptions  the  order  is  the  same  as  that 
given  by  Pfeiffer  and  as  that  found  when  glycine,  8-hydroxyquinoline,  or 
ethylenediamine  is  the  chelating  group  in  aqueous  solution. 

50.  VanUitert  and  Fernelius:  J.  Am.  Chem.  Soc,  76,  375,  379  (1954). 

51.  Pfeiffer,  Thielert,  and  Glaser:  J.  prakt.  Chew.,  152,  145  (1939). 
62.  Mellor  and  Maley:  Nature,  159,  370  (1047;;  161,  136    1 048) . 

53.  Bjerrum:  "Metal  Ammine  Formation  in  Aqueous  Solution,"  Copenhagen,  P. 
Haase  and  Son,  1941. 


176  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Calvin  and  Melchior35a  used  the  method  of  Bjerrum  to  study  the  stability 
of  the  salicylaldehyde  chelates  in  water  solution,  using  a  sulfonated  salicyl- 
aldehyde  to  obtain  water  solubility.  A  similar  set  of  data  was  accumulated 
for  o-formylnaphthols.  In  all  cases  the  stability  of  the  series  was: 

Cu++,  Ni++,  Co++,  Zn++ 
Decreasing  Stability 


of  Chelate 

Since  the  order  is  essentially  the  same  as  that  of  Mellor,  the  role  of  the  sol- 
vent seems  to  be  small. 

VanUitert,  Fernelius,  and  Douglas56,  using  a  modification  of  the  Bjerrum 
titration  method,  studied  the  stabilities  of  the  metal  chelates  of  several 
substituted  /3-diketones.  They  found  that  the  general  order  of  stability  in 
75  per  cent  dioxane-water  solution  is: 

Hg++,  (Cu++,  Be++)  Fe++,  Ni++,  Co++,  Zn++,  Pb++,  Mn++,  Cd++,  Mg++,  Ca++,  Sr++  Ba++. 

> 

Decreasing  stability 

Similar  series  using  other  ligands  have  also  been  given54- 55, 56.  Results  show 
some  deviation  from  the  above  lists,  but  certain  features  are  recurrent. 

In  general,  the  stability  of  the  complexes  of  the  alkali  and  alkaline  earth 
metals  decreases  as  the  charge  on  the  cation  decreases  or  as  the  size  of  the 
cation  increases.  Lumb  and  Martell57  found  that  the  stabilities  of  alkaline 
earth  complexes  of  glutamic  and  aspartic  acids  fall  in  the  order  Mg++  > 
Ca++  >  Sr++  >  Ba++  >  Ra++.  The  stability  of  the  citric  acid  complexes 
of  the  alkaline  earths  falls  in  the  order  Ca++  >  Sr++  >  Ba++58.  A  similar 
order  has  been  reported  for  the  complexes  of  a  number  of  alkali  and  alka- 
line earth  metal  ions  with  N-acetic  acid  substituted  amines  and  with 
poly  amines.59  All  data  on  the  complexes  of  the  rare  earth  ions  are  also 
consistent  in  showing  a  decrease  in  complex  stability  with  increasing  size 
of  the  rare  earth  ion60-  *■ 62- 63- 64.  (See  Fig.  4.4) 

54.  Merritt,  "Frontiers  of  Science  Outline,"  Wayne  University,  Spring,  1949; 

55.  Chabarek  and  Martell: J.  Am.  Chem.  Soc,  75,  2888  (1953). 

56.  VanUitert,  Fernelius,  and  Douglas:  /.  Am.  Chem.  Soc,  75,  457,  2736,  2739,  3577, 

(1953);  VanUitert  and  Hass,  J.  Am.  Chem.  Soc,  75,  451   (1953);  VanUitert, 
Hass,  Fernelius,  and  Douglas,  J.  Am.  Chem.  Soc,  75,  455  (1953). 

57.  Lumb  and  Martell,  J.  Am.  Chem.  Soc,  75,  690  (1953). 

58.  Hennig,  Schmahl,  and  Theopold,  Biochem.  Z.,  321,  401  (1952). 

59.  Martell  and  Calvin,  "Chemistry  of  the  Metal  Chelate  Compounds,"  (a)  p.  192; 

(b)  p.  190,  New  York,  Prentice-Hall,  Inc.,  1952. 

60.  Spedding  and  Powell,  J.  Am.  Chem.  Soc,  76,  2545,  2550  (1954)  and  earlier  papers 

of  Spedding  on  ion  exchange  separation  of  rare  earths  with  citrate. 

61.  Fitch  and  Russell,  Can.  J.  Chem.,  29,  363  (1951);  Anal.  Chem.,  23,  1469  (1951); 

Beck,  Chem.  Acta,  29,  357  (1946). 

62.  Moeller,  Record  Chem.  Progress,  14,  69  (1953). 


ELECTRON  PAIR  BOND  AND  STRUCTURE  111 

Irving  ami  William.-""  summarized  the  results  of  many  investigators  in  an 
excellent  review  of  available  stability  data.  They  recognized  thai  compari- 
sons of  the  stabilities  of  complexes  of  different  ligands  are  mosl  effective 
when  metals  of  the  same  type  are  used.  Reversals  found  in  the  earlier  lists 
arise  because  comparisons  were  drawn  between  complexes  of  disshnilai 
metals.  When  comparisons  wen1  restricted  to  bivalent  metals  of  the  first 
transition  series  they  found  thai  the  order  Mn  <  (poorer  than)  Fe  <  Co  < 
Ni  <  Cu  >  (better  than)  Zn  is  valid  irrespective  of  the  nature  of  the 
coordinated  ligand  or  the  Dumber  of  ligands  involved.  Since  the  ability  of 
metals  to  coordinate  with  nitrogen,  oxygen  or  sulfur  varies,  depending  upon 
the  type  oi  metal  considered,  no  single  series  involving  all  metal  ions  with 
all  ligands  can  ever  be  expected.  Irving  and  Williams  correlated  their  series 
with  the  reciprocal  of  the  ionic  radii  and  the  second  ionization  potential- 
of  the  metal-  as  suggested  by  Irving  and  Williams65b  and  by  Calvin  and 
Melchior35a. 

Such  a  correlation  finds  justification  in  that  the  second  ionization  poten- 
tial for  ions  of  comparable  size  can  be  used  as  an  estimate  of  the  strength 
of  the  a  bond  between  metal  and  ligand.  The  ion  type  is  important  in  that 
it  determines  the  extent  of  secondary  interactions  such  as  multiple  bond 
formation  (p.  191).  The  data  for  the  alkaline  earth,  alkali  metal,  and  rare 
earth  metal  ions  can  best  be  considered  in  terms  of  predominantly  ionic 
bonds  (Chapter  3). 

Martell  and  Calvin59b  indicated  the  general  relationship  between  the  for- 
mation constants  of  metal  chelates  and  the  second  ionization  potentials  of 
the  metals  by  means  of  the  plot  shown  in  Fig.  4.3.  The  relationship  between 
the  stability  constants  of  the  rare  earth  chelates  of  ethylenediamine  tetra- 
acetate and  the  reciprocal  of  the  radius  of  the  rare  earth  ions  is  shown  in 
Fig.  4.4.  In  both  of  these  cases  the  ions  are  sufficiently  similar  so  that  the 
method  chosen  to  estimate  the  field  strength  around  the  ion  is  reasonably 
good  for  all  members  of  the  series.* 

The  Role  of  the  Ligand.  If  one  accepts  the  definition  of  G.  N.  Lewis  that 
a  base  is  an  electron  pair  donor,  the  process  of  coordination  is  an  acid-base 
phenomenon  in  which  the  coordinated  ligand  acts  as  a  base  and  the  metal 
ion  acts  as  an  acid.  The  point  is  illustrated  by  comparing  the  typical  acid- 

*  Wheelwright,  Spedding  and  Schwarzenbach64b  suggested  that  the  rare  earth 
ethylenediaminetetraacetate  complexes  change  from  hexadentate  to  pentadentate 
structures  at  Gd'"^  because  of  steric  effects  due  to  decreasing  size  of  cation. 

63.  Dissertations,  University  of  Illinois,  Brantley  (1949);  Moss  (1952). 

64.  Martell  and  Plumb,  J.  Phys.  Chem.,  56,  993  (1952);  Wheelwright,  Spedding,  and 

Schwanenbach,  ./.  Am.  ('hit,,.  Soc,  75,  4100  (1953);  Spedding,  Powell,  and 
Wheelwright,  ./.  Am.  Chi  m.  8oc.s  76,  2557  (1954);  Templeton  and   Dauben,  J. 
Chem.  Soc.  76,5237  (1951  . 

65.  Irving  and  Williams,  ./.  Chem.  Soc,  1963,  3192;  Nature  162,  746  (1948). 


178 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

22 


Fig,  4.3.  Relationships  between  formation  constants  of  metal  chelates  and  the 
second  ionization  potentials  of  the  metals.  #  Ethylenediamine;  O  8,  8',  8"  tri- 
aminotriethylamine;  ■  salicylaldehyde. 

base  reaction  between  ammonia  and  hydrogen  ion  with  the  similar  reaction 
between  ammonia  and  copper  (I)  ion. 
H  H 

(1)  H++  :N:H-*H:N:H+ 

U  H 

Acid       Base 

H  H 

(2)  Cu+  +  :N:H  ->  Cu:N:H+ 

H  a 

Acid       Base 

The  formal  analogy  is  apparent,  though  even  elementary  considerations 
suggest  that  the  ability  of  the  positive  ion  to  attract  electrons  will  be  in- 
fluenced by  many  characteristics  of  the  cation  such  as  charge,  size,  polariza- 


Acid-base  process 


Coordination  process 


ELECT  Hits  PAIR  BOND  AND  STRUCTURE 


L79 


9.5  100  1 1.0  12  0 

1/7  X  10-1 

Fig.  4.4.  Log  of  the  Stability  constants  of  the  rare  earth  complexes  of  ethylene- 
diamine  tetraacetate(64b)  as  a  function  of  reciprocal  of  the  empirical  radius84*. 
O  —  Potentiometric  data 
•   —  Polarographic  data  in  KNOj  n  =  0.1 
A  —  Polarographic  data  in  KC1  /i  =  0.1 

The  potentiometric  data  are  most  accurate  for  the  ions  La-Eu. 

The  polarographic  data  are  most  accurate  for  the  ions  Gd-Lu. 

bility,  screening  constants  and  other  properties  as  well  as  by  properties  of 
the  ligand.  In  view  of  the  formal  analogy,  a  correlation  between  the  basic 
strength  of  a  ligand  and  its  coordinating  ability  is  not  unexpected,  although 
one  could  hardly  hope  for  a  strict  parallelism. 

In  1928  Riley66  suggested  that  any  factor  which  increases  the  localiza- 
tion of  negative  charge  in  the  base  (coordinating  ligand)  makes  the  elec- 
trons more  readily  available  and  thus  increases  the  ability  of  the  base  to 


180  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

coordinate.  These  ideas  were  used  to  explain  a  number  of  phenomena.  It 
has  been  observed  that  sulfate  and  sulfite  ions  each  tend  to  occupy  a  single 
coordination  position  while  carbonate  preferentially  forms  a  four  membered 
chelate  ring  involving  two  coordination  positions.  Steric  factors  cannot 
explain  this  difference.  Riley  attributed  the  difference  to  a  tighter  binding 
of  the  electrons  on  the  sulfate  because  of  the  higher  nuclear  charge  on  the 
central  sulfur  atom.  Carbonate  ion  with  a  lower  nuclear  charge  on  the 
central  carbon  atom  supposedly  can  contribute  the  four  electrons  necessary 
to  form  two  coordinate  bonds  more  readily  than  can  the  sulfate  ion. 

Many  attempts  to  establish  a  linear  relationship  between  the  basic 
strength  of  a  ligand  (as  measured  by  its  pKH+  value)  and  the  complex  form- 
ing ability  of  the  ligand  (as  measured  by  the  logarithm  of  the  formation 
constant  of  its  metal  complexes)  have  been  recorded.  One  of  the  first  at- 
tempts was  that  of  Larsson67.  The  relationship  was  disputed  by  later 
workers68, 69,  but  it  now  seems  well  established  that  when  systems  of  suffi- 
cient structural  similarity  are  compared,  a  linear  relationship  between 
pKCompiex  and  pKbase  is  obtained.  Bruehlman  and  Verhoek70  found,  for 
example,  that  when  the  logarithm  of  the  first  association  constants  of 
silver-amine  complexes  are  plotted  against  the  pK  values  for  the  correspond- 
ing substituted  ammonium  ions,  two  straight  lines  are  obtained:  one  for 
the  pyridines  and  primary  aliphatic  amines  and  one  for  the  secondary 
amines.  Data  from  the  literature  indicate  that  tertiary  aliphatic  amines  lie 
on  a  third  curve.  The  slope  of  the  curves  (Fig.  4.5)  is  approximately  one- 
fourth,  indicating  a  much  smaller  range  of  basic  strengths  when  measured 
against  hydrogen  ion,  a  not  unexpected  observation. 

Bjerrum71  confirmed  the  linear  relationship  for  cyclic  amines  and  primary 
amines  and  extended  the  data  to  include  mercury (II)  complexes  as  well. 
The  data  of  Schwarzenbach  and  his  co-workers  on  the  stability  constants 
of  the  complexes  of  the  alkaline  earths  with  aminopolycarboxylic  acids 
show  a  similar  relationship  if  the  number  of  chelate  rings  formed  in  the 
structure  is  taken  into  account  (page  229). 

Calvin  and  Bailes72  in  1946  studied  polarographically  the  stability  of 
copper  chelates  of  the  form 


66.  Riley, ./.  Chem.  Soc,  1928,  2985;  Ives  and  Riley,  /.  Chem.  Soc,  1931, 1998. 

67.  Larsson,  Z.  physik.  Chem.,  A169,  215  (1934). 


ELECTROS  PAIR  HOM)  AM)  STRl  <  !  I  RE 


1S1 


Fig.  4.5.  Relationship  between  strength  of  the  base  and  its  ability  to  form  co- 
ordination compounds  with  silver  (I)  ion  (From  Ref.  70). 

(1)  Pyridine 

(2)  a-Picoline 

(3)  7-Picoline 

(4)  2,4-Lutidine 

(5)  /3-Methoxyethylamine 

(6)  Ethanolamine 

(7)  Benzj'lamine 

(8)  Isobutylamine 

(9)  Ethylamine 

(10)  Morpholine 

(11)  Diethylamine 

(12)  Piperidine 

where  A  represents  an  electron  attracting  group.  They  found  that,  in 
general,  the  greater  the  electron  attracting  power  of  A  the  greater  the  tend- 
ency to  remove  electrons  from  the  nitrogen  and  hence  the  lower  the  basic 
.strength  of  the  amine  and  the  stability  of  the  copper  complex.  The  stabilil  y 
of  the  compounds  varied  as  A  was  changed,  the  order  being 


\<>_  <  — S03Xa  < 


<  — H  <  — CH3  <  —OH  <  — OCH3 


More  recently  Fernelius  and  co-workers66  have  carried  out  intensive  in- 
vestigations on  the  coordinating  ability  of  diketones  of  the  type  EtCOCHr 
OOR'  where  the  nature  of  the  R  groups  was  varied  systematically.  They 
report  a  linear  relationship  between  the  Logarithm  of  the  formation  con- 

urgh  and  Cogswell,  J.  Am.  Chem.  Soc,  66,  2412    1943 
Britten  and  Williams,  •/.  CJu  m.  Soc.,  1935,  796. 

70.  Bruehlman  and  Verhoek,  /.  Am.  Chem.  80c.,  70,  1401  (1948  . 

71.  Bjerrum,  Chem.  ft      .  46,  381     I960  . 

72.  Calvin  and  Hail--.  /.An    <  >  ■       S        68,  953    1946 


182 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


stants  of  the  sodium  complexes  of  the  /3-diketones  with  aromatic  R  groups 
and  the  pK  values  of  the  diketones.  If  one  assumes  that  the  /3-diketones 
are  largely  in  the  enol  form,  the  following  represents  the  influence  of  the 
aromatic  R  groups  in  decreasing  the  basic  strength  of  the  diketone  ion. 


o->Lt>  cy) 


So- 


decreasing  BASIC  STRENGTH 


It  was  found  also  that  the  relationship  holds  for  many  metals,  the  slope 


13 


O    12 
O 


10 


3 

-f           7 

,-•''' 

12 


13 
PKHA 


Fig.  4.6.  Relationship  between  the  acid  pK  of  various  /3-diketones  and  the  ability 
of  the  diketones  to  coordinate  with  copper (II).  (Data  from  Ref.  56a) 
O  log  Ki  for  process  M++  +  Ke~  ->  MKe+ 
•  log  K2  for  process  MKe+  +  Ke~  -»  M(Ke)2 
Solid  lines  represent  ketones  with  two  aromatic  groups;  dotted  lines,  those  with 
one  alkyl  group. 


Compound 
1 

2 
3 

4 
5 
6 

7 
8 
9 


R 
phenyl 
2,  thenyl 
2,  furyl 
2,  thenyl 
2,  furyl 
2,  thenyl 
methyl 
phenyl 
silyl 


Ri 

phenyl 

phen}rl 

phenyl 

2,  thenyl 

2,  thenyl 

methyl 

methyl 

methyl 

methyl 


ELECT  la  i\  PAIR  BOND  AND  STRUCT!  rRE 


1  s:; 


Tablk  4.4.  Stability  Constants  01  Substituted 
M  vlonatb  Complexes  of  Copper 


Acid  Constants  of  Ma  Ionic  Acids 

—log 

R 

R' 

k oomplex  ilissoc 

pK, 

pKi 

pKi  - 

(larger  value  = 
:  stability) 

H 

II 

2.75 

5.36 

8.11 

8 

Me 

H 

2.97 

5.46 

8.43 

8 

Et 

H 

2.90 

5.55 

8.45 

8 

n-Pr 

H 

2.97 

5.68 

8.65 

8 

i-Vv 

11 

2.93 

5.80 

8.73 

9 

Me 

Me 

3.08 

5.82 

8.90 

5.4 

Et 

Et 

2.24 

7.23 

9.47 

5 

n-Pr 

n-Pr 

2.06 

7.48 

9.54 

5 

A  similar  situation  was  observed  qualitatively  by  Bailar  and  Work  when  they 
observed  that  neopentanediamine,  NHoCHoCCCHs^CHoNF^  ,74  coordinates  more 
readily  and  gives  more  stable  compounds  than  its  unsubstituted  analog,  trimethyl- 
enediamine. 


of  the  line  becoming  greater  with  more  electronegative  metals.  A  second 
linear  relationship  was  obtained  for  those  ligands  in  which  R  is  an  aromatic 
group  and  R'  an  alkyl  group.  Representative  data  for  their  copper  com- 
plexes are  shown  in  Fig.  4.6.  In  general  the  /3-diketones  containing  two  aro- 
matic rings  form  more  stable  chelate  compounds  than  those  containing  one 
aliphatic  group.  This  difference  is  greater  for  the  second  ligand  than  for 
the  first.  Among  the  alkyl  groups  studied  were  CH3 — ,  C2H5 — ,  (CH3)2CH — , 
(CH3)3Si(:iI,CII2— ,  (CH3)3C—  and  F3C— .  As  might  be  expected  from 
the  inductive  effect,  the  trifluoromethyl  group  reduces  the  basic  strength 
of  the  ligand  very  markedly. 

Two  rather  anomalous  observations  on  electronic  effects  merit  brief 
consideration.  Riley66  studied  the  stability  of  copper  complexes  of  substi- 
tuted malonic  acids  of  the  type  CHR(COOH)2 .  He  found  that  if  R  is 
methyl, 'ethyl,  or  normal  propyl,  the  resulting  complex  is  slightly  less  stable 
than  if  R  is  hydrogen.  On  the  other  hand,  if  R  is  an  isopropyl  group  the 
complex  is  reportedly  much  more  stable  than  when  R  is  hydrogen.  For 
malonate  ions  of  the  type  RR'C(COO)2=  the  resulting  complex  is  much  less 
-table  than  when  only  a  single  group  is  present.  Ethyl  and  propyl  have  a 
bigger  effect  than  methyl  in  reducing  stability,  which  implies  steric  factors 
or  solvation  factors  with  the  disubstituted  compound. 

A-  the  data  in  Table  1.1  show,  the  stability  constants  for  the  copper 
complexes  do  not  parallel  the  pK  values  tor  these  acids71.  The  role  of  the 
isopropyl  group  appears  to  be  anomalous  in  this  case. 

73.  Gane  and  Ingold,  ./.  Cht        -         1929,  1698. 

74.  Bailar  and  Work,  ./.  Am.  Chem.  80c. ,  68,  232  (1946). 


184  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Stabilization  of  Valence  by  Coordination.*  A  number  of  attempts 
have  been  made  to  justify  by  the  atomic  orbital  theory  the  fact  that  co- 
ordination can  stabilize  both  common  and  uncommon  valence  states  of  a 
metal.  For  example,  the  relative  stabilities  of  the  2+  and  3+  states  of 
cobalt  have  been  explained  repeatedly  in  terms  of  atomic  orbital  theory. 
The  cobalt(II)  cyanide  complex  is  so  unstable  that  it  reduces  water  with 
i  he  liberation  of  hydrogen,  while  the  hydrated  cobalt(III)  ion  is  so  unstable 
that  it  will  liberate  oxygen  from  water  (see  page  185).  Pauling  suggested27*1 
that  these  facts  may  be  explained  from  a  consideration  of  orbitals  available 
in  the  cobalt  cyanide  complexes.  The  hexacovalent  cobalt(II)  cyanide  is 
represented  as: 

fco(CN)T|  '  _J*«LIL~ '{fJApIT\i± 

In  order  to  free  two  d  orbitals  for  complex  formation,  the  seventh  d  elec- 
tron in  the  cobalt (II)  ion  is  promoted  to  a  higher  energy  level  where  it  is 
easily  lost  to  give  the  cobalt(III)  complex.  Two  arguments  may  immedi- 
ately be  raised  against  such  a  simple  explanation.  First,  it  is  known  that 
the  hydrated  cobalt(III)  ion  is  also  diamagnetic;  hence,  it,  like  the  cyanide, 
should  have  little  tendency  to  pick  up  the  electron  in  the  excited  level  to 
permit  reversion  to  cobalt (II).  This  is  a  contradiction  of  fact.  Second, 
Adamson77  has  presented  evidence  to  indicate  that  the  cobalt(II)  cyanide 
complex  is  actually  pentacovalent  [Co(CN)5]~;  hence,  the  necessity  to  free 
the  sixth  orbital  by  promotion  of  an  electron  is  eliminated.  Without  the 
promoted  electron,  the  argument  loses  its  validity. 

In  general  terms,  one  can  say  that  the  oxidation  state  which  is  lowest  in 
energy  will  be  most  stable.  Obviously,  then,  any  comprehensive  explanation 
must  involve  consideration  of  all  of  the  terms  which  contribute  to  the  energy 
of  different  oxidation  states.  Sufficient  data  to  make  such  a  study  meaning- 
ful are  not  available,  but  a  number  of  empirical  rules  which  systematize 
many  oxidation  states  can  be  employed.  Usually,  more  than  one  factor 
must  be  considered  because  of  the  large  number  of  energy  terms  involved 
in  even  the  energy  of  coordination.  For  this  reason  the  treatment  is  only 
an  approximation. 

The  "anomalous"  oxidation  states  of  the  rare  earths  can  be  systematized 
by  assuming  that  certain  electronic  configurations  such  as  an  empty /level, 
a  half  full ./'  level,  or  a  full  /  level  will  be  stable.  The  same  type  of  argument 
is  utilized  here.  The  following  postulates  are  made:f 

*  Sec  also  ( !hapter  2. 

I  The  authors  :ir<'  indebted  to  Dr.  Daryle  Busch  for  many  helpful  suggestions  in 
outlining  this  set  of  generalizations. 
77.  Adamson,  ./.  Am.  Chem.  Soc.,  73,  5710  (1951). 


ELECTRON  PAIR  BOND  AND  8TRUCTI  RE  L85 

(1)  Stable  electronic  configurations  for  the  central  metal  ion  are: 

a.  a  half  filled  shell,  as  in  the  iron(III)  ion. 

1).  a  completely  filled  d  level,  as  in  covalent  ironi  1 1 1  and  cobalt  I  III). 

c.  halt*  filled,  unused  d  orbitals  which  are  left  afterhybridization  to  obtain 

the  bonding  orbital-   e.g.,  I  'r  '  ■  and  Y++). 


Cr+  +  +  ! 

<-<<,'.       U       IHHI 
UNUSED    d|               2      3 
ORBfTALS    : A  £& ' 

2    [f  the  elect  ronegativity  of  the  bonding  atom  in  the  ligand  is  high*  so 

that  ionic  bonds  are  favored,  thai  valence  of  the  central  metal  which  involves 
the  half  filled  d  shell  (postulate  la)  or  the  ionic  state  of  maximum  multi- 
plicity is  usually  favored. 

(3)  If  the  electronegativity  of  the  bonding  atom  in  the  ligand  is  lowf  so 
that  covalent  bonds  are  favored,  the  valence  with  either  completely 
filled  d  levels  or  half  filled  unused  d  levels  is  stabilized  (postulate  lb  or  lc). 

A  number  of  examples  may  be  used  to  illustrate  the  applications  of  the 
above  postulates: 

(a)  Co(XH3)6++  ^±  Co(XH3)6+++  +  e~  E°  =  -0.1  v. 

Since  Co+++  cannot  achieve  a  half  filled  d  shell  and  since  \H.;  forms  bonds 
which  are  quite  covalent  in  character,  stable  struct  tire  lb  (completely  filled 
d  levels)  is  obtained  to  stabilize  the  Co+++  state. 

(b)  Co(H20)6++  ^  Co(H20)6+++  +  e-  E»  =  -1.84  v. 

Water  has  little  tendency  to  enter  the  covalent  state,  so  the  ionic 
cobalt (II)  state  is  obtained  without  achieving  any  of  the  preferred 
structures. 

(c)  [CuI2(H20)2l-  ;=±  [CuI2(H20)2l  +  e~     E°  =  lakge  negative  value 

+                  3d            J   4s          4p 
Cu  — , 

innnnu  i 

i j 

Since  iodide  forms  strongly  covalent  bonds  (easily  polarized),  the  -tinc- 
ture giving  a  full  d  shell  or  a  Cu+  state  is  favored. 

(d)  [Cu(H20)4l+^  [Cu(H20)4]++  +  e-  E»  =  -0.153  v. 

Water  is  too  electronegative  to  form  strongly  covalent  bonds;  hence,  (  u" 
is  more  stable.  The  same  appears  to  be  true  of  ammonia  and  ethylenedi- 

*  An  alternative  statement  is:  "If  the  Ligand  is  of  low  deformability  .  .  ."  See 
125  for  discussion. 

t  An  alternative  statement  is:  "It"  the  ligand  is  of  high  deformability  .  .  ." 


186  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

amine  complexes,  the  Cu — NH3  bond  being  less  covalent  than  the  Co — NH3 
bond. 

(e)  [Cu(CN)2]-  ;=±  Cu++  +  2CN~  +  e~  E°  =  -1.1  v. 

Since  cyanide  prefers  covalent  bond  formation,  the  Cu+  state  is  favored. 
Coordination  of  nitrites  or  sulfur  compounds  appears  to  give  a  similar  re- 
sult. 

(f)  [Cr(H20)6]++  ^  [Cr(H20)6]+++  +  e"  E°  =  0.414  v. 

Although  water  should  not  be  expected  to  form  strong  covalent  bonds  on 
the  basis  of  electronegativity,  the  strong  covalent  character  of  the  water- 
Cr  bond  is  indicated  by  the  slow  rate  of  exchange  between  coordinated  and 
solvent  water.  On  this  basis  the  covalent  state  Cr4-1"1"  defined  by  lc  is 
favored. 

(g)  [Cr(CN)6]4-  ^  [Cr(CN)6]3  +  e"  E°  =  1.28  v. 

Qr+++  should  be  stabilized  here  more  than  in  the  corresponding  case  of 
water  since  CN~  forms  bonds  of  greater  covalent  character.  This  is  ob- 
served. 

(h)  [Cr(NH3)6]++  ^±  [Cr(NH3)6]+++  +  e"  #°  =  ? 

Although  the  potential  for  this  couple  is  not  known,  one  would  predict 
that  it  lies  at  about  0.7  v,  between  that  for  the  aquated  chromium  system 
and  the  cyanide  system. 

(i)  V(h2o)++  ^  V(H2o)+++  +  e-  E«  =  0.255  v. 

The  structure  of  Y++  is : 

v++ 3d_  4f_  _4p 

111 

Since  water  does  not  form  strongly  covalent  bonds  with  V2+  there  is  no 
advantage  to  the  2+  state  as  opposed  to  the  3+  state. 

(J)  [V(CN)6r  ^  tV(CN),r  +  e-  E°  =  ? 

The  E  value  for  this  system  is  not  known,  but  the  possibility  of  stabilizing 
a  half  filled,  unused  d  shell  by  covalent  bond  formation  on  V2+  would  sug- 
gesl  that  [V(CX)6]4_  should  be  stabilized  with  respect  to  the  [V(CN)6]S 
state.  The  above  potential  would  be  more  negative  than  that  for  the  aquated 
system;  qualitative  data  indicate  that  such  a  potential  is  reasonable78. 
The  ability  of  the  vanadium (II)  ion  to  form  a  stable  complex,  [V(dipy)3]~H~, 
as  against  a  less  pronounced  tendency  by  the  vanadium(III)  ion  would  also 
be  expected  from  the  above  treatment.  King  and  Garner79  report  that  this 

7s.  Reference  21,  p.  806;  Taube,  Chem.  Revs.,  50,  69  (1952). 
79.   King  and  Garner, ./.  Am.  Chem.  Soc,  74,  3709  (1952). 


ELECT lio\  PAIR  BOND  AND  STRUCTURE 


187 


difference  is  so  strong  thai  Y++  and  V+++  can  be  separated  quantitatively 

in  aqueous  solution  by  complexing  the  V  '  ;  and  then  precipitating  the  V+++. 
(k)  Arguments  similar  to  these  have  been  employed  quite  successfully  in 
correlating  the  oxidation  states  of  nickel.  Nickel  IS  normally  divalent,  but 
Jensen80  found  that  the  complex  [XiBiv  JP^H^hH  can  be  oxidized  to  give 
pentacovalent  nickel(III).  The  electronic  configuration 


3d 


4s        4p 


correlates  with  the  observed  paramagnetism  equivalent  to  one  unpaired 
electron. 

To  form  hexacovalent  nickel  by  d2sp*  hybridization,  a  d  electron  would 
have  to  be  promoted  to  a  4d  or  5s  level.  The  5s  level  has  been  suggested  as 
the  preferred  lower  energy  level42*  ■ 81.  Such  promotion  would  lead  to  easy 
oxidation  of  nickel  to  the  4+  state  if  the  six  covalent  bonds  were  very 
strong.  Xyholm42  reports  the  complex, 


r^> 


^^ 


CH^    XCH3 


CI, 


containing  tetracoordinate  nickel(II);  this  can  be  oxidized  to  NiClr 
2  diarsine.  The  structure  proposed  for  the  latter  compound  is 

[Ni(diarsine)2Cl2]Cl. 

Since  hexacoordinate  covalent  nickel  is  present,  one  d  electron  was  probably 
promoted  to  a  five  s  level  and  should  be  easily  lost.  Such  a  hypothesis  re- 
ceives support  from  the  fact  that  the  complex  may  be  oxidized  to  nickel  (IV) 
complexes;  furthermore,  the  magnetic  moment  of  the  nickel(III)  compound 
corresponds  to  one  unpaired  electron  with  little  spin  contribution,  a  fact 
expected  from  an  odd  electron  in  an  s  state.  Xyholm42a  has  pointed  out 
that  if  use  is  made  of  a  ligand  of  low  electronegativity  which  forms  very 
stable  covalent  complexes  and  if  the  metal-ligand  bonds  are  sufficiently 
strong,  other  examples  of  nickel  (III)  and  nickel(IV)  compounds  might  be 
observed,  provided  the  coordination  number  is  expanded  to  five  or  six. 

It  is  evident  that  the  metal  is  as  important  as  the  ligand  in  determining 
the  degree  of  covalent  character  and  the  strength  of  the  metal-ligand  bond. 
(This  is  also  evident  from  the  field  splitting  treatment  of  magnetism  using 

80.  Jensen,  Z.  anorg.  nllgem.  Chem.,  229,  265  ,1936). 

81.  Burstall  and  Xvholm,  ./.  Chem.  Soc,  1952,  3570. 


188  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

the  ionic  model,  page  132.)  For  example,  water  is  able  to  form  stable  co- 
valent bonds  with  chromium,  unstable  covalent  bonds  with  cobalt,  and 
apparently  very  unstable  bonds  with  copper(II).  An  even  more  striking 
example  is  found  in  the  case  of  complex  fluorides.  Klemm  and  Huss82  pre- 
pared the  following  complex  fluorides:  K3FeF6 ,  K3C0F7 ,  K2NiF6 ,  and 
KsCuFe .  The  magnetic  moments  of  the  iron  and  cobalt  complexes  indicate 
an  ionic  type  of  bond,  the  ionic  structure  for  the  iron(III)  and  cobalt (IV) 
being  particularly  favored  by  the  half  filled  d  shell.  It  is  surprising,  how- 
ever, that  the  nickel  (IV)  in  K2NiF6  was  found  to  be  diamagnetic,  indicat- 
ing covalent  Ni — F  bonds.*  The  corresponding  K2PtF6  is  also  diamagnetic. 
It  is  of  interest  that  fluorine  and  oxygen  can  stabilize  unusually  high  valence 
states  such  as  Co4+,  Ni4+,  and  Fe6+. 

As  might  be  expected  with  a  topic  of  this  complexity,  any  set  of  valence 
generalizations  is  apt  to  produce  inconsistencies.  For  example, 

(a)  [FeF6]4"  ^±  [FeF6]s  +  e~  E°  =  >  -0.4  v. 

The  half  filled  d  level  in  Fe(III)  and  the  small  tendency  for  covalent 
character  in  Fe — F  bonds  should  stabilize  the  Fe+++  state.  This  is  roughly 
true. 

(b)  [Fe(H20)6]++  ^  [Fe(H20)6]+++  +  e"  #°  =  -0.771  v. 

Water-metal  bonds  are  likewise  ionic  but  less  so  than  fluorine-metal 
bonds;  so  the  trivalent  state  here  should  be  somewhat  less  stable  than  in 
the  case  of  the  fluoride  complex.  This  is  also  roughly  true;  but 

(c)  [Fe(CN)6]4-  ^±  [Fe(CN)6]3  +  er  E°  =  -0.36  v. 

The  metal-cyanide  bonds  should  be  strongly  covalent  and  should  favor 
the  Fe(II)  state  with  the  structure 


IHMHUU      U      IHiO 


as  compared  to  the  Fe(III)  state  with  the  structure 


i  - 

nOT!7T77    T7   rTTTu 


The  electrode  potential  indicates  that  the  ferricyanide  is  more  stable  (i.e., 
poorer  oxidizing  agent)  than  the  corresponding  [Fe(H20)6]+++  ion,  in  direct 

82.  Klemm  and  Huss,  Z.  anorg.  allgem.  Chem.,  258,  221  (1949);  262,  25  (1950) ;  Natur- 
wissenschaften,  37,  175  (1950). 
*  An  alternative  treatment  of  these  facts  can  be  given  by  the  crystal  field  splitting 
method  described  on  page  132. 


ELECTRON  IWlli  HuSD  AM)  STL'CCTURE  189 

contradiction  to  theory: 

[Fe(H20)6]+++  +  [Fe(CN),]*-  ;=±  [Fe(H20)6]++  +  [Fe(CN),]-        E»  =  +.41  v. 

Pauling81  has  attempted  to  explain  this,  bu1  he  appears  to  have  the  facte 
reversed.  He  states,  "The  interesting  tact  that  the  ferrocyanide  ion  is  less 
easily  oxidized  to  the  ferricyanide  ion  than  is  the  hydrated  ferrous  ion  to 
the  hydrated  ferric  ion  can  now  be  explained."  His  explanation,  based  on 
double  bonds,  attributes  enhanced  stability  to  the  ferrocyanide.  From  the 
potentials  given  by  Latimer84,  it  is  apparent  that  ferrocyanide  is  more 
easily  oxidized  to  ferricyanide  than  hydrated  ferrous  ion  is  to  ferric  ion. 

(d)  [Fe(ophen)3]++  ^±  [Fe(ophen)3]+++  +  e~  E*  -  -1.12  v. 

If  it  is  assumed,  as  seems  logical,  that  the  Fe — ophen  bonds  are  strongly 
covalent,  the  iron (II)  state  would  be  expected.  (See  electron  diagram 
above.)  This  is  an  agreement  with  fact.  Similar  arguments  explain  the 
system 

[Fe(dipy)3]++  ^  [Fe(dipy)3]+++  +  e~  E°  =  -1.096  v. 


JN     N     . 

ORTHOPHENANTHROLINE  tf,  C*-DIPYRIDYL  CONJUGATED    SYSTEM 

INVOLVING     METAL - 
LIGAND    DOUBLE   BOND 

STABILIZE      Fe(H) 


a. 

Fe'  Ve  —  NH2 


N^  *N-NHN 


0<  -  PYRIDYLPYRROLE  0<-PYRIDYLHYDRAZINE 

STABILIZE     Fe(ir) 
PlO.  4.7.  Heteroc\clic  coordinating  agents  and  the  oxidation  states  of  iron 

On  the  other  hand,  the  cases  of  the  tris  a-pyridylhydrazine,  the  tria  a- 
pyridylpyrrole,  and  the  /3-diketone  complexes  of  iron  are  not  so  obvious. 
Electrode  potential  data  for  these  systems  are  not  available,  but  the 
iron(III)  state  is  supposedly  stabilized  strongly  by  these  ligands.  The  reason 
for  a  big  difference  in  the  ability  of  the  nitrogen  in  these  ligands  to  form 
covalent  bonds  as  compared  to  the  nitrogen  in  orthophenanthroline  and 
dipyridyl  is  not  immediately  obvious.  The  possibility  of  forming  multiple 
metal-ligand  bonds  with  the  nitrogens  of  both  aromatic  rings  is  probably 

83.  Pauling,  /.  Ch       So      1948,  1461. 

84.  Latimer,  "Oxidation  Potentials,"  2nd  Ed.,  New  York,  Prentice-Hall,  1952. 


190  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

important  in  the  orthophenanthroline  and  dipyridyl  systems.  In  a-pyridyl- 
hydrazine  and  a-pyridylpyrrole  only  one  nitrogen  is  part  of  an  aromatic 
ring  system,  so  the  possibility  of  resonating  metal-ligand  double  bonds  on 
both  nitrogens  is  reduced.  This  is  seen  by  reference  to  the  structural  formu- 
las in  Fig.  4.7.  On  the  other  hand,  the  /3-diketones  might  logically  be  ex- 
pected to  form  more  ionic  bonds  than  orthophenanthroline  since  coordina- 
tion is  through  the  more  electronegative  oxygen  atom  and  the  complex  is 
paramagnetic.  The  stability  of  the  3+  state  here  is  not  surprising. 

(f)  Several  unusual  oxidation  states  of  silver  pose  rather  vexing  prob- 
lems, particularly  in  view  of  the  conclusions  about  the  strong  covalent 
bond-forming  power  of  orthophenanthroline.  Silver  has  an  outer  electronic 
structure  similar  to  that  of  copper;  hence,  strongly  covalent  ligands  might 
be  expected  to  give  a  stable  silver  (I)  state  for  tetracoordinate  or  bicoordi- 
nate  covalent  derivatives. 


+  __4d [5s     5p^ 

A9  THHHtfij J 

AVAILABLE  FOR 
Sp3OH  LINEAR 
HYBRIDIZATION 

Actually  orthophenanthroline  and  dipyridyl,  which  form  very  stable  co- 
valent bonds  in  the  iron  system,  give  stable  complexes  of  silver(II)  such  as 
[Ag(ophen)2]++  and  [Ag(dipy)2]++.  The  reason  why  such  divalent  tetra- 
coordinate silver  complexes  should  be  stable  is  not  immediately  obvious 
from  the  preceding  set  of  rules. 

Ionic  and  Multiple  Bonds  Between  the  Metal  and  the  Ligand. 
The  Principle  of  Electro-neutrality.  The  concept  of  the  coordinate  bond 
appears  simple  enough,  yet  more  careful  scrutiny  of  the  nature  of  these 
bonds  from  the  standpoint  of  electron  distribution  and  bond  polarities  led 
to  difficulties85,86,87,88  in  interpretation  which  are  not  yet  entirely  re- 
solved. 

In  normal  covalent  bond  formation  in  which  each  of  two  atoms  shares 
one  electron  with  the  other,  no  considerable  electrical  disturbance  should 
result;  if  the  pair  of  electrons  were  equally  shared,  there  should  be  no  re- 
sulting dipole.  However,  the  situation  is  somewhat  altered  in  the  case  of 
coordinate  bond  formation.  In  this  instance,  one  atom  gains  and  the  other 
atom  loses  a  share  in  two  electrons;  consequently,  the  acceptor  atom  gains 
in  nH  negative  charge  and  the  donor  atom  gains  in  net  positive  charge85. 

85.  Sidgvvick,  Chemistry  &  Industry,  46,  803  (1927).  Reference  5b,  pp.  71  and  122. 

86.  Lou  rv,  Chemistry  &  Industry,  42,  412  (1923). 

87.  Sidgwick,  Trans.  Faraday  Soc,  19,  473  (1923). 

88.  Lowry,  Chemistry  &  Industry,  42,  715  (1923);  Sidgwick,  Ann.  Reports,  1934,  38; 

Hunter  and  Samuel,  Chemistry  &  Industry,  1935,  635;  Mathieu,  Compt.  rend., 
215,  325  (1942);  Reference  5b,  p.  121. 


ELECTRON  PAIR  BOND  AND  STRUCTURE  L91 

This  is  implied  by  Sidgwick's  arrow.  .1  — >  />,  where  .1  is  the  donor  atom 

and  B  is  the  acceptor.  Lowry88  indicated  this  by  plus  and  minus  > i ll 1 1 - .  as 

- 
A— B. 

Of  direct  interest  is  the  fact  that  the  above  logic  would  seem  to  call  for 
an  accumulation  of  negative  charges  on  the  central  atom  of  coordination 

compounds  an  unaccustomed  concept  for  metallic  element-  traditionally 
considered  as  electropositive  in  character. 

In  modern  theory  the  problem  has  been  considered  in  two  more  or  Less 
complementary  ways:  (1 1  by  assuming  the  formation  of  double  (or  triple) 
bonds  in  which  unused  (/  electron  pairs  of  the  metal  are  donated  hack  to 
the  ligand,  and  (2  I  by  assuming  an  ionic  contribution  to  the  bond  such  that 
the  negative  charge  on  the  ion  is  reduced.  Pauling  has  expressed  the  opinion 
that  this  charge  transfer  takes  place  until  each  atom  has  essentially  zero 
residual  charge.  He  has  expressed  this  formally83- 90  as  the  postulate  of  th 
electrical  neutrality  of  atoms;  namely,  "that  the  electronic  structure  of  suit- 
stances  is  such  as  to  cause  each  atom  to  have  essentially  zero  resultant 
electrical  charge,  the  amount  of  leeway  being  not  greater  than  about  =b  V2, 
and  these  resultant  charges  are  possessed  mainly  by  the  most  electroposi- 
tive and  electronegative  atoms,  and  are  distributed  in  such  a  way  as  to 
correspond  to  electrostatic  stability."  Data  on  x-ray  K  absorption  edges 
for  complexes  of  Cr,  Mn,  Fe,  and  Ni91  have  been  interpreted  as  supporting 
the  principle  of  electrical  neutrality. 

Multiple  Bonds.  Multiple  bonds  can  arise  in  those  cases  in  which  the 
entering  ligand  can  act  as  an  electron  acceptor  as  well  as  an  electron  donor. 
Cyanides,  carbonyls,  and  other  groups  containing  first  period  elements 
joined  to  other  atoms  by  multiple  bonds  can  serve  as  such  acceptors  by 
virtue  of  their  own  double  bonds.  In  addition,  recent  work  suggests  that 
second  period  elements  such  as  phosphorus  and  sulfur  may  be  joined  to 
the  metal  by  double  bonds  if  3d  orbitals  in  these  atoms  are  used  to  receive 
the  electrons  from  the  metal92.  The  carbonyls  and  cyanides  have  been  ex- 
tensively considered  by  many  workers.  On  the  basis  of  the  hybridized  orbi- 
tal treatment  as  applied  to  Ni(CO)4 ,  the  nickel  atom  contains  5  unshared 
3<7  electron  pairs: 


3d           '4s         4p 
in        Ni(CO).  J    — i 

4       ihumhi-v^^     i 


Lowry,  Tram.  Faraday  Soc. ,  19,   188    I 

Pauling  in  Victor  Henri  Memorial  Volume,  "Contribution  to  the  Study  of  Mo 
lecular  Structure,"  p.  1.  Liege,  Desoer,  1947. 

Mitchell  and  Beeman,  J.C)  ,  20,  1298    1952 

Ch.-itt  and  William.^  /•  Chem.  Soc,  1951,  3061;  Chatl  L66,  L9fi0  ; 

-    :kin  and  Dyatkina,  ./.  Gen.  Chem  .  I    8  8  S  .  16,  345  (1946). 


194  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

platinum.  The  unusual  stability  of  Cr(CN)6~  (comparable  to  iron  cyanides) 
is  not  amenable  to  such  a  treatment.  Since  the  chromic  ion  has  three  un- 
paired electrons, 


Cr+^+      _£f_|__      4s     _±P_ 


1    K    \\ 


d  sp 


the  formation  of  double  bonds  is  improbable,  and  the  entire  elimination  of 
charge  from  the  central  metal  is  usually  assumed  to  take  place  through 
resonance  with  ionic  forms100  (page  208).  Similarly,  the  stability  of  the  com- 
plex ions  Mo(CN)s~  and  Mo(CN)84-  and  their  tungsten  analogs  cannot  be 
attributed  to  double  bond  formation  because  of  the  small  number  of  d 
electrons.  Pauling  suggests  that  these  complexes  likewise  involve  single 
covalent  bonds  with  some  ionic  character  which  transfers  the  negative 
charge  from  the  central  atom  to  the  attached  groups. 

It  is  interesting  that  many  of  the  donor  atoms  which  show  strong  com- 
plexing  tendencies  and  which  stabilize  unusual  oxidation  states  are  potential 
electron  acceptors  as  well  as  electron  donors.  Among  these  are  the  tertiary 
phosphines  and  arsines,  cyanide,  nitrite,  and  molecules  containing  aromatic 
nitrogen  such  as  orthophenanthroline  and  a,a'-dipyridyl.  On  the  other 
hand,  it  is  difficult  to  seriously  attribute  the  stability  and  other  properties 
of  their  complexes  to  double  bond  formation,  since  available  data  indicate 
that  these  same  properties  are  displayed  by  structures  such  as  the  chromi- 
cyanide  and  molybdenum  cyanide  in  which  the  possibility  of  double  bond 
formation  is  absent.  Further,  electrode  potential  data  indicate  that  the 
ferricyanide,  in  which  only  two  double  bonds  are  possible,  is  more  stable 
than  the  corresponding  ferrocyanide  in  which  three  double  bonds  can  be 
formed.  This  lack  of  correlation  between  the  properties  of  these  complexes 
and  the  ability  of  the  donor  metal  to  form  double  bonds*  must  be  regarded 
as  a  serious  weakness  in  the  concept. 

*  In  this  same  connection  Chatt92b  has  pointed  out  that  boron,  which  can  form  no 
double  bonds,  gives  much  weaker  complexes  with  carbon  monoxide  than  does  plati- 
num, which  can  form  double  bonds.  Qualitative  data  obtained  by  Lutton  and 
Parry1018  indicate  that  under  comparable  conditions  this  difference  is  not  as  large  as 
usually  assumed  since  even  [PtCl2CO]2  will  lose  carbon  monoxide  under  reduced 
pressure  at  room  temperature  to  give  black  residues;  hence  apparent  stability  differ- 
ences reflect  only  rates  of  decomposition.  Further,  the  stable  compound,  H3BP(Me)3  , 
has  been  reported102  to  melt  at  106°C  without  decomposition  and  to  withstand  tem- 
peratures  up  to  200°C,  indicating  a  stability  comparable  to  that  of  the  platinum 
phosphine  complexes.  On  the  other  hand,  Chatt  points  out  that  PF3  will  not  form 
complexes  with  boron  or  aluminum  compounds  but  will  complex  with  platinum — a 
fact  which  is  interpreted  as  offering  strong  support  for  his  argument.  Recently,  how- 
ever, the  compound  H3BPF3  has  been  prepared101b. 
100.  Ref.  95,  p.  375. 


ELECTRON  PAIR  BOND  AND  STRUCTV/:/:  L95 

In  a  separate  treatment  of  charge  distribution  in  complexes,  Syrkin  and 
Dyatkina101,  I"1  ' "■'■  started  with  somewhat  different  philosophical  assump- 
tions and  arrived  at  the  sain*1  picture  as  Pauling.  It  lias  been  suggested  that 
their  ideas  might   be  helpful  in  estimating  electronic  transitions  in  the 

molecule18.  The  concept  has  definite  limitations. 
Ionic    Structure.   For   complexes   containing   ammonia,   derivatives   of 

ammonia,  water,  hydroxy!  ion,  and  the  like,  it  is  not  possible  to  in\  <>ke  the 
double  bond  to  reduce  the  negative  charge  on  the  metal  ion  and  to  explain 
complex  stability,  tor  these  coordinating  groups  cannot  act  as  acceptors  of 
electrons.  Here,  the  2s  and  2p  orbitals  an1  full,  and  the  3s,  3p,  and  3d  orbi- 
tals  are  of  too  high  energy  for  bond  formation.  Paulmg  pointed  out  that  the 
usual  coordinating  groups  of  this  type  which  commonly  form  complexes 
with  the  iron  group  transition  elements  are  in  the  main  strongly  electro- 
negative in  character,  and  suggested  that,  because  of  this  property,  they 
are  able  to  remove  most  or  all  of  the  negative  charge  from  the  central  atom 
and  thus  stabilize  the  complex  without  converting  the  essentially  covalent 
structure  to  an  essentially  ionic  structure.  He  has  cited  as  possible  evidence 
for  this  argument  the  fact  that  the  iron  group  elements  tend  to  form  less 
stable  halide  complexes  as  the  electronegativity  of  the  halogen  decreases. 
For  example,  the  iodide  complexes  of  the  3d  elements  are  very  unstable. 

According  to  Pauling,  the  electropositive  character  of  the  4d  palladium 
and  5d  platinum  transition  elements  is  less  than  that  for  the  3d  series.  This 
difference  is  reflected  in  the  type  of  complexes  they  form.  The  metals  of 
the  palladium  and  platinum  series  not  only  enter  into  combination  with  all 
the  coordinating  groups  mentioned  in  connection  with  the  iron  group  ele- 
ments, but  they  also  form  stable  complexes  with  less  electronegative  groups 
such  as  iodide.  Since  it  is  assumed  that  the  metals  of  these  two  groups  have 
little  or  no  tendency  to  form  positive  ions,  but  prefer  to  remain  neutral  or 
even  become  negative,  some  of  the  negative  charge  may  actually  be  left  on 
the  central  metal  of  the  complex.  It  becomes  less  essential,  therefore,  ac- 
cording to  Pauling,  to  search  for  conditions  which  can  bring  about  reduction 
of  the  negative  charge  on  the  central  atom. 

The  Trans  Effect  in  Resonance  Theory.  An  explanation  of  the  trans 
effect  (page  146)  in  terms  of  the  ion-polarization  theory  was  given  in 
(  hapter  3.  It  was  noted  that  the  magnitude  of  the  trans  effect  in  a  series 

101.  button  and  Parry,  ./.  Am.  Chem.  Soc.,  76,  4271   (10.54);  Bissot  and  Parry,  un- 
published  results. 

I  Wagner,  /.  A  ,  75,  3872  (195.3). 

103.  Syrkin  and  Dyatkina..  Acta  /  n.  U.  /.'.  8.  8.,  20,  137,  273    1945 

Chap.  14. 

104.  VanVleck,/.  1, 177  (1933);2t20  (1934);Mullikan,  J  *hys., 

2,  7v.'  (1934);  Mofl  -Ion),  A202,  534,  548  '1050). 

105.  Ref.  95,  (a)  p.  371,  (b)  p.  383. 


L96  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

of  donors  increases  in  the  direction  of  decreasing  electronegativity106,  which 
parallels  the  direction  of  expected  increase  in  covalent  character  of  a  bond. 
Syrkin1"7  proposed  an  explanation  of  the  phenomenon  based  on  the  concept 
of  resonating  ionic  and  covalent  forms.  In  the  case  of  platinum(II)  com- 
plexes, Syrkin  suggested  that  the  actual  state  of  the  platinum  might  be 
intermediate  between  those  represented  by  structures  (A),  (B),  (C),  and  (D). 


X               X 

X              X 

x          x- 

X              X 

\_/ 

\-/ 

\ 

\+ 

Pt 

Pt 

Pt 

Pt 

/  \ 

/ 

/ 

X              X 

x          x- 

x          x- 

X             X 

(A) 

(B) 

(C) 

(D) 

Structure  (A)  involves  covalent  dsp2  hybridized  bonds;  structure  (B)  in- 
volves three  covalent  dsp  bonds  and  a  single  ionic  bond  (four  such  struc- 
tures would  contribute  toward  the  bonding  in  the  resultant  species) ;  struc- 
ture (C)  represents  two  covalent  ds  hybridized  bonds  and  two  ionic  bonds 
(four  structures  assumed);  and  (D)  represents  a  single  covalent  d  bond 
along  with  three  ionic  bonds  (four  structures).  When  all  the  coordinated 
groups  are  identical  (as  in  this  example)  the  various  permutations  of  bonds 
for  a  single  contributing  structure,  such  as  (B)  are  of  equal  weight.  How- 
ever, in  the  case  where  one  of  the  groups,  X,  is  replaced  by  a  group  Y, 
which  forms  bonds  of  a  higher  degree  of  covalent  character,  certain  of  the 
permutations  are  enhanced  or  minimized  in  importance.  Thus,  in  the 
complex  PtX3Y,  structure  (B)  has  three  of  its  forms  approximately  equiva- 
lent while  the  fourth,  that  involving  covalent  bonds  to  the  three  X  groups 
and  an  ionic  bond  to  the  Y  group,  is  minimized.  Similarly,  for  structure 
(D),  the  form  in  which  Y  is  bound  covalently  while  the  three  X  groups  are 
ionic  would  be  enhanced  in  its  importance.  According  to  the  changes  in 
importance  of  the  canonical  forms  represented  by  structures  (B)  and  (D), 
the  effect  of  substituting  Y  for  X  to  produce  PtX3Y  is  merely  to  weaken 
the  bonds  holding  the  X  groups.  However,  similar  treatment  of  the  struc- 
ture (C)  indicates  that  the  group  trans  to  Y  is  weakened  to  a  greater  extent. 
The  four  forms  of  structure  C  considered  are : 

x  x-       x  x        x-        x       x-        x- 

\  \  /  / 

Pt  Pt  Pt  Pt 

/  \       /  \ 

X  Y~         X-  Y"         X-  Y  X  Y 

E  F  G  H 

106.  Quagliano  and  Schubert,  Chem.  Rev.,  50,  246  (1950). 

107.  Syrkin,  Bull.  acad.  sci.  U.  R.  S.  S.,  Classe  set.  chim.,  1948,  69. 


ELECTRON  PAIR  BOND    iND  STRUCTl  RE  L97 

Since  V  tends  to  form  covalent  bonds  to  a  greater  extent  than  does  X,  forms 
( ;  and  II  will  be  favored.  From  this  picture,  it  is  apparent  that  the  bonds  of 

the  groups  X  which  are  cis  to  Y  are  strengthened  by  the  presence  of  Y, 

while  the  group  X  which  is  trans  to  Y  has  lost  in  covalent  character. 

Such  a  model  does  not  justify  the  strong  trans  effect  attributed  to  PFj  by 

Chattw,  since  three  fluorines  attached  to  the  phosphorus  might  be  expected 
to  increase  its  electronegativity  enough  to  minimize  its  strong  covalent 

bond-forming  tendencies.  In  addition,  if  such  ionic  resonance  forms  make  a 
major  contribution  to  the  structure,  the  rationalization  of  the  planar  ge- 
ometry becomes  more  difficult  in  atomic  orbital  theory.  Finally  the  reason 
for  neglecting  sp  hybridization  and  the  contributing  .structure 

X  x- 

\ 
Pt 

_         \ 
X  Y 

is  not  obvious.  Inclusion  of  this  structure  would  invalidate  the  argument. 
( )n  the  other  hand,  the  general  concept  of  charge  distribution  indicated 
by  all  structures  does  give  an  explanation  of  most  cases  of  trans  labiliza- 
tion  and  cis  stabilization.  The  unexpected  trans  influence  of  PF:5  mentioned 
above  has  not  been  proved  without  question  (see  p.  148);  hence,  it  cannot 
be  cited  as  a  completely  valid  objection.  Furthermore,  the  ability  of  fluorine 
to  reduce  the  covalent  bond-forming  power  of  phosphorus  has  not  been 
considered  on  a  quantitative  basis,  so  such  arguments  are  equivocal.  This 
then  represents  an  additional  approach  to  the  trans  effect. 

The  Molecular  Orbital  Approximation 

The  method  of  molecular  orbitals  was  conceived  and  developed  in  its 
early  years  largely  by  Hund,  Mulliken,  and  Lennard-Jones.  Though  the 
method  itself  is  as  old  as  the  Heitler-London-Pauling-Slater  atomic  orbital 
approximation,  its  extensive  application  to  coordination  compound.-  has 
occurred  only  in  very  recent  years,  largely  as  a  result  of  the  work  of  Len- 
nard-Jones, Coulson,  and  their  associates.  From  this  work  have  emerged 
valuable  ideas  relative  to  such  problems  as  the  structure  of  the  carbonyls 
Chapter  16),  coordination  through  the  ethylenic  double  bond  (Chapter 
1")  .  and  the  structure  of  the  metal  cyclopentadiene  complexes.*  An  ex- 
cellent non-mathematical  resume  of  the  results  of  the  molecular  orbital 
method  up  to  1947  was  given  by  Coulson108.  other  Qonmathematical  treat- 

*  The  coordination  Dumber  eight  for  Zr,  Mo,  Ru,  Ce,  Bf,  W,  I  >s,  and  Th  baa  been 

treated  l.v  Penney  and  And'  :iK  the  method  «»f  molecular  orbitals. 

108.  Coulson,  Quart.  Revs.,  1,  144  'HJ47j. 


198  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

incuts  have  been  given  by  Palmer109,  Bowen110,  Walsh111,  Emeleus  and  An- 
derson112  ,114,  and  by  later  workers  applying  the  ideas  to  specific  prob- 

lemgi;       ■     92,   LIS, 115, 

Probably  the  best  comparison  of  the  two  methods  is  in  Coulson's  out- 
standing book,  "Valence"23.  The  essential  mathematical  methods  as  well 
as  the  chemical  results  of  the  theory  are  summarized  in  a  fashion  which 
can  be  understood  by  both  the  mathematical  and  non-mathematical  reader. 
Mathematical  methods  are  available  in  books  on  quantum  mechanics25. 

In  general,  the  atomic  orbital  theory  assumes  that  through  the  hybrid- 
ization of  atomic  orbitals  a  new  set  of  directed  orbitals  is  obtained  (page 
164).  The  bond  between  groups  then  arises  from  the  overlap  of  one  of  the 
orbitals  of  this  set  and  the  bonding  orbital  of  the  coordinated  ligand.  In 
short,  a  highly  localized  bond  is  formed  involving  only  a  bonding  function 
from  each  of  the  two  groups  which  are  joined.  In  the  molecular  orbital 
theory  the  situation  is  quite  different.  The  bonding  orbitals  for  the  entire 
complex  group  (e.g.,  Ni(CN)4=)  are  involved  in  the  formation  of  each  bond. 
For  instance,  in  the  bonding  of  four  cyanide  ions  to  a  central  nickel(II)  ion, 
a  nonlocalized  set  of  molecular  orbitals  may  be  obtained  from  the  four  nickel 
orbitals  (c?sp2-hybridized,  if  necessary)  and  all  four  cyanide  groups.  It  is 
true  that  usually  the  orbital  of  one  cyanide  group  will  contribute  much 
more  heavily  to  a  given  bond  than  the  other  three  cyanides,  but  the  im- 
portant point  is  that  provision  is  made  for  all  to  contribute.  From  the 
physical  standpoint,  the  original  atomic  orbital  theory*  pictured  the  bond 
as  being  restricted  to  the  interaction  of  a  single  electron  pair;  in  contrast, 
the  molecular  orbital  method  assumes  that  a  pair  of  bonding  electrons  is 
not  confined  to  a  single  bond  but  participates  in  all  bonds.  A  necessary  con- 
sequence of  the  molecular  orbital  picture  is  that  the  bonds  will  all  be  inter- 
related and  changes  in  one  bond  will  be  propagated  to  all  other  links  in  the 
compound.  The  effect  produced  by  altering  one  bond  in  the  complex  is 
illustrated  by  "trans  elimination"  (page  204). 

One  may  also  consider  that  the  simple  atomic  orbital  representation  and 

*  The  above  description  of  the  Pauling  theory  is  not  representative  of  the  present 
day  version.  More  recent  modifications  introduce  ionic  contributions  and  resonance 
among  several  canonical  structures  to  account  for  nonlocalization  of  electrons273-  83. 
In  this  form,  it  approaches  the  original  molecular  orbital  treatment.  See  the  section 
on  ionic  structures  and  double  bonds  (pages  191  and  195). 

I  (Hi.  Palmer,  "Valency,  Classical  and  Modern,"  pp.  179-196,   London,  Cambridge 
University  Press,  1944. 

110.  Bowen,  Endeavor,  4,  75  (1945). 

111.  WatehtQuart.Rev8.,2f  73  (1948). 

112.  Ref.  L5c,  pp.  .",1-59. 

113.  Jaffe,  J.  Phys.  Chem.,  58,  185  (1954). 

ill.  Van  Yleck  mii. I  Sherman,  Rev.  Mod.  Physics,  7,  167  (1935). 

LIS.   I. <niK.nl  .Jones  and  Pople,  Proc.  Roy.  Soc.  (London),  210,  190  (1951). 


ELECTRON  PAIR  BOND  AND  8TRI  CT\  RE  L99 

the  extreme  ionic  viewpoint  are  really  .special  cases  of  the  molecular  orbital 

theory.  For  instance,  the  complex  ion  [Fe(CN)e]  may  be  represented  in 
molecular  orbital  theory  as  the  ionic  [Fc'  (  \  ,;  or  the  covalent 
[Fe"(CN)§]  or  as  any  structure  in  between,  depending  upon  the  relative 
sizes  of  three  arbitrary  coefficients  in  the  wave  equation.  The  intermediate 
state  is  achieved  in  the  atomic  orbital  system  by  introducing  the  concept 
of  "resonance."  That  is,  the  molecule  may  he  represented  by  the  super- 
position of  a  number  of  canonical  structures,  each  of  which  corresponds  t<> 
a  chemical  picture  of  localized  bonds  or  ions.  The  state  of  the  molecule  has 
properties  which  are  different  from  those  of  the  individual  canonical  struc- 
tures, but  can  be  represented  in  terms  of  a  set  of  structures.  Ionic  structures 
and  double  bonded  structures  are  utilized  to  remove  charge  from  the  cen- 
tral metal  (pages  191  and  195).  The  same  end  is  achieved  in  the  ionic  model 
by  the  introduction  of  polarization  terms  and  the  concept  of  the  crystal 
field  splitting  of  the  degenerate  d  levels  in  the  central  ion.  (See  Chapter  3.) 

Coulson103  has  differentiated  between  "localized"  molecular  orbitals 
which  resemble  the  atomic  orbital  picture,  and  the  "nonlocalized"  molecu- 
lar orbitals  described  above.  The  nonlocalized  orbitals  have  been  particu- 
larly useful  for  simple  systems  such  as  the  oxygen  and  nitrogen  molecules 
and  systems  of  conjugated  double  bonds  such  as  benzene.  On  the  other 
hand,  most  complex  systems  usually  demand  some  bond  localization  as  a 
simplifying  approximation. 

The  <r,  7r,  b  Designation  of  Molecular  Orbitals.  Bonding,  Anti- 
bonding  and  Nonbonding  Orbitals.  The  designation  of  molecular  or- 
bitals as  a,  7r,  or  5  has  arisen  in  both  atomic  and  molecular  orbital  theories. 
The  symmetry  of  bonds  with  these  designations  is  most  easily  seen  from 
a  brief  consideration  of  the  methods  for  combining  atomic  orbitals  to  give 
molecular  levels.  The  symmetry  of  the  individual  s,  p,  and  d  orbitals  has 
already  been  indicated  (Fig.  4.1).  It  is  usually  assumed  in  molecular  orbital 
theory  that  suitable  localized  molecular  orbitals  can  be  obtained  by  a  com- 
bination of  the  appropriate  atomic  functions.  Thus,  two  s  orbital  functions 
may  be  added  to  give  a  molecular  orbital  which  is  symmetrical  around  the 
intrmuclear  axis  and  which  concentrates  the  electronic  charge  between  the 
two  nuclei.  Such  an  orbital  is  known  as  a  a  bonding  orbital,  the  a  <1> 
nation  indicating  bond  symmetry  around  the  Internuclear  axis.  Alterna- 
tively, two  a  functions  may  be  subtracted  to  give  an  orbital  which  is  still 
symmetrical  about  the  internuclear  axis,  but  which  concentrates  the  charge 
away  from  the  space  between  the  two  nuclei  I  Fig.  L8  i.  This  is  known  as  a 
a  antibonding  level. 

In  contrast  to  a  bond.-,  the  combination  of  two  pt  or  two  pu  orbital-  to 
give  a  bonding  molecular  orbital  results  in  a  concentration  of  charge  in 
ribbon-shaped  streamer  above  and  below  the  internuclear  axis  (Fig.   I 


200 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


ATOMIC 
ORBITAL  1 

ATOMIC 
ORBITAL  2 

COMBINA- 
TION OF 
FUNC- 
TIONS 

APPROXIMATE 

FORM  OF  MOLECULAR 

ORBITAL 

: 
M.O. 
CLASSIFI- 
CATION 

O 

s 

© 

s 

Ys'Ys 

(         A        *        A         ) 

<5-s 

BONDING 

Ys"  Ys 

©CD 

<r  S 

ANTI- 

BONDINC 

OR 

Fig.  4.8.  Bonding  and  antibonding  a  molecular  orbitals  between  two  atoms- 
localized  bonds. 


Fig.  4.9.  Bonding  and  antibonding  x  orbitals  between  2  atoms — localized  bonds 


Since  such  an  orbital  is  not  symmetrical  around  the  bond  axis  and  since  it 
represents  a  component  of  angular  momentum  around  the  bond  direction 
equal  to  one,  it  is  known  as  a  w  orbital.  It  is  the  molecular  analog  of  the 
atomic  p  state.  (See  end  view,  Fig.  4.11,  for  analogy  to  atomic  p  orbital.) 
7r  bonds  can  also  be  of  antibonding  character  as  illustrated  in  Fig.  4.9. 

5  orbitals  are  of  relatively  rare  occurrence  in  most  systems.  The  formation 
of  a  8  bond  by  combination  of  two  dxy  bonds  along  the  z  axis  is  showrn  in 
Fig.  4.10.  From  the  end-on  view,  Fig.  4.11,  this  orbital  is  seen  to  have 
symmetry  similar  to  that  of  the  atomic  dxy  orbital,  and  hence,  has  a  com- 
ponent of  angular  momentum  equal  to  two  around  the  bond  direction.  This 
then  justifies  the  5  designation.  In  short,  molecular  orbitals  are  designated 


ELECTRON  PAIR  BOND  AND  STRUCTl  RE 

x 


201 


dxy     *■    dyy    ATOMIC  ORBITALS    IN    POSITION    TO    FORM 
J    MOLECULAR    ORBITAL     BY    APPROACH    DOWN     Z     AXIS 

Fig.  4.10.  8  Orbital  formation 


0~~    BOND 

Fig.  4.11.  View  of  <r, 
Unity  to  atomic  s,p,d. 


W  BOND  </     BOND 

and  5  molecular  orbitals  down  internuclear  axis.  Note  simi- 


as  a,  it,  8,  etc.,  accordingly  as  the  component  of  angular  momentum  around 
the  bond  direction  is  0,  1,  2,  . . .  etc.  If  the  electrons  in  a  given  orbital 
spend  most  of  their  time  between  the  nuclei,  the  orbital  is  termed  bonding; 
if  the  electron  is  restricted  in  its  movement  so  that  only  a  small  percentage 
of  its  time  is  spent  between  the  nuclei,  the  orbital  is  termed  antibonding; 
and,  finally,  if  the  electron  in  an  atom  is  not  disturbed  seriously  by  the 
presence  of  the  second  nucleus  (i.e.,  inner  core  electrons),  the  orbital  is 
termed  nonbonding. 

Application  of  Molecular  Orbital  Theory  to  Complex  Compounds. 
The  Compound  KJtu^ClvO-H-jO.  The  diamagnetism  of  the  compound 
K4Ru2ClioO-H20  which  contains  two  atoms  of  formally  tetravalent  ru- 
thenium has  already  been  mentioned  as  a  point  of  difficulty  in  the  atomic 
orbital  interpretation  (page  167  and  Fig.  4.2).  Dunitz  and  Orgel17  showed 
by  a  molecular  orbital  treatment  that  an  earlier  suggestion  of  Pauling  (men- 
tioned in  Ref.  32)  to  the  effect  that  "seven  orbitals  of  each  ruthenium  are 


202 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


used  in  bond  formation  of  which  two  on  each  ruthenium  are  used  in  double 
bond  formation  with  the  central  oxygen"  can  be  understood  from  a  molecu- 
lar orbital  treatment.  Actually,  however,  all  available  remaining  spd 
orbitals  of  ruthenium  must  be  considered  rather  than  just  seven.  Dunitz 
and  Orgel  assumed,  in  essence,  that  each  of  the  ten  chlorine  atoms  is  bound 
to  the  ruthenium  ion  by  a  a  bond.  They  then  obtained  non-localized  molecu- 
lar orbitals  for  the  Ru — O — Ru  system  involving  the  w  oxide  levels  and 
the  remaining  available  orbitals  of  the  ruthenium  ion.  The  transformation 
of  atomic  orbitals  into  the  appropriate  molecular  forms  is  indicated  sche- 
matically in  Fig.  4.12.  Each  molecular  orbital  may  be  made  from  half  of 
two  atomic  orbitals  (Eu  from  Px  and  Pxy)  or  from  a  single  atomic  orbital 
(Eg  from  Pyz).  The  total  number  of  molecular  orbitals  must  be  equal  to  the 
number  of  atomic  orbitals  used.  (The  symbolism  of  Eyring,  Walter,  and 
Kimball25  is  used.)  After  the  five  a  bonds  to  chlorine  and  one  <r  bond  to 
oxygen  are  formed  by  each  ruthenium  ion,  the  four  remaining  electrons  on 
each  ruthenium  ion  and  the  four  unused  tt  electrons  on  the  oxide  must  be 
placed  in  molecular  orbitals  which  are  shown  inside  the  dotted  line  in 
Fig.  4.12.  When  these  levels  are  filled  by  the  twelve  electrons  in  ac- 


Antibonding  Eu° 


Molecular  Ru-0-Ru  Levels 


Linear   SP 

FlO.   1.12.  Molecular  orbital  representation  of  diamagnetism  in  K2RU2CI10OH2O 


cordance  with  the  principle  of  maximum  multiplicity,  diamagnetism  is 
obtained.  The  double  bond  to  oxygen  from  each  ruthenium  is  then  con- 
tributed by  a  o-  Ru — O  bond  and  an  Eub  molecular  orbital  level.  The  Eub 


ELECTRON  PAIR  BOND    IND  STRl  CTl  RE  203 

orbital  may  be  described  as  a  double  degenerate  bonding  w  orbital.  The 
actual  extent  of  the  it  bonding  will  be  sensitive  to  the  relative  electronega 
tivities  of  the  atoms  concerned.  but  the  observed  Ru  0  distance,  L.80  A. 
is  close  to  the  value  L.74AfoundinRu04  ,a  fact  which  has  been  interpreted 
as  indicating  considerable  double  bond  character  in  the  Ru  0  interaction. 
It  is  also  clear  that  the  degree  of  bonding  and  hence  the  stability  of  the 
anion  would  be  diminished  by  any  departure  from  linearity  for  the  Ru  ( I 
Ru  system. 

The  molecular  orbital  explanation  of  diamagnetism  in  this  case  is  remi- 
niscent  of  its  similar  success  in  interpreting  the  paramagnetism  of  the 
oxygen  molecule116. 

In  Chapter  3  it  was  stated  that  the  quanticule  theory  of  Fajans  (page 
L32)  bears  a  resemblance  to  the  molecular  orbital  interpretation.  This  can 
now  be  seen  since  in  quanticule  terms  the  [Ru — O — Ru]64"  grouping  would 
be  considered  as  a  quanticule  to  which  ten  Cl~  ions  could  be  bound  through 
the  polarized  ion  concept.  After  considering  appropriate  polarization  terms, 
the  end  result  would  approach  quite  closely  the  above  molecular  orbital 
treatment,  even  though  the  starting  points  are  very  different. 

The  Compounds  [(XH3)5Co— 02— Co(NH3)5]X5  and  [(NH3)5Co— 02— 
Co(XH3)5]X4 .  The  linear  Co — 02 — Co  group  can  be  treated  analogously 
to  the  ruthenium  compound  except  that  the  peroxide  ion  now  has  both 
internally  bonding  Eu(\yt)  and  antibonding  Eg(vr)  orbitals  which  follow 
directly  from  the  treatment  for  molecular  oxygen.  It  follows  that  there  are 
twenty  electrons  after  a  bonding  to  place  in  molecular  levels  (i.e.,  six  elec- 
trons from  each  cobalt  and  eight  it  electrons  from  02=).  The  order  of  the 
molecular  levels  is: 

(EJ>)[(B2g)(B2u)(Eu«)(E0t>)](Ea°) 

The  relative  order  of  levels  inside  the  square  brackets  is  not  known.  The 
bonding  Egb  and  antibonding  Ega  levels  now  arise  from  interaction  of  the 
previously  described  Eg  metal  levels  (see  the  case  of  [Ru — 0 — Ru]6+)  with 
the  extra  -k  levels  of  the  02=  ion.  The  oxide  ion  had  only  two  unused  p  levels 
for  interaction  with  the  metal,  whereas  the  peroxide  ion  now  has  four  un- 
used 7r  levels,  giving  additional  interaction  possibilities.  Placing  the  twenty 
electrons  in  appropriate  levels  gives: 

(Ej>mB2g)HB2uy(Eu°y(E0by](E0«y 

Since  all  orbitals  are  filled,  diamagnetism  follow.-.  The  filling  of  both  the 
bonding  and   the   corresponding  antibonding   levels   indicates   that    the 

metal-Oi  bond  and  the  O — O  bond  Bhould  have  HO  double  bond  character. 

The  oxidation  of  [(XH3)5Co — O2 — Co  NH  ;XS  to  the  corresponding 
[(NHa)sCo — 02 — Co(NHs)i]Xf  must  involve  removal  of  an  electron  from 

116.  Lennard-Jones,  Trans.  Faraday  80c.,  25,  668    l *»29; . 


204  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

the  least  stable  orbital,  which  is  Ega,  and  presumably  centered  mainly  on 
the  — 02 —  grouping.  It  is  in  this  sense  that  one  would  attribute  the  electron 
loss  to  the  02=  rather  than  to  cobalt  (III).  The  0 — 0  group  would  then  re- 
semble the  superoxide  ion,  O2-;  the  preparation  of  the  compound  by  means 
of  alkali  metal  superoxides  might  be  suggested. 

The  Fe — Fe  interaction  in  metal  carbonyls  has  also  been  justified  by 
the  molecular  orbital  theory17. 

The  Paramagnetic  Resonance  of  IrCl6=.  Stevens117  has  recently  applied  the 
molecular  orbital  theory  to  a  discussion  of  details  in  the  paramagnetic 
resonance  absorption  spectrum  of  IrCl6=.  The  paramagnetic  absorption 
data  are  usually  interpreted  in  terms  of  an  ionic  model.  His  work  represents 
an  initial  attempt  to  formulate  orbitals  that  describe  some  deviations  from 
an  ionic  model  which  seem  to  be  required  by  details  of  the  spectrum. 

On  an  ionic  model,  the  complex  is  considered  to  be  a  central  iridium (IV) 
with  five  5d  orbital  electrons,  surrounded  by  a  regular  octahedron  of  Cl~ 
ions.  The  complex  shows  s  =  }^  and  g  =  1.8  and  is  a  typical  (de)b  com- 
pound. According  to  the  Stevens'  modification,  an  electron  which  is  on  one 
of  the  chlorine  ions  migrates  to  the  iridium.  It  will  presumably  go  into  the 
(de)b  shell  which  then  has  six  electrons  and  is  closed.  The  chloride  ion  be- 
comes a  chlorine  atom  with  one  unpaired  spin,  so  that  as  far  as  the  mag- 
netic properties  are  concerned,  the  process  looks  like  the  transfer  of  a  mag- 
netic hole  from  the  iridium  to  a  chlorine.  Adopting  this  sort  of  an  approach, 
the  next  step  was  to  fit  it  into  the  self-consistent  field  model  and  set  up  a 
wave  function  which  has  the  required  symmetry  and  allows  the  electron  to 
spend  part  of  its  time  near  the  chlorine.  Such  a  molecular  orbit  was  con- 
structed from  a  dxy  type  of  metallic  function  and  a  p  type  function  from  the 
ligands. 

Double  Bonds  and  the  Trans  Effect.  The  possibility  of  double  bond  for- 
mation arising  from  the  donation  of  central  cation  d  electrons  to  acceptor 
levels  in  the  coordinated  ligand  has  been  considered  extensively  in  molecu- 
lar orbital  theory.  Craig,  Maccoll,  Nyholm,  Orgel,  and  Sutton28  have  sum- 
marized the  evidence  for  the  existence  of  dT — p*  bonding  using  a  penulti- 
mate c^-orbital  as  follows : 

(1)  Complexes  in  which  this  could  occur  (i.e.,  cyanide,  carbonyl,  and 
nitrosyl)  are  formed  with  elements  which  have  suitable  penultimate  d 
orbitals  such  as  the  transition  metals,  copper  or  silver,  and  even  the  group 
I  IB  elements.  These  compounds  are  not  formed  by  elements  which  lack 
penultimate  d  orbitals  such  as  aluminum. 

(2)  Such  complexes  are  more  stable  than  the  corresponding  ones  formed 
with  ( 51~,  and  Br~,  which  have  no  pw  orbitals  free  to  accept  a  bond  from  the 
metal  atom. 

I  17    Stevens,  Proc.  Roy.  Soc.  (London),  A219,  542  (1953). 


ELECTRON  PAIR  BOND  AND  STRUCTURE  205 

(3)  The  bond  lengths,  where  known,  are  less  than  would  be  expected  for 

o-  bonding  alone. 

All  three  of  these  points  are  subject  to  criticism.  Points  (1)  and  (2)  be- 
come Less  impressive  when  the  stabilities  of  Mo(CN)g"  and  Cr(CN)6E  are 
recalled.  The  latter  stable  complex  cannot  be  stabilized  by  dK-  -pw  bonds 
unless  one  assumes  the  participation  of  unpaired  electrons  in  such  a  bond. 
In  the  former  case,  no  electrons  are  available.  Further,  the  extreme  sta- 
bility of  certain  of  the  phosphorus-boron  bonds  in  compounds  between 
boron  hydrides  and  the  alky]  phosphines  would  require  the  postulation  of 
a  source  of  double  bonding  electrons  other  than  the  d  orbitals102.  In  connec- 
tion with  point  (3),  Wells  has  criticized  the  use  of  bond  lengths  as  a  cri- 
terion of  double  bond  character93. 

Additional  evidence  cited  for  double  bond  character  is  that  for  those 
metals  in  which  no  double  bonding  is  possible  the  coordinating  power  for 
a  Belies  of  amines  runs  parallel  to  the  basic  constants;  so,  if  only  a  bonds 
were  formed,  ethylenediamine  would  always  be  a  stronger  coordinating 
agent  than  dipyridyl.  Since  the  reverse  is  true  with  the  transition  metals, 
it  is  concluded  that  double  bonding  occurs  with  the  transition  metal  com- 
plexes. Since  molecular  orbital  calculations28,  113  indicate  the  theoretical 
feasibility  of  dr-pT  bonds,  the  principal  remaining  problem  is  to  obtain 
proof  that  such  bonds  produce  the  results  attributed  to  them. 

The  stability  of  PF3  complexes  such  as  (PF3)2PtCl292a  and  Ni(PF3)4118  has 
been  attributed  to  dr-pT  double  bonding.  Because  the  x  bond  wrould  tend  to 
neutralize  the  formal  charges  set  up  by  the  formation  of  the  a  bond,  the 
latter  might  be  strengthened. 

Since  two  of  these  x  bonds  could  be  formed  at  right  angles,  the  cis  form 
of  compounds  L2MX2  would  be  favored  if  only  L  could  form  such  bonds 
with  M.  Such  cis  stabilization  would  then  provide  a  reasonable  basis  for 
trans  weakening  and  would  thus  explain  the  trans  effect  or  trans  elimi- 
nation of  PF3  .  Chatt92a  has  treated  the  trans  effect  along  these  lines;  his 
explanation  of  the  trans  effect  for  PF3  is  cited  as  one  of  the  major  advan- 
tages of  his  treatment  as  compared  to  the  two  previous  explanations  (pp. 
1  17  and  195). 

The  argument  can  be  illustrated  by  following  the  explanation  of  Chatt 
and  Wilkins119  for  tin-  cis-trans  conversion  of  [P(Et)3}2PtCl2  .  They  esti- 
mated from  a  t  hermochemical  study  that  the  conversion  of  trans 
[P(E1  PtCli  to  the  cis  form  results  in  an  increase  of  about  L2  kcal  in 
bond  energy.  Since  both  phosphorus  and  chlorine  have  vacant  '/  orbitals, 
(L-il.  bond-  could  be  expected  for  Pi     P  and  Pt     CI.  li  is  assumed  that  the 

118.  Irvine  and  Wilkinson,  Science,  113,  71-'    1951). 

119.  Chatt  and  Wilkins, ./.  Chem.  Soc  ,  1952,  273,  1300;  1953,  70. 


206  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Pt — P  bond  has  greater  double  bond  character  than  the  Pt — CI  bond  be- 
cause P  is  higher  in  the  trans  influence  series. 

The  dotted  lines  in  Fig.  4.13  represent  the  x  or  other  bond  components  in 
which  electron  pairs  from  the  filled  d  orbitals  of  the  metal  atom  contribute 
in  some  manner  to  the  strength  of  the  Pt — P  and  Pt — CI  bonds.  The 
strengths  of  these  components  are  represented  by  the  size  of  the  dots.  In 
the  trans  complex  (I)  both  the  Pt — P  bonds  must  use  the  same  d  orbitals 
in  the  ir  component;  hence  the  x  components  are  weaker  than  in  the  cis 

Pv     .CI  P.       .CI 

V  X 

cf/%V  p"y  xci 

(I)  (H) 

Fig.  4.13.  Bond  components  in  Pt— P  and  Pt— CI  bonds 

complex  where  each  Pt — P  bond  has  available  a  different  d  orbital.  On  the 
other  hand,  the  chlorine  atoms  in  the  cis  complex  (II)  are  now  competing 
with  the  phosphorus  atoms  for  electrons  from  d  orbitals  of  the  platinum 
atom,  so  will  get  a  smaller  share  than  they  had  in  the  trans  isomer.  The 
chlorine  bonds  in  the  trans  position  are  thus  weakened,  as  the  trans  effect 
indicates. 

The  argument  has  an  interesting  application  to  complexes  containing 
PF3 .  Only  the  cis  form  of  PtCl2(PF3)2  is  stable,  as  this  argument  suggests92*. 
Further,  the  weakening  of  the  a  bond  between  phosphorus  and  platinum 
due  to  the  inductive  effect  of  the  fluorine  would  be  partially  compensated 
by  the  increased  strength  of  the  t  bond,  since  the  electronegative  fluorine 
attached  to  phosphorus  should  make  the  phosphorus  d  levels  contract  to  a 
point  where  they  would  be  more  capable  of  w  bond  formation28.  This  line 
of  argument  would  then  suggest  that  in  (C2H5)3P — Pt  bonds,  where  w 
bonds  are  somewhat  less  effective*  than  in  F3P — Pt,  one  might  expect  a 
more  polar  bond  than  in  the  latter  case.  Estimates  of  bond  dipole  moments 
by  Chatt  and  Williams92a  bear  out  this  prediction.  In  such  a  circumstance, 
strong  B — P(C2H5)3  bonds  might  occur  with  less  -k  bonding  contribution 
than  would  be  required  to  stabilize  the  B — PF3  bond.  Hence,  Chatt92a  cites 
the  nonexistence  of  X3B — PF3  complexes  as  strong  support  for  his  double 
bond  postulate  since  boron  does  not  have  d  electrons  available  for  donation 
to  the  phosphorus  in  PF3  .  The  compound  H3B — PF3  is  now  known,  how- 
ever101b. 

A  variation  of  this  dT-dT  treatment  of  the  trans  effect  using  dv  and  dp* 
hybrid  orbitals  has  been  given  by  Jaffe113. 

*  The  less  electronegative  (C2H5)  groups  would  not  be  as  effective  as  F  in  making 
I  he  phosphorus  orbil  als  contract  to  a  point  where  strong  w  bonds  could  form28. 


ELECTRON  PAIR  BOND  AND  STRUCTl  RE  L'07 

Bonding  of  Metals  to  Double  Bonds  in  Terms  of  the  Molecular  Orbital 
Theory.  Coordination  of  metals  to  the  double1  bond  of  ethylene  and  related 
olefins  has  been  treated  by  several  investigators  (e.g.,  Ref.  UD)  using  the 
molecular  orbital  theory  and  is  discussed  elsewhere  (page  506).  A.  E.  A. 
Werner1-1  postulated  a  tt  electron  bond  between  carbon  and  nitrogen  in  the 
azobenzene  platinum (IV)  chloride  described  by  Kharasch  and  Ashford98: 

CI  CI 

CeHs-N        \    /         N-C6H5 

\\—/X II 


■N       /    \ 
CI  CI 


In  order  to  represent  the  difference  between  the  t  and  a  electrons  of  the 
double  bond,  he  suggested  that  the  bond  might  be  formulated  as 


N-^-N 


where  xx  represents  the  electrons  in  the  tt  orbital.  However,  it  is  quite 
possible  that  the  unshared  pair  of  electrons  of  one  or  both  of  the  nitrogen 
atoms122  in  the  azo  group  contributes  to  the  bonding. 

The  metal  cyclopentadiene  complexes  such  as  M(cyclopentadiene)2  with 
their  interesting  sandwich  structure  are  obvious  compounds  for  a  molecular 
orbital  treatment.  Such  treatments  have  been  given  by  Dunitz  and  Orgel123, 
Jaffe124,  and  Moffitt125. 

Bond  Classification — Ioxic  and  Covalent  Bonds — Inner 
and  Outer  Orbital  Complexes 

Throughout  this  and  the  preceding  chapter  the  idea  that  there  are  two 
limiting  types  of  complexes  has  been  recurrent.  The  discussions  based  on 
the  electron-pair  bond  have  dealt  with  complexes  of  the  type  which  might 
most  unambiguously  be  called  penetration  complexes.  They  are  distin- 
guished from  the  normal  or  "ionic"  complexes  by  gross  properties  such  as 
stability  in  the  solid  state,  slow  rates  of  reaction  and  dissociation,  irre- 
versible electrode  and  dissociation  behavior,  and  almost  complete  masking 

120.  Dewar,  Bull.  Soc.  chim.,  18,  C79  (1951);  Chatt  and  Duncanson,  /.  Chem.  Soc, 

1949,  3340;  1952,  2622;  1953,  2939. 

121.  Werner,  Nature,  160,  644  (1947). 

122.  Callis,  Nielsen,  and  Bailar, ./.  Am.  Chen,.  Soc,  74,  3461  (1952) ;  Bailai  and  Callis, 

./.  .1//'.  ('hem.  Soc,  74,  6018  (1952);  Liu,  Thesis,  University  of  Illinois,  1951. 

123.  Dunitz  and  Orgel,  Nature,  171,  121  (19.53). 

124.  Jaffe,  ./.  Chem.  Phys.,  21,  156  (1953). 

125.  Moffitt,  ./.  Am.  Chem.  Soc,  76,  3386  (1954). 


208  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

of  the  constituent  groups.  The  marked  differences  in  the  properties  of  the 
two  types  of  complexes  have  commonly  been  attributed  to  a  distinct  dif- 
ference in  their  bond  types.  The  penetration  complexes  are  often  tacitly 
assumed  to  be  predominately  covalent  while  the  normal  complexes  are 
considered  to  be  ionic.  The  designations  covalent  and  ionic,  however,  ap- 
pear to  depend  in  large  measure  upon  the  individual  using  the  terms,  since 
no  unequivocal  experimental  test  is  available  as  a  means  of  classification. 
With  this  in  mind  it  appears  to  be  profitable  to  review  the  experimental 
parameters  considered  in  the  classification  and  then  to  try  to  relate  these 
parameters  to  electronic  structure  or  other  fundamental  characteristics  of 
the  complex. 

The  Magnetic  Criterion  for  Bond  Type 

Reference  has  already  been  made  to  the  interesting  observation  that  in 
the  formation  of  typical  coordination  compounds  from  paramagnetic  metal 
ions  the  magnetic  susceptibility  of  the  resulting  complexes  is  frequently 
changed  from  that  of  the  simple  ions.  This  is  usually  interpreted  in  terms 
of  the  atomic  or  hybridized  orbital  theory  as  meaning  that  unpaired  d  elec- 
trons in  the  simple  ion  have  become  paired  in  the  complex  and  that  the  d 
orbits  thus  made  available  have  formed  covalent  bonds  with  the  coordi- 
nated groups  or  ions.  In  some  cases,  however,  the  full  paramagnetism  of  the 
central  ion  is  unchanged  when  this  ion  is  made  part  of  a  complex.  For 
example,  the  compounds  [Fe(NH3)6]Cl2 ,  [Co(N2H4)2]Cl2 ,  (NH4)3[FeF6],  and 
K3[CoF6]  appear  to  possess,  respectively,  the  same  number  of  unpaired 
electrons  as  the  gaseous  metal  ions  in  the  ground  state.  It  would  seem  that 
in  these  instances  there  has  been  no  fundamental  reorganization  of  the 
electrons  about  each  component  of  the  complex. 

Pauling,  following  the  lead  of  earlier  workers,  considered  the  bonding 
forces  in  the  "ionic"  126,  m  *  or  normal  complexes  to  be  essentially  electro- 
static in  character.  He  did  not  believe,  however,  that  a  complex  ion,  such 
as  [FeF6]=,  which  contains  five  unpaired  electrons,  should  be  considered  to 
be  of  the  extreme  ionic  type127.  Use  could  be  made  of  the  4s  and  4p  orbitals 
to  form  as  many  as  four  covalent  bonds  without  disturbing  the  3d  shell,  the 
magnetic  moment  of  the  complex  being  unchanged  by  this  amount  of  co- 
valent character  of  the  bonds. 

In  considering  resonance  possibilities  it  is  important  to  realize  that  the 
ion  [FeF6]-  cannot  have  an  intermediate  structure  corresponding  to  reso- 

*  The  terms  "covalent"  and  "ionic"  are  purely  comparative,  but  their  use  in  this 
connection  is  somewhat  confusing.  For  example,  the  fluoride  complex  [FeF6]=  is  not 
ionized  in  water  and  the  Fe — F  bond  is  not  at  all  "ionic"  as  compared  with  the 
Na — F  bond  in  sodium  fluoride. 

126.  Pauling,  ./.  Am.  Chem.  Soc,  54,  1002  (1932). 

127.  Ref.  27a,  pp.  37,  38  and  115. 


ELECTROS  PAIR  BOND  AND  STRUCTURE  209 

Dance  between  the  ionic  type  (containing  five  unpaired  electrons)  and  the 
(/'-Vp:i  covalent  type  (containing one  unpaired  electron)*  since  the  conditions 
for  resonance  require  that  the  resonating  structures  have  the  same  number 
of  unpaired  electron.-'-7.  Since  there  can  be  DO  intermediate  type,  the  mag- 
netic criterion  should  be  capable  of  distinguishing  between  the  predomi- 
nantly covalent  and  predominantly  ionic  complexes  as  defined  above.  In 

each  of  the  examples  cited  above,  the  number  of  unpaired  electrons  for  the 
covalent  type  of  structure  is  different  from  that  for  the  ionic  type,  and 
measurement-  of  magnetic  moments  can  be  used  conveniently  to  determine 
which  type  exists.  This  criterion  fails,  however,  in  those  cases  where  the 
Dumber  of  unpaired  electrons  is  the  same  for  either  extreme  structural  type. 
For  example,  the  number  of  unpaired  electrons  is  three  in  a  complex  of 
chromium(III)  of  coordination  number  six,  assuming  either  a  covalent 
<l-sp;  structure  or  an  essentially  ionic  structure.  Similarly,  it  has  been  sug- 
gested35a-  m  that  a  complex  of  cobalt(II)  and  four  associated  groups  may 
contain  three  unpaired  electrons  for  an  ionic  structure  or  a  covalent  tetra- 
hedral  configuration. 

Xo  distinction  can  be  made  by  means  of  magnetic  moment  measure- 
ments between  covalent  tetrahedral  (spz  hybridization)  and  ionic  structures 
for  complexes  of  copper(I),  silver(I),  and  gold(I);  nor  between  covalent 
planar  (dsp2  hybridization  with  promotion  of  one  d  electron  to  a  p  orbital), 
covalent  tetrahedral  (spz  hybridization),  and  ionic  structures  for  copper(II) 
and  silver(II). 

In  a  similar  manner,  magnetic  susceptibility  measurements  fail  to  serve 
as  a  criterion  for  distinguishing  between  bond  character  in  the  compounds 
of  the  nontransition  elements,  all  of  the  simple  ions  of  these  elements — as 
well  as  their  complex  ions — being  uniformly  diamagnetic. 

The  outstanding  example  in  which  measurements  of  magnetic  suscepti- 
bility have  been  of  value  in  assigning  stereochemical  configurations  is  in 
connection  with  the  complexes  of  tetracoordinate  nickel (II).  This  case  has 
been  discussed  on  page  171.  Figgis  and  Xyholm35h  have  also  considered  the 
for  cobalt (II)  complexes  and  have  suggested  the  size  of  the  orbital 
component  as  an  additional  variable  with  stereochemical  significance. 

Koolution  of  Optical  Isomers  as  a  Criterion  for  Bond  Type 

Some  attempts  have  been  made  to  employ  the  results  of  resolution 
studies  as  an  additional  key  to  the  character  of  bonds  in  compounds. 
Mann129,  for  example,  considered  his  isolation  of  the  dextro  form  of  tetra- 

*  See  Table  4.3. 
128.  Calvin  and  Barkelew,  J.  Am.  Chem.  Soc,  68,  2267  (1946). 
120.  Mann,  J.  Chem.  Soc,  1930,  1745. 


210  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

chloro  (fi ,  /3'-diaminodiethylsulfide  monohydrochloride)  platinum (IV), 

CI     ?  7NH2 

Pt      CH2 

/     V 

CI       I       XS-CH2CH2NH2-HCI 

as  decisive  evidence  for  the  presence  of  a  coordinate  bond  between  the 
sulfur  and  platinum  atoms.  In  this  compound  the  valence  bonds  of  the 
sulfur  atom,  which  has  apparently  become  asymmetric  by  the  process  of 
coordination,  presumably  possess  space  directions  similar  to  those  of  the 
sulfur  atom  in  the  asymmetric  sulfoxides130  and  sulfinates131.  Johnson132 
went  so  far  as  to  propose  a  connection  between  the  existence  or  nonexistence 
of  stable  optical  isomers  and  the  bond  character  of  the  coordination  com- 
pounds. He  indicated  that  stable  optical  isomers  are  possible  only  in  those 
cases  in  which  the  coordinated  groups  are  attached  to  the  central  metal  ion 
by  covalent  bonds.* 

Johnson132  cited  the  apparently  good  correlation  between  resolvability  of 
complexes  and  the  magnetic  criterion  for  bond  type.  The  following  diamag- 
netic  ions,  for  example, 

[Co(C204)3]3  136,  [Rh(C204)3]3  137,  [Co(en)3]+++  138,  and  [Rh(en)3]+++  139 

have  been  resolved  into  stable  optical  isomers,  whereas  [Mn(C204)3]~  and 
[Fe(C204)3]~,  which  contain  four  and  five  unpaired  electrons,  respectively, 
have  resisted  all  attempts  at  unequivocal  resolution132, 140.  Failure  to  resolve 
complexes  of  this  type,  in  which  configurational  dissymmetry  almost  cer- 
tainly exists,  is  probably  due  to  a  rapid  rate  of  racemization  of  the  optical 
isomers.  The  assumption  made  by  Johnson  implies  that  this  rate  is  too  rapid 
to  allow  separation  and  identification  of  the  isomers  when  the  bonds  be- 
tween the  central  metal  atom  and  the  attached  groups  are  essentially  ionic, 
but  is  sufficiently  slow  for  resolution  to  be  effected  when  the  attached 

130.  Harrison,  Kenyon,  and  Phillips,  /.  Chem.  Soc,  1926,  2079. 

131.  Phillips,  J.  Chem.  Soc,  127,  2552  (1925). 

132.  Johnson,  Trans.  Faraday  Soc,  28,  845  (1932). 

*  Essentially  the  same  suggestion  had  been  made  earlier  by  Sidgwick133. 

133.  Ref.  5b,  p.  86. 

134.  Hunter  and  Samuel,  Chemistry  and  Industry,  1935,  34. 

135.  Orgel,  /.  Chem.  Soc,  1952,  4756. 

136.  Jaeger  and  Thomas,  Proc  Acad.  Sci.  Amsterdam,  21,  693  (1919);  Johnson  and 

Mead,  Trans.  Faraday  Soc,  29,  626  (1933). 

137.  Werner,  Ber.,  47,  1954  (1914);  Jaeger,  Rec  Trav.  Chim.,  38,  256  (1919). 

138.  Werner,  Ber.,  45,  121  (1912). 

139.  Werner,  Ber.,  45,  1228  (1912). 

140.  Thomas,  /.  Chem.  Soc,  119, 1140  (1921);  Jaeger,  Rec  Trav.  Chim.,  36,  242  (1919). 


ELECTRON  PAIR  BOND  AND  STRl  CT\  RE  211 

groups  are  bound  by  covalenl  bonds.  Inherent  in  all  of  the  foregoing  argu- 
ments is  the  assumption  that  a  covalenl  bond  is  of  necessity  stronger  than 

an  ionic  one  or  is  slower  in  reaction.  This  point,  has  been  jusl  lv  cril  ici/cd' 

It  is  significant  in  support  of  Johnson's  arguments  that  Ci  •(<    '  >        d 
been  resolved141  while  All  ('■_■<  V:i   could  not  l>e  resolved15-  ni  despite  earlier 
claims  for  resolution11-. 

Exchange  Studios  as  a  Criterion  for  Bond  Type.  There  appears  to 
I>e  a  rough  parallelism  between  the  conclusions  obtainable  from  exchange 
experiments,  magnetic  susceptibility  data,  and  studies  involving  the  iso- 
lation of  stable  isomers.  That  is  to  say,  those  complexes  which,  on  the  basis 
of  magnetic  moment  measurements,  appear  to  satisfy  the  criterion  for 
covalent  binding  are  also  usually  resolvable  into  optical  isomers  or  separable 
into  cis  and  trans  isomers  and  do  not  undergo  rapid  exchange  between  the 
central  metal  atom  of  the  complex  and  a  radioactive  isotopic  ion  of  this 
metal78**-  144  145.  To  illustrate,  bis(methylbenzylglyoxime)nickel(II)  is  dia- 
magnetic,  has  been  separated  into  two  stable  geometric  isomers41a,  and  does 
not  exchange  with  radioactive  nickel(II)  ions144a.  Similarly,  the  diamag- 
netic  ion  [Copn-2Cl2]+  was  found  not  to  exchange  with  radioactive  cobalt(II) 
ions146.  Further,  the  diamagnetic  ion  [Co^O^]",  which  has  been  resolved136 
into  stable  d  and  I  forms,  does  not  exchange147  its  bonded  oxalate  radicals 
with  uncombined  oxalate  ions  containing  radioactive  carbon. 

Exchange  experiments  carried  out  by  Long147,  148  between  uncombined 
oxalate  ions  containing  radioactive  carbon  and  the  complex  oxalato  ions  of 
aluminum(III),  iron(III),  cobalt(III),  and  chromium(III)  appear  to  be  in 
agreement  with  the  resolution  studies.  The  oxalate  complexes  of  alumi- 
num(III)  and  iron(III)  undergo  rapid  interchange  while  those  of  cobalt(III) 
and  chromium(III)  show  none. 

The  results  of  exchange  experiments  between  radioactive  cobalt  and 
complexes  of  cobalt (II)  and  cobalt(III)  containing  bidentate  ligands  led 
\Yest144c  to  the  general  conclusion  that  slow  exchange  can  be  associated 
with  strong  covalent  bonds  in  the  complex  and  rapid  exchange  with  weak 

HI.  Werner,  Be,.,  45,  3061  (1912). 

142.  Wahl,  Ber.,  60,  399  (1927);  Burrows  and  Lauder,  ./.  Am.  Chem.  8oc.,  53,  3600 

(1031). 

143.  Johnson,  Trans.  Faraday  Soc,  31,  1612  (1935). 

144.  Johnson  and  Hall,  ./.  Am.  Chem.  Soc.,  70,  2344  [1948);  Hall  and  Willeford,  ./. 

Am.  Caem.  Soc.,  78, 5419  (1951);  West, ./.  Chem.  Soc.,  1958, 3115;  Libby,  "The- 
ory  of  Electron  Exchange  Reactions  in  Aqueous  Solutions."  p.  :;■•,  Preprint, 

posium  on  Electron  Transfer  and  Esotopic  Reactions,  Division  of  P] 
and  Inorganic  Chemistry,  American  Chemical  Society,  and  Division  of  Chemi- 
cal Physics,  American  Physical  Society,  Notre  Dame,  .tunc  11   I 

145.  Adamson,  Welker,  and  Volpe,  ./.  An,.  <  .  72,  1090    1950  . 

146.  Flagg,  J   Am.  Chi  m.  Soc.,  63,  557    i'.»41). 

147.  Long,  J.  Am.  Chem.  Soc,  63,  1353  (1941;. 

148.  Long,  J.  Am.  Chem.  Soc,  61,  570  (1939). 


212  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

covalent  or  ionic  bonds.  Oxalato  and  malonato  complexes  of  iron(III) 
which  have  magnetic  susceptibilities  corresponding  to  five  unpaired  elec- 
trons are  reported  to  exchange  rapidly  with  carbon-14  labeled  ligands, 
whereas  K3Fe(CN)6 ,  wThich  has  a  moment  corresponding  to  one  unpaired 
electron,  shows  negligible  exchange149. 

The  above  facts  support  the  general  consistency  of  the  three  experimental 
criteria  used  for  bond  classification  (i.e.,  magnetic  moment,  resolution,  ex- 
change) ;  however,  some  cases  of  apparent  disagreement  have  been  reported 
and  should  be  considered.  According  to  Johnson150,  the  ion  [Ni  en3]++  could 
not  be  resolved  into  its  optical  iosmers,  and  on  this  basis  the  bonds  between 
the  nickel  and  nitrogen  atoms  would  be  termed  ionic  in  character.  In  the 
case  of  [Ni  dipy3]++,  there  seems  no  obvious  reason  for  expecting  a  funda- 
mentally different  type  of  binding  between  nickel  and  the  nitrogen  atoms, 
yet  this  complex  ion  has  been  resolved151  and  so  would  be  classed  as  covalent 
in  character.  Claims152  have  also  been  made  for  the  resolution  of 
[Ni  en2(H20)2]++.  This  would  require  the  highly  improbable  conclusion  that 
the  binding  in  [Ni  en2(H20)2]++  is  covalent  in  character,  whereas  the  tris- 
(ethylenediamine)  complex  is  ionic.  Magnetic  moment  measurements 
obviously  can  supply  no  clue  in  these  cases  inasmuch  as  both  the  ionic  and 
covalent  structures  involve  twro  unpaired  electrons. 

Further  disagreement  in  classification  is  observed  between  the  resolution 
method  and  the  exchange  method78b.  Neogi  and  Dutt153  have  reported  the 
resolution  of  [Ga(C204)3]s;  however,  the  general  exchange  behavior  of  gal- 
lium (I  II)  makes  it  seem  almost  certain  that  the  complex  would  exchange 
oxalate  rapidly.  Resolution  of  [Zn  en3]++  and  [Cd  en3]++  has  been  reported154, 
yet  formation  and  dissociation  of  these  complexes  is  instantaneous.  Such 
resolution  seems  improbable. 

The  complexes  of  iron (II)  with  o-phenanthroline  and  a ,  a'-dipyridyl  are 
diamagnetic155  and  the  tris  complex  of  the  latter  coordinating  molecule  has 
been  resolved  into  its  stable  optical  isomers156.  Accordingly,  the  iron- 
nitrogen  bonds  in  these  complexes  are  generally  conceded  to  be  mainly 
covalent  in  character157.  Thus,  exchange  between  radioactive  iron (II)  and 
these  complex  ions  might  not  be  anticipated.  However,  Ruben  and  co- 
workers158 demonstrated  that  these  ions  experience  exchange  at  a  slow  but 

149.  Clark,  Curtis  and  Odell,  J.  Chem.  Soc,  1954,  63. 

150.  Johnson,  Trans.  Faraday  Soc,  28,  854  (1932). 

151.  Morgan  and  Burstall, ./.  Chem.  Soc,  1931,  2213;  Nature,  127,  854  (1931). 

152.  Wahl,  Acta  Sci.  Fennicae,  Comm.  Phys.  Math.  4,  1  (1927). 

153.  Neogi  and  Dutt,  J.  Indian  Chem.  Soc,  15,  83  (1938). 

1.">1    Xeogi  and  Mukherjee,  J.  Indian  Chem.  Soc,  11,  225  (1934). 

155.  Ref.  22b. 

156.  Werner,  Per.,  45,  433  (1912). 

157.  Ref.  27a,  p.  117. 

158.  Ruben,  Kamen,  Allen,  and  Nahinsky,  J.  Am.  Chem.  Soc,  64,  2297  (1942). 


ELECTRON  PAIR  BOND  AND  STRUCTURE  213 

easily  measurable  rate  in  aqueous  solution.  On  the  contrary,  the  iron(III) 
in  ferrihemcgLobirj  and  ferriheme,  which  is  considered  to  l>e  held  by  ionic 

or  electrostatic  forces  on  the  basis  of  magnetic  data15'*,  did  not  exchange 
with  radioactive  iron(III)  ions  after  two  months.  These  workers  concluded 
that  the  rate  oi  exchange  appears  to  depend  more  on  the  structural  features 
of  the  complex  ion  than  on  bond  type.  It  has  been  suggested158'  160  that  in 
those  complexes  with  a  fused  ring  structure,  such  as  ferrihemoglobin,  there 
may  be  considerably  greater  stereochemical  resistance  to  exchange  than  in 
tlu1  case  of  dipyridyl  and  similar  complexes  simply  because  of  the  necessity 
of  breaking  the  four  metal-nitrogen  bonds  without  bond  reformation  in  the 
former  as  against  a  "stepwise"  exchange  in  the  latter.  On  the  basis  of  prob- 
ability considerations,  then,  exchange  in  the  dipyridyl  type  complexes  may 
be  favored  over  that  in  the  fused  ring  type  in  spite  of  predictions  to  the  con- 
trary based  on  magnetic  data. 

The  diamagnetic  Xi(CX)4=  undergoes  rapid  exchange  in  direct  contra- 
diction to  the  expected  result. 

Other  Criteria  for  Bond  Type 

X-ray  analyses,  electron  diffraction  studies,  and  optical  methods  have 
supplied  extremely  useful  information161'  162  regarding  complex  molecule- 
and  ions,  but  such  information  usually  yields  clues  as  to  the  nature  of  the 
bonds  between  the  constituent  parts  of  these  complexes  only  as  it  can  be 
interpreted  in  the  light  of  other  data  and  current  theories  of  binding.  Some 
information  regarding  the  force  constants  of  the  bonds  in  coordination 
compounds  has  been  obtained  from  a  study  of  the  Raman  spectra  of  these 
substances.  From  these  studies  has  come  the  rather  unexpected  result150 
that  the  force  constants  for  typical  coordinate  bonds  are  of  the  same  order 
of  magnitude  though  somewhat  smaller  than  that  for  ordinary  single  bonds. 

The  "Inner  and  Outer  Orbital'*  Complexes  of  Tanbe 

The  entire  field  of  substitution  reactions  in  complex  ions,  including  both 
radioactive  exchange,  racemization,  and  chemical  substitution  reactions 
was  considered  in  an  excellent  review  by  Taube78b.  He  pointed  out  that  a 
useful  classification  of  complexes  can  be  based  on  differences  in  their 
adjustment  to  equilibrium  with  respect  to  substitution  reactions  (chemical 
basis  of  bond  type).  On  the  other  hand,  he  emphasized  that  a  slower  rate 
for  substitution  doe-  not  necessarily  mean  greater  bond  stability  and  thai 
rates  of  reaction  will  not,  of  necessity,  correlate  with  factors  related  to  bond 

159.  Pauling  and  Coryell,  Proc.  Natl.  Acad.  Sri.,  22,  150.  210    1931 
ICO.  Ikler,  J.  Am.  Chem.  Soc,  69,  724  (1947; ;  Reference  22,  p.  171 . 

161.  Fernelius,  "Chemical  Architecture"   (diurk  and  Grummitt,   Eds.),  Chap.    Ill 

York,  Interscience  Publishers,  Inc.  1948;  Ref.  15c,  p.  it'»7;  Chap.  V. 

162.  Szabo,  Acta  Univ.  Szegediensis,  Acta  Chem.  et  Phys.  (A\  S.),  1,  52  (1942;. 


214  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

strength.  (On  this  basis  Bjerrum's  term  "robust"  complexes  was  criticized, 
since  it  implies  greater  stability.)  As  a  case  in  point,  Taube  noted  that  the 
complex  CrCl++  is  more  dissociated  at  equilibrium  than  the  corresponding 
FeCl4"1"  ion,  yet  the  iron  (III)  complex  is  in  labile  equilibrium  with  its  sur- 
roundings while  the  chromium(III)  ion  is  not. 

Taube's  summary  of  the  data  relative  to  the  lability  of  various  complexes 
with  respect  to  substitution  reactions  is  made  in  Table  4.5.  Inert  and  labile 
groups  may  be  readily  distinguished. 

The  electron  structures  for  the  complexes  of  coordination  number  six  fall 
quite  naturally  into  two  classes:  in  one  class,  which  will  be  designated  as 
the  "inner  orbital"  type,  relatively  stable  d  orbitals  of  lower  principal 
quantum  number  are  combined  with  the  sp3  set  of  orbitals  of  higher  quan- 
tum number;  in  the  other,  designated  as  the  "outer  orbital"  type,  the  d 
orbitals  have  a  considerably  lower  stability,  since  they  are  of  the  same 
principal  quantum  number  as  the  s  and  p  orbitals  with  which  they  are 
hybridized.*  The  subdivision  of  the  inner  and  outer  orbital  complexes  into 
the  labile  and  inert  classifications  is  indicated  in  Table  4.6.  The  important 
point  indicated  by  the  classification  is  the  discontinuity  in  rates  which  ap- 
pears at  the  point  at  which  the  last  available  inner  d  orbital  is  occupied  by 
an  unshared  electron.  For  example,  reactions  of 

1  \     ! I 

are  rapid,  while  those  of 

CrM"*-       _2tf_L 45  -*£—\ 

<  <   <L j 

are  slow. 

Mo5+  (dW°D2£P3)  complexes  are  labile;  those  of  Mo+++  (dWWSP*) 
are  relatively  inert. 

Taube  pointed  out  that  this  factor  appears  to  be  of  major  significance, 
and  it  cannot  be  attributed  to  a  sudden  change  in  degree  of  covalent  char- 
acter of  the  bonding  since  evaluation  of  degree  of  covalent  bond  character 
by  independent  means  shows  no  sudden  discontinuity  at  the  appearance  of 
fchis  particular  configuration.  As  independent  indices  of  covalent  character 

*  Huggins163  first  proposed  the  use  of  inner  and  outer  orbitals  for  coordinate  bond 
formation.  Pauling  rejected164  the  idea  on  the  grounds  that  such  bonds  are  too  weak 
to  be  of  importance.  More  recent  calculations28  of  bond  strength  from  the  overlap 
integral  indicate"  that  such  outer  orbital  complexes  are  justifiable,  particularly  under 
i  he  conditions  outlined  by  Huggins  (i.e.,  with  groups  of  high  electronegativity). 
163.  Buggins, ./.  Chem.  Phys.,  5,  527  (1937). 
I  til    Ref .  27a,  p.  115. 


Tabi.f.  -1.5.  Lability  of  Hexacoordinated  Complex  Ions 
(From  Reference  78b) 
Complex  ions  of  the  following  are  Labile  with  respect  to  simple  substitution: 
aluminum  (III),  Boandium(III),  yttrium  (III),  tripositive  rare  earth  ions,  titani- 
um (IV),    zireonium(IV),    thorium(IV),    U02++,    plutonium(III),    plutonium(IV), 

PU02++. 


Element 

Lability  of  Complex  Ions 

V(II) 

V(CN)64~  is  inert;  no  definite  evidence  on  other  com- 

plex ions 

van) 

F",  CNS",  CN",  SOr,  C2Or,  citrate,  and  pyrophos- 

phate complex  ions  are  "labile";  V(CX)6a  appears 

to  be  more  labile  than  V(CN)64" 

Nb(II) 

Only  polynuclear  complexes  known  in  solution 

Nb(III) 

SO4"  complex  probably  labile 

Nb(V) 

Cl~,  Br",  and  H20  complexes  labile 

Ta(II) 

Only  polynuclear  complexes  known  in  solution 

Ta(III) 

Xo  definite  information;  CN-  complex  probably  labile 

Ta(V) 

CN~  complex  labile;  F"  and  C204=  complexes  probably 

labile 

Cr(II) 

Cl~  complex  reported  inert 

Cr(III) 

H2Or  F",  CI",  CN",  CNS",  NH3  ,  etc.  complexes  inert 

Mo(II) 

Only  polynuclear  complexes  known  in  solution 

Mo(III) 

CI",  Br~,  and  CNS"   complexes  inert;  replacement  of 

NHi  slow  in  acid 

Mo  (IV) 

Mo(CNy-  inert 

Mo(V) 

CI"  and  Br"  complexes  labile;  CNS"  complex  may  be 

measurably  slow  in  substitution;  Mo(CN)g-  inert 

Mo(VI) 

F",  CI",  and  HOO"  complexes  labile 

W<II) 

Only  polynuclear  complexes  in  solution 

W(III) 

W2C19-  characterized  as  inert 

W(IV) 

Cl~  complex  probably  labile;  W(CN)84~  inert 

W(V) 

Cl~  and  C204"  complexes  probably  labile;  W(CN)8S 

inert 

W(VI) 

F"  and  Br-  complexes  labile;  CI"  complex  doubtful 

Mn(II) 

En  and  pyrophosphate  complexes  labile;  Mn(CN)«a 

inert 

Mn(III) 

F",  CI",  C204~,  and  pyrophosphate  complexes  labile; 

Mn(CN)63  inert 

Mn(IV) 

F~  and  C204~  complexes  inert 

Re(III) 

CI"  complex  inert;  NH3  complex  probably  inert 

Re  (IV) 

CI",  Br~,  and  I"  complexes  inert 

Re(V) 

CI"  and  CNS"  complexes  labile;  Re02(CN)4s  indeter- 

minate, may  be  inert 

Re  (VI) 

F~  complex  labile 

Fe(II) 

En  and  C204"  complexes  labile;  Fe(CN)«4"    (and  sub- 

stitution derivatives),  Fe(ophen)3++,  and  Fe(dipy)i++ 

inert 

Fe  HI) 

F",  CI",  Br-  CN-  .  NH,,S«Or,  SO,"  and  CtOr  com- 

plexes labile;  Fe(CN)»"   (and  substitution  deriva 

tives)  and  Fe(ophen)j+++  inert 

21.5 


Table  4.5 — Continued 


Element 

Lability  of  Complex  Ions 

Ru(II) 

Cl~,  CN",  and  NH3  complexes  inert 

Ru(III) 

CI",  Br",  and  C204"  complexes  inert;  complex  ammines 

and  derivatives  inert 

Ru(IV) 

CI"  complex  inert 

Ru(VI) 

Cl~  complex  labile 

Os(II) 

Cl~  complex  inert;  CN"  complex  probably  inert 

Os(III) 

Cl~  complex  inert 

Os(IV) 

Cl~  complex  inert 

Os(VI) 

F"  complex  labile;  C204~,  N02~,  and  Cl~  complex  on 

Os02++  undergo  rapid  substitution 

Co  (II) 

CI",    Br-,    I"",    CNS~,    and    NH3    complexes   labile; 

Co(CN)64~  may  be  inert 

Co  (III) 

H20  in  presence  of  Co++  labile;  CN",  S03™,  N02~,  and 

C204=  complexes  inert;  complex  ammines  and  de- 

rivatives inert 

Rh(II) 

Br"  in  Rhpy5Br+  slow  in  substitution 

Rh(III) 

Cl~,  CN",  S04~,  and  NH3  complexes  inert 

Ir(III) 

CI",  Br~  probably,  and  CN"  complexes  inert;  complex 

ammines   and   derivatives   inert;   S04~   and   C204" 

complexes  inert 

Ir(IV) 

Cl~  and  py  complexes  inert 

Ni(II) 

NH3  ,  en,  C204=,  tartrate,  and  CN~  complexes  labile, 

dipyridyl  complex  inert 

Pd(II) 

Coordination  number  4  only  in  complex  ions  and  de- 

rivatives; some  reactions  measurably  slow 

Pd(IV) 

No  definite  conclusions 

Pt(II) 

Coordination  number  4  only;  Cl~  and  NO 2"  complexes 

inert;  ammines  and  derivatives  inert;  complexes  less 

labile  than  those  of  palladium (II) 

Pt(IV) 

Halide  and  CNS~  complexes  inert;  ammines  and  de- 

rivatives inert 

Cu(I),  Cu(II) 

Cl~,  Br~,  NH3  ,  and  SO3"  complexes  labile 

Ag(I) 

NHj  ,  CN~,  and  S03"  complexes  labile 

Au(I) 

CI" ,  Br ,  CN"  and  CNS"  complexes  probably  labile 

Au(III) 

S04°",  Cl~,  and  NH3  complexes  inert;  NOr  complex 

hydrolyzed  rapidly 

Zn(II),Cd(II),Hg(II) 

Labile 

Ga(III) 

F~,  CI-,  and  C204=  complexes  labile 

In(III) 

Probably  labile 

Tl(III) 

C204~  complex  labile;  CI"  and  Br-  complexes  not  cer- 

tain 

Si  (IV) 

F"  in  SiF6=  measurably  slow  in  substitution 

Ge(IV) 

No  conclusions  for  coordination  number  6 

Sn(IV) 

No  conclusions  for  coordination  number  6 

P(V) 

PF6~  inert 

As(V) 

AsF6"  and  As(C6H402)3~  inert 

Sb(V) 

SbF6~  and  SbCU"  inert 

SF6  ,  SeF6  ,  TeF6 

Inert 

216 


ELECTRON  PAIR  BOX  J)  AX  J)  STRUCTl  RE  217 

Table  4.6.  Inner  and  Outer  Orbital  COMPLEXES  Inert  and  Labile  Forms 

(From  Reference  78b) 

I.  Inner  orbital  complexes 

A.  Labile  members 

(1)  d°d°d°D^SP3    Sc(III),  Y(III),  rare  earths(III),  Ti(IV),  Zr(IV),  Hf(IV), 

Ce(IV),     Th(IV),     Nb(Y),     Ta(V),     Mo(VI),     W(VI), 
Np(III),  Np(IV),  Pu(III),  Pu(IV). 

(2)  d*d<>d0D7SP3    Ti(III),  Y(IV),  Mo(V),  W(V),  Re(VI). 

{3)  dhPd°D*SP*    Ti(II),  V(III),  Nb(III),  Ta(III),  W(IV),  Re(V),  Ru(VI). 

B.  Inert  members 

(1)  dWdWSP*     V(II),  Cr(III),  Mo(III),  W(III),  Mn(IV),  Re(IV). 

(2)  d'dWD*SP*     Cr(CN)64-,  Mn(CN)6s,  Re(III),  Ru(IV),  Os(IV). 

(3)  d*-dWD*SP3     Mn(CN)6-,  Re(II),  Fe(CN)68S,  Fe(ophen)3+++,  Fe(dipy)3+++ 

Ru(III),  Os(III),  Ir(IV). 

(4)  dHNNPSP*     Fe(CN)64-,    Fe(ophen)3++,    Fe(dipy)3++,    Ru(II),    Os(II), 

Co(III)  in  all  but  F  complexes,  Rh(III),  Ir(III),  Pd(IV), 
Pt(IV). 
II.  Outer  orbital  complexes 

Lability  tends  to  decrease  slowly  as  charge  on  central  cation  increases.  Typical 
"outer  orbital  ions":  A1+++,  Mn++  Fe++,  Fe+++,  Co++,  Ni++,  Zn++,  Cd++,  Hg++, 
G&+++,  In+++,  and  T1+++. 

he  used  the  acid  dissociation  constants  of  the  hydrated  ions,  the  hydration 
energies  of  the  metal  ions,  and  theoretical  arguments  from  size  and  charge 
of  the  ion. 

This  is  not  to  imply  that  the  degree  of  covalent  character  in  the  bond 
may  not  exercise  an  influence  on  the  rate  of  substitution  reactions;  on  the 
contrary,  the  variation  in  the  degree  of  covalent  character  is  an  important 
factor  in  determining,  for  those  ions  for  which  both  possibilities  exist, 
whether  the  complex  ion  will  be  of  the  inner  orbital  or  outer  orbital  elec- 
tronic  type.  But  it  is  particularly  significant  that  under  some  circumstances, 
complexes  of  the  outer  orbital  type  which  are  described  as  "ionic"  may  have 
bonds  of  more  covalent  character  than  some  of  the  inner  orbital  complexes 
which  are  classified  as  "covalent".  For  example,  there  is  reason  to  believe 
that  [Ga(H20)6]"f++  is  more  covalent  in  its  bonds  than  is  [Cr(H20)6]+"H", 
yet  from  exchange  studies  [Ga(H20)6]++"f  is  classed  as  "ionic"  while 
[Cr(H20)6]+~H"  is  classed  as  "covalent".  It  is  in  this  sense  that  Taube's 
classification  seems  much  superior  to  the  conventional  ionic-covalent  de- 
scription.  The  terms  "ionic"  and  "covalent"  must  remain  indefinite  be- 
cause  they  are  not  defined  unambiguously. 

On  the  other  hand,  the  experimental  classification  of  complexes  into 
inert  and  labile  compounds  is  usually  definite  and  the  theoretieal  descrip- 
tion of  these  complexes  is  quite  unambiguous  except  in  a  relatively  small 
number  of  eases  where  either  the  inner  or  outer  orbital  designation  may 
apply  (i.e.,Cu++,Ni(A),++f  etc.). 

The  implication  that  all  bonds  involving  a  change  in  magnetic  moment 


218  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

are  stronger  than  bonds  in  which  no  such  change  is  observed  (i.e.,  "co- 
valent"  bonds  by  magnetic  criterion  are  stronger  than  "ionic")  has  been 
shown  to  be  untrue  in  an  earlier  discussion  (Chapter  3,  p.  136).  One  then 
looks  to  a  factor  other  than  "bond  strength"  to  explain  the  rapid  exchange 
in  the  labile  complexes  and  the  slow  exchange  in  the  inert  complexes.  Since 
one  is  dealing  with  a  problem  in  kinetics  in  all  exchange  and  racemization 
studies,  it  would  appear  that  there  is  a  sharp  discontinuity  in  the  energy 
required  to  form  the  activated  complex  as  soon  as  the  last  d  orbital  gets  at 
least  one  electron. 

Taube  interpreted  these  facts  as  indicating  that  substitution  proceeds  by 
an  intermediate  species  of  coordination  number  seven  which  can  be  stabi- 
lized through  utilization  of  the  empty  d  orbital  on  the  central  metal  ion. 
If  the  inner  d  orbitals  are  completely  occupied,  electrons  must  be  promoted 
or  paired  to  make  a  d  level  available.  Either  process  would  require  energy 
which  would  appear  as  an  activation  energy.  The  alternative  path,  in  which 
a  ligand  is  lost  in  the  rate  determining  steps,  can  also  be  supposed  to  re- 
quire a  high  activation  energy,  since  there  is  no  factor  which  compensates 
effectively  for  the  energy  required  to  remove  the  group. 

In  outer  orbital  complexes  lability  is  observed  if  the  central  ion  has  low 
charge,  while  increasingly  inert  character  is  observed  as  the  charge  on  the 
central  ion  builds  up.  Substitution  by  dissociation  mechanism  seems  reason- 
able when  the  charge  is  low  (i.e.,  1,  2,  or  3).  It  has  been  suggested  that  the 
energy  required  to  remove  one  of  the  groups  is  compensated  in  part  by 
rerrybridization  of  the  lower  orbitals  (i.e.,  sp3  or  sp2d  to  a  lower  coordination 
number).  The  observation  that  many  of  the  metals  of  these  complexes 
readily  assume  a  coordination  number  of  four  was  cited  in  support  of  such 
an  argument.  Increasing  charge  on  the  central  ion  is  bound  to  produce 
bonds  of  more  covalent  character  which  are  stronger  and  harder  to  dis- 
sociate or  substitute  by  any  mechanism.  This  is  illustrated  by  the  fact  that 
the  rate  of  hydrolysis  decreases  in  the  series  of  the  hexafluoro  complexes: 
AlF«r  >  SiF6=  >  PF<r  »  SF6 . 

An  exception  to  the  above  rules  is  found  in  the  case  of  [Co(H20)6]+++. 
This  ion  exchanges  water  rapidly,  much  more  rapidly  than  replacement  of 
NH3  by  H20  in  [Co(NH3)6]+++.  The  electronic  structure  as  determined  by 
its  diamagnetism  is  d2d2d2D2SP3,  which  should  lead  to  slow  exchange  on  the 
basis  of  the  above  considerations  for  inner  orbital  complexes.  It  is  probable, 
however,  that  the  paramagnetic  labile  state  for  Co+++  (d2dldldl  SP3D2)  is 
only  slightly  above  the  diamagnetic  ground  state  in  energy.  This  relation 
is  expected  from  the  fact  that  in  the  complex  with  fluoride  the  paramagnetic 
state  is  lowest  while  in  the  hexammine  the  diamagnetic  state  is  lowest. 
Since  water  is  intermediate  between  fluoride  and  ammonia  in  polarizability, 
one  might  expect  on  the  basis  of  crystal  field  splitting  arguments  (Chapter 


ELECTRON  PAIR  BOND  AND  STRUCTURE  219 

3)  that  the  two  states,  paramagnetic  and  diamagnetic,  would  lie  close  to- 
gether in  the  water  complex  (i.e.,  near  to  the  point  (A)  of  intersection  of 
the  two  lines  in  Fig.  3.3  (p.  L36)).  Od  this  basis  a  small  activation  energy 
would  suffice  to  give  the  outer  orbital  paramagnetic  structure,  which  could 
undergo  exchange  more  readily  than  the  closed  shell  type  of  structure. 

Taube's  work  emphasizes  a  point  which  should  be  obvious  but  which 
none  the  less  results  in  much  confusion.  Criteria  based  on  rate  are  de- 
pendent upon  mechanism  and  as  such  are  frequently  much  less  dependent 
upon  bond  strength  than  is  commonly  supposed.  In  this  sense  all  expla- 
nations of  the  trans  effect  are  inadequate,  since  it  has  never  been  fully 
established  that  the  result  is  due  to  bond  strength  rather  than  rate  and 
mechanism.  Taube's  postulates  would  suggest  that  mechanism  might  be  of 
major  importance  in  explaining  these  substitution  processes,  yet  all  expla- 
nations of  the  effect  are  based  on  the  concept  of  bond  strength.  In  fact,  one 
must  conclude  with  Taube  that  our  knowledge  of  reaction  mechanisms  of 
coordination  compounds  is  still  very  meager. 


O.  Chelation  and  the  Theory  of  Hetero- 
cyclic Ring  Formation  Involving  Metal  Ions 

Robert  W.  Parry 

University  of  Michigan,  Ann  Arbor,  Michigan 

The  term  "chelate"  was  proposed  by  Morgan1  to  designate  those  cyclic 
structures  which  arise  from  the  union  of  metallic  atoms  with  organic  or  in- 
organic molecules  or  ions.  The  name  is  derived  from  the  Greek  word  chela 
which  means  the  claw  of  a  lobster  or  crab.  Chelate  ring  systems  can  be 
formed  only  by  ligands  which  have  more  than  one  point  of  attachment  to 
the  metal.  For  example,  unidentate  NH3  cannot  form  a  ring,  but  bidentate 
ethylenediamine  can  form  chelate  structures.  Ligands  with  three  points  of 
attachment  are  known  as  tridentate,  those  with  four,  as  tetradentate,  and 


so  on: 


H, 


H2 


Hi 


M  <-  NH3  N 

/    \ 
M  CH2 

T  I 

N CH2 

I 
H2 

Monodentate  Bidentate 

Ligand  Ligand 

No  Chelation        One  Chelate  Ring 


N  N 

/    \     •    \ 

CH2        M  CH2 

I         T         I 

CH2 N CH2 

I 
H 

Tridentate  Ligand 

Two  Interlocked 

Chelate  Rings 


A  comprehensive  review  of  the  chelate  rings  was  given  by  Diehl2  in  1937 
and  a  more  recent  treatment  by  Martell  and  Calvin3  in  their  book,  "The 
Chemistry  of  the  Metal  Chelate  Compounds. " 

Many  widely  divergent  chemical  and  biological  problems  are  intimately 
related  to  the  formation  of  chelate  rings.  For  example,  metals  which  are 

1.  Morgan  and  Drew,  J.  Chem.  Soc,  117,  1456  (1920). 

2.  Diehl,  Chem.  Rev.,  21,  39  (1937);  (a)  p.  84. 

3.  Martell  and  Calvin,  "Chemistry  of  the  Metal  Chelate  Compounds,"  New  York, 

Prentice-Hall,  Inc.  1952. 


220 


THEORY  OF  HETEROCYCLIC  RING  FORMATION  221 

essential  for  plant  and  animal  nutrition  form  chelate  rings  m  the  organism 
(Chapter  "J P.  Thus,  hemin  is  an  iron  chelate  and  chlorophyll  Lfl  a  magne- 
sium ring  compound.  Also,  metals  play  an  important  role  in  the  functioning 

of  enzymes  apparently  through  chelate  ring  format  ion  in  the  inter- 
mediates. 

Another  point  of  biological  interest  is  the  use  of  metal  ion  buffers.  By 
selecting  a  proper  completing  agent,  free  metal  ion  concent  ration  can  be 
maintained  at  a  relatively  constant  level  in  a  predetermined  range  just  as  a 
constant  hydrogen  ion  concentration  is  maintained  in  conventional  buffer 
systems. 

A  novel  use  of  chelating  agents  for  the  direct  titration  of  metals  has  been 
suggested  by  Schwarzenbach4.  He  points  out  that  many  chelating  agents 
change  color  according  to  the  metal  ion  concentration  in  a  manner  com- 
pletely analogous  to  the  pH  dependent  color  changes  observed  with  acid- 
base  indicators.  This  makes  direct  metal  titrations  possible. 

The  Stability  of  Chelate  Structures 
Extra  Stability  Due  to  Chelation— The  "Chelate  Effect" 

One  of  the  most  striking  properties  of  chelate  ring  compounds  is  their 
unusual  stability.  In  this  respect  they  resemble  the  aromatic  rings  of  or- 
ganic chemistry.  As  an  illustration,  one  may  compare  the  relatively  stable 
chelate  [Xi(en)3]++  with  the  analogous,  but  less  stable  non-chelate  com- 
pound [Xi(XH2CH3)6]++.  The  ethylenediamine  complex  is  stable  in  solution 
at  high  dilution,  but  the  methylamine  compound  dissociates  under  the  same 
conditions  to  precipitate  nickel  hydroxide2a.  Data  on  formation  constants 
in  solution5  indicate  that  the  chelate  complexes  of  ethylenediamine  and 
other  polydentate  amines  are  usually  much  more  stable  than  the  corre- 
sponding ammonia  complexes. 

An  illustration  involving  compounds  of  a  different  type  is  found  in  the 
£-diketones  which  may  enolize  and  form  stable  six-membered  rings  with 
metal  atoms.  Representative  acetylacetonates  are  shown  in  Fig.  5.1.  The 
stability  of  the  metal  acetylacetonates  is  indicated  by  the  fact  that  they 
may  be  heated  without  decomposition  to  temperatures  well  above  that  at 
which  acetylacetone  itself  is  decomposed2.  This  remarkable  stability  con- 
trasts sharply  with  the  low  stability  of  coordination  compounds  containing 
simple  ketones  such  as  acetone. 

The  formation  of  fused  rings  around  the  metal  seems  to  confer  an  even 

4.  Schwarzenbach,  Chimin,  3,  1  (1949);  Schwarzenbach  and  Gysling,  Helv.  Chim. 

Acta,  32,  1314  (1949) ;  Schwarzenbach  and  Willi,  HeUf.  Chim.  Ada,  34,  528  (1951); 
and  other  papers  in  the  series  on  metal  indicators. 

5.  Schwarzenbach,  Helv.  Chim.  Acta,  35,  2344  (1952). 


222 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


H— C 


/ 


CH3 
I 

C= 


V 


CH- 


\ 


.o- 


Be 


</      V 


CH- 
I 
C 


H3C-C 


A 


<\H3 

J* 

>-HH-C 

I  / 


CH, 


CH- 


C-CH3 


o        o 

Al 


CH- 


N 


XCH. 


C-H 


[Xl  (a  c  30)3] 


M.P.=  192 


[Be  (acac^J 

B.P.=  270°  B.P.=3I4V 

Fig.  5.1.  Acetylacetone'complexes  of  beryllium  and  aluminum 

greater  stability  than  the  formation  of  single  rings.  For  instance,  copper(II) 
ethylenediamine-bis-acetylacetone,  which  contains  three  interlocked  rings, 


CH3 

s 

/ 

CH3 


Cu 

/    \ 

■N  N: 


/ 

X 


CH, 


./ 


CH 


CH- 


■CH. 


CH- 


may  be  heated  nearly  to  redness  without  suffering  decomposition6.  Calvin 
and  Bailes7  made  a  polarographic  study  of  the  compounds  (A)  and  (B) 
(Fig.  5.2)  and  reported  that  the  reduction  potentials  indicate  much  greater 


CH3    CH- 


Ei 

2 


_A_ 

+  0.02  (reduction 


6.  Morgan  and  Smith,  J.  Chem.  Soc,  127,  2030  (1925). 

7.  Calvin  and  Bailes,  J.  Am.  Chem.  Soc,  68,  953  (1946). 


THEORY  OF  HETEROCYCLIC  RING  FORMATION 


223 


o      o- 

Cu 
=N         N=C 


\ 


I      I 

CH2-CH2 


B 


Ej_    -  -0.75 

2 

Fig.  5.2.  Polarographic  comparison  of  chelated   and   nonchelated  structures 

stability  for  the  interlocked  three  ring  system,  (B),  than  for  the  comparable 
two  ring  system,  (A).  Other  examples  have  also  been  cited. 

Of  even  more  interest  are  the  biologically  important  metal  porphyrin  de- 
rivatives which  are  constituents  of  chlorophyll  X  and  hemin  (Chapter  21). 
These  have  completely  interlocked  ring  systems  (Fig.  5.3).  Such  materials 
and  the  structurally  similar  phthalocyanines  (Chapter  22) 


<6  R5 

Fig.  5.3.  The  porphyrin  ring  system 

are  very  .-table  in  acid  solution.  In  fact,  the  copper  phthalocyanine  complex 
i-  reported  to  be  .-table  in  the  vapor  phase  at  500°C. 

The  stability  of  multiple  ring  systems  has  been  utilized  extensively  in 
the  commercial  applications  of  ethylenediaminetetraacetic  acid,  salts  of 
which  are  sold  under  such  trade  names  as  "Yersene,"  "Sequestrene,"  and 
"Nullapon."  Schwarzenbach  has  published  an  outstanding  series  of  papers 
on  the  stability  of  such  system-,  varying  a  number  of  structural  factors 
in  the  ligand.  The  enhanced  stability  conferred  on  a  complex  as  a  result  of 
ring  formation  has  been  termed  the  "chelate  effect"  by  Schwarzenbach'.  A 


■2-2-2 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


H— C 


/ 


CH- 

I     ' 


'C— 

I 

CH- 


Be 


CH- 
I 
C 


\ 


H3C-C 


/\ 


C-CH, 


CH- 


I 


c-hh-c 


V 


/ 

CH3 


O  O 

Al 

0/  % 


CH- 


X 


C-H 


[be  (acac)2J 


[ai  (acac)^ 

M.P.=  192° 
B.P.=  270°  B.P.=3I4° 

Fig.  5.1.  Acetylacetone_complexes  of  beryllium  and  aluminum 


greater  stability  than  the  formation  of  single  rings.  For  instance,  copper(II) 
ethylenediamine-bis-acetylacetone,  which  contains  three  interlocked  rings, 


CH3 
\ 

/ 

CH3 


Cu 

/    \ 

■N  N: 


CH: 


■CH- 


\ 

\ 


CH- 


CH 


CH- 


may  be  heated  nearly  to  redness  without  suffering  decomposition6.  Calvin 
and  Bailes7  made  a  polarographic  study  of  the  compounds  (A)  and  (B) 
(Fig.  5.2)  and  reported  that  the  reduction  potentials  indicate  much  greater 


CH3     CH3 


El    =  +0.02    (  REDUCTION  \ 


6.  Morgan  and  Smith,  /.  Chem.  Soc,  127,  2030  (1925). 

7.  Calvin  and  Bailes,  J.  Am.  Chem.  Soc,  68,  953  (1946). 


THEORY  OF  HETEROCYCLIC  RING  FORMATION 


223 


CH2-CH2 


B 


E|     =  "0.75 

2 


Fig.  5.2.  Polarographic   comparison  of   chelated   and   nonchelated   structures 

stability  for  the  interlocked  three  ring  system,  (B),  than  for  the  comparable 
two  ring  system,  (A).  Other  examples  have  also  been  cited. 

Of  even  more  interest  are  the  biologically  important  metal  porphyrin  de- 
rivatives which  are  constituents  of  chlorophyll  X  and  hemin  (Chapter  21). 
These  have  completely  interlocked  ring  systems  (Fig.  5.3).  Such  materials 
and  the  structurally  similar  phthalocyanines  (Chapter  22) 


Re  R5 

Fig.  5.3.  The  porphyrin  ring  system 

are  very  -table  in  acid  solution.  In  fact,  the  copper  phthalocyanine  complex 
is  reported  to  be  stable  in  the  vapor  phase  at  500°C. 

The  stability  of  multiple  ring  systems  has  been  utilized  extensively  in 
the  commercial  applications  of  ethylenediaminetetraacetic  acid,  salts  of 

which  are  sold  under  such  trade  names  as  "Versene,"  "Sequesi  rene."  and 

"Nullapon."  Schwarzenbach  has  published  an  outstanding  series  of  papers 

on  the  stability  of  such  Bystems,  varying  a  number  of  structural  factors 
in  the  ligand.  The  enhanced  stability  conferred  on  a  complex  as  a  result  of 
ring  formation  has  been  termed  the  "chelate  effect''  by  Schwarzenbach1.  A 


224  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

review  of  the  factors  contributing  to  the  stability  of  complexes  will  be  a 
useful  starting  point  in  the  consideration  of  the  chelate  effect. 

Factors  Involved  in  Chelate  Stability 

Since  chelate  compounds  are  merely  a  special  class  of  coordination  com- 
pounds, all  factors  outlined  in  Chapters  3  and  4  are  important  in  deter- 
mining their  stability.  In  addition,  a  few  factors  assume  special  importance 
as  a  result  of  ring  formation  and  will  be  considered  specifically  here.  The 
question  of  solvation  effects  is  of  particular  importance  in  the  study  of 
chelate  compounds  since  many  of  the  large  organic  ligands  are  only  very 
slightly  soluble  in  water  so  their  complexes  have  been  studied  in  mixed 
solvents11,  or  in  organic  solvents9-  12.  If  solvation  terms  (p.  138)  were  truly 
negligible,  the  choice  of  solvent  would  be  of  minor  importance.  That  such 
is  not  always  the  case  is  shown  by  a  number  of  investigations  (e.g.,  Refs. 
12,  13).  In  fact,  in  organic  solvents,  a  metal  cation  and  its  anion ^are  usually 
associated.  An  interesting  correlation  of  observations  in  mixed  solvents  and 
in  water  was  given  by  Van  Uitert  and  Haas12b.  Van  Uitert,  Fernelius, 
Douglas,  and  their  co-workers12, 13  have  applied  data  from  mixed  solvents 
to  the  study  of  many  different  chelate  systems.  Trotman  and  Dickenson10 
suggest  that  solvation  energy  terms  may  even  be  of  major  importance  in 
determining  the  relative  stabilities  of  some  non-chelated  complexes,  such 
as  the  silver  ammines. 

It  is  important  to  note  that  in  the  thermochemical  (p.  138)  cycle  en- 
tropy effects  have  been  neglected  and  the  change  in  heat  content,  AH,  is 
taken  as  an  approximate  measure  of  the  change  in  free  energy,  AF,  which 
determines  the  stability  of  the  compound.  In  a  consideration  of  the  " chelate 
effect"  the  entropy  terms  are  so  large  that  they  can't  be  neglected,  even  as 
a  first  approximation.  These  effects  are  discussed  in  more  detail  in  a  later 
section.  Since  AF  =  AH  —  TAS,  a  consideration  of  factors  influencing 
both  AH  and  AS  is  appropriate.  It  will  be  convenient  as  a  conventional 
simplification  to  assume  that  AH  is  determined  in  large  measure  by  the 
energy  of  coordination  (see  p.  138)  (i.e.,  the  energy  for  the  processes): 

Jlf(f)*  +  yABigf  ->  [M(AB),]l0)*™ 

8.  Bjerrum,  Chem.  Revs.,  46,  381  (1950). 

9.  Burkin,  ./.  Chem.  Soc,  1954,  71;  Jonassen,  Fagley,  Holland,  and  Yates,  J.  Phys. 

Chem.,  58,286  (1954). 

10.  Trotman  and  Dickenson,  ./.  Chem.  Soc,  1949,  1293. 

11.  Calvin  and  Wilson,  ./.  Am.  Chem.  Soc,  67,  2003  (1945). 

12.  VanUiterl ,  Fernelius,  and  Douglas, ./.  Am.  Chem.  Soc,  75,  3577  (1953);  VanUiterl 

and  Baas,  •/.  Am.  Chem.  Soc,  75,  451  (1953);  VanUitert,  Fernelius,  and  Doug- 
las; ./.  Am .  Chem.  Soc,  75,  457  (1953);  VanUitert .  Haas,  Fernelius,  and  Doug- 
las, ./.  Am.  Chem.  Soc,  75,  455  (1953). 
L3.  VanUitert,  Fernelius,  and  Douglas,  J.  Am.  Chem.  Soc,  75,  2736,  2739  (1953). 


THEORY  OF  HETEROCYCLIC  RING  FORMATION 

The  energy  of  coordination  may  then  be  considered  iii  terms  of  Steric  lac- 
tors  for  both  the  central  ion  and  the  ligand  which  arise  from  chelation  and 
electronic  factors  for  both  components  of  the  complex,  which  are  peculiar 

to  chelate  Bystems. 

Steric  Factors  in  Chelate  Ring  Formation 

Ring  Size 

Bonds  in  coordination  compounds  may  arise  from  two  general  types  of 
groups:  (1  I  primary  acid  groups  in  which  the  metal  ion  replaces  an  acid 
hydrogen  and,  (2)  neutral  groups  which  contain  an  atom  with  a  free  elec- 
tron pair  suitable  for  bond  formation.  If  two  groups  from  either  class  1  or 
2  or  from  classes  1  and  2  are  present  in  the  same  molecule  in  such  positions 
that  both  groups  can  form  bonds  with  the  same  metal  ion,  a  chelate  ring 
may  be  formed.  When  the  groups  are  present  in  such  positions  as  to  form  a 
five-  or  si.\-membered  ring,  the  resulting  complex  is  most  stable,  although 
4-,  7-,  S-  and  even  larger  rings  are  known  (Chapter  6).  The  existence  of 
three-membered  rings  has  not  been  established. 

Evidence  on  Three -Membered  Rings 

In  a  review  of  the  coordination  compounds  of  hydrazine,  Audrieth  and 
Ogg14  point  out  the  interesting  fact  that  in  a  surprisingly  large  number  of 
cases,  the  number  of  hrydrazine  groups  coordinated  to  a  metal  ion  is  one-half 
the  normal  coordination  number  of  the  metal.  Since  no  structural  determi- 
nations have  been  made,  the  possibility  of  a  three-membered  chelate  ring 
cannot  be  definitely  eliminated;  however,  the  low  solubilities  of  most  of 
these  compounds  suggest  polynuclear  structures  involving  hydrazine 
bridges  rather  than  chelate  structures.  The  complexes  [PtCl2(N2H4)]  and 
[PdBi v  X2H4)]15  are  probably  dimers  of  the  type: 

("1  CI  CI 

\    /     \     / 
Pt  Pt 

/      \     /     \ 
\JU  CI  N0H4 

In  one  of  the  few  cases  ID  which  hydrazine  complexes  have  been  studied 
in  solution,  Rebertus,  Laitinen,  and  Bailar16  found  that  the  zinc(II)  ion  will 
coordinate  four  hydrazine  molecules  with  only  small  differences  between 
the  separate  dissociation  constants;  this  indicates  >t rongly  that  hydrazine  is 

14.  Audrieth  and  Ogg  "The  Chemistry  of  Hydrazine,"  p.  181,  New  York,  John  Wiley 

as,  Inc.,  1961 . 

15.  Goremykin  and  GladyahevBkaya,  •/.  Oen.  Ck  m.  [UJ3M  R.)  13,  762  (1943);  14,  13 

(1944. . 

16.  Rebertus,  Laitinen,  and  Bailar,  •/.  Am.  «.,  75,  3051  (1953). 


226  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

monodentate  with  the  normally  four  coordinate  zinc (II)  ion.  A  similar 
study  conducted  by  Schwarzenbach  and  Zobrist17  indicated  that  four  hy- 
drazine molecules  are  bound  to  zinc(II)  and  six  to  nickel (II)  in  a  manner 
comparable  to  the  binding  of  ammonia  to  these  metals.  They  concluded 
thai  no  three-membered  chelate  rings  were  ever  formed. 

Finally,  no  well  authenticated  case  of  optical  isomerism  which  might  be 
used  as  evidence  for  a  chelate  ring  structure  has  been  observed  with 
hydrazine  complexes.* 

Four-Membered  Rings 

The  stereochemistry  of  metal  chelate  rings  differs  from  that  of  carbon 
ring  systems  in  that  all  of  the  atoms  in  the  ring  are  not  the  same  size  and 
some  of  the  bond  angles  normally  vary  from  109°  (or  120°)  as  a  result  of  the 
directed  valences  of  the  metal  ion.  These  two  factors  may  relieve  the  in- 
stability of  four-membered  ring  systems.  For  example,  the  carbonate 
group  in  [Co  en2  C03]+  occupies  two  positions  to  give  a  rather  stable  four- 
membered  ring.  Scale  drawings  of  this  ring,  using  Pauling's  covalent  radii, 
indicate  that  the  steric  strain  is  much  less  than  in  a  corresponding  four- 
membered  carbon  system.  Similarly,  sulfate,  sulfite,  thiosulfate,  thiocar- 
bonate,  selenate,  selenite,  molybdate;  and  chromate  can  each  occupy  two 
positions  in  the  coordination  sphere2, 18a-  19.  (See  also  p.  180  for  electronic 
interpretations.)  Four-membered  rings  are  very  common  in  bridged  mole- 
cules such  as: 

XXX  R3P  CI  PR3 

\    /   \    /  \    /    \    / 

Al  Al  and  Pt  Pt 

/     \    /     \  /     \    /     \ 

XXX  CI  CI  CI 

(p.  18  and  22).  The  formation  of  four-membered  oxo-bridges  in  basic  solutions 
of  chromium(III)  is  of  great  importance  in  the  leather  tanning  industry 
(Chapter  13). 

Unusual  four-membered  rings  have  been  reported  by  Dwyer  and  Mel- 
lor21  • 22,  who  found  that  copper,  nickel,  palladium,  and  silver  ions  form 
complexes  with  triazene  derivatives  which  are  much  more  stable  than  the 

*  A  report  that  [Co(N2H<)3]Br3  has  been  resolved  into  optical  isomers  is  a  typo- 
graphical error.  The  ligand  should  be  ethylenediamine,  not  hydrazine.  (Wells,  "Struc- 
tural Inorganic  Chemistry,"  p.  530). 

17.  Schwarzenbach  and  Zobrist,  Helv.  Chim.  Acta,  35,  1291  (1952). 

18.  Riley,  /.  Chem.  Soc,  1928,  2985;  1929,  1307;  1930,  1642. 

19.  Briggs,  /.  Chem.  Soc,  1929,  685. 

20.  Yoe  and  Sarver,  "Organic  Analytical  Reagents,"  New  York,  John  Wiley  &  Sons, 

Inc.,  1941. 

21.  Dwyer,  J.  Am.  Chem.  Soc.,  63,  78  (1941). 

22.  Dwyer  and  Mellor,  J.  Am.  Chem.  Soc,  63,  81  (1941);  Dwyer,  /.  Am.  Chem.  Soc, 

63,  78  (1951). 


THEORY  OF  HETEROCYCLIC  RING  FORMATION  227 

parent  triazene.  They  withstand  the  action  of  boiling  hydrochloric  acid  and 
concentrated  alkali;  some  of  them  are  stable  at  temperatures  above  300°C. 
The  following  structure  has  been  suggested: 

N 

/    \ 
R— N  N— R 

\     / 
M 

/     \ 
R— N  N— R 

\    / 

N 

One  would  expect  a  ring  of  this  type  to  be  somewhat  strained,  but  the  un- 
usual stability  of  the  compounds  gives  no  indication  of  this.  It  is  observed, 
however,  that  at  low  temperatures  the  compound  dimerizes,  a  process 
which  could  relieve  strain  by  opening  the  rings  and  crosslinking  the  metal 
atoms.  Four-membered  diamagnetic  nickel  chelate  rings  of  ethylxanthoge- 
nate 

S  S 

/\    /  X 

C2H5— O— C  Ni  C— 0— C2H5 

\   /     \/ 

s  s 

and  nickel  ethyl  dithiocarbamate, 

H  S  S  H 

I         /\     /  \         I 
C2H5— N— C  Ni  C— N— C2H6 

\   /      \/ 

s  s 

have  been  described23. 
Five-Membered  Rings 

Five-  and  six-membered  rings  are  very  common.  Hundreds  of  examples 
of  each  type  have  been  described2- 20, 24, 25>  26.  In  general,  it  is  observed  that 
saturated  compounds  tend  to  form  five-membered  structures  whereas  those 
ligands  which  give  rings  with  two  double  bonds  tend  to  form  six-membered 
rings.  The  evidence  for  a  five-membered  saturated  ring  arises  from  Beveral 
unrelated  types  of  experiments.  For  example,  1,2,3-triaminopropane, 

NH2     NH2     NH2 

I            I  I 

II— c c C— II, 

I      I      I 

II       II       II 

23.  Cambi  and  Szego,  Ber.,  64,  2591  (1931). 


228  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

can  react  with  a  metal  so  as  to  occupy  only  two  coordination  positions,  the 
third  amine  group  then  being  capable  of  salt  formation.  The  compound  of 
this  type  formed  with  platinic  chloride  will  then  be  either  disymmetric  (A) 
or  symmetrical  (B),  according  as  a  five-  or  six-membered  ring  is  formed 
preferentially  by  chelation.  Mann27  was  able  to  resolve  the  complex,  estab- 
lishing the  existence  of  the  five-membered  ring  A. 


NH2— CH2 

NH2- 

-CH2 

/ 

/ 

\ 

Cl4Pt 

CUPt 

CH— NHaHX 

\ 

\ 

/ 

NH2— CH 

NH2- 

-CH2 

CH2— NH2HX 

A.  Resolvable  B.  Nonresolvable 

Five-membered  Ring  Six-membered  Ring 

Fig.  5.4.  Chelation  of  1,2,3-triaminopropane 

Another  example  is  found  in  the  fact  that  ethylenediamine  forms  very 
stable  five-membered  chelate  rings.  The  presence  of  substituents  on  the 
carbon  does  not  disturb  the  five-membered  ring  and  thus  has  only  a  minor 
effect  on  the  color  and  stability  of  the  coordination  compound.  The  co- 
balt(III)  compounds  containing  propylenediamine  and  2 , 3-butylene- 
diamine  are  similar  to  their  ethylenediamine  homologs  in  ease  of  formation, 
stability  and  color.  Other  substituted  ethylenediamines  such  as  meso- 
stilbenediamine,  isobutylenediamine28,  cyclopentanediamine29,  and  cyclo- 
hexanediamine29  form  very  stable  coordination  compounds  comparable  to 
their  ethylenediamine  parent.  On  the  other  hand,  a  very  different  effect  is 
produced  by  increasing  the  number  of  carbon  atoms  between  the  amine 
groups,  since  this  expands  the  ring.  Trimethylenediamine  forms  six- 
membered  chelate  rings  with  cobalt30,  nickel31,  platinum31, 32,  and  iron33; 

24.  Flagg,  "Organic  Reagents  in  Gravimetric  and  Volumetric  Analysis,"  New  York, 
Interscience  Publishers,  Inc.,  1948. 

26.  Mellan,   "Organic  Reagents  in  Inorganic  Analysis,"  p.  53,  Philadelphia,  The 

Blakiston  Co.,  1941;  Freudenberg,  "Stereochemie,"  Vol.   3,   p.    1200,  Franz 
Deuticke,  Leipzig  and  Wien,  1932. 

27.  Mann,  J.  Chem.  Soc,  129,  2681  (1926). 

28.  Mills  and  Quibbell,  J.  Chem.  Soc,  1935,  839;  Lidstone  and  Mills,  J.  Chem.  Soc., 

1939,  1754. 

29.  Jaeger  and  terBerg,  Proc.  Acad.  Sci.  Amsterdam,  40,  490  (1937) ;  Jaeger  and  Bijerk, 

Proc.  Acad.  Sci.  Amsterdam,  40,  12, 116,  316  (1937) ;  Z.  anorg.  allgem.  Chem.,  233, 
97  (1937);  earlier  articles  by  Jaeger. 

30.  Werner,  Ber.,  40,  61  (1907). 

31.  Tschugaeff,  Ber.,  39,  3190  (1906);  /.  prakt.  Chem.  [2]  75,  159  (1907);  12]  76,  89 

(1907). 


THEORY  OF  HETEROCYCLIC  RINQ  FORMATION  229 

available  evidence  indicates  that  such  compounds  are  less  stable  and  more 
difficult  to  prepare  than  the  analogous  propylenediamine  compounds  con- 
taining five-membered  rings.  Bailar  and  Work*  found  thai  neopentane- 
diamine,  NHsCHsC(CHi)sCHiNHs ,  coordinates  more  readily  and  gives 
more  stable  compounds  than  docs  trimethylenediamine,  HA'  CHj 
CHi  CHj  NHi.  This  unexplained  observation  contrasts  sharply  with 
the  fact  that  propylenediamine,  2,3-but  yleiicdiamiiie  and  many  other 
2,3-diamines  strongly  resemble  ethylenediamine  in  their  complexing  be- 
havior.  In  the  latter  case,  substitution  on  the  carbon  docs  not  greatly  alter 
the  complexing  properties. 

A  second  line  of  evidence  has  been  obtained  by  Schwarzenbach5  from  a 
consideration  of  the  formation  constants  of  metal  complexes  related  to 
ethylenediaminetetraacetates,  and  of  the  general  type: 


(I) 


The  value  of  n  varied  from  2  to  5,  giving  five-,  six-,  seven-,  and  eight- 
membered  chelate  rings  involving  the  nitrogen  atoms.  The  corresponding 
imino  diacetate  complexes 

O 

/ 

CH2— C— O— 

/ 
IIX  (II) 

\ 

CH2— C— 0— 
\ 

o 

were  studied  as  standards  in  which  no  chelate  ring  formation  involving  only 
nitrogen  atoms  was  possible.  Data  indicate  that  when  n  =  2  the  stabiliza- 
tion due  to  the  chelate  ring  formation  is  a  maximum.  As  the  chain  length 
(value  of  n)  increases,  the  stabilizing  effect  due  to  chelation  disappears  and 
en  replaced  by  a  slight  destabilizing  effect.  It  was  also  observed  that 

■>l.  Drew  and  Tress,  ./.  Chem.  8oc.,  1933,  1335. 

Breuil,  Campt.  rend.,  199,298  (1931,. 
34.  Pfeiffer  and  Bainmann,  Ber.t  36,  10G4  (1903). 

Wilar  and  Work,  J.  Am.  Chem.  Soc,  68,  232  (1946). 
36.  Schwarzenbach  and  Ackerman,  Hclv.  Chim.  Acta,  32,  1682  (1949). 


0 

0 

\        H 

H 

/ 

0— c— c 

C- 

-C— 0 

H\ 

/H 

N— (CH2)n- 

-X 

H/ 

\H 

0— c— c 

C- 

-C— 0 

/        H 

11 

\ 

0 

0 

230 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


as  the  chain  length  is  increased  the  tendency  of  the  ligand  to  bind  two 
separate  metal  ions  increases  rapidly  so  the  formation  of  polynuclear  com- 
plexes takes  place.  Similar  results  were  reported  by  Schwarzenbach  and 
Ackerman36  from  their  study  of  the  isomeric  diaminocyclohexane-N,N'- 
tetraacetates  (Fig.  5.5)  coordinated  with  the  alkaline  earth  ions.  The  cal- 


P 

CH2—  C-O- 


CH2-C-0- 

,P 

CH2~C-0- 


Cf-U-C-0 


Fig.  5.5.  l,2-Diaminocyclohexane-N,N'  tetraacetate 


cium  chelate  compound  of  the  1,2  isomer,  which  contains  a  five-membered 
chelate  ring,  is  even  more  stable  (K  =  1012*5)  than  the  ethylenediamine 
tetraacetate  complex  (K  =  1010-5).  On  the  other  hand,  the  1,3  and  1,4 
derivatives  which  would  give  badly  strained  ring  structures  in  the  metal 
complexes  are  much  less  stable  and  show  a  strong  tendency  to  coordinate 
with  two  metal  cations  rather  than  to  form  a  ring. 

Schwarzenbach6  also  reports  formation  constants  for  complexes  of  ethyl- 
enediamine and  trimethylenediamine  which  confirm  the  greater  stability  of 
the  five-membered  metal-nitrogen  ring. 

The  stability  of  five-membered  rings  is  not  restricted  to  the  coordination 
of  amines.  Dey37  compared  the  efficacy  of  dicarboxylic  acids  in  the  forma- 
tion of  coordination  compounds  with  tin.  He  found  the  order  of  decreasing 
complexing  power  to  be  oxalic,  malonic,  and  succinic  acids.  This  corresponds 
to  a  decrease  in  chelate  stability  as  one  goes  from  a  five-  to  a  seven-mem- 
bered  ring. 

Similar  observations  were  made  by  Riley18.  He  found  that  the  stability 
of  complexes  formed  between  the  Cu+2  ion  and  the  oxalate,  malonate,  and 
succinate  ions  decreased  in  the  order  listed.  Electronic  effects  cannot  justify 
this  observation  since  succinate  ion  is  a  stronger  base  than  oxalate38.  Re- 
cently Courtney,  Chabarek,  and  Martell39  found  that  if  the  acetate  groups 
of  ethylenediaminetetraacetate  are  replaced  by  propionate  groups  to  give 
terminal  rings  of  six  rather  than  five  members,  the  stability  of  the  chelate 
is  reduced. 

37.  Dey,  Univ.  Allahabad  Studies,  Chem.  Sect.,  1946,  7;  [Chem.  Abs.,  41,  6169  (1947)]. 

38.  Hixon  and  Johns,  /.  Am.  Chem.  Soc,  49,  1786  (1927). 

39.  Courtney,  Chaberek,  and  Martell,  J.  Am.  Chem.  Soc.,  75,  4814  (1953). 


THEORY  OF  HETEROCYCLIC  RING  FORM  \TI<>\ 


233 


Rings  of  Six  or  More  Members 

In  general  it  is  found  that  stable  chelate  rings  involving  two  double  bonds 
are  usually  six-membered  structures.  Thus  acetylacetone  and  salicylalde- 

hyde  and  their  derivatives  coordinate  readily  to  give  very  stable  six- 
membered  chelate  complexes: 


CH3 
C 

■c      o 


CH3   \/ 


/\/\ 


v 


SAL1CYLALDEHYDE 
ACETYLACETONE  CHELATE  CHELATE 

If  only  one  double  bond  is  present  in  the  ring,  both  five-  and  six-membered 
structures  are  common,  with  the  five-membered  unit  appearing  somewhat 
more  frequently  in  the  usual  descriptions.*!  Heller  and  Schwarzenbach40 
examined  iron(III)  complexes  of  pyrocatechindisulpho  acid  and  chromo- 
tropic  acid.  In  the  former  case  (A)  a  five-membered  ring  involving  one 
resonating  double  bond  is  formed  and  in  the  latter  case  (B)  a  comparable 

+++ 


X" 


Fe 


+++ 


REMOVE 


SO: 


o(h- 


SQ- 


PYROCATECHIN  COMPLEX 

OF  Fe+++ 

A 


REMOVE 


CHROMOTROPIC  ACID 

COMPLEX  OF  Fe  "^ 

B 


Fig.  5.6 


*  Lowry41  attempted  to  justify  the  stability  of  six-membered  rings  on  the  basifl  of 
alternating  polarity  with  the  metal  atom  as  a  negative  group.  Using  this  hypothesis, 
he  concluded  that  six-memhered  rings  are  more  stable  than  those  containing  live 
members.  The  limitations  of  this  concept  are  obvious  from  the  discussion  on  ring 
size. 

t  Bobtelsky  and  Bar-Gadda"  conclude  that  a  double  bond  in  a  ring  is  apparently 
effective  in  stabilizing  even  a  seven-memben-d  ring. 
40.  Heller  and  Schwanenbaeh,  HeUf.  Ckim.  Ann,  34,  1876  (1951). 


232  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

six-membered  ring  is  produced.  The  values  of  the  formation  constants  were 
given  as: 

FeX  +  A"4  ->  [FeXA]~4  log  K  =  15.7  ±  0.4 

FeX  +  B"4  ->  [FeXB]~4  log  K  =  17.0  ±  0.5 

where  X  =  anion  of  nitrilotriacetic  acid.  The  differences  in  the  formation 
constants  are  smaller  than  the  differences  in  the  acid  constants  of  the  parent 
compounds,  thus  indicating  little  influence  due  to  ring  size. 

The  problem  of  ring  size  also  arises  in  the  discussion  of  citrate  and  tar- 
trate complexes.  A  variety  of  formulas  has  been  proposed  which  involve 
rings  of  various  sizes20, 43, 44, 45> 46.  It  has  been  established  that  the  citrate  ion 
can  lose  its  hydroxyl  hydrogen  as  well  as  the  carboxyl  hydrogens  and  can 
coordinate  with  a  bivalent  metal  such  as  copper  even  in  acid  solution46, 47. 
This  suggests  the  possibility  of  the  formation  of  both  six-  and  seven-mem- 
bered  rings  in  the  citrate  complexes,  the  six-membered  ring  probably  form- 
ing preferentially46  • 47 : 


/CH2  CH2 

o=cr  <r  c=o 

The  fact  that  tartrate  complexes  are  in  general  more  stable  than  the  analo- 
gous succinate  complexes  and  that  citrate  complexes  are  more  stable  than 
tricarballylate  complexes  also  indicates  the  involvement  of  the  OH  groups 
in  the  chelation  process. 

Rings  of  seven  or  more  members  are  comparatively  uncommon,  but  are 
well  established  (Chapter  6).  As  the  length  of  the  chain  between  the  two 
donor  atoms  increases,  so  does  the  tendency  to  form  polymetallic  complexes. 

A  few  interesting  exceptions  to  the  foregoing  generalizations  are  known. 
Thus,  the  dimethyl  glyoxime  chelate  ring  with  nickel  involves  twTo  double 

*  Alternatively  both  rings  may  form  on  the  same  metal  to  give  the  ion  [MCi]~. 

41.  Lowry,  Chemical  &  Industrial,  42,  715  (1923). 

42.  Bobtelsky  and  Bor-Gadda,  Bull.  soc.  chim.  France,  1953,  382. 

43.  Paulinova,  /.  Gen.  Chem.  (U.S.S.R.)  17,  3  (1947);  [Chem.  Abst.,  42,  53  (1948)]. 

44.  Bobtelsky  and  Jordan,  J.Am.  Chem.  I Soc.  ,67,  1824  (1945);  69,  2286  (1947);  75,  4172 

(1953). 

45.  Harada,  Sci.  Papers  Inst.  Phys.  Chem.  Research  (Tokyo)  41,  68  (1943),  [Chem. 

Abs.,  41,  6206  (1947)]. 

46.  Parry  and  DuBois,  /.  Am.  Chem.  Soc,  74,  3752  (1952). 

47.  Warner  and  Weber,  J.  Am.  Chem.  Soc,  75,  5086  (1953). 


THEORY  OF  H  FTFh'OCYCUC  RING  FOR M  A  Tin  A 


233 


bonds  and  may  be  formulated  as  a  five-  or  six-membered  structure: 


R 

1 

C 

/    \ 
R— C             N- 

II               1 
N            Ni 

\   / 

0 

OH 

Ni 

/     \ 
HON              N— >0 

II                II 

r— c c     n 

Six-membered 

ring 

Five-membered  ring 

The  original  formulation48  of  the  structure  as  a  five-membered  ring  was 
based  on  the  fact  that  the  anti-glyoxime  is  the  only  isomer  which  gives  the 
characteristic  red  nickel  salt. 


OH 


R— C=N 


R— C=N 


R— C=N 


anti 


R— C=N 


R— C=N     OH 


OH 


OH 


OH 


OH 


amphi 


R— C=N 

syn 


These  stereochemical  deductions  have  been  supported  completely  by  recent 
x-ray  data49.  Examination  of  the  structure  of  the  entire  molecule  makes  the 
choice  of  five-membered  rings  reasonable  even  though  two  double  bonds 
are  involved.  As  Fig.  5.6  shows,  the  formation  of  five-membered  rings  gives 


/oh\ 

C=N  N=C-R 

Ni 
C=N  ^N=C-R 

\>HO 

Alultiple        ring 
formation        with 
five-membered 
ring  and  hydrogen 
bonds. 


R-C 


N- 


0\       N  = 

:N  O 


•" 


// 


C-R 


Only      two      possible 

rings  if  ring  is  six-mem- 
bered. 


Fig.  5.7.  Possible  structures  of  nickel  dimethylglyoxinn' 
the  possibility  of  multiple  ring  formation  through  hydrogen  bonding.  I 
48.  Pfeiffer,  Ber.,  63,  1811  (1930). 


234  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

dci ice  cited  earlier  indicates  a  marked  increase  in  stability  arising  from  the 
presence  of  multiple,  interlocked  rings.  It  is  of  some  interest  to  note  that 
the  hydrogen  bond  in  this  complex  is  the  shortest  yet  reported49. 

Another  interesting  exception  is  found  in  the  complexes  of  silver. 
Schwarzenbach  and  his  co-workers50  report  that  the  complexes  of  silver  (I) 
with  trimethylenediamine,  tetramethylenediamine,  and  pentamethylene- 
diamine  (six-,  seven-,  and  eight-membered  rings)  are  all  more  stable  (log 
K  =  5.85,  5.90,  5.95,  respectively)  than  the  corresponding  silver  complex 
with  ethylenediamine  (log  K  =  4.7).  This  is  attributed  to  the  fact  that  the 
two  bonds  of  silver  are  linear  and  the  longer  membered  chains  are  better 
able  to  form  rings  than  are  the  shorter  chains.  Such  an  interpretation  re- 
ceives further  support  from  the  fact  that  the  complex  [Ag2  en2]+2  is  formed 
and  was  isolated  as  the  crystalline  sulfate.  The  molecular  weight  was  con- 
firmed by  cryoscopic  measurements. 

Polydentate  Ligands — Multiple  Ring  Systems 

In  recent  years  ligands  capable  of  occupying  as  many  as  six  coordination 
positions  on  a  single  metal  ion  have  been  described.  Studies  on  the  for- 
mation constants  of  coordination  compounds  with  these  ligands  have  been 
reported39- ».».«.«.  In  general  it  is  observed  that  the  stability  of  the 
complex  goes  up  with  an  increase  in  the  number  of  groups  available  for  co- 
ordination. Other  studies,  particularly  those  involving  the  preparation  of 
penetration  complexes  of  cobalt,  are  of  considerable  interest.  Three  types 
of  chelating  agents  have  been  placed  around  all  six  of  the  coordination 
positions  of  cobalt  (III).  They  are: 


— ooc- 

-CH2 
\ 

CH2COO— 
/ 

N- 

-CH2— 

-CH2- 

-N 

/ 

\ 

— ooc- 

-CH2 

(A) 

CH2COO— 

49.  Godycki  and  Rundle,  Acta  Cryst.,  6,  487  (1953). 

50.  Schwarzenbach,  Maissen,  and  Ackermann,  Helv.  Chim.  Acta,  35,  2333  (1952); 

Schwarzenbach,  Ackermann,  Maissen,  and  Anderegg,  Helv.  Chim.  Acta,  35, 
2337  (1952). 

51.  Jonassen,  LeBlanc,  and  Rogan,  J.  Am.  Chem.  Soc,  72,  4968  (1950). 

52.  Chaberek  and  Martell,  J.  Am.  Chem.  Soc,  75,  2888  (1953);  Lumb  and  Martell, 

J.  Am.  Chem.  Soc,  75,  690  (1953). 

53.  Chaberek,  Courtney,  and  Martell,  J.  Am.  Chem.  Soc,  74,  5052  (1952);  75,  2185 

(1953);  Courtney  and  Martell,  J.  Am.  Chem.  Soc,  74,  5057  (1952);  Chaberek 
and  Martell,  J.  Am.  Chem.  Soc,  74,  6021,  6228  (1952). 


THEORY  OF  HETEROCYCLIC  RING  FORMATION  235 

\ll  CH  CB  I  ii  CH1NH1 

\  / 

NT —  CHj —  CH       \ 

/  "  \ 

MI,CII2CH2  CM, (Ml  All 

(B) 

ethylenediaminetetraacetate  (A),  tetrakis(2-aminoethyl)ethylenediamine 
(  B),  and  compounds  of  the  general  form: 

H 

^n-(ch,)x-s-(chA-s-(ch2)z-n/ 

OH  HO^-^ 

(C) 
(X,Y,  ANDZ   HAVE    BEEN2  0R3) 

of  which  3,6-dithia,l)8-bis(salicylideneamino)octane,  (C),  is  an  example. 
Schwarzenbach54  showed  that  cobalt(II)  may  fill  only  five  of  its  coordi- 
nation positions  with  ethylenediaminetetraacetate  and  the  sixth  with  an 
auxiliary  ligand  such  as  Br~,  H20,  or  CNS~.  The  stable  penetration  com- 
plex of  cobalt  (III),  [Co(Y)Br]=,  can  be  prepared  from  the  cobalt(II)  salt 
by  oxidation.  On  the  other  hand,  the  cobalt  (III)  ion  can  satisfy  all  of 
its  coordination  positions  with  ethylenediaminetetraacetate  to  give  the 
sexicovalent  complex,  [Co(Y)]~.  This  ion  can  be  produced  by  complete 
substitution  of  the ligands  from  other  cobalt(III)  complexes:  [Co(XH3)6]+++ 
+  H4Y  ->  4XH4+  +  2XH3  +  [CoY]-.  Cis-  and  trans-[Co  en2Cl2]+  and 
[Co (ox) 3]-  behave  in  the  same  way.  No  intermediates  have  yet  been  iden- 
tified. Bailar  and  Busch55  confirmed  the  sexidentate  character  of  the  salt 
by  examination  of  its  infrared  spectrum  and  by  the  resolution  of  the  com- 
plex into  optical  isomers.  They  also  reported  that  the  elimination  of  the 
extra  substituent  (i.e.,  Br)  in  the  pentadentate  complex  [Co(Y)Br]=  pro- 
ceeds without  complete  loss  of  optical  activity. 

Schwarzenbach  and  A  loser56  have  also  prepared  complexes  of  Fe+++, 
Co"1-1-1",  and  Ni++  with  the  amine  analog  (B)  of  ethylenediaminetetraacet  ic 
acid;  these  appear  to  be  sexidentate  structures. 

Dwyer,  Lions,  Gill,  and  Gyarfas57  have  synthesized  main-  ligands  of  the 
third  type  (C),  and  have  formed  sexidentate  complexes  using  Co(III). 
Such  complexes  have  been  resolved  and  show  the  highest  optical  activity 

54.  Schwarzenbach,  Helo.  Chim.  Acta,  32,  841  (1949). 
56.  Bailar  and  Busch,  ./.  Am.  Chem.  Soc.,  75,  4574  (1953). 

56.  Schwarzenbach  and  Moeer,  Helv.  Chim.  Acta,  36,  681  (1963). 

57.  Dwyer,  Lions,  Gill,  and  Gyarfaa,  Nature,  168,  29  <  1  < »5 1  * ;  ./.  .1///.  Chem.  Soc,  69, 

2917  (1947);  72,  1545,5037  (1950);  74,  4188  (1952);  75,  2443  (1953);  76,  383  (19J 


236 


(  UKMISTRY  OF  THE  COORDINATION  COMPOUNDS 


yet  recorded.  They  can  be  represented  schematically  as: 

-  + 


(CH2)Y 

A  ligand  containing  one  oxygen  in  place  of  a  sulfur  also  serves  as  a  sexi- 
dentate  group ;  this  is  most  remarkable  in  that  an  ethereal  oxygen  is  coordi- 
nated firmly  to  cobalt  in  a  penetration  complex.  This  ability  of  stable 
terminal  groups  to  stabilize  unstable  ring  arrangements  in  the  complex  is 
interesting  but  not  unique  (Ref.  3  p.  142). 

Steric    Factors   Within    the    Complex.    Interference    by    Attached 
Groups:  F -Strain 

In  some  cases  the  clashing  of  groups  on  two  coordinated  ligands  will  re- 
sult in  a  distortion  of  bond  angles  and  a  decrease  in  stability.  This  is  the 
phenomenon  of  F-strain,  described  by  Brown58,  as  applied  to  coordination 
compounds.  A  number  of  experimental  observations  on  complex  compounds 
can  be  reasonably  interpreted  in  terms  of  steric  strain.  The  thermodynamic 
stability  of  N  and  N,N'-alkyl  substituted  ethylenediamines  has  been 
studied  by  a  number  of  investigators59, 60- 61, 62.  The  data  clearly  show  re- 
duction in  the  stability  of  the  complex  with  substitution  of  alkyl  groups  for 
hydrogen  atoms  on  the  nitrogen.  This  is  indicated  by  the  instability  con- 
stants for  the  nickel  complexes  in  Table  5.1  and  the  thermodynamic  values 
in  Table  5.2.  Steric  strain  or  F  strain  appears  to  offer  a  logical  though  not 
unique  interpretation  of  these  data. 

Data  of  Smirnoff63  and  Willink  and  Wibaut64  on  complexes  of  iron  (II)  sug- 

58.  Brown,  Bartholomay,  and  Taylor,  J.  Am.  Chem.  Soc,  66,435  (1944);  Brown  and 

Barbaras,  ./.  Am.  Chem.  Soc,  69,  1137  (1947),  and  other  papers,  H.  C.  Brown. 

59.  Keller  and  Edwards,  /.  Am.  Chem.  Soc,  74,  215  (1952);  74,  2931  (1952) ;  Edwards 

dissertation,  University  of  Michigan,  1950. 

60.  Irving  and  Griffiths,  J.  Chem.  Soc,  1954,  213. 

61.  Basoloand  Murmann,  J\  Am.  Chem.  Soc,  74,  5243  (1952);  76,  211  (1954). 
62    Mc  In  tyre,  dissertation,  Pennsylvania  State  College,  1953. 

63.  Smirnoff,  Helv.  Chim.  acta,  4,  802  (1921). 

64.  Willink  and  Wibaut,  Rec  Trav.  Chim.,  54,  275  (1935). 


THEORY  or  HETEROCYCLIC  RING  EORMATIOX 


237 


Table  5.1.  Stability  Constants  it  26°  oi    ran  Ni<  cbl  Complexes  of  Bomj 
Diamines  oi  the  Type  NRR'CH»CHtNHR* 

(Collected  by  [rving  and  ( iriiliths60) 


R 

R' 

R* 

log  a:, 

log  K* 

log  A,/A% 

i-K|,ir 

11 

U 

II 

7.60 

6.48 

1.12 

10.18 

Mo 

11 

H 

a 

g  IS 

7.36 

5.74 

1.62 

10.40 

i:t 

11 

II 

U  3d 

6.78 

5.30 

1.48 

10.56 

Pr 

11 

H 

Q 

5.17 

3.47 

1.70 

10.62 

Me 

H 

Me 

6.65 

3.85 

2.80 

10.16 

Table  5.2.  Thermodynamic  Data  (0°) 

M  B,0)J  „■*  +  n(AA)  aq  ^  [M(AA)„]  an+2  +  *H,0 

(Collected  by  Basolo  and  Murmann61) 


Xickel(II) 

Copper(II) 

n 

3 

AF° 

AH° 

AS0 

n 

AF° 

AH° 

AS0 

Ethylenediamine 

-25.1 

-24.9 

+  1 

2 

-26.6 

-24.6 

+7 

Ethylenediamine 

2 

-18.1 

-16.3 

+7 

X-Methvlcthvlenediamine 

2 

-17.2 

-17.0 

+  1 

2 

-25.3 

-23.0 

+8 

X ,  N  '-Diet  hylethylenedi- 

2 

-15.3 

-7.8 

+27 

2 

-23.3 

-17.5 

+21 

amine 

gest  reduced  stability  when  interference  of  groups  arises.  It  is  reported 
that  a,a-dipyridyl  coordinates  with  iron  whereas  the  6,6-disubstituted 
dipyridyl  does  not.  The  low  coordinating  ability  is  attributed  to  clashing 
of  the  methyl  or  amino  groups  in  the  6 , 6-substituted  complex.  Merritt65 
reports  an  analogous  case  with  8-hydroxyquinoline  and  its  derivatives 


COORDINATES  WITH  Fe^- 


B 

DOES  NOT  COORDINATE  WITH  Fe+* 

R  =  CH3or-NH2 


0<-  0<  -  DIPYRIDYL 


6,  6-  SUBSTITUTED     DIPYRIDYL 


and  has  proposed  the  use  of  selected  steric  factors  to  obtain  selective  or 
specific  analytical  reagents.  His  work  is  described  in  more  del  ail  under 
the  use  of  coordination  compounds  in  analytical  chemistry  (see  p.  t»78). 

85.  Merritt,  "Frontiers  of  Science  Outline,"  Lecture  Wayne  University,  Spring  1949; 
Merritt  and  Walker,  Ind.  I  ■•<.,  Anal.  Ed.,  16,  387  (1944);  Phillips,  El- 

binger,  and  Merritt,./  rn.  Soc,  71,  3986  (1949);  Phillips,  Buber,  Chung, 

and  Merritt,  J.  Am.  Chem.  Soc,  73,  630  (1951). 


238 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Irving,  Cabell,  and  Mellor66  have  used  the  same  type  of  arguments  to 
justify  reduced  stability  of  the  copper(II)  and  iron(II)  complexes  of  2,9- 
dimethyl-1 ,  10-phenanthroline. 


4^ 

3 


n 


I,  10  -  PHENANTHROLINE 

As  noted  in  Chapter  3,  the  coordination  number  is  inadequately  treated 
if  size  alone  is  considered,  but  size  factors  can  be  understood  if  the  interac- 
tion energy  or  bond  energy  at  a  permitted  distance  of  approach  is  taken  into 
account.  It  is  thus  apparent  that  the  interaction  energy  of  metal  and  ligand 
at  the  permitted  distance  is  important  in  determining  compound  stability. 
Recognizing  this  important  restriction,  Irving  and  his  co-workers  justified 
the  fact  that  ions  only  slightly  larger  than  aluminum(III),  such  as  gal- 
lium (III)  and  iron(III),  can  give  precipitates  while  aluminum(III)  cannot. 
In  view  of  such  differences,  Irving  and  his  co-workers66,  as  well  as  Berg67,  have 
also  suggested  the  possibility  of  designing  selective  chelating  agents  based 
on  stereochemical  differences.  Irving,  Butler,  and  Ring66a  have  prepared  a 
number  of  methyl  and  phenyl  substituted  8-hydroxyquinolines.  They  found 
that  substitution  only  in  the  2  position  always  prevented  formation  of  the 
Al+++  complex,  but  permitted  chelation  with  chromium  (III),  iron(III), 
gallium(III),  copper(II),  and  zinc(II)  and  that  the  acridines,  which  in- 
volve ring  formation  on  the  2  position,  also  fail  to  yield  complexes  with 
aluminum  (III),  but  give  precipitates  with  the  other  cations  listed. 


OH 

I  -HYDROXY   ACRIDINE 


OH 

9 -HYDROXY- 1:2:3 :4-TETRAHYDR0    ACRIDINE 


Figure  5.8,  taken  from  Irving,  Butler,  and  Ring,  shows  the  interference  of 
the  2-methyl  groups  with  the  oxygen  and  nitrogen  atoms  in  the  chelate 
rings  of  the  tris-2-methyl-8-hydroxyquinoline  complex  of  aluminum  (III). 
Phillips,  Huber,  Chung  and  Merritt65d  report  that  the  ultraviolet  absorption 
spectrum  of  the  copper  chelate  of  2-methyl-8-hydroxyquinoline  gives  no 
evidence  of  steric  hindrance  and  that  the  unhindered  aluminum  complex 

86    [rving,  Butler,  and  Ring,  ./.  Chem.  Soc,  1949,  1489;  Irving,  Cabell,  and  Mel- 
lor, /.  Chem.  Soc.,  1963, 3417. 

67.  Berg,  Z.  anorg.  Chem.,  204,  208  (1932). 


THEORY  OF  HETEROCYCLIC  IIISC  l<)l!\l AVION  239 


O  OXYGEN  •  N,TROCEN    Q  SMOTHER  TER- 

^-^  VALENT  METAL 

Fig.  5.8.  Steric  hindrance  in  the  tris-2-methyl-8-hydroxyquinoline  chelate  of 
aluminum.  Points  of  interference  are  indicated  by  double  arrows. 

involving  only  one  2-methyl-8-hydroxyquinoline  could  be  identified  in 
solution  by  the  method  of  continuous  variations,  yet  no  hindered  bis-  or 
tris-complexes  of  aluminum  could  be  found.  These  facts  are  consistent  with 
the  proposed  steric  effect. 

Steric  Factors  Determined  by  the  Metal  Ion 

Elementary  theory  indicates  that  the  most  stable  structures  arise  when 
the  bonds  of  the  metal  are  so  directed  in  space  that  they  overlap  the  orbitals 
of  the  ligand  without  serious  distortion  of  either  set  of  orbitals. 

An  interesting  problem  arises  when  the  bonds  of  the  metal  ion  and  the 
bonds  of  the  coordinating  group  do  not  have  the  same  basic  geometry. 
A  case  of  this  type  is  the  divalent  platinum  complex  of  /3,/3',0"-triamino- 
triethylamine  which  was  studied  by  Mann68.  The  base  is  a  quadridentate 
molecule  in  which  the  four  nitrogen  atoms  can  be  expected  to  occupy  the 
corners  of  a  tetrahedron  bul  not  the  corners  of  a  square.  The  bonds  of  the 
platinum(II)  are  normally  directed  to  the  corners  of  a  square,  but 
they    are    apparently    forced    into    the    tetrahedral    configuration     in 

68.  Mann,./.  Chew..  Soc,  1926,  482;  Mann  and  Pope,  •/.  Chem.  8oc.t  1926,  2675. 


240 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


[PtN(CH2CH2NH2)3]++  (p.  363).  The  complex  could  also  be  octahedral 
if  the  two  anion  groups  were  coordinated  to  the  platinum  (see  Figs.  5.9  and 
5.10).  A  crystal  structure  analysis  of  this  complex  is  needed.  There  are  no 


CH2— CH2— N 


NH2^CT___ 


~^>NH, 


Fig.  5.9.  Tetrahedral  coordination  of  jS,^',/?  triaminotriethylamine 


+  + 


NH 


Fig.  5.10.  Octahedral  coordination  of  P,(3',p  triaminotriethylamine  and  two  other 
groups. 


data  to  indicate  that  this  complex  is  any  less  stable  because  of  the  steric 
strain.  Data  are  available,  however,  for  the  copper(II)  complex  which 
should  also  be  planar,  and  it  is  indeed  less  stable  than  one  would  expect 
from  trends  in  the  periodic  table.  In  Fig.  5.11  the  log  of  the  formation  con- 
stants for  a  number  of  metal  amines  are  plotted  for  the  metals  from  manga- 
nese to  zinc. 


THEORY  OF  HETEROCYCLIC  RING  FORMATION 


24] 


II!'! 
O 

20 

//V 

18 

7          \ 

16 

/       9     \\  Mtren 
A         /      \     \  •   6 

14 

- 

/    ?        7         \    ^ 

Locj  K 

/   ■''           \  \  • 

12 

10 

/        ,*       /             /      \          y 

8 

T  °  /     \ 

6 

■V 

t 

s 

4 

s 
s 

2 

- 

1         1          1          1 

Mn 


Fe 


Co 


Ni 


Cu 


Z^n 


Fig.  5.11.  Logarithms  of  the  formation  constants  for  complexes  of  polyamines 
with  transition  metals.  (Data  from  Ref.  5). 

en  =  ethylenediamine  NH2CH2CH2NH2 
dien  =  3,/3'diaminodiethylamine  NH(CH2CH2NH2)2 
trien"=  triethylenetetraamine  NH2CH2CH2NHCH2CH2NHCH2CH2XH2 
tren  =  j3,/3',£"triaminotriethylamine  N(CH2CH2NH2)3  (forced  tetrahedral  config- 
uration) 


For  those  metals  which  have  no  strong  planar  preference,  the  P ,  (3' ,  0" - 
triaminotriethylamine  complex  (M-tren)  is  more  stable  than  the  bis- 
ethylenediamine  complexes  because  of  the  entropy  associated  with  the 
completely  interlocked  ring  system.  On  the  other  hand,  the  copper(II)  com- 
plex, [Cu-tren],  is  less  stable  than  the  bis-ethylenediamine  complex 
[Cu(en)J.  This  phenomenon  has  been  associated  with  the  steric  strain 
arising  from  the  tetrahedral  structure  around  the  normally  planar  cop- 
per(II)  ion5.  It  is  interesting  to  note  that  the  nickel  complex  [Ni-tren]  shows 


242 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


no  reduced  stability  as  a  result  of  the  tetrahedral  configuration,  but  this  is 
not  unexpected  since  even  the  Ni(NH3)4++  ion  is  normally  tetrahedral 
rather  than  planar. 

Another  much  quoted  though  unproved  case  of  steric  hindrance  is  that 
cited  by  Porter69,  who  has  shown  with  molecular  models,  that  when  bis- 
S^^jS^'-tetramethyl^^'-dicarbethoxypyrromethene    (Fig.    5.12)    func- 


F.tOOC- 


EtOOC — C 


C— COOEt 


C— COOEt 


Fig.  5.12.  Overlapping    of    S^^S^'-tetramethyl^^'-dicarbethoxypyromethane 


groups1 


tions  as  a  bidentate  chelate  group,  the  chelate  is  prevented  from  assuming 
a  planar  configuration  by  steric  hindrance.  The  a  methyl  groups  (marked 
by  asterisks)  overlap  seriously  as  is  seen  in  Fig  5.12.  Complexes  with 
Fe+2,  Ni+2,  Co+2,  Cu+2,  Zn+2,  Cd+2,  Pd+2  69  and  Pt+2  70  have  been  prepared. 
Both  the  palladium71  and  platinum70  complexes  are  diamagnetic,  indicating 
"covalent"  bonding;  the  nickel  complex  is  paramagnetic  indicating  an 
"ionic"  bond71.  Since  the  normal  covalent  bonds  of  palladium (II)  and  plati- 
num (II)  are  planar,  one  would  expect  that  steric  inhibition  to  the  planar 
arrangement  would  lower  the  complex  stability.  Actually,  little  evidence 
is  available  to  indicate  that  such  is  the  case.  In  fact,  limited  data  on  com- 
plexes of  3,3'-dimethyl-4,4'-dicarbethoxydipyrromethene,  in  which  there 
are  no  a  methyl  groups  to  overlap,  indicate  that  the  metal  complexes  are 

69.  Porter,  J.  Chem.  Soc,  1938,  368;  Mellor,  Chem.  Revs.,  33,  171,  175  (1943). 

70.  Mellor  and  Willis,  /.  Proc.  Roy.  Soc.  N.  S.  Wales,  79,  141  (1945). 

71.  Mellor  and  Lockwood,  J.  Proc.  Roy.  Soc.  N.  S.  Wales,  74,  141  (1940). 


THEORY  OF  HETEROCYCLIC  HINQ  FORMATIOh 


243 


Table  5.3,  Magnetk    Moments  of  Phthaloctaninb  Complexes 

Metal  in  Complex 

\Mg    Moment 
in  Bohr  Magnetons 

Theoretic  .il  Moment 

.   />-  Bonds 

Theoretii  al  Moment 
for  Planar  />-; 

Theoretical 
Moment  for 
tonic  Binding 

Cu+J 

Fe+I 
Md 

1.73 

i) 
2.16 

3.96 
4.55 

1.73 
0 

1.7:5 
2.83 
3.87 

3.87 
2.83 
1.73 

0 
1.73 

1.73 
2.83 
3.87 
4.90 
5.92 

actually  less  stable  than  the  fully  methylated  compound69  in  which  steric 
hindrance  supposedly  occurs.* 

The  converse  problem  of  fitting  a  normally  tetrahedral  ion  to  a  planar 
quadridentate  molecule  has  also  received  attention.  The  phthalocyanine 
molecule  (p.  73)  is  rigidly  coplanar,  and  its  complexes  with  the  divalent 
ions  of  copper,  nickel,  platinum,  cobalt,  iron,  manganese,  magnesium  and 
beryllium  have  been  shown  by  x-ray  studies  to  be  planar72.  The  appearance 
of  magnesium  and  beryllium  with  planar  coordination  is  indeed  surprising, 
since  these  metals  normally  assume  a  tetrahedral  structure.  It  is  noteworthy 
that  both  beryllium  and  magnesium  phthalocyanins  readily  form  hydrates; 
Buch  behavior  may  be  indicative  of  lower  stability  in  the  forced  configura- 
tion. Two  molecules  of  water  would  allow  octahedral  coordination. 

The  magnetic  properties  of  the  remaining  phthalocyanines  have  been 
studied  by  Klemm  and  his  students73,74.  Their  data  permit  an  answer  to 
the  problem:  "Does  assumption  of  a  forced  planar  configuration  by  the 
metal  ion  require  the  use  of  planar  dsp2  or  d2p2  bonds?"  Data  in  Table  5.3 
indicate  that  it  does  not,  since  the  observed  moments  do  not  correspond  to 
those  expected  for  dsp2  bonds.  Selwood75  suggested  that  the  magnetic  data 
actually  indicate  a  transition  from  covalent  to  ionic  bonds  in  the  iron  and 
manganese  complexes  with  forced  configurations. 

Schwarzenbach  and  Ackerman8  have  invoked  favorable  steric  and  en- 
tropy factor-  as  an  argument  to  justify  their  observation  that  1,2-cyclo- 
hexanediamine-N.N'-tetraacetate    forms    a    more    stable    chelate    with 


*  It  is  interesting  that  none  of  the  pyrromethene  complexes  even  approach  I  In- 
analogous  porphyrins  or  phthaloyamins  in  stability,  because  of  multiple  ring  effects 
in  the  latter69. 

72    Robertson,/.  I  Stoc.,  1935,  615;  1936,  1195;  [instead  and  Robertson,  ./. 

S     ..  1936, 
Oemm  and  Klemm,  ./.  prakt.  Chem.,  143,  82  (1935). 
74.  Senff  and  Klemm,  J.  prakt.  ('hem.,  154,  73  (!'• 

ood,  "Magnet  ochfini-t  r\."  p.  163,  Ne*  York,  Interscience  Publishers,  Inc., 
1943. 


244  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Ca++  (K  =  1012-5)  than  does  the  related  ethylenediaminetetraacetate 
(K  =  1010-5).  It  is  assumed  that  this  difference  exists  because  the  coordinat- 
ing  groups  in  the  cyclohexanediamine  derivative  are  fixed  in  position  while 
those  in  the  ethylenediamine  derivative  are  free  to  rotate  about  the  ethy- 
lene group.  The  smaller  magnesium  ion  and  the  larger  barium  ion  are  less 
able  to  utilize  this  stereochemical  advantage,  so  there  are  smaller  differences 
for  these  ions  between  the  complexes  of  the  cyclohexanediamine  and  ethy- 
lenediamine derivatives. 

Irving,  Cabell,  and  Mellor66  also  suggest  that  the  apparent  relative  stability 
of  the  ferrous  tris-orthophenanthroline  complex  may  be  due  in  part  to  the 
fact  that  the  ferric  tris-orthophenanthroline  structure  is  destabilized  by 
steric  hindrance.  Evidence  for  this  is  obtained  from  the  observation  that 
when  iron(III)  ions  react  with  orthophenanthroline  directly,  the  binuclear 
complex 

H 
O 

/   \ 

(ophen)2Fe  Fe(ophen). 

\    / 

O 

H 

is  formed,  rather  than  the  tris-complex. 

In  summary,  there  is  some  evidence  to  indicate  that  the  stereochemistry 
of  metal  cations  is  important  in  determining  the  stability  and  type  of 
complex  formed.  However,  exceptions  are  known.  Present  data  indicate 
that  the  stereochemical  properties  of  the  metal  ion  are  much  more  flexible 
in  chelate  ring  formation  than  the  stereochemical  properties  of  the  ligand. 

Electronic  Effects  Peculiar  to  Chelate  Rings 

Effects  Due  to  Ring  Closure 

A  few  unusual  electronic  effects  seem  to  arise  in  chelate  systems  as  a  re- 
sult of  ring  formation.  Such  effects  are  as  yet  incompletely  understood. 
Spike  and  Parry77  measured  indirectly  the  enthalpy  and  entropy  associ- 
ated with  reactions  of  the  type 

M(NH3)2+  en  ->  Men  +  2NH3 

In  some  cases  similar  studies  were  made  using  methylamine  in  place  of 
ammonia.  If  the  formation  of  chelate  rings  produced  no  increase  in  the 

76.  Sidgwick, ./.  Chem.  Soc,  433  (1941);  "The  Electronic  Theory  of  Valency,"  Oxford 

Univ.  Press,  1927. 

77.  Spike  and  Parry,  ./.  Am.  Chem.  Soc,  75,  2726,  3770  (1953);  Spike,  PhD  Disserta- 

tion, University  of  Michigan,  1952. 


THEORY  OF  HETEROCYCLIC  RING  FORMATION  245 

stability  of  the  metal-ligand  bond.  All  for  the  above  process  Bhould  be 
utially  zero  ami  the  Increased  stability  of  the  chelated  system  should 

arise  as  a  result  of  entropy  factors.  If,  however,  ring  formation  results  in  a 
stronger  metal  ligand  bond.  All  for  the  above  process  should  be  negative. 
When  zinc  and  cadmium  were  used,  AH  for  the  process  was  found  to  be  es- 
sentially zero,  hut  when  copper(II)  was  the  metal  ion  the  AH  term  was  afi 
large  as  the  entropy  term,  indicating  a  much  si  ronger  metal-ligand  bond  as 
a  result  of  ring  formation.  The  absence  of  double  bonds  in  the  ethylene- 
diamine  makes  the  usual  resonance  interpretations  (see  below)  difficult. 

Resonance  Effects 

In  194-3  Cabin  and  Wilson11,  using  the  method  of  Bjerrum82,  found  a 
straight  line  relationship  between  the  basic  strength  of  enolate  0-diketones 
and  the  stability  of  copper(II)  complexes  (see  also  p.  178).  Their  work  also 
indicated  the  necessity  for  subdividing  the  ligands  into  similar  groups  in 
order  to  establish  a  correlation.  The  data  shown  in  Fig.  5.13  were  classified 
into  four  groups  (A),  (B),  (C),  and  (D),  A  and  C  giving  linear  plots  with 
considerable  scatter,  and  B  and  D  giving  one  point  lines.  The  structural 
types  associated  with  the  four  lines  are: 
CH3 

H  c>*-°-  rO°~ 


*=/       c=0 
H 


A  B. 

ENOLATE  TYPE  OF  NAPTHOLATE  ION  OF 

ACETYLACETONE        2-HYDROXYNAPTHALDEHYDE - I 


\    hr 


o- 

/ 
H 


o— 


C  D. 

PHENOLATE     ION  NAPTHOLATE     ION     OF 

OF    SALICYLALDEHYDE  2- HYDROXYNAPTHALDE HYDE  -  3 

According  to  Calvin  and  Wilson,  the  most  important  difference  in  these 
structures  is  the  nature  of  the  double  bond  between  the  two  carbon  atoms 
of  the  three  carbon  Bystem  which  forms  the  conjugated  chain  between  the 
two  oxygen  atom-.  These  bonds  are  marked  with  asterisks  in  the  above 
formulas.  In  structure  (A)  only  a  methyl  group  and  a  hydrogen  are  at- 


246 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

T 


4.0  5.0  60  7.0 

INCREASING  COMPLEX  STABILITY-* 
LOG  Kav. 
Fig.  5.13.  Relationship  between  the  basic  strength  of  enolate  /3-diketones  and  the 
stability  of  their  copper(II)  complexes.  (From  Ref.  11). 

Line  A: /3-Diketones  and  0-keto  ester:  (16)  trifluoro  acetylacetone  (17)  furoylace- 
tone;    (18)   acetylacetone;    (19)   benzoylacetone;    (12)   acetoacetic  ester; 
(14)  C-Methyl  acetylacetone 
Line  B:  2-hydroxynaphthaldehyde-l 

Line  C:  Substituted  salicylaldehydes :  (2)  4-Nitro;  (3)  3-Chloro;  (4)  5-Chloro;  (5) 
3-Fluoro;  (6)  Salicylaldehyde;  (7)  5-Methyl;  (8)  3-Methoxy;  (9)  4-Meth- 
oxy;   (10)  3-n-Propyl;   (11)  3-Ethoxy;   (13)  4,6  Dimethyl;   (20)  3-Nitro; 
(21)  5-Nitro 
LineD:  2-hydroxynaphthaldehyde-3 

Kav  =  equilibrium  constant  for: 


c— o-  C— O 

/  /  \ 

— C  +  |  Cu++  ±=;  — C  Cu 


c=o 

/ 

Kd  =  equilibrium  constant  for: 

\ 
0—0 


c=o 


c— o- 

— C  II  ^±  — C  +  H+ 


C=0 


C=0 


THEORY  OF  HETEROCYCLIC  RING  FORMATION  247 

bached  to  this  bond.  In  structures  (H),  (C),  and  (D)  the  double  bond  is 

also  part  of  a  resonating  aromatic  ring.  According  to  the  met  hod  used  by 
Pauling78  and  by  Branch  and  Calvin79,  the  double  bond  A  which  docs  not 
resonate  with  any  single  bonds  in  attached  rings  is  given  an  arbitrary  bond 
order  of  2.  In  the  case  of  structure  (C)  the  double  bond  must  resonate  in 
the  benzene  ring,  hence  it  may  be  regarded  as  only  half  of  a  double  bond 
for  the  enolate  system.  It  is  assigned  the  value  1.5.  Similarly,  structure  (B) 
is  assigned  the  double  bond  order  1.1)7  and  (D),  1.33.  It  can  be  seen  from 
Fig.  5.12  that  the  stability  of  the  copper  complex  at  constant  acidity  of  the 
chelating  agent  decreases  in  the  same  order  as  the  decrease  in  this  double 
bond  character.  In  short,  the  greater  the  double  bond  character  of  the 
bond  in  the  enolate  system,  the  more  stable  is  the  complex.  It  is  reported 
that  these  observations  on  stability  of  complexes  of  different  types  have 
also  been  supported  by  polarographic  studies7  and  by  exchange  studies 
involving  radioactive  copper(II)  ions80. 

The  observations  led  to  the  following  conclusion,  "Resonance  in  the 
enolate  (or  chelate)  ring  plays  a  far  greater  part  in  the  bonding  of  copper 
than  it  does  in  the  bonding  of  hydrogen."  Calvin  suggested  two  possible 
explanations  for  this.  The  first  is  represented  electronically  as  follows: 


According  to  the  second  suggestion,  a  completely  conjugated  six-membered 
chelate  ring  analogous  to  pyridine  is  formed: 

\    /2  v  J- 

,C=Q  /C-Q 

The  second  hypothesis  assumes  considerable  double  bond  character  for  the 
metal-oxygen  bond.  Although  double  bonds  between  metal  and  ligand  have 
been  extensively  postulated  (see  p.  191,  Chapter  X)  the  suggestion  in  this 
runs  into  rather  serious  difficulty.  An  electron  balance  shows  that  the 
electron  pair  used  to  form  the  metal-oxygen  double  bond  came  from  the 
oxygen  rather  than  from  the  metal  ion  as  is  normally  postulated.  A  double 

>    Pauling,  "Nature  of  the  Chemical  lion.]/'  pp.  L79,  182,  L87,  L39,  Cornell  Uni- 
versity Press,  1942. 
Branch  and  Calvin,  "The  Theory  of  Organic  Chemistry,"  p.  113,  New  York, 
Prentice-Hall,  Inc.  1941. 

80.  Duffield  and  Calvin,  ./.  ,1///.  Chem.  Soc,  68,  557  (1946). 


248  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

l)oii(l  of  this  type  is  diametrically  opposed  to  the  usual  assumption  that 
the  metal  ion  donates  the  electrons  and  the  ligand  accepts  them  (see  p. 
191).  Such  a  double  bond  would  increase  the  residual  negative  charge  on 
the  copper  rather  than  decrease  it  as  is  normally  postulated.  To  assume 
that  the  copper(II)  ion  behaves  in  a  normal  fashion  and  uses  d  electrons  to 
form  a  double  bond  with  the  oxygen  is  equally  distasteful  since  oxygen  has 
no  low  level  orbit als  which  permit  it  to  serve  as  an  acceptor  without  destroy- 
ing the  conjugated  double  bond  system  in  other  parts  of  the  ring. 

Marked  deviations  between  fact  and  prediction  have  been  attributed  to 
steric  inhibition  of  resonance11  although  the  supporting  evidence  for  this 
postulate  is  still  extremely  sketchy  in  many  cases.  One  of  the  more  convinc- 
ing illustrations  is  the  copper  complex  derived  from  salicyl aldehyde  and 
1 , 8-diaminonaphthalene 


Q_0-Cu-0-/--) 

~Vh=n    n=hc' 


Since  this  complex  is  a  multiple  ring  type  involving  a  highly  conjugated 
system,  we  would  expect  it  to  be  more  stable  than  comparable  complexes  in 
which  the  entire  chelating  system  is  not  fused  together.  Actually,  the  com- 
plex is  only  slightly  more  stable  than  the  open  ring  structures.  Duffield 
and  Calvin80  attributed  this  unexpected  behavior  to  the  fact  that  steric 
factors  prevent  the  complex  from  assuming  a  coplanar  structure  about  the 
copper  atom.  It  is  suggested  that  such  nonplanarity  prevents  or  reduces  the 
benzenoid  chelate  resonance  and  thus,  the  stability  of  the  complex.  It  is 
possible  that  steric  factors,  independent  of  resonance  effects,  could  also 
account  for  the  reduced  stability  since  Cu-1"4"  is  normally  a  planar  ion. 

The  opposite  situation,  in  which  stability  of  a  strained  structure  is  at- 
tributed to  resonance  has  been  described  bj^  Dwyer  and  Mellor22.  A  metallic 
triazine  complex  such  as 

N 

/   \ 
R'— N  N— R 

\    / 
M 

forms  a  four-membered  ring  which  is  unexpectedly  stable.  This  stability 
has  been  attributed  to  resonance  of  the  following  type. 

N  N 

•    \  /    % 

R— N  X— II'  <->  R— N  X— R' 

\    /  \    / 

M  M 


THEORY  OF  HETEROCYCLIC  RING  FORMATION  249 

Chelates  Involving  Conjugated  Double  Bonds 

Finally,  an  interesting  compound  described  by  Chatl  and  Wilkins81  should 
be  mentioned.  This  stable  complex  appears  to  be  a  chelate  si  ructure  involv- 
ing coordination  to  two  double  bonds  of  pentadiene.  The  molecular  formula 

of  the  complex  is  PtCb(C5H8)2,  the  monomeric  nature  of  the  compound 
having  been  established  by  molecular  weight  measurements.  Butadiene, 
which  would  make  a  small  and  highly  strained  ring,  does  not  chelate  under 
the  conditions  used  by  Chatl  and  Wilkins  but  reacts  independently  with 
different  platinum  atoms. 

Entropy  Effects  in  Chelation 

Sidgwick76  suggested  in  1941  that  the  stability  of  chelate  systems  as 
compared  to  similar  nonchelate  structures  may  be  due  to  a  statistical  factor 
which  he  pictured  as  follows.  If  one  of  the  two  metal-ligand  bonds  of  a 
chelate  system  is  broken,  the  remaining  bond  will  hold  the  molecule  in 
place  so  that  the  broken  link  can  be  reformed,  whereas  an  atom  or  group 
that  is  bound  by  a  single  link  will  drift  away  if  the  bond  is  broken.  Since 
this  is  a  question  of  probability,  it  should  appear  in  the  entropy  term.  The 
relationship  is  somewhat  more  apparent  if  one  writes  a  typical  equation 
denning  the  chelate  effect: 

M(NH2Me)2++  +  en  ->  Men++  +  2NH2Me 

The  equation  suggests  an  increase  in  the  disorder  of  the  system  on  chela- 
tion or  an  increase  in  the  translational  entropy  of  the  system. 

Concurrent  with  Sidgwick's  1941  paper,  J.  Bjerrum82  published  one  of 
the  most  important  experimental  papers  to  appear  in  the  field  of  coordina- 
tion chemistry  since  the  early  work  of  Werner.  In  a  classical  theoretical 
and  experimental  analysis  of  metal  ammine  formation,  he  considered  two 
factors  which  are  important  in  determining  the  ratio  between  successive 
dissociation  constants  for  a  metal  ammine  such  as  the  ethylenediamine 
complex  of  a  metal.  These  are:  (1)  a  statistical  effect,  and  (2)  a  ligand  effect. 
The  statistical  effect  is  defined  as  the  joint  contribution  to  the  ratio  of  the 
dissociation  constants  which  is  attributable  to  purely  statistical  causes  plus 
the  stereochemical  effects  of  dissimilar  coordination  positions.  For  example, 
if  a  given  metal  can  coordinate  a  maximum  of  N  Uganda  and  at  a  particu- 
lar time  has  bound  only  n  ligands,  then  the  statistical  probability  thai  the 
complex  will  lose  a  ligand  should  be  proportional  to  n  whereas  the  proba- 
bility that  it  can  pick  up  another  ligand  should  be  proportional  to  the  num- 
ber of  stereoehemically  satisfactory  sites  remaining  in  the  coordination 
sphere,  or  for  a  nonchelate  ligand.  (JV-n).  For  a  chelate  ligand  the  two  sites 

81.  Chatt  and  Wilkins,  ./.  Chi  •■     fl         1952,  2622. 

82.  Bjerrum,  J.,   "Metal   Ammine   Format  ion   in   Aqueous  Solution/'   Copenhagen, 

P.  Haase  and  Son,  1**41 . 


250  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

must  be  adjacent  in  order  to  meet  the  sterochemical  requirements  of  the 
donor  molecule.  It  is  apparent  that  this  factor  should  appear  in  entropy 
terms.  The  ligand  effect  includes  the  joint  contribution  to  the  ratio  of  the 
dissociation  constants  which  is  attributable  directly  or  indirectly  to  the 
ligands  taken  up.  This  would  be  an  enthalpy  term.  The  work  of  Bjerrum 
and  others  was  admirably  summarized  by  Burkin83. 

In  1952,  Schwarzenbach5  and  Spike77  utilized  the  model  suggested  by 
Sidgwick  as  the  basis  for  independent  kinetic  treatments  of  the  chelate 
effect.  Following  the  suggestion  of  Bjerrum82,  the  formation  and  dissociation 
of  the  nonchelated  complex  MA2  and  the  chelated  complex  M(AA)  are 
considered  to  be  step  processes.  It  is  then  logical  to  assume  that  the  chelate 
molecule  (AA)  reacts  with  or  dissociates  from  the  metal  ion  in  two  steps. 
The  intermediate  form  is  a  complex  in  which  the  chelating  ligand  is  bound 
by  only  one  donor  atom.  By  application  of  simple  collision  theory  of  reac- 
tion rates,  by  assuming  a  comparable  energy  of  activation  for  the  reaction 
of  chelate  and  nonchelate  structures,  and  by  using  the  best  available  data 
on  sizes  of  molecules,  one  can  estimate  the  order  of  magnitude  of  the  en- 
tropy term  in  the  chelate  effect776.  It  appears  from  the  above  models  that 
the  rate  of  the  reaction 

MA++  +  A  -+  MA2++ 

can  be  related  to  the  size  of  the  volume  element  containing  one  free  amine 
molecule  and  the  rate  of  the  comparable  reaction 


[M— AA— ]++ 


M 


can  be  related  to  that  volume  inside  the  sphere  of  radius  r'  which  is  available 
to  the  second  end  of  the  chelating  ligand. 

The  above  model  suggests  that  the  stabilization  due  to  chelation  should 
decrease  rapidly  as  the  chain  of  the  ligand  is  lengthened.  Schwarzenbach 
has  shown  that  the  difference  in  free  energy  of  formation  between  chelate 
and  nonchelate  structures  decreases  rapidly  and  even  reverses  in  sign  as 
the  chain  is  lengthened.  One  also  arrives  at  a  justification  for  the  stability 
of  five-membered  rings.  As  a  result  of  steric  strain  the  energy  of  bond  forma- 
tion is  low  for  small  rings  but  increases  as  increasing  size  of  the  ring  re- 
lieves strain.  On  the  other  hand,  the  stabilizing  influence  of  chelation,  which 
appears  in  the  entropy  term,  is  greatest  for  small  rings.  These  two  terms, 
working  in  opposite  directions,  produce  a  stability  maximum  in  a  five- 

83.  Burkin,  Quarterly  Revs.,  5,  1  (1951). 


THEORY  OF  HETEROCYCLIC  RISC  FORMATION 


251 


Table  5.4.  Thebmodyn  lmic  Constan  re  in  -M  Qnii  mini  s  m.i  Solution  \  t  25°C 
for  Reaction  MAt4"*  +  en  — ♦  Men++  +  2  \ 


IF 

MI 

CdiXII.UI    r+— Cd(en)++ 
Cd  XH3)2++— Cd(en)++ 

Zn(XH3)2++— Zn(en)++ 
Cu(XH3)2++— Cu(en)++ 

-1.  Hi 
-1.20 
-1.55 
-4.30 

0.0 

+  .1 
+.1 

-2.6 

1.7 
t.:i 
5.3 
5.7 

membcred  saturated  ring  and  in  a  six-membered  unsaturated  ring,  the 
stereochemistry  of  which  is  further  restricted  by  double  bond  formation. 

The  model  also  indicates  that  further  restriction  on  the  mobility  of  the 
second  ligand  should  enhance  the  stability  of  the  complex  if  the  size  of  the 
metal  ion  is  such  as  to  fit  into  the  space  between  the  binding  atoms.  Schwar- 
zenbach  and  Aekerman37  found  that  1 ,2-cyclohexanediamine  tetraacetate 
forms  a  more  stable  chelate  with  calcium(II)  than  does  ethylenediamine 
tetraacetate.  They  attribute  this  to  such  steric  stabilization. 

The  model  also  suggests  that  multiple  ring  formation  should  result  in 
enhanced  chelate  stability,  a  fact  which  has  already  been  well  established. 
Schwarzenbach5  reports  that  the  chelate  effect  in  a  bidentate  ligand  is 
about  half  of  that  in  a  tridentate  ligand  which  can  form  two  interlocking 
rings  and  is  about  one  third  of  that  in  a  tetrafunctional  ligand  which  can 
form  three  rings. 

The  preceding  model  would  indicate  that  the  chelate  effect  should  be 
quite  independent  of  the  metal  except  insofar  as  special  steric  requirements 
of  the  metal  are  concerned  (e.g.,  a  linear  structure  of  silver).  Schwarzen- 
bach5 noted  the  low  chelate  effect  for  the  [Zn(en)]++  complex  and  suggested 
that  this  may  indicate  a  tendency  of  the  zinc(II)  ion  toward  linearity.  He 
interpreted  the  data  on  copper(II)  complexes  as  being  more  representative 
of  the  chelate  effect. 

Spike  and  Parry77  measured  the  entropy  and  enthalpy  changes  for  reac- 
tions of  the  type  M(XH3)2  +  en  — »  Men  +  2XH3 .  Their  data  for  the 
changes  at  25°  in  a  solution  of  2  molar  univalent  nitrate  salt  (i.e.,  KNO|  or 
NH4NO1)  are  summarized  in  Table  5.4. 

All  entropy  differences  are  roughly  of  the  same  size,  as  might  be  expected, 
and  the  chelate  effect  for  zincfll)  and  cadmium(II)  is  definitely  an  en- 
tropy effect.  On  the  other  hand,  it  is  significant  that  in  the  case  of  copper 
there  is  a  marked  enthalpy  contribution  to  the  chelate  effect  (i.e.,  bond- 
are  stronger  in  the  chelate  structure.)  The  basis  for  this  effect  is  still  ob- 
scure. Irving54  has  confirmed  the  enthalpy  contribution  for  the  copper  Bys- 
tem  by  calorimetric  measurement  a. 

84.  Irving,  private  communication. 


252  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

The  entropy  term  in  chelate  formation  can  also  be  considered  qualita- 
tively  in  terms  of  the  number  of  particles  on  each  side  of  the  equation.  For 
the  reaction  Ni(NH3)6++  +  3en(aq)  ->  [Ni(en)3]++  +  6NH3j  Calvin  and 
Bailes7  reported  the  thermodynamic  values:  AF  =  —13.2;  AH  =  —6; 
AS  =  24.  Another  factor  of  importance  is  the  relative  orientation  of  water 
molecules  around  the  simple  and  chelated  ions.  Such  a  factor  is  of  major 
importance  when  large  organic  ligands  serve  as  the  chelating  ligands.  The 
importance  of  such  hydration  effects  has  been  considered  by  Cobble85 
in  a  series  of  useful  empirical  relationships. 

Adamson86  has  recently  suggested  a  new  approach  to  the  chelate  effect 
in  which  the  standard  state  of  the  ligand  is  changed  to  give  a  condition  of 
minimum  translation  entropy.  He  proposes  to  use  mole  fraction  unity  as 
the  standard  for  the  ligands  rather  than  the  conventional  molarity  unity. 
Using  this  approach,  the  data  are  comparable  to  those  using  the  conven- 
tional standard  states  if  a  comparable  series  of  reactions  is  considered; 
however,  comparisons  between  reactions  involving  different  numbers  of 
ligands  will  be  altered. 

85.  Cobble,  /.  Chem.  Phys.,  21, 1443  (1953). 

86.  Adamson,  J.  Am.  Chem.  Soc,  76,  1578  (1954). 


Large  Rings 


Thomas  D.  O'Brien* 

University  of  Minnesota,  Minneapolis,  Minnesota 

The  more  stable  ring  sizes  among  coordination  compounds  are  analogous 
to  those  occurring  among  organic  compounds.  The  coplanar  five-  and  six- 
membered  carbon  rings  are  the  most  stable,  according  to  the  Baeyer  strain 
theory,  because  of  the  smaller  requisite  deviation  from  the  natural  tetra- 
hedral  bond  angle  of  1Q9°  28'.  However,  organic  ring  compounds  which  are 
thought  to  be  strainless  and  which  contain  more  than  thirty  members  have 
been  prepared.  These  compounds  are  quite  possible  if  the  atoms  are  not 
forced  to  be  coplanar. 

Stable  chelate  rings  of  five  and  six  members  containing  metallic  atoms 
are  numerous  and  well  known,  but  rings  of  seven  or  more  members  are 
comparatively  uncommon.  This  is  illustrated  by  early  failures  to  prepare 
chelate  rings  with  polymethylenediamines1, 2-  3> 4- 5.  Only  recently  has 
Pfeiffer6  reported  the  preparation  of  seven-membered  chelates  of  tetra- 
methylenediamine  and  nine-membered  chelates  of  hexamethylenediamine. 
These  were  prepared  in  alcohol  or  ether  solution,  and  are  immediately 
hydrolyzed  by  water.  The  studies  of  Schwarzenbach  (p.  229)  on  tetra- 
acetic  acid  derivatives  of  such  amines  indicate  that  polymetallic  com- 
plexes are  to  be  expected  as  chain  length  increases.  Duff7  prepared  com- 
plexes such  as  [(NH3)5CoOOCRCOOCo(NH3)5]4+  and  Macarovici8  reported 

*  Now  at  Kansas  State  College,  Manhattan,  Kansas. 

1.  Werner,  Ber.,  40,  61   (1907). 

2.  Tschugaeff,  Ber.,  39,  3190  (1906);  J.  prakt.  Chem.  [2],  75,  159  (1907). 

3.  Drew  and  Tress,  ./.  Chem.  Soc.,  1933,  1335. 

4.  Pfeiffer  and  Baimann,  Ber.,  36,  1064  (1903). 

5.  Pfeiffer  and  Lubbe,  ./.  prakt.  Chem.  [S],  136,  321  (1933). 

6.  Pfeiffer,  Bohm  and  Schmita,  Naturwissenschaften,  35,  190  (1948). 

7.  Duff,  ./.  Chem.  Soc.,  1923,  560. 

8.  Macarovici,  Bull  sect.  set.  acad.  roumaine,  23,  61   (1940);  Chem.  Abs.,  37,  6642 

(1943). 

253 


254 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


,N  i  -  NH2-<^> <^>- 


NH: 


NH2-Ni 


/ 
\ 


CI 


CI 


NH2-<^> <^]>-NH2 

(Water  may  complete  the  coordination  sphere.) 

This  structure,  however,  is  based  only  on  analysis.  Dwyer  and  Mellor' 
formulated  the  nickel  triazine  complex  as  a  dimer 

R  R 

I  I 

N— N=N 

\     /  \     / 

Ni  Ni 

/     \  /     \ 

N=N— N 


R 


R 


This  raises  an  interesting  question  about  the  benzidine  complexes  of  di- 
valent metals,  the  formulas  of  which  are  frequently  written 


NH. 


t  + 


Such  complexes  are  possibly  polymeric,  since  benzidine  does  not  chelate 
with  cobalt  in  [Co  en2  benzidine  Br]Br2  but  is  monodentate10. 

Seven-Membered  Rings 
Duff11  found  that  the  dibasic  acids  meso-tartaric,  maleic,  dibromsuccinic, 
itaconic  and  citraconic,  when  added  to  carbonatobis(ethylenediamine)co- 
balt(III)  bromide  yielded  crystalline  compounds,  which  he  supposed  con- 
tained the  ion 


R— CH— C 


• 


O 


0 


\ 
/ 


Co  en2 


<> 


R— CH— C 


\ 


LARGE  RING8 


255 


It  i>  possible  that  the  o-hydroxy  acids  form  five-membered  rings  involving 
the  metal,  the  carboxy  and  the  hydroxy  groups.  A  series  of  related  dibasic 
acids  in  which  the  carbonyl  groups  arc  in  the  trans  positions  give  only  vis- 
cous syrupe  which  have  not  been  identified,  [t  is  possible  that  the  acid  mole- 
cules Berve  as  bridge  groups  in  building  polymers.  Several  analogous  com- 
pounds between  cobalt  and  phthalic  acid  and  some  sulfur  derivatives  of 
phthalic  acid  were  also  reported  by  Duff11.  Be  assigned  the  following  struc- 
tures on  the  basis  of  analytical  data  alone: 


+  r- 


S-0N, 


(X;>-s  («>-.   £ 


S-O' 

o/xo 


o  o 

\/ 

s 


-1  + 


Co  en- 

C~0'  2 

M 

o 


Tetrachlorodimethylphthalatotitanium(IV)  has  been  imported  by  Scagli- 
arini  and  Tartarini13  who  proposed  the  following  structure,  again  on  the 
of  analytical  data  alone: 

0-CH3 


CI4Ti 


,o=c 


0-CH3 

Shuttleworth"  states  that  for  the  chromium  chelate  derivatives  of  di- 
basic  acid-.  se\  en-membered  ring  structures  are  intermediate  in  stability 
between  four-  and  six-membered  rings.  He  reports  complexes  of  the  type 
AA  with  maleic,  malonic,  glutarie,  adipic,  suberic,  phthalic  and  azelaic 
acids,  remarking  that  the  acids  which  do  not  form  five-,  six-  or  seven-mem- 
bered  rings  tend  to  form  polymers. 

Brady  and  Hughes15  investigated  the  reaction  of  2,2'-biphenol  with  a 
number  of  metallic  ions  and  complexes,  and  proposed  seven-membered 
ring  structures  for  two  of  the  compounds  prepared.  When  thallium(I)  ace- 
tate, in  ammoniacal  methanol  solution,  was  treated  with  2,2'-biphenol,  a 
precipitate  was  formed,  which,  from  analysis,  was  assigned  the  structure 


:  >-<  : 

1            1 

\    / 

Tl 

> 


87  2291*    1943). 


►wyerand  Mellor,  ./.    1  63.  Bl     1941   . 

I.  s-ri.  acad.  ri,          ne,  23,  181    1940  ;  Cfu  n 

11.  Duff,  J.  CI                  119,  "    1921   . 

12.  Duff,  ./.  (  ■         j        119,  21  . 
Scagliarini  and  Tartarini.  AUi  aeead.  Lincei,  4,  318 

14.  Shuttleworth,  J.  A       I.  45,  ISO    I 

15.  Brady  and  Hugfa  1988,1227. 


256 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


When  this  substance  was  treated  with  aqueous  alkali,  a  less  soluble  com- 
pound was  formed  along  with  the  liberation  of  an  equivalent  amount  of 
biphenol. 


o  o 

i  i 

CD 


/ 


NH2CH3 


Cu 


(M) 


NH2CH3 


NH 


2  H2N 


Structure  (I)  was  assigned  on  the  basis  of  these  observations.  Another 
srven-membered  ring  structure  proposed  by  the  same  authors  was  that 
of  the  copper  complex  showrn  in  (II). 

Para-aminophenol  yields  blue-violet  insoluble  compounds  with  copper(II) 
and  iron(II).  From  the  composition  of  the  compounds  and  their  insolubility 
in  water,  Augusti16  proposed  the  unlikely  structure  (III)  for  the  copper 
complex. 

Seven-membered  rings  have  been  reported  in  which  the  central  atom  co- 
ordinates to  two  nitrogen  atoms17  of  a  diamine.  Middleton  reports  cobalt 
complexes  with  the  structures 


I  —  NH. 


Co  en2 


NH, 


CU    AND 


CI 


The  correctness  of  these  formulas  is  indicated  by  analysis  and  by  the  colors 
of  the  compounds,  the  first  having  the  orange  color  of  luteo  salts,  and  the 
second,  the  green  color  of  the  praseo  salts. 

Rings  Containing  Eight  or  More  Members 

The  first  ci<;ht-membered  ring  was  reported  by  Price  and  Brazier18  who 
treated  carbonatobis(ethylenediamine)cobalt(III)  bromide  with  sulfonyl- 
diacetic  acid.  They  assigned  the  structure 

O 
OH2—  c— o 


o 


1 1 


cir— c— o 

\ 

o 


/ 


Co    en2 


X 


LARGE  RINGS 

Under  different  conditions  the  two  carboxy]  groups  are  attached  to  two 
different  cobalt  atoms,  giving  rise  to  polynuclear  structures.  Moreover,  if 
the  sulfone  group  Is  replaced  by  sulfide,  do  compounds  are  obtained  analo- 
gous to  those  for  which  the  eight-membered  ring  structure  was  proposed. 

This  suggests  that  the  chelation  may  involve  the  oxygen  atoms  of  the 
sulfone  group  rather  than  the  carboxy]  groups. 

Schmitz-Dumont  and  Motzkus18  obtained  an  insoluble  compound  when 
copper(I)  ion  was  treated  with  bis-a-methyl-0-indyl  methene,  to  which  they 
s&  _  ted  the  structure 


Triethanolamine  has  been  used  as  a  coordinating  agent  with  a  number  of 
metallic  ions.  Tettamanzi  and  Carli20  found  that  coordination  compounds 
rather  than  basic  salts  are  formed  with  nickel,  cadmium,  calcium  and 
magnesium.  They  proposed  the  alternative  structures  (IV)  and  (V). 


ch2-ch2-oh 

N-CH2-CH2-OH 
CH2-CH2-0H 


HO-CHo-CH 


HO-CH- 


\ 


CHo-N 


H0-CH2-CH2 


/ 


(C2H40H)3NN(I)      X 

M 
(C2H40H)3l\r        xx 


K 


The  blue  color  of  the  nickel  salt  furnishes  evidence  for  structure  (V).  Since 
magnesium  does  not  form  stable  magnesium  to  nitrogen  coordinate  bonds 
with  other  amines,  structure  (IV)  is  favored  for  the  magnesium  salt.  Fur- 
ther work  by  Tettamanzi  and  Garelli21  showed  that  when  cobalt,  copper, 
or  zinc  was  used  as  the  central  atom,  a  hydrogen  of  one  hydroxyl  group  was 
replaced  by  the  metal  giving  a  compound  which  they  formulated  as 


0-CH2CH2 


\. 


H0-CHo-CH 


0 


HO-CH2-CH2 


Millet--  has  prepared  some  crystalline  derivatives  of  bismuth  triethanol- 


16.  August  i;  M  ie,  17,  11^    1935  . 

17.  Middleton,  Thesis.  University  of  Illinois,  IS 

18.  Price  and  Brazier,  ./.  Chem.  Soc.,  107,  1367     l'e 

19.  Schmits-Dumonl  and  Motzkus,  Ber.,  61,  581     l  IS 

H    Tettamanzi  and  Carli.  Gazz.  rhim.  it<il.,  63,  566  (1033  . 

21.  Tettamanzi  and  Garelli,  Gazz.  chim.  ital.,  63,  ."(>    1933) 

22.  Miller,../.  A  8oc.t  62,  2707    l'Un  . 


258  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

amine  and,  from  analytical  data,  has  assigned  the  formulas 

O— CIT2—  Cir,  O— CH2— CH2 

/  \  /  \ 

X:i— O— Hi  N— CH2— CH2OH     and     Bi— O— -CH2— CH2— N 

\  /  \  / 

O— CH2-  CII,  O— CH2—  CH2 

A  rather  odd  addition  compound  of  thallium  acetoacetic  ester  and  carbon 
disulfide  has  been  reported  by  Feigl  and  Backer23.  Because  of  the  color, 
insolubility,  and  stability  (even  toward  acids  and  bases)  the  authors  have 
proposed  the  following  eight-membered  ring  structure: 

CH3— C=C=C— O  C2H5 

I  I 

o         o 

I       I 

Tl  Tl 

\    / 

cs2 

The  double  enolization  of  the  methylene  group  is  experimentally  indicated 
by  the  fact  that  compounds  of  this  type  are  not  formed  if  one  or  both  of 
the  methylene  hydrogens  are  replaced  by  an  alkyl  or  aryl  radical.  It  seems 
hard  to  conceive  of  the  carbon  atom  in  the  carbon  disulfide  as  the  donor 
atom  because  it  has  no  available  electrons;  however,  each  of  the  sulfur  atoms 
has  electron  pairs  available,  so  it  seems  more  logical  for  the  structure  to  be 

CH3— C=C=C— OC2H5 , 


Tl  Tl 

I  I 

s=c=s 

thus  giving  rise  to  a  ten-member ed  ring. 

Some  early  work  by  Schlesinger24,  who  was  attempting  to  span  trans 
positions  with  a  bidentate  group,  resulted  in  the  preparation  of  a  number 
of  complexes  of  copper  with  polymethylene  bis-a-amino  acids: 

o=c — o  o  —  c=o 

1  X  1 

I        H^(cH2)n^H  I 

R  —  R 

Compounds  were  prepared  in  which  n  has  the  values  2,  3,  5,  7,  and  10; 
thus,  if  these  structures  are  correct,  the  rings  contain  5,  6,  8,  10,  and  13 


23.  Feigl  and  Backer,  Monatsh.,  49,  401  (1928). 

24.  Schlesinger,  Ber.,  58,  1877  (1925). 


LARGE  RINGS 


259 


members.  For  n  =  2  or  3,  deep  blue  compounds  are  formed,  for  n  =  10,  the 

product  is  red-violet,  and  for  ;/  =  5  or  7,  both  the  blue  and  violet  forms 
are  obtained.  These  products  are  nonelectrolytes  and  monomolecular  so 
that  cis-t rans  isomerism  was  suspected,  with  the  methylene  groups  span- 
ning trans  positions  in  the  red-violet  compound 


R' 

I 
NH-C-R 


0=0-0^  ^ 

I       A^CuCl 

R-C-NH  ^^-O  —  C=0 

R' 


Mat  tern-5  prepared  an  interesting  compound  in  which  an  eight-mem- 
bered  ring  apparently  spans  the  trans  positions  in  the  coordination  sphere 
of  a  platinum(II)  ion.  The  substance  was  produced  by  the  series  of  reactions 
shown  below. 


Pi 


ci 


Cl- 


Cl 
NH3      1  CI 


CI 


^Z-NH-CH,  - 
^\ 


FH 


H2N^ 1 NH3 

-  CI 


+  +  + 

REDUCE 


ELEC. 


P+ 


1 NH3 


CI 


(NH2CH2CH2)2NH 

ACTIVATED        *~ 
CHARCOAL 


^2-NH-ch2 
C*NH3  \ 


HoN NH3 


NH; 


++ 


The  structure  of  the  end  product  was  deduced  from  the  mode  of  preparation, 
analysis,  titration  of  available  chlorine,  and  preparation  of  the  dichloro- 
diammineplatinum(II)  complex  as  a  derivative.  This  dichloro  derivative 
was  shown  to  be  the  trans  isomer,  indicating  that  the  original  ion,  containing 
diethylenetriamine  hydrochloride,  was  also  trans  in  configuration.  When 
recrystallized  from  water,  the  compound  tended  to  rearrange,  liberating 
ammonium  chloride  according  to  the  equation 


•HCI 
,CH2-NHCH2C 

&*>  \ 

NH2 NH2 

\  /     pt      /  • 

H^  NH3 


a*' 


<**' 

fZ 


-CH2-CH2 

\ 
NH2 


Pt 


7 


+  nh4ci 


HoN 


Pfeiffer  and  co-workers26  have  investigated  the  reactions  of  various  metal 
ions  with  condensed  systems  of  salicylaldehyde  and  several  diamines.  With 
deeamethylenediamine  salicylaldehyde  and  copper(II)  ion  they  obtained 
a  compound  to  which  they  assigned  a  thirteen-membered  ring  structure, 

25.  Mattern,  thesis,  University  of  Illinois,  1946. 

26.  Pfeiffer,  et  al.,  Ann.,  503,  84  (1933);  J.  prakt.  Chem.  [2]  145,  243  (1936). 


260  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

(struct ure  (VI)  with  n  =   10) 

H-C       /  \       .CH  HC^N  N  =  CH 


N^/.,,   >  ^N 


(CH2)n- 
(2T) 


o  o 

I  I 

HC=NL  N=CH 

|  CU  I 


YE 


N-CH 


0-v'0 


VTTT 

Calvin  and  Barkelew27  have  also  reported  compounds  of  copper  with 
condensed  systems  of  salicylaldehyde  and  diamines  of  the  general  type 
shown  in  structure  (VI).  Penta-,  hexa-,  and  heptaamines  were  prepared  giv- 
ing rings  of  eight,  nine  and  ten  members  respectively.  As  shown  in  the 
structural  formulas,  these  molecules  also  involve  two  six-membered  rings. 
The  stability  of  these  smaller  rings  and  the  flexibility  of  the  di-,  tri-,  penta-, 
hexa-,  hepta-,  and  decamethylene  groups  probably  account  for  the  forma- 
tion of  these  complexes.  The  latter  factor  is  emphasized  in  the  cases  where 
ortho-,  meta-,  and  paradiamino  benzene  and  benzidine26  were  substituted 
for  the  decamethylenediamine  in  the  condensed  ring  system.  Monomeric 
compounds  were  first  reported.  However,  Pfeiffer  later  showed,  on  the  basis 
of  cryoscopic  measurements,  that  these  were  actually  dimers,  so  the  meta 
phenylenediamine  salt  would  have  structure  (VII)  which  contains  a  twelve- 
membered  ring  and  four  six-membered  rings.  The  corresponding  para- 
phenylenediamine  derivative  would  contain  a  fourteen-membered  ring, 
while  the  benzidine  dimer  would  contain  a  twenty-two-membered  ring  as 
shown  in  (VIII). 

It  is  quite  evident  that  the  proposed  structures  of  complexes  with  chelate 
rings  containing  more  than  six  atoms  are  not  firmly  established.  Lack 
of  x-ray  and  other  conclusive  data,  the  several  possible  linkages,  and  the 
possibility  of  polymerization,  all  tend  to  make  the  proposed  structures 
highly  speculative. 

27.  Calvin  and  Barkelew,  J.  Am.  Chem.  Soc,  68,  2267  (1946). 


/  .    General  Isomerism  of  Complex 
Compounds 

Thomas  D.  O'Brien* 

University  of  Minnesota,  Minneapolis,  Minnesota 

A  consideration  of  the  number  of  different  isomeric  forms  in  which  a 
relatively  simple  inorganic  coordination  compound  can  exist  makes  it  ap- 
parent that  the  study  of  the  isomerism  of  coordination  compounds  may 
become  extremely  complicated.  As  simple  a  compound  as  Co(en)2(H20)- 
I  NOj)C1j  can  exist  in  eighteen  different  isomeric  forms,  twelve  of  which  are 
optically  active.  Whereas  stereoisomerism  has  probably  been  the  most 
widely  investigated  of  the  different  types  of  isomerism,  the  others  are 
equally  important. 

Solvate  Isomerism 

The  classic  example  of  solvate  isomerism  is  concerned  with  the  three 
hydrate  isomers  of  the  compound,  CrCl3-6H20.  The  green  form,  which  is 
obtained  from  fairly  concentrated  solutions  of  hydrochloric  acid,  has  been 
assigned  the  formula  [Cr(H20)4Cl2]Cl-2H20  on  the  basis  of  conductivity 
measurements  and  relative  ease  of  precipitation  of  the  chlorides  with  sil- 
ver(I)  ion1.  Upon  dilution,  stepwise  aquation  takes  place.  The  resulting  solu- 
tions yield  the  blue-green  [Cr(H20)5Cl]Cl2  •  H20  and  the  violet  [Cr(H20)6]Cl3 . 

Britton2  reports  that  the  decrease  in  conductivity  and  the  decrease  in 
the  amount  of  chloride  precipitated  with  silver  nitrate,  in  going  from  the 
violet  to  the  green  form,  are  due,  not  to  the  transition  between  the  two 
forms  proposed  by  Werner1,  but  to  the  formation  of  a  green,  highly  ag- 
gregated, basic  chromium(III)  chloride  which  is  virtually  a  colloidal  elec- 
trolyte. If  this  explanation  were  correct,  the  green  solutions  should  be  more 
viscous  than  those  containing  the  violet  form  of  the  chromium  compound. 
However,  Partington  and  Tweedy3  measured  the  viscosities  and  found  that 

*  Now  at  Kansas  State  College,  Manhattan,  Kansas. 

1.  Werner  and  Gubser,  Ber.,  34,  1601  (1901);  Bjerrum,  Ber.,  39,,  1599  (1906);  Bjer- 
rum:  "Studier  over  Kromiklorid,"  Kopenhagen,  1907;  Bjerrum,  Z.  phys.  Chem., 
59,  336,  581  (1907). 

2.  Britton,  J.  Chem.  Soc,  127,  2128  (1925). 

3.  Partington  and  Tweedy,  Nature,  117,  415  (1926). 

261 


262  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

the  violet  solutions  are  more  viscous  than  the  green.  This  is  in  agreement 
with  Werner's  postulate  since  the  tervalent  hexaquochromium(III)  ion 
should  form  solutions  in  which  the  pseudolattice  is  more  stable  than  would 
be  the  case  with  the  singly  charged  dichlorotetraquo  ion. 

Some  doubt  has  been  cast  on  the  simple  interpretation  of  Werner  by  the 
results4  obtained  in  the  preparation  of  tris(ethylenediamine)chromium(III) 
chloride.  The  reaction  of  hexaquochromium(III)  chloride,  [Cr(H20)6]Cl3 , 
with  anhydrous  ethylenedimaine  in  toluene  solution  gives  a  yield  of  about 
25  per  cent  of  yellow  tris(ethylenediamine)chromium(III)  chloride.  Similar 
treatment  of  ordinary  hydrated  chromium(III)  chloride,  which  contains 
[Cr(H20)4Cl2]Cl-2H20  and  [Cr(H20)5Cl]Cl2-H20,  yields  none  of  the  tris- 
(ethylenediamine)  complex.  Normally,  ethylenediamine  replaces  coordi- 
nated chlorides  more  easily  than  it  replaces  coordinated  water.  Marchi  and 
McReynolds4  state  that  [Cr(H20)4Cl2]Cl-2H20  should  not  result  in  a  differ- 
ent product  than  that  obtained  with  [Cr(H20)6]Cl3 ,  and  that  the  system 
is  more  complex  than  is  implied  by  Werner. 

Further  evidence  that  the  equilibria  are  complex  has  been  reported  in 
connection  with  the  study  of  the  transformation  of  [Cr(H20)4Cl2]Cl-2H20 
to  [Cr(H20)6]Cl3  by  warming  in  dilute  solutions521  and  by  conductometric 
titration513.  In  the  dark,  equilibrium  was  reached  in  six  and  one-half  hours, 
but  in  ultraviolet  light  the  reaction  was  much  faster.  Also,  if  the  equi- 
librium mixture  obtained  in  the  dark  was  subsequently  exposed  to  ultra- 
violet light,  there  was  a  considerable  shift  in  the  equilibrium  point.  After 
measuring  the  pH,  conductance,  and  extinction  coefficient,  the  authors 
concluded  that  the  equilibrium  is  very  complex,  that  the  conversions  take 
place  in  steps  and  that  each  isomeric  change  is  preceded  by  rrydrolysis. 
This  evidence  does  not  appear  to  show  anything  about  the  nature  of  the 
isomerism.  The  shift  in  equilibrium  simply  indicates  that  the  different 
compounds  contain  different  amounts  of  energy.  Conductometric  titration 
of  chromium(III)  solutions  shows  nonstoichiometric  ratios  of  bound 
chloride,  the  breaks  occurring  at  1.54,  2.1,  and  3.0  equivalents  of  silver  ion. 

The  equilibria  are  probably  still  best  represented  by  the  simple  explana- 
tions  given  above.  Recent  kinetic  studies  support  this  conclusion50.  As  in 
any  other  chemical  reaction,  the  equation  is  not  intended  to  represent  a 
mechanism,  but  only  the  starting  materials  and  final  products. 

Fremy6  first  prepared  nitratopcntamminecobalt(III)  nitrate  1-hydrate, 
and  converted  it  to  the  solvate  isomer,  aquopentamminecobalt(III)  ni- 
trate. The  reverse  reaction  was  carried  out  by  Benrath  and  Mienes7. 

4.  Marchi  and  McReynolds,  J.  Am.  Chcm.  Soc,  65,  481  (1943). 
."».  Data!  and  Quershi,  •/.  Osmania  Univ.,  8,  6  (1940);  Law,  Trans.  Faraday  Soc,  32, 
1  Mil  |  1936) ;  llannn,  ./.  Am.  Chcm.  Soc,  73,  1240  (1951). 

6.  Fremy,  Arm.  chim.  phys.,  [3]  25,  296  (1852). 

7.  Benrath  and  Mienes,  Z.  anorg.  Chcm.,  177,  289  (1929). 


ISOMERISM  OF  COMPLEX  COMPOUNDS  263 

When  a  solution  of  sulfatopentaimninecobalt(III)  hydrogen  sulfate  2- 
hydrate,  [Co(NH,)sS04]HS04-2H20,  is  treated  with  chloroplatinic  acid, 
orange-red  crystals  of  sulfatopentanuninecobalt(III)  chloroplatinate  2- 
hydrate  precipitate.  An  isomeric  red-yellow  aquopentammine  complex, 
[Co(NH,)8H20]2(S04)2[PtCl6]J  is  obtained  when  sulfuric  acid  and  chloro- 
platinic acid  arc  added  to  an  aqueous  solution  of  sulfatopentammineco- 
balt(III)  sulfate8. 

Because  water  is  by  far  the  most  widely  used  solvent,  t  be  above  examples 
show  hydrate  isomerism,  but  this  by  no  means  precludes  the  possibility  of 
other  solvate  isomers  such  as  might  be  formed  by  alcohols,  amines,  or  am- 
monia. 

Coordination  Isomerism 

Two  salts  with  the  empirical  formula  CoCr(XH3)6(CX)6  are  known.  Both 
are  yellow  and  relatively  insoluble  in  water.  One  is  prepared  by  treating 
aqueous  hexamminecobalt(III)  chloride,  [Co(NH3)6]Cl3 ,  with  potassium 
hexaeyanochromate(III),9  K3[Cr(CX)6],  while  the  other  is  prepared  by 
treating  aqueous  hexamminechromium(III)  chloride,  [Cr(XH3)6]Cl3 ,  with 
potassium  hexacyanocobaltate(III),  K3[Co(CX)6].10  The  differences  be- 
tween them  can  easily  be  shown  by  treating  solutions  of  each  with  silver 
nitrate.  In  each  case  an  insoluble  silver  salt  is  obtained,  but  hexammine- 
cobalt(III)  nitrate  is  present  in  one  nitrate  and  hexamminechromium(III) 
nitrate  in  the  other.  It  follows  that  the  formulas  of  the  original  compounds 
are  [Co(XH3)6]  [Cr(CX)6]  and  [Cr(XH3)6]  [Co(CN)J.  Another  example  is 
found  in  the  isomerism  of  the  violet  tetramminecopper(II)  tetrachloro- 
platinate(II)  [Cu(XH3)4]  [PtCl4],  and  the  green  tetrammineplatinum(II) 
tetrachlorocuprate(II),  [Pt(XH3)4]  [CuCl4]. 

It  is  not  necessary  that  coordination  isomers  contain  two  different  cen- 
tral atoms,  as  in  the  examples  above.  Atoms  of  the  same  metal  can  appear 
in  both  the  cation  and]  the  anion  as  in  [Co(X"H3)4(X02)2]  [Co(XH3)2(X02)4] 
and  [Co(XH3)6]  [Co(X02)6]-H  A  similar  example  of  this  type  of  isomerism 
is  found  in  the  orange-yellow  [Co(XH3)6]  [Co(XH3)2(X02)4]3  which  is  iso- 
meric with  the  orange-red  salt  [Co(XH3)4(X02)2]3[Co(X02)6].n 

The  reversible  transformation  at  45°  of  a  double  silver-mercury  iodide 
from  a  red  to  a  yellowT  form12  has  been  explained  by  the  hypothesis  that  the 
change  is  due  to  a  change  in  function  of  the  metal  atoms  according  to  the 
equation  AgHg|AgI4]  ^±  Ag2[HgI4].  The  crystal  structures  are  showrn  in 

8.  Jorgensen,  J.  prakt.  Che?n.,  31,  271  (1885). 

9.  Braun,  Ann.,  125,  183  (1863). 

10.  Jorgensen,  ./.  prakt.  Chem.,  [2]  30,  31  (1884);  PfeifTer,  Ann.,  346,  42  (1906). 

11.  Jorgensen,  Z.  anorg.  Chem.,  5,  177  (1894);  ibid.,  5,  182  (1894);  ibid.,  7,  287  (1894); 

ibid.,  13,  183  (1897);  Werner  and  Miolati,  Z.  physik.  Chem.,  14,  514  (1894). 

12.  Roozeboom,  Proc.  K.  Akad.  Wctensch  ;//>,  3,  84  (1900). 


264 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Fig.  7.113.  X-ray  and  conductivity  measurements  show  that  in  the  alpha 
form  the  mercury  and  silver  atoms  statistically  fill  three  out  of  the  four 
equivalent  positions  in  the  crystal  lattice. 


fi+sz"sK 


cC-ASzHs\< 


•  =Ag,  0--Hg-,O  =  l 
Fig.  7.1.  Crystal  structures  of  the  two  forms  of  Ag2HgI4 

Polymerization  Isomerism 

The  word  "polymerization"  as  applied  to  polymerization  isomerism  in 
coordination  chemistry  has  a  different  connotation  from  that  in  modern 
usage  in  organic  chemistry.  In  organic  chemistry,  polymerization  denotes 


Table  7.1 

Formula 

Molecular 
Weight 

Properties11 

[Co.(NH,)6]  [Co(N02)6] 
[Co(NH3)4(N02)2]  [Co(NH3)2(N02)4] 

[Co(NH3)5N02]  lCo(NH3)2(N02)4l2 

[Co(NH3)c]  [Co(NH3)2(N02)4]3 

lCo(NH3)4(N02)2]3[Co(N02)6] 
[Co(NH3)BN02]3[Co(N02)6]2 

Double 
Double 

Triple 

Quadruple 

Quadruple 
Quintuple 

Yellow.  Insoluble  in  water. 

Yellow-brown.  Four  forms 
possible;  cis-cis,  trans- 
trans,  cis-trans,  trans- 
cis. 

Orange.  Difficulty  soluble. 
Anion  can  exist  in  cis  or 
trans  form. 

Yellow-orange.  Anion  can 
exist  in  cis  or  trans 
form. 

Orange-red.  Cation  can 
be  either  cis  or  trans. 

Brown-yellow. 

13.  Ketelaar,  Z.  Kriat.,  87,  436  (1934)  Fig.  7.1  is  taken  from  Clark,  Applied  X-rays, 
3rd  Edition,  p.  364.  McGraw-Hill  Book  Co.,  New  York,  1940. 


ISOMERISM  OF  COMPLEX  COMPOUNDS 


265 


Table  7.2 


Formula 

Molecular 
Weight 

Comments 

PI   NH.)tCl»]" 

Single 

Yellow.   The   cis   isomer  is   com- 
monly called  Peyrone's  chloride, 
while    the    trans    is    known    :is 
Reiset's  chloride. 

113)4]  [PtCl4]" 

Double 

Green.  Known  as  Magnus'  salt. 

DPt(NH,),Cl]  [Pt(NH3)Ci3] 

Double 

lPt(NH,)4]  [Pt(NH,)Cl,]«" 

Triple 

Orange-yellow. 

[Pt(NH,),Cl]«[PtCl4]1" 

Triple 

the  union  of  a  large  number  of  separate  units.  The  implications  associated 
with  the  term  in  coordination  chemistry  can  probably  best  be  illustrated 
by  the  following  examples,  which  were  originally  reported  by  Werner.  Six 
polymerization  isomers  of  trinitrotriammine-cobalt(III)?  [Co(NH3)3(N02)3], 
are  shown  in  Table  7.1.  Examples  are  known  in  the  chromium  series, 
also14, 15  and  a  platinum  series  is  included  in  Table  7.2. 

There  are  two  salts,  one  green  and  the  other  red,  with  the  empirical 
formula  (CH3)2Tel2 .  For  some  time  it  was  thought  that  the  compound 
had  a  planar  configuration  and  that  these  two  were  the  cis  and  trans  forms20. 
It  is  now  believed  that  the  green  salt  has  a  molecular  weight  double  that 
of  the  red  one,  and  that,  in  the  green  isomer,  one  tellurium  atom  is  associ- 
ated with  a  cation  while  a  second  is  a  part  of  an  anion ;  the  true  formula  is 
[(CH3)3TeI]  [CH3TeI3]21. 

Polymerization  isomers  of  complex  ions  in  which  diallylamine  behaves 
as  a  bidentate  group  have  been  prepared22,  [Pt{(CH2=CHCH2)2NH)Cl2] 
[Pt{(CH2=CHCH2)2NH}2]  [RCI4]. 

This  type  of  isomerism  is  also  known  in  cases  where  bridging  occurs. 
Octammine-^-diol-dicobalt(III)  bromide  2-hydrate  is  a  polymerization 
isomer  of  hydroxyaquotetramminecobalt(III)  bromide.  The  formulas  of 


14.  Werner  and  Jovanovits,  unpublished  work;  cf.  Werner:  "New  Ideas  on  Inorganic 

Chemistry,"  2nd  Ed.,  p.  232,  New  York,  Longmans,  1911. 

15.  Christensen,  J.  prakt.  Chem.,  45,  371  (1892). 

16.  Peyrone,  Ann.,  51,  1  (1844),  55,  205  (1845),  61,  178  (1847).  Gerstl,  Ber.,  3,  682 

(1870);  Odling,  Phil.  Mag.,  4,  No.  38,  455  (1870). 

17.  Magnus,  Pogg.  Ann.,  14,  242  (1828). 

18.  Cossa,  Ber.,  23,  2503  (1890). 

19.  King,  J.  Chem.  Soc,  1948,  1912. 

20.  Vernon,  J.  Chem.  Soc,  117,  86,  889  (1920);  119,  687  (1921);  Knaggs  and  Vernon, 

J.  Chem.  Soc,  119,  105  (1921). 

21.  Drew,  J.  Chem.  Soc,  1929,  560. 

22.  Rubinstein  and  Derbisher,  Dohladij  Akad.  Nauk  S.S.S.R.,  74,  283  (1950). 


266 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


these  compounds  are 


OH 

/   \ 
| Ml  )4Co  Co(NH3)4 

\    / 

OH 


Br4-2I120        and         fcoCNH,)*^ ]  Br2  , 


respectively2'.  Another  interesting  example  is  to  be  found  in  the  isomeric 
hexammine-ju-triol-dicobalt  (III), 

OH 

/      \ 

(NH3)  3  Co— OH— Co  (NH3) , 

\       / 
OH 

and   dodecammine-/x-hexol-tetracobalt (III) , 


Co 


HO, 


*HO' 


C0(NH3)4 


6  + 


ions24.  The  second  compound  may  be  considered  to  be  a  dimer  of  the  first, 
though  its  structure  is  quite  different.  The  structure  of  this  tetracobalt 
complex  was  proven  by  resolution  into  optical  enantiomers25.  The  structural 
formula  of  another  dodecammine-/x-hexol-tetracobalt(III)  ion  may  be 
written26 


NH3 


(NH3)4Co 


OH 


OH 


OH 


Co 


NH3 
OH 

,    /       \ 
Co  Co(NH3), 


OH 


OH 


NH3 


NH, 


However,  there  is  no  indication  that  this  compound  has  been  prepared. 
An  odd  type  of  polymerization  isomerization  is  implicit  in  the  work  of 
Rubinstein27,  who  reports  the  formation  of  a  new  compound  by  the  following 
reaction: 

[(NH3)4NH2ClPt]Cl2  +  [Pt(NH3)4Cl2]Cl2-+  [(NH3)4NH2ClPtCl2]  [(NH3)4Cl2PtCl2] 

23.  Werner,  Ber.,  40,  4434  (1907). 

24.  Birk,  Z.  anorg.  allgem.  Chem.,  175,  405  (1928);  Werner,  Ber.,  40,  4836  (1907). 

25.  Werner,  Ber.,  47,  3087  (1914). 

26.  Hiickel,  "Structural  Chemistry  of  Inorganic  Compounds,"  Vol.  I,  p.  166,  New 

York,  Elsevier  Publishing  Co.,  1950. 

27.  Pubinstein,  Izvest.  Seklora  Platini  i  Drug,  Blagorod  Metal.  Inst.  Obschei  i  Neorg. 

Khim.  Akad.  Nauk.  S.S.S.R.,  20,  53  (1947). 


ISOMERISM  OF  COMPLEX  COMPOl  NDS  267 

According  to  Rubinstein,  the  new  compound  is  characterized  by  the  differ- 
ences in  its  chemical  and  physical  properties  as  compared  bo  those  associ- 
ated with  mixtures  of  the  reactants.  The  author  did  not  indicate  any 
probable  structure  for  this  new  compound,  bul  it  might  be  formulated  as 
a  dinuclear  complex: 

NH2 


CI, 


CI  CI 

L  Nir3)«Pt— NH2— Pt(NH3)4JCl5 


(NHs)4Pt  Pt(NH3) 

\    / 
CI 

Ionization  Isomerism 

Bromopentamminecobalt(III)  sulfate28  is  dark  violet  in  color;  its  solu- 
tions give  no  precipitate  upon  the  addition  of  silver  nitrate  but  give  a 
precipitate  immediately  when  barium  chloride  is  added.  If  this  dark-violet 
salt  is  heated  with  concentrated  sulfuric  acid  and  then  cooled,  the  addition 
of  dilute  hydrogen  bromide  produces  a  violet-red  compound  of  the  same 
empirical  formula29.  This  violet-red  salt,  however,  gives  no  precipitate  when 
barium  chloride  is  added  but  silver  bromide  precipitates  immediately  with 
silver  nitrate.  From  these  experimental  facts  it  is  concluded  that  the  iso- 
mers are  bromopentamminecobalt(III)  sulfate,  [Co(XH3)5Br]S04 ,  and 
sulfatopentamminecobalt(III)  bromide,  [Co(XH3)5S04]Br. 

A  similar  set  of  isomers  consists  of  the  green  2rans-dichlorobis(ethylene- 
diamine)cobalt(III)  nitrite30,  trans-[Co  en2  C12]X02 ,  and  the  red  trans- 
nitrochlorobis(ethylenediamine)cobalt(III)  chloride31,  [Co  en2  C1X02]C1. 
Still  another  example  is  furnished  by  dihydroxytetrammineplatinum(IV) 
sulfate32,  [Pt(XH3)4(OH)2]S04 ,  which  yields  neutral  solutions,  and  sulfa- 
totetrammineplatinum(II)  hydroxide,  [Pt(XH3)4S04](OH)2 ,32  solutions  of 
which  are  strongly  basic. 

Xyholm33  has  reported  a  compound  having  the  formula 

PdBr2{As(C2H5)  (C6H5)2}3, 

that  exists  in  two  forms  which  might  be  considered  to  be  ionization  isomers. 
The  compound  is  soluble  in  organic  solvents  both  at  room  temperature  and 
at  low  temperatures.  Molecular  weight  determinations  indicate  that  it  is 
dissociated  over  the  entire  temperature  range  studied.  However,  it  under- 

28    Jorgensen,  J.  prakt.  Chem.,  [2]  19,  49  (1879);  Z.  anorg.  Chem.,  17,  463  (1898); 
Diehl,  Clark,  and  Willard,  Inonjanic  Syntheses,  1,  186  (1939). 

29.  Jorgensen,  J.  prakt.  Chem.,  [2]  31,  262  (1885). 

30.  Jorgensen,  J.  -prakt.  Chem.,  [2]  39,  1   (1889). 

31.  Werner,  Ber.,  34,  1773  (1901). 

32.  Cleve,  K.  Sv.  Vet.  A  had.  Handl.,  10,  No.  9  (1871). 

33.  Xyholm,  J.  Chem.  Soc,  1960,  848. 


268  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

goes  a  change  in  character  as  the  temperature  is  lowered  considerably  be- 
low zero.  The  color  of  the  solutions  changes  at  —  78°C,  and  a  distinct  in- 
crease in  the  conductivity  is  observed.  The  equilibria  proposed  to  explain 
this  behavior  are  shown  below. 

(solid)  [PdBr(AsR,),]Br    b^nlU") )  [PdBr2(AsR,)2]  +  AsR3 

[PdBr2(AsR3)2]  +  AsR3  ^— >  [PdBr(AsR3)3]+  +  Br~ 

Structural  Isomerism 

The  existence  of  this  type  of  isomerism  is  based  almost  exclusively  on 
the  nitro  and  nitrito  compounds.  Jorgensen34  prepared  two  compounds 
in  the  following  manner: 


Cool 


[Co(NH3)sCHCl2  -Mi,       Jff*cl    >  -^ 


heat/    conc-  HC1 


stand 


*  [Co(NH3)5N02]Cl2 
brown-yellow 


in  cold 


->  [Co(NH3)5N02]Cl2 
red 


The  red  form,  believed  to  contain  the  nitrito  group,  is  converted  to  the 
brown-yellow  nitro  form  quite  rapidly  by  heating  in  solution  or  by  adding 
concentrated  hydrochloric  acid.  It  changes  slowly  even  in  the  solid  state. 
Lecompte  and  Duval35  prepared  these  two  salts  according  to  the  method 
of  Jorgensen34  and  determined  the  Debye-Scherrer  patterns,  the  infrared 
absorption,  and  the  ultraviolet  absorption  bands.  The  Debye-Scherrer 
patterns  were  "rigorously  identical."  By  comparison  with  organic  nitro 
and  nitrito  compounds,  Lecompte  and  Duval  concluded  that  there  were 
no  — O — N=0  links  in  the  red  cobalt  compound,  but  only  those  of  the 
O 

— N  type.  The  isoxantho  or  red  compound,  in  addition  to  having  the 

\ 
O 

same  two  strong  absorption  bands  at  0.5  and  7..")  M  as  the  xantho  or  yellow7 

-.ill,  showed  an  additional  band  at  7.65  fi.  This  was  shown  to  be  the  same 

as  the  maximum  absorption  band  of  chloropentamminecobalt(III)  chloride, 

the  starting  material  in  the  preparation  of  the  nitro  complex.  Lecomte  and 

Duval  conclude  that  the  red  color  is  due  to  the  presence  of  some  unreacted 

starting  material.   These  results  are  in  accord  with  the  earlier  work  of 

Piutii'1  who  reported  thai  the  absorption  spectra  of  the  two  forms  are 

34.  Jorgensen,  Z.  anorg.  ('hem.,  5,  168  (1894). 

;;.").  Lecompte  and  Duval,  Bull.  soc.  chirn.,  12,  678  (1945). 

36.  Piutti,  Ber.,  46,  1832  (1912). 


ISOMERISM  OF  COMPLEX  COMPOl   VDS  269 

identical.  Shibata17,  however,  claimed  that  the  two  forms  had  quite  different 
absorption  spectra. 

Adell*  measured  the  rate  of  conversion  of  the  "nitrito"  to  the  nitro  form 
photometrically  and  concluded  thai  it  followed  the  law  for  a  first  order 

reaction.  These  results  can  be  considered  to  be  only  indirect  structural 
evidence;  however,  it  should  be  pointed  out  thai  the  conversion  of  the 
highly  ionized  salt,  chloropentanuninecobalt(III)  chloride-nitrite, 
[Co(NHj)iCl]ClNOj ,  to  the  nitro  complex  (assuming  the  conclusion  of 
Lecomte  and  Duval  to  be  correct)  in  solution  should  follow  a  second  order 
rati1  law,  unless  the  rate-determining  step  is  a  slow  rearrangement  which 
takes  place  subsequent  to  the  collision  of  a  ehloropentamminecobalt(III) 
ion  and  a  nitrite  ion.  This  would  imply  a  mechanism  of  substitution  involv- 
ing a  temporary  coordination  number  of  seven  for  the  cobalt  ion. 

Yalman  and  Kuwanawb  have  confirmed  the  results  of  Adell38  and  have 
shown  that  the  conversion  of  the  cis  dinitritotetrammine  salt  to  the  cor- 
responding cis  dinitro  salt  is  also  first  order.  However,  they  were  unable 
to  show  by  spectrophotometric  means,  the  existence  of  the  cis  nitritonitro- 
tetramminecobalt(III)  salt,  a  logical  intermediate  in  the  isomeric  trans- 
formation. Neither  were  they  able  to  synthesize  the  cis  nitritonitro  salt 
from  the  corresponding  cis  nitroaquo  salt. 

Basolo,  Stone,  Bergman,  and  Pearson38c,  however,  report  the  existence  of 
the  analogous  cis  nitritonitro-bis  (ethylenediamine)  cobalt(III)  compound, 
but  state  that  it  is  relatively  unstable  and  undergoes  an  intramolecular  re- 
arrangement to  the  stable  cis  dinitro  compound.  The  nitritonitro  isomer 
could  be  isolated  only  when  stabilized  by  high  concentrations  of  nitrite 
ion. 

The  strongest  evidence  for  the  existence  of  the  two  structural  isomers 
comes  from  the  work  of  Taube  and  Murmann  (private  communication), 
who  studied  the  reaction 

[Co(XH3)50*H]++  +  HOXO  ->  [Ck)(NH,)60*NO]++  +  H20, 

(where  0*  is  oxygen  enriched  with  O18).  Their  results  show  all  of  the  heavy 
oxygen  isotope  is  retained  in  the  nitritopentamminecobalt  ion,  indicating- 
no  rupture  of  the  cobalt-oxygen  bond  in  the  transformation.  When  the  pink 
nitrito  sail  was  heated  either  in  the  solid  state  or  in  solution,  the  yellow 
nitro  isomer  was  formed.  This,  when  treated  with  excess  NaOH  to  reform 
the  hydroxypentammine  cobalt  salt,  released  all  the  heavy  oxygen  in  the 
nitrite  ion. 

/.  Coll.  Sri.  Imp.  Univ.  Tokyo,  37,  15  (1915  . 

38.  A<lell,  Sicnsk.  Km,.  Tvi.,  56,  :;>ls    1944   \Z. anorg.  C hem., 25%,  272  (1944  . 

Yalman  and  Kuwana,  paper  presented  before  Physical  and  Inorganic  Division, 

American  Chemical  Society,  Kansas  City,  April,  1954. 
38c.  Basolo,  Stone,  Bergman,  and  Pearson,  ./.  .1///.  Chem.  Soc,  76,  3079,  5920  (1<)">  1 


270  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Aquopentamminecobalt  salts  gave  the  same  results  when  treated  with 
nitrite  ion,  as  did  diaquotetramminecobalt  ion. 

Tracer  experiments  further  showed  that  in  going  from  the  nitrito  to  the 
nitro  isomer,  there  was  no  oxygen  exchange  of  the  coordinated  nitrite  with 
the  solvent  water  or  with  nitrite  ions.  After  isomerization  was  completed 
there  was  no  exchange  of  the  nitro-oxygen  with  nitrite  ion. 

These  results  indicate  that  the  isomerization  must  occur  by  an  intra- 
molecular rearrangement  in  which  the  nitrite  ion  is  never  free  and  in  which 
the  oxygen  first  linked  to  the  cobalt  is  completely  transferred  to  the  nitro- 

gen  :  Co;; 

N— O 

The  dithiocyanatobis(ethylenediamine)cobalt(III)  halides  were  reported 
by  Werner  to  exist  in  two  forms39,  one  red  and  the  other  blue-red.  These 
two  forms  were  thought  to  differ  in  the  manner  in  which  the  thiocyanate 
groups  are  linked  to  the  central  atom.  However,  Werner40  later  showed  that 
in  both  forms  the  thiocyanate  group  is  attached  to  the  cobalt  through  the 
nitrogen  and  so  concluded  that  these  are  cis-trans,  rather  than  structural 
isomers. 

Ray  and  Maulik41  report  isomerism  associated  with  the  compound 
H4[(CN)5Co(S203)].  These  investigators  suggested  that  it  is  possible  that 
coordination  takes  place  through  oxygen  in  one  case  and  through  sulfur 
in  the  other,  thus  giving  rise  to  structural  isomerism.  This  suggestion  is 
supported  by  the  fact  that  the  solid  salts  of  the  normal  form  are  gold  in 
color  while  those  in  which  the  thiosulfate  ion  is  supposedly  coordinated 
through  a  sulfur  atom  are  brown. 

Other  Types  of  Isomerism 
Coordination  Position  Isomerism 

Another  type  of  isomerism  is  encountered  in  the  polynuclear  compounds. 
Werner42  calls  this  "Coordination  Position  Isomerism";  it  is  illustrated  by 
symmetrical  dichlorohexammine-ju-diol-dicobalt(III)  chloride 


CNH3)3         /0H^         (NH3)3 
Co              Co^ 

Cl2 

and  the  unsymmetrical 

"CNH3)2              OH 

^Co               CO  =(NH3)4 

Cl2 

39.  Werner  and  Braunlich,  Z.  anorg.  Chem.,  22,  127,  141  (1899), 

in.  Werner,  Ann.,  386,  22,  41,  192  (1912). 

ll.  Ray  and  Maulik,  Z.  anorg.  Chem.,  199,  353  (1931). 

42.  Werner,  Ann.,  375,  7,  39,  32,  107,  111  (1910). 


ISOMERISM  OF  COMPLEX  COMPOUNDS 


271 


Werner-  also  studied  Baits  containing  the  symmetrical  and  imsymmetrical 
forms  of  dicUorohexammine-Ai-amino-peroxo-cobalt(III)-cobalt(rV  I   ions, 


CI. 
(NH3V 


NH- 


,co: 


co; 


x 


and 


ci2. 


lCnh3)2 


^ 


Co 


0E 


,NH2< 


Oi 


CI 


'(NHO 


3'3 


+  + 


AND 


C0=CNH3)4 


+  + 


The  first  isomer  forms  gray-black  salts  which  are  difficulty  soluble  in  water, 
while  the  second  is  green-brown  in  color  and  is  easily  soluble  in  water. 

Jensen43  described  a  second  type  of  coordination  position  isomerism  in  the 
rhodo  and  er3rthrochromic  complex  ions.  The  two  isomers  differ  in  the 
nature  of  the  bridge  group  connecting  the  two  cobalt  atoms.  The  rhodo 
and  erythro  complex  ions  are  reported  by  Jensen  to  have  the  formulas, 

H20 


[(XH3)oCrOHCr(XH3)5]5+        and 


[(XH; 


)5CrXH2Cr 


H20      ~\ 

(XH,)J 


respectively.  Recent  work44  indicates  that  these  ions  are  not  isomeric  but 
that  they  have  the  formulas,  rhodochromic, 

[(XH3)5CrOHCr(XH3)5p+;  eiythrochromic,  [(XH^CrOHCr^^4]5*, 

Isomers  Resulting  from  Isomerism  of  Ligands 

S  eral  types  of  isomerism  met  in  organic  chemistry  are  also  found  in 
the  inorganic  field.  For  example,  Ablov45  has  studied  the  reaction  of  chloro- 
aniline  with  ?ra^s-dichlorobis(ethylenediamine)cobalt(III)  chloride  and 
found  that  the  reaction  involves  only  rearrangement  to  the  cis  form.  How- 
ever, under  the  proper  conditions,  it  is  entirely  possible  that  chloroaniline 
could  replace  a  coordinated  chloride  to  give 


Co  en2 


<_> 


XH2  \C1 


CI 


Isomers  of  this  ion  could  exist,  depending  on  whether  the  chloroaniline  were 
ortho,  meta,  or  para.  The  action  of  toluidine  on  dichlorobisfethylenedi- 

43.  Jensen,  Z.  anorg.  Chem.,  232,  257  (1937). 

44.  Wilmarth,  Graff,  Gustin,  and  Dharmatti,  "The  Structure  and  Properties  of  the 

Rhodo  and  Erythro  Complex  Compounds,"  preprint,  Symposium,  Division  of 
Physical  and  Inorganic  Chemistry,  American  Chemical  Society,  1952. 

45.  Ablov,  Bull.  soc.  chim.  [5]  3,  2270  (1936). 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

amine)cobalt(III)  chloride  has  been  reported  to  result  in  the  compound, 
[Co  en2(CH3C6H4NH2)Cl]Cl2 ,46  which  can  exist  in  forms  containing  either 
ortho,  meta,  or  para  toluidine.  Similarly,  Kats47  and  Griinberg  have  re- 
ported dichlorobis(aminobenzoic  acid)platinum(II), 

[Pt(NH2C6H4COOH)2Cl2], 

in  which  ortho,  meta,  or  para-aminobenzoic  acid  is  present  in  the  coordina- 
tion sphere. 

Ring  Size  Isomerism 

The  isomerism  of  the  many  diamines  used  as  coordinating  groups  may 
lead  to  different  types  of  isomerism  in  the  coordination  compounds  formed. 
One  of  these  is  dependent  on  ring  size.  Tris(propylenediamine)  cobalt  (III) 
chloride  and  tris(trimethylenediamine)cobalt(III)  chloride  illustrate  this 
phenomenon48.  The  trimethylenediamine  compound  is  less  stable,  more 
soluble,  and  different  in  color  from  the  propylenediamine  complex. 

Summation  Isomerism 

Another  type  of  isomerism  which  might,  for  want  of  a  better  name,  be 
called  "summation  isomerism"  includes  those  instances  in  which  entirely 
different  groups  are  coordinated  to  the  central  atom,  but  the  sum  of  all 
the  atoms  is  constant.  An  example  is  to  be  found  in  the  identical  empirical 
formulas  of  the  complex  ions,  dichloro(tetramethylenediamine)  (ethylene- 
diamine)cobalt(III)  and  dichlorobis  (trimethylenediamine)  cobalt  (III).  Al- 
though the  following  pairs  of  complexes  have  not  actually  been  prepared, 
they  serve  to  exemplify  the  type  of  isomerism  under  consideration: 

[Co(NH3)4Cl(Br03)]+,         [Co(NH3)4(C103)Br]+; 

lCo(NH3)4(S03)(SCN)],        [Co(NH3)4(S203)(CN)]; 

[Co(NH3)4(C103)(N03)]+,        [Co(NH3)4(C104)(N02)]+. 

Electronic  Isomerism 

The  cations  of  nitrosylpentamminecobalt  salts49  may  be  obtained  in  two 
forms  which  are  strikingly  different  in  their  physical  and  chemical  proper- 
ties, though  their  stoichiometrics  and  structural  formulas  are  identical, 
[Co(NH3)5NO]++.50  The  chloride  of  one  series  is  black  and  paramagnetic 

46.  Bailar  and  Clapp,  /.  Am.  Chem.  Soc,  67,  171  (1945). 

47.  Kats  and  Griinberg,  Zhur.  Obshchei  Khim.,  20,  248  (1950). 

48.  Bailar  and  Work,  /.  Am.  Chem.  Soc,  68,  232  (1946). 

49.  Moeller,  J.  Chem.  Ed.,  23,  542  (1946). 

50.  Sand  and  Genssler,  Ber.,  36,  2083  (1903);  Werner  and  Karrer,  Helv.  chim.  acta., 

1,  54  (1918) ;  Milward,  Wardlaw,  and  Way, ./.  Chem.  Soc.,  1938,  233;  Ghosh  and 
Ray,  J.  Indian  Chem.  Soc,  20,  409  (1943). 


ISOMERISM  OF  COMPLEX  COMPOUNDS  273 

while  the  corresponding  Bait  of  the  second  scries  is  pink  and  dia- 
ma^netic500,  50d-  51.  It  is  believed  that  dipositive  cobalt  and  neutral  nitro- 
gen(II)  oxide  are  present  in  the  black  salt  and  that  tripositive  cobalt  and 
\<  I     ions  are  present  in  the  pink  complex60,  5,!l. 

51.  Frazer  and  Long,  J.  Chem.  Phys.,  6,  462  (1938);  Mellor  and  Craig,  J.  Proc.  Roy. 
Soc.,  N.S.  Wales,  78,  25  (1944);  Ray  and  Bhar,  J.  Indian  Chem.  Soc.,  5,  497 
(1928). 


8 


Stereoisomerism  of  Hexacovalent  Atoms 


Fred  Basolo 
Northwestern  University,  Evanston,  Illinois 

Introduction 
Werner's  Coordination  Theory 

Shortly  after  Tassaert1  discovered  the  compound  C0CI3  •  6NH3 ,  it  was 
noticed  that  some  of  the  complex  compounds  with  the  same  chemical 
composition  had  very  markedly  different  properties.  It  was  known,  for  in- 
stance, that  CoCl3-4NH3  could  exist  as  a  dark  purple  or  a  bright  green 
crystalline  salt.  In  terms  of  the  structure  of  the  molecule,  this  implies  that 
the  two  forms  differ  in  the  arrangement  of  the  atoms  in  the  molecule. 
Numerous  theories  (Chapter  2)  were  proposed  in  an  attempt  to  explain 
the  experimental  facts;  at  the  turn  of  the  century  there  were  three  popular 
theories.  Jorgensen2  modified  the  chain  theory  of  Blomstrand3  and  repre- 
sented what  we  now  call  the  cis  and  trans  isomers  of  [Co  en2  C12]C1  as 
shown  in  Fig.  8.1.  Friend4  designated  the  structures  by  means  of  a  "shell" 

CI  CH2-CH2    CH2  ~CH2  C(  CH2  CH2 

Co-NH2-NH2— NH2— NH2— C!  NCo-NH2-  NH2—  NH2~ NH2—  CI 

CI  CI  CH2 CH2 

trans  cis 

Fig.  8.1 

surrounding  the  central  atom  (Fig.  8.2).  In  his  coordination  theory,  Werner 

1.  Tassaert,  Ann.  chim.  phys.,  28,  92  (1798). 

2.  Jorgensen,  Z.  anorg.  Chem.,  5,  147  (1894). 

3.  Blomstrand,  Ber.,  4,  40  (1871). 

1.  Friend,  Trans.  Chem.  Soc,  93,  260  (1908). 


274 


STEREOISOMERISM  OF  HEXACO}    VLENT  ATOMS  275 


0^ 

CI 

t>*\ 

/ 

NH2        / 

NHp- 

-CH? 

I        Xo 

1 

1 

cr      ;    nh2- 

-CH2 

TRANS  CIS 

Fig..  8.2 

postulated  that  there  must  be,  in  addition  to  the  primary  valence  bond,  a 
secondary  valence  bond.  Unlike  Friend,  he  said  the  coordination  groups  are 
connected  to  the  metal  and  not  to  each  other  (Fig.  8.3). 

/CH2 
CI  CH2       ^  NH2 

CHp^  I  NH2/ .CI 

/      2^NH2 1 NhU     rM  2/      1 

I  /       CO       /  2  -CH2  /     Co        / 


/         UO      / 


Cb2 NH^        I         NH2-CH2  £"2 


NH?  T-      CI 


CHz 


CI 


-^CH2-NH2 


TRANS  £!§ 

Fig.  8.3 

Werner  predicted  that  cis-[Co  en2  CUJCl  would  be  found  to  be  optically 
active;  this  could  be  accounted  for  on  the  basis  of  the  octahedral  structure 
which  he  proposed.  Jorgensen  mentioned,  however,  that  his  structure  like- 
wise permitted  a  symmetrical  trans  form  and  an  asymmetrical  cis  form. 
After  the  accumulation  of  more  experimental  data,  Werner  was  able  to 
convince  his  contemporaries  that  the  structure  he  had  proposed  was  cor- 
rect. Of  course,  with  the  present-day  knowledge  of  atomic  structure,  the 
configuration  proposed  by  Jorgensen  can  be  ruled  out  immediately,  since 
it  involves  five  covalent  bonds  attached  to  one  nitrogen  atom. 

Proof  of  Octahedral  Structure  of  Hexacovalent  Elements 

Three  of  the  more  symmetrical  arrangements  of  six  equivalent  groups 
about  a  common  center  are:  (a)  plane  hexagonal,  (b)  trigonal  prismatic 
and  (c)  octahedral  (Fig.  8.4).  If  these  groups  differ  in  composition  they  can 
be  arranged  in  different  ways  depending  on  the  structure  or  spatial  ar- 
rangement of  the  system.  The  number  of  possible  arrangements,  or  of 
stereoisomers,  will  suggest  the  geometric  configuration  involved.  Each  of 
the  three  models  under  consideration  allows  only  one  possible  form  for  the 
compound  [Ma5b];  for  the  compound  [Ma4b2],  (A)  and  (B)  Lead  to  three 


276 


(HhMISTRY  OF  THE  COORDINATION  COMPOUNDS 


isomer 


rs  while  (C)  allows  only  two  forms;  for  the  compound  [Ma3b3],  (A) 
and  (B)  again  give  three  forms  while  (C)  gives  only  two  isomers. 

Stereoisomers  Theoretically  Possible 


Com- 
pounds 

Ma5b 

Ma4b2 

Ma3b3 


(A)  Plane  hexagonal 

one 

three(l,2;l,3;l,4) 
three  (1,2,  3;  1,  2,4;1 
3,  5) 


M 

3 


(B)  Trigonal  prismatic 

one 

three  (1,  2;  1,  3;  1,  4) 

three  (1,  2,  3;  1,  2,  5;  1, 

2,  6) 

Fig.  8.4 


(C)  Octahedral 

one 

two  (1,  2;  1,  6) 

two  (1,2,  3;  1,2,  6) 


Many  compounds  of  the  types  [Ma^]  and  [Ma3b3]  have  been  prepared 
and  in  no  case  has  it  been  possible  to  isolate  more  than  two  isomers.  This 
would  indicate  that  the  octahedral  arrangement  is  correct,  but  it  should 
be  remembered  that  failure  to  isolate  a  third  form  does  not  necessarily 
prove  its  nonexistence. 

Much  more  conclusive  evidence  on  the  spatial  arrangement  of  the  groups 
can  be  obtained  by  considering  the  symmetry  of  the  entire  complex.  If  it  is 
assumed  that  bidentate  groups  span  only  adjacent  positions,  then  the 
compound  [M(AA)3]  may  exist  in  one  form  if  the  structure  is  plane  hex- 
agonal and  two  forms  if  it  is  either  trigonal  prismatic  or  octahedral  (Fig. 
8.5).  The  trigonal  prismatic  arrangement  yields  two  geometrical  isomers, 


AA 


AA 


AA  ( 


AA 


kP 


lAA        AAl 


A"A 
AA 


AA 


AA 


AA 


.A  A 


\J 


V 


AA 


(a)  Plane 

Hexagonal 


(b)  Trigonal  Prismatic 
(Geometrical  Isomers) 

Fig.  8.5 


(c)  Octahedral 

(Optical  Isomers) 


each  of  which  has  a  plane  of  symmetry,  but  an  asymmetric  molecule  re- 
sults if  the  arrangement  is  octahedral.  Werner5  prepared  the  purely  inor- 

/     /0HV 

ganic  compound  [Co(AA)3]6+,  in  which  AA  =      (NH3)4Co  ,  and 

\  OH/ 

5.  Werner,  Ber.,  47,  3087  (1914). 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS  277 

resolved  it  by  means  of  the  dextro-a-bromocamphor-T-sulfonate  into  dex- 
tro  and  levo  forma  (see  page  323).  This  proved  conclusively  the  octahedral 

structure  of  hexacovalent  cobalt  (III)  and  it  is  now  realized  that,  almost 
without  exception,  this  is  the  correct  structure  for  compounds  containing 
atoms  which  are  hexacovalent. 

The  Stereochemistry  of  Inorganic  Complex  Compounds  Compared 
to  That  of  Organic  Compounds 

The  octahedral  configuration  of  hexacovalent  metals  is  now  as  generally 
accepted  as  the  tetrahedral  configuration  of  carbon.  It  presents  many  more 
possibilities  for  isomerism  and  intramolecular  rearrangement  than  does 
the  tetrahedral  configuration  of  carbon.  There  are  numerous  questions 
which  have  not  yet  been  answered,  largely  because  the  syntheses  for  these 
complex  compounds  are  often  based  on  empirical  knowledge  alone  and  it  is 
frequently  impossible  to  make  a  molecule  of  known  configuration.  The  num- 
ber of  possible  isomers  becomes  extremely  large  as  the  degree  of  complexity 
of  the  molecule  increases;  a  compound  of  the  type  [Mabcdef]  may  exist  in 
thirty  different  forms  (fifteen  pairs  of  mirror  images).  It  is  not  surprising, 
therefore,  that  very  little  is  known  of  compounds  more  complex  than 
[M(AA)a2b2]. 

Geometrical  Isomerism  (cis-trans  Isomerism) 

The  octahedral  structure  of  hexacovalent  atoms  wTas  first  indicated  by 
the  fact  that  only  two  stereoisomers  could  be  isolated  for  compounds  of 
the  types  [Ma4b2]  and  [Ma3b3].  On  the  basis  of  this  structure,  the  number 
of  position  isomers  theoretically  possible  for  any  complex  can  be  easily  de- 
termined; in  some  cases  all  of  the  predicted  isomers  have  been  isolated, 
but  many  instances  are  known  in  which  only  the  most  stable  form  has 
been  obtained. 

Chelating  Molecules  Occupy  cis  -Positions 

The  principle  that  chelating  groups  span  adjacent  cis  and  not  remote 
trans  valence  bonds  of  the  central  atom  has  been  widely  used  to  determine 
the  configuration  of  complex  compounds  and  to  prepare  compounds  of 
known  configuration.  This  principle  was  derived  by  comparing  chelate  ring 
formation  with  the  formation  of  maleic,  but  not  fumaric  anhydride,  and 
from  the  similarity  of  metal  and  carbon  atoms  in  forming  five-  and  six- 
membered  rings  more  readily  than  those  containing  larger  numbers  of 
atoms.4  Tic--7  points  oul  that  this  principle  can  also  be  deduced  from  the 
isomerism  of  certain  types  of  complex  compounds.  In  tin-  complex 
[M(AA)2bo],  if  the  chelating  group  -pan-  only  cis  positions  the  compound 

6.  Wen,  40,  51  (1907  . 

7.  Tress,  Chemistry  <fe  Industry,  1938,  1234. 


278 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


can  exist  in  a  racemic  mixture  and  one  inactive  trans  form;  however,  if  it 
spans  trans  positions,  only  a  racemate  is  possible  (Fig.  8.6).  A  point  which 


AA 


^J 


AA 


AA 


AA 


^AA 

CIS  (dl) 

Group  AA  spanning  cis-positions 


b 

RANS 


AA 


Group  AA  spanning  trans-positions 
Fig.  8.6 

was  not  mentioned  by  Tress  is  that  this  assumes  the  trans  spanning  groups 
are  not  free  to  rotate  around  the  corners.  If  this  rotation  were  possible  then 
only  one  optically  inactive  form  would  exist.  Numerous  compounds  of  this 
type,  which  exist  in  racemic  and  inactive  forms,  are  known.  In  addition, 
several  compounds  of  the  type  [M(AA)a2b2]  have  been  resolved  into  their 
optically  active  antipodes.  Optical  activity  can  exist  in  these  compounds 
only  if  the  chelate  ring  spans  cis  positions;  (Fig.  8.7). 


a,.  -  I 

/ 


n 


b- - 


(d.0 

Group  AA  spanning  cis  positions 


,y. 


X 

__        b 

/ 

M           / 

a  -  -   H 

nVb 

AA 

V_ 

(OPTICALLY    INACTIVE) 

Group  AA  spanning  trans 

P< 

»8] 

fcions 

Fig.  8.7 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS 


279 


Although  it  is  generally  agreed  that  chelating  groups  such  as  ethylene- 
diamine  are  stem-ally  incapable  of  spanning  trans  positions  in  the  coordi- 
nation sphere  of  a  metal,  there  is  no  reason  to  suppose  that  a  chelating 
group  of  sufficient  size  cannot  do  so  under  the  proper  conditions.  However, 
except  for  recent  work  by  Pfeiffer8,  all  attempts  to  prepare  simple  chelate 
rings  oi  seven  or  more  members  have  given  inconclusive  or  negative  results 
(Chapter  (>).  A  new  approach  has  been  studied9  using  2-chloro-l  ,6-diam- 
mine-3,4.r)-diethylenetriamineplatinum(IV)  chloride  (see  page  259). 

Various  Types  of  Cis-trans  Isomers 

Cat  ionic  Complex  Compounds.  The  method  of  preparation  of  both 
the  cis  and  trans  isomers  of  a  complex  depends  upon  the  compound  in  ques- 
tion and  no  general  rules  for  the  preparation  of  these  isomers  can  be  laid 
down.  It  must  also  be  remembered  that  molecular  rearrangements  are  com- 
mon in  reactions  of  coordination  compounds  and  that  the  expected  isomer 
may  not  always  be  the  one  isolated.  The  fact  that  bidentate  groups  span 
cis  positions  suggests  the  possibility  of  preparing  a  cis  salt  by  the  displace- 
ment of  groups  occupying  cis  positions.  This  technique  has  been  employed. 

A  very  common  starting  material  for  the  preparation  of  diacidotetra- 
minecobalt(III)  complexes  is  carbonatotetramminecobalt(III)  nitrate10.  The 
carbonate  radical  is  coordinated  firmly  to  the  cobalt  as  is  illustrated  by 
the  fact  that  it  does  not  precipitate  upon  the  addition  of  barium  chloride. 
However,  it  does  liberate  carbon  dioxide  when  acid  is  added  (Fig.  8.8). 


NH 


NH3 


NH 


Fig.  8.8 


Assuming  that  no  rearrangement  lakes  place  during  this  reaction,  one  can 

expeci  to  obtain  the  corresponding  cw-diacido  compound.  Rearrangement 

to  th<-  trans  -alt  can  be  kept  al  a  minimum,  if  the  solid  complex  is  allowed 
to  react  with  an  alcoholic  solution  of  the  desired  acid. 

8.  Pfeiffer,  Bohn,  and  Schmitz,  Natururissenschaften,  35,  190    1948  . 

9.  M:ittf ■ni.  thesis,  University  of  Illinois,  1947. 

10.  Biltz  and  Hiltz,  "Laboratory  Methods  of  Inorganic  Chemistry,"  translated  by 
Hall  and  Blanchard,  p.  171.,  New  York,  John  Wilej 


280  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

The  procedure  described  above  is  adaptable  to  the  preparation  of  cis- 
[Co(NH3)4(N02)2]+,  which  is  yellow-brown.  The  orange-yellow  isomeric 
ion,  frans-[Co(NH3)4(N02)2]+,  is  readily  obtained  by  the  oxidation  of  co- 
balt (II)  chloride  6-hydrate  in  the  presence  of  ammonium  chloride,  am- 
monia and  sodium  nitrite11.  These  stereoisomers  react  differently  with 
concentrated  hydrochloric  acid;  the  cis  salt  dissolves  completely  in  the 
boiling  acid,  forming  the  green,  crystalline  £rans-[Co(NH3)4Cl2]Cl,  whereas 
the  trans  salt  forms  a  red  precipitate  of  /rans-[Co(NH3)4N02Cl]Cl. 

The  analogous  compound  containing  ethylenediamine  has  been  thor- 
oughly studied  by  Werner12  and  his  findings  are  illustrated  by  means  of  a 
flow  sheet  (Fig.  8.9). 


[Co  en2  (N02)2]+ 

concentrated 
1  HNOs 
4 

[Co  en2  (N03)2]+ 

JH20 

[Coen2  (H20)2]+++- 

KOH 

->  [Co  en2  (H20)OH]++  - 

dilute  HNOj         r^              ,TT  --.n   1 1  1  i. 

>  [Co  en2  (H20)2]+++ 

1  NaN02  + 
1  HC2H3O2 

NaN02 
HC2H3O2 

*[Coen2(ONO)2]+ 

*[Co  en2  (ONO)2l+ 

J  stand 
I    (warm) 

stand 
(warm) 

[Co  en2  (N02)2]+ 

[Co  en2  (N02)2]+ 

Cis  -Series 

Trans  -Series 

Fig.  8.9 
*  Recently,  some  conflicting  reports  have  appeared  in  the  literature  with  regard 
to  the  actual  existence  of  nitrito  complexes  (page  268). 

Although  the  cis  isomer  can  sometimes  be  obtained  by  the  displacement 
of  a  bidentate  group,  the  procedures  employed  to  produce  the  trans  isomer 
are  almost  entirely  empirical.  There  is  some  reason  to  believe,  however, 
that  when  a  planar  tetracovalent  compound  changes  to  an  octahedral 
structure,  the  two  groups  added  occupy  trans  positions13.  This  procedure 

11.  Biltz  and  Biltz,  ibid.,  p.  178. 

12.  Werner,  Ber.,  44,  2445  (1911). 

13.  Werner,  "New  Ideas  on  Inorganic  Chemistry,"  translated  by  Hedley,  p.  261, 

London,  Longmans,  Green  and  Co.,  1911;  Jorgensen,  Z.  anorg.  Chem.,  25,  353 
(1900). 

14.  Basolo,  Bailar,  and  Tarr,  J.  Am.  Chem.  Soc,  72,  2433  (1950);  Heneghan  and 

Bailar,  J.  Am.  Chem.  Soc,  75,  1840  (1953) 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS 

CI 


28 1 


PtCI2-h  en 


-i  +  + 


en/  P*   /en 


+  CI. 


Fig.  8.10 


L 


Pt 


^ 


was  recently  applied11  in  the  preparation  of  /ratts-dichlorobis(ethylene- 
diamine)platinnm(IY)  chloride  (Fig.  8.10). 

Anionic  Complex  Compounds.  There  are  fewer  examples  of  cis-trans 
isomerism  in  anionic  complexes  and  these  have  not  been  studied  as 
extensively  as  fche  corresponding  cationic  compounds.  The  ion15 
[Co(XH3)2(X02)4]_  should  exist  in  cis  and  trans  forms,  but  only  one  isomer 
is  known  and  there  are  conflicting  reports  as  to  its  structure  (page  292). 

Delepine16  has  shown  that  potassium  hexachloroiridate(III),  K3[IrCl6], 
and  potassium  oxalate  react  to  form  potassium  a's-dichlorobis(oxalato)iri- 
date(III),  KsIIi^CoO^Clo].  The  cis  configuration  of  this  complex  was  es- 
tablished by  its  resolution,  using  strychnine.  Prolonged  boiling  of  a  solution 
of  the  potassium  salt  yielded  the  corresponding  trans  isomer.  The  complex, 
K[Ir  py2  (C204)2],  (and  its  rhodium(III)  analog17)  was  prepared  by  various 
methods  and  in  every  case  the  trans  salt  was  isolated. 

Ammonium  disulfitotetramminecobaltate(III),  NH4[Co(NH3)4(S03)2], 
was  first  prepared18  by  the  reaction  of  carbonatotetramminecobalt(III) 
chloride  and  ammonium  sulfite.  The  cis  configuration  was  assigned  to  this 
salt19  on  the  basis  of  the  fact  that  ethylenediamine  replaces  two  of  the  am- 
monia molecules  much  more  readily  than  the  other  two.  If  the  sulfite 
groups  are  trans  to  each  other,  the  four  ammonia  molecules  are  equivalent, 
and  all  of  them  would  be  replaced  by  ethylenediamine  with  equal  ease. 
However,  if  the  sulfite  groups  are  cis  to  each  other,  the  introduction  of 
ethylenediamine  may  follow  either  of  two  paths;  the  path  which  allows  the 
replacement  of  only  two  ammonia  molecules  would  be  expected  because, 
according  to  the  principle  of  trans  elimination,  the  two  ammonia  molecules 
which  are  trans  to  the  negative  sulfite  groups  should  be  labilized  (Fig.  8.11). 

15.  Erdmann, ./.  prakt.  Chem.,  97,  385    1S66);  Biltz  and  Biltz,  "Laboratory  Methods 

of  Inorganic  Chemistry,"  translated  by  Hall  and  Blanchard,  p.  150,  New 
York,  John  WUey  &  Sons,  Inc.,  1909. 

16.  Delepine,  Ann.  chim.,  19,  149  (1923). 

17.  Delepine.  Soc.  Espanola  Fi*.  y  Quim,  27,  485  (1929). 

18.  Hofmann  and  Jenny,  her.,  34.  01). 

19.  Klement,  Z.  ano  g.  aUgem.  Chem.,  150,  117  (1925);  Bailar  and  Peppard,  ./. 

62,105(1940). 


282 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

NH3 


NH 


Fig.  8.11 


Nonionic  Complex  Compounds.  Complex  compounds  in  which  the 
charge  on  the  central  atom  is  neutralized  by  the  coordinating  groups  are 
nonionic.  Compounds  of  this  type  are  usually  capable  of  existing  in  stereo- 
isomeric  modifications,  and,  in  some  cases  both  isomers  have  been  obtained. 
However,  satisfactory  proofs  of  their  structures  have  not  been  possible. 
Much  of  the  difficulty  encountered  results  from  the  fact  that  suitable 
solvents  are  not  known  for  some  of  these  substances. 

A  very  strong  argument  against  the  Blomstrand-Jorgensen  chain  theory 
and  in  favor  of  Werner's  coordination  theory  was  the  fact  that 
[Co(NH3)3Cl3]  did  not  give  a  silver  chloride  precipitate  readily.  Werner 
interpreted  this  to  mean  that  all  of  the  chlorine  was  held  firmly  by  the 
central  metal  atom.  The  analogous  nitro  compound20,  [Co(NH3)3(N02)3], 
is  believed  to  have  the  trans,  (1,2,6)  configuration.  Duval21  has  prepared 
[Co(NH3)3(N02)3]  by  five  different  methods  and  the  five  products  showed 
identical  absorption  spectra  and  similar  electrical  conductivities,  but  the 
x-ray  diagrams  of  some  of  the  powders  differed  slightly.  It  was  concluded 
that  this  evidence  was  insufficient  to  establish  the  existence  of  different 
geometric  structures  for  any  of  the  five  products.  On  the  other  hand, 
Sueda22  claims  to  have  prepared  cis,  (1 ,2,3)-[Co(NH3)3(N02)3]  by  starting 
with  as-[Co(NH3)3(H20)3]+++22- 23' 24. 

20.  Biltz  and  Biltz,  "Laboratory  Methods  of  Inorganic  Chemistry,"  translated  b}r 

Hall  and  Blanchard,  p.  182,  New  York,  John  Wiley  &  Sons,  Inc.,  1909. 

21.  Duval,  Compt.  rend.,  206,  1652  (1938). 

22.  Sueda,  Bull.  Chem.  Soc,  Japan,  13,  450  (1938). 

23.  Matsuno,  J.  Coll.  Sci.  Imp.  Univ.  Tokyo,  41,  10  (1921). 

24.  Sueda,  Bull.  Chem.  Soc,  Japan,  12,  188  (1937). 


STEREOISOMKh'/SM  OF  1IFX \<  < M  .1  LENT  ATOMS 


283 


The  nonelectrolyte  complexes  do  not  necessarily-  have  to  contain  equal 
numbers  of  neutral  groups  and  anions  [MajbJ,  but  may  also  be  of  the  type 
[Ma4b2].  This  particular  type  is  realized  with  hexacovalent  metals  having 
oxidation  states  of  two  or  four.  A  good  example  is  shown  by  cifl  and  trans 
isomers  of  |Pt  (  XUAjCli],  which  may  be  obtained  1>\  the  oxidation  of  the 
corresponding  dichlorodiammine  platinum (II)  compounds1'5;  this  also 
illustrates  that  the  two  groups  added  to  the  planar  tetracovalent  compound 
occupy  trans  positions  in  the  resulting  octahedron  (Fig.  8.12). 


NH 


NH- 


NH3 


Fig.  8.12 


Still  another  type  of  nonelectrolyte  complex  is  possible  if  the  neutral 
group  and  acid  radical  are  united  in  the  same  molecule,  as  is  the  case  in 
the  amino  acid,  glycine,  XH2CH2COOH.  These  are  termed  inner  complexes 
and  are  important  in  analytical  chemistry  and  mordant  dyeing.  Cobalt  (III) 
oxide  reacts  with  a  solution  of  glycine  to  form  a  mixture  of  two  com- 
pounds, both  of  which  have  the  composition  [Co(NH2CH2COO)3],  and 
which  can  be  separated  because  of  a  slight  difference  in  their  solubilities25. 
They  are  extremely  stable  and  may  be  dissolved  without  change  in  con- 
centrated sulfuric  acid;  their  aqueous  solutions  have  practically  no  elec- 
trical conductivity;  cryoscopic  measurements  show  that  they  are  undis- 
Bociated  in  solution.  They  are  believed  to  represent  geometrical  isomers  in 
which  all  of  the  same  groups  ( — NH2  or  — COO)  of  the  glycine  molecules 
occupy  adjacent  positions,  or  in  which  two  of  these  are  opposite  to  each 
other  (Fig.  8.13). 

2.5.  Ley  and  Winkler,  Ber.,  42,  3894  (1909). 


284  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

,CH2-NH2 


NH2 


CH2 


0  =  C 


CH2-NH2 

TRANS   OR    1,2,6 


Fig.  8.13 


The  absorption  spectra  suggest  that  the  more  soluble  form  is  the  trans 
isomer26.  Examination  of  the  diagrams  will  reveal  that  in  neither  case 
does  the  compound  possess  a  plane  of  symmetry,  so  there  exists  the  possi- 
bility of  mirror  image  isomerism  in  each  case,  bringing  the  total  number  of 
stereoisomers  to  four.  Since  this  compound  is  a  nonelectrolyte  it  does  not 
lend  itself  to  the  formation  of  salts  and  has  not  been  resolved.  Evidence  has 
been  obtained,  however,  for  the  existence  of  the  four  isomers  of  the  analo- 
gous complex  formed  between  d-alanine  and  cobalt(III)27. 

Complex  Compounds  Containing  Unsymmetrical  Bidentate 
Donor  Molecules.  The  same  type  of  isomerism  which  has  just  been  dis- 
cussed can  also  be  realized  when  only  one  or  two  unsymmetrical  molecules 
are  introduced  into  the  coordination  sphere  of  a  complex.  The  compound 
[Co(DMG)2  NH3C1]  has  been  resolved  (page  313)  by  Tsuchida,  Koboyaski, 
and  Nakamura28.  They  said  this  means  the  ammonia  and  chloro  groups 
occupied  cis  positions.  If  this  is  true,  the  two  molecules  of  dimethylgly- 
oxime  are  in  different  planes,  which  is  contrary  to  the  structure  of  analogous 
compounds  of  the  types  [Co(DMG)2A2]X  and  [Co(DMG)2X2]-29.  A  more 
recent  study  of  the  ultraviolet  absorption  spectrum  of  this  complex  indi- 
cates that  the  negative  portions  of  the  dimethylglyoxime  ions, 

26.  Kuroya  and  Tsuchida,  Bull.  Chem.  Soc,  Japan,  15,  429  (1940);  Basolo,  Ball- 

hausen,  and  Bjerrum,  Acta.  Chem.  Scand.,  9,  810  (1955). 

27.  Lifschitz,  Z.  physik.  Chem.,  114,  485  (1925). 

28.  Tsuchida,  Kobayoski  and  Nakamura,  Bull.  Chem.  Soc.,  Japan,  11,  38  (1936). 

29.  Nakatsuka,  Bull.  Chem.  Soc,  Japan,  11,  48  (1936) ;  Thilo  and  Heilborn,  Ber.,  64, 

1441  (1931). 


STEREOISOMERISM  OF  IIEXACOVALENT  ATOMS 


285 


CH3 

\ 
C 
1 

0 

=  N 

1 

c 

=  N 

/ 

-H3 

0 

H 

occupy  trans  positions  (page  295).  It  is  therefore  suggested  by  Tsuchida 
and  Koboyashi10  that  the  dimethylgloximes  may  be  in  the  same  plane  and 
the  optical  activity  of  the  compound  [Co(DMG)2NH3Cl]  is  caused  by  the 
unsymmetrical  oximes  (Fig.  8.14).  No  case  of  optical  isomerism  of  this  type 
has  been  definitely  established.  Furthermore,  there  is  reason  to  believe 
that  hydrogen  bonding  would  occur31  and  that  the  trans  complex  is  not 
optically  active  as  represented  in  Fig.  8.14(a  and  b)  but  is  instead  sym- 
metrical, as  shown  in  Fig.  8.14c. 


ci 


ci 


DMG 


CO 


DMG 


DMG 


CI 


0_ 


/ 


I 


Co 


DMG 


CHo-C-hW 


/      /Co 

r.H3-r.=NJ  4  N  =  r-r.H3 


NH3 

C<5J 


NH3 
Cb) 


Fig.  8.14 


L 


I 

N-C-CH- 

/ 


I 


NH3 
CO 


A  very  striking  example  of  isomerism  resulting  from  the  coordination  of 
an  unsymmetrical  molecule  has  been  clearly  demonstrated  with  the  com- 
pound dinitro(ethylenediamine)  (propylenediamine)  cobalt  (III)  bromide32. 
Since  propylenediamine,  NH2(CH3)CHCH2NH2 ,  is  not  symmetrical,  the 
methyl  group  (represented  in  Fig.  8.15  by  the  symbol  *)  can  be  placed  in 
the  cis  complex  ions  either  near  to  the  plane  of  the  two  nitro  groups,  or 
far  from  this  plane. 

30.  Tsuchida,  and  Kobayoski,  Bull.  Chem.  Soc,  Japan,  12,  83  (1937). 

31.  Rundle  and  Parasol,  J.  Chem.  Phys.,  20,  1489  (1952). 

32.  Werner  and  Smirnoff,  Helv.  chim.  Acta.,  1,  5  (1918). 


286  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

NOz 


Co 


N02 

TRANS 


en 


N02 


PL  -ISOMERS 
C/3) 


en^ 


N02 


Co 


'N02 


NO2' 


f> 


DL-ISOMERS 

(V; 
Fig.  8.15 


[d-pn]D 

[d-pn]D 

[rf-pnj 

(fi) 

[Z-pn]D 

(7)     J 

[Z-pn]D 

[*-pn  } 

[d-pn]L 
.[^-pn]L 

[d-pn]L 
il^-pn  ]L 

These  geometrical  isomers  will  be  distinguished  as  a,  (3,  and  7  compounds. 
In  addition  to  being  unsymmetrical,  propylenediamine  contains  an  asym- 
metric carbon  atom  and  may  exist  in  both  the  dextro  and  levo  modifi- 
cations; therefore,  the  total  number  of  isomers  possible  is  twice  that  shown 
in  Fig.  8.15. 


(«) 


All  of  the  predicted  isomers  were  isolated. 

Complex  Compounds  Containing  Polydentate  Donor  Molecules. 

The  most  extensively  studied  chelate  groups  attached  to  a  central  atom 
are  bidentate,  but  compounds  are  known  which  can  fill  three  (tridentate), 
four  (tetradentate),  five  (pentadentate)  or  six  (hexadentate)  positions  in 
the  coordination  shell.  The  presence  of  six  functional  groups  in  the  ethyl- 
enediaminetetraacetic  acid  (EDTA)  molecule  first  provided  the  possibility 
of  forming  compounds  in  which  a  substance  acts  as  a  hexadentate  chelating 
agent.  The  salts  of  the  complex  ions  formed  by  this  substance  are  usually 


STEREOISOMERISM  OF  HEXACOY ALEXT  ATOMS 


287 


hydrated;  however,  Brintzinger,  Thiele  and  Mtiller88  prepared  anhydrous 

Xa[Co(EDTA)]  by  drying  the  -4-hydrate  at  150°.  Schwarzenbach84  pre- 
pared the  anhydrous  salt  [Co  en2  Cl2][Co(EDTA)].  Complex  ions  containing 
pentadentate  ethylenediaminetetracetic  arid  have  also  been  prepared. 
Schwarzenbach  reports  several  salts  of  the  ions  [Co(IIY)Br]~  and 
[Co(HY)N02]-  (Y  represents  the  EDTA4-  ion).  The  pK  of  the  free  car- 
boxyl  group  is  approximately  3  in  both  cases.  The  infrared  studies  of  Busch 
and  Bailar  con  (inn  the  hypothesis  that  EDTA  may  behave  as  either  a 
pentadentate  or  hexadentate  donor35.  The  hexadentate  Co(III)  complex 
has  been  resolved  into  optical  isomers35, 36.  Recently,  Dwyer  and  Lions37 
have  reported  a  cobalt(III)  complex  cation  containing  a  new  hexadentate 
chelate  (Fig.  8.16);  they  report37b  the  extremely  high  molecular  rotation  for 
this  compound  of  over  50,000°  at  the  mercury  green  line  (5461  A.).  Models 


d_[-  1,8- BIS  CSALICYLIDENEAMINO)- 
3,6-  DITHIAOCTANECOBALT  (III) 

Fig.  8.16 


show  that  this  compound  can  exist  in  only  one  strainless  geometrical  form 
in  which  the  nitrogen  atoms  are  in  trans  positions  and  the  sulfur  atoms  and 
oxygen  atoms  are  in  cis  positions  to  each  other.  The  resulting  compound 
is  asymmetric  and  the  two  enantiomorphs  of  the  cobalt  (III)  complex  were 
isolated.  These  investigators38  have  successfully  extended  the  group  of 
hexadentate  compounds  to  several  analogs  of  1,8  bis-(salicylideneamino) 
3,6  dithiaoctane  cobalt(III).  Dwyer  and  his  co-workers3713, 39  have  continued 

33.  Brintzinger,  Thiele,  ami  Muller,  Z.  anorg.  allgem.  Chem.,  251.  285  (1943). 

34    Schwarzenbach,  //</>■.  chim.  Acta,  32,  K.V.)  (1949). 

35.  Busch  and  Bailar,  ./.  .1///.  Chem.  Sac,  75,  1574  (1953). 

arfas,  and  Mellor,  ./.  Phys.  Chan.,  59,  296    L955). 

37.  Dwyer  and  Lions,  •/.  .1///.  Chem.  Soc.,  69,  2917  <  1917);  72,  1645    I960 

38.  Dwyer  and  Gyarfas,  •/.  Proc.  Roy.  Soc.  A.  8.  Wales,  83,  170    1949). 

Dwyer,  Lions  and  Mellor,  •/.  Am.  Chem.  Soc.,  72,  .r)0:57  (1950).  Dwyer,  Gill, 
Gyaifasand  Lions,  ibid. ,74,  U88  (1952).  Collins,  Dwyer,  and  Lions,  //>/»/..  74, 
3134  1952).  Dwyer,  Gill,  Gyarfas  and  Lions,  J.Am.  Chem.  So,..  75,  1526,  2443 
(1953). 


288 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


their  investigations  of  these  hexadentate  chelate  compounds  utilizing 
different  central  metal  atoms  and  different  ligands.  Other  hexadentate 
chelates  were  prepared  in  which  one  and  then  both  of  the  sulfur  atoms  in 
some  of  the  above  ligands  were  replaced  by  oxygen  atoms39b.  The  authors 
also  reported  a  resolution  of  the  cobalt(III)  complex  containing  the  hexa- 
dentate chelate  in  which  one  sulfur  was  replaced  by  an  oxygen  atom.  Mag- 
netic studies39a  supported  their  conclusions  that  the  central  atom  is  octa- 
hedral in  configuration  and  that  the  bonds  are  of  the  hybridized  d2sp3  type. 
Tridentate  groups,  such  as  tripyridyl41b  and  a ,  /3 , 7-triamino  propane410, 
form  very  stable  compounds  with  hexacovalent  metals  of  the  types 
[M(tripy)2]  and  [M{NH2CH2CH(NH2)CH2NH2l2],  respectively.  It  is  be- 
ieved  that  in  some  of  these  compounds  the  coordinated  group  is  attached 
in  the  1,2,6  positions  along  an  edge  of  the  octahedron  and  not  solely  in 
the  1,2,3  positions  bounding  an  octahedral  face.  That  this  is  probably 
correct  is  indicated  by  the  ease  with  which  these  tridentate  groups  fill  three 
coordination  positions  in  the  planar  tetracovalent  complex,  [Pt  tripy  CI]  CI. 
This  cannot  be  taken  as  conclusive  evidence  and  certainly  it  is  possible 
for  some  tridentate  groups  to  be  attached  on  an  octahedral  face  as  wTell 
as  along  the  central  plane  of  an  octahedron.  Models  show  that  complexes 
in  which  triaminopropane  is  tridentate  should  have  only  the  1,2,3  con- 
figuration. Diethylenetriamine,  NH2CH2CH2NHCH2CH2NH2 ,  is  also 
knowTn  to  behave  as  a  tridentate  donor  molecule40  and  should  be  capable 
of  forming  three  geometrical  isomers  with  a  hexacovalent  atom  (Fig.  8.17). 


The  two  1,2,3  isomers,  (B)  and  (C)  would  form  optical  isomers.  Only 
one  isomer  of  this  type  has  been  isolated  and  its  configuration  has  not  been 
definitely  established. 

Tetradentate  donors  are  known  to  be  possible  and,  recently,  numerous 


!(>.   Mann,  J.  Chew.  Soc,  1934,  466;  1930,  1745. 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS 


289 


coordination  compounds  of  this  type  have  been  prepared41.  Mann  used 

/3,0'vJ''-tnaminotiictliYlainine  and  obtained  CW-[Co  trill  (NCS)j]NCS. 
Because  of  the  structure  of  this  amine,  the  corresponding  trans  Bait  does 
not    exist.    Morgan   and   BurstaU   investigated   2,2/,2*,2'"-tetrapyridy] 

and  reported  it  to  yield  tran8-[Co  tetrpy  C1JC1.  The  salt  had  the  charac- 
teristic  green  color  of  ^n«^cUorotetrajninecobalt(III)  cations  (p.  294). 
Basolo41"  has  isolated  coordination  compounds  of  cobalt(III)  with  tri- 
ethylenetetramine,  NHiCHrf)HjraCH,CH2NHCH2CH2NH2 ,  behaving  as 
a  tetradentate  group.  The  complex  [Co  trien  C12]C1  was  isolated;  theo- 
retically, it  can  exist  in  three  geometrical  forms  (Fig.  8.18):  one  isomer  in 
which  the  chloro  groups  occupy  trans  positions,  and  two  isomers,  both 
optically  active,  with  the  chloro  groups  adjacent  to  each  other.  Only  one 
isomer  was  obtained  and  the  cis  configuration  of  this  salt  was  established. 


ci 

TRANS 
CSYMMETRICAL) 


N 
CIS 

COPTICALLY    ACTIVE) 

Fig.  8.18 


ct 

CIS 
COPTICALLY    ACTIVE) 


Since  cis-trans  rearrangements  are  known  to  occur  readily  in  cobalt  com- 
plexes, it  may  be  that  such  a  change  in  configuration  always  resulted  in 
favor  of  a  more  stable  cis  modification.  However,  the  fact  that  geometrical 
isomers  are  possible  for  coordination  compounds  containing  certain  poly- 
dentate  groups  has  been  demonstrated3911. 

Poly  nuclear  Complex  Compounds.  Numerous  polynuclear  complexes 
of  hexacovalent  elements  have  been  isolated  and  properly  identified.  The 
majority  of  these  compounds  are  binuclear  and  result  from  the  fact  that 
Borne  groups  are  capable  of  donating  two  pairs  of  electrons  and,  in  so  doing, 
can  form  a  bridge  between  two  metal  atoms.  A  consideration  of  the  octa- 
hedral structure  reveals  that  this  bridge  can  be  formed  in  three  different 
ways:  1 1)  one  donor  group  joining  two  corners  of  the  octahedron,  (2)  two 
donor  groups  occupying  one  edge  of  each  octahedron  or  (3)  three  donor 
group-  occupying  one  face  of  each  octahedron  (Fig.  8.19). 

41.  Bailes  and  Calvin.  ./.  .1       Cfo      .  Soc.,  69,  1886  (1947);  Morgan  and  Buret  all,  ./. 
.  Soc,  1934,  1498;  Pope   and  .Mann,  ibid.,  1926,  2675,  2681;  ibi<L,  1927, 
1224;   Basolo,  ./.  Am.  ('hem.  Soc,  70,  2634  (1948);  Morgan  and  Buret  all,  ./. 
-'  .,  1938,  1672;  Mann,  J  Soc.,  1929,  409. 


290 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

x 


[a6M— X— Ma5] 
(CORNK10 


|      X             |        i 

A                    A 

X            "1 

/  \ 

a4M              Ma4 

\    / 

X 

(EDGE) 

Fig.  8.19 

X 

/  \ 

d3M— X— Ma3 

\    / 
X 

(FACE) 


The  number  of  possible  geometrical  isomers  of  these  polynuclear  hexa- 
covalent  complex  compounds  is  extremely  large.  Even  the  very  simplest 

X 

/    \ 
compounds  of  the  types  [ba4M — X — Ma4b]  and  [ba3M  Ma3b]  mav 

\    / 
X 

exist  in  three  and  five  different  geometric  forms,  respectively  (Fig.  8.20). 

One  of  the  latter  (E  and  F)  is  optically  active. 


f 


A 

CU') 


B 

CI, 2') 


J: 

*  1 

; 

/;/ 

d 
b 

*         1 

L      D       a. 
X    |a 

7 

/'/ 

,h  a, 


c 

(2,2') 

b  a 

X 


t% — r 

a.      E       a 


fQ.  Oi< 


t* — r 

a      F        a 
0,3') 


,a  a, 


a  b 


b  b 


G 
0,6') 


H 
C2,20 


I 
(2,4') 


Fig.  8.20 


STEREOISOMERISM  OF  II  EX  ACOV  A  LENT  ATOMS 


291 


The  rather  involved  stereochemistry  of  the  polynuclear  cobalt(III)  and 
chromium (II I)  ammines  was  investigated  extensively  by  Werner42.  His 
study  was  undertaken  for  the  purpose  of  preparing  mononuclear  compounds 
of  known  structure  and  to  "establish"  the  configuration  of  mononuclear 
complexes. 

Determination  of  Configuration 

Chemical  Methods.  Bidentate  Group.  The  mosl  commonly  used  method 
of  determining  configurations  depends  on  the  fact  thai  bifunctional  groups 
ran  span  only  coordination  positions  which  are  adjacent  to  each  other. 
Hence,  provided  that  no  rearrangement  of  configuration  occurs  during 
the  reaction,  the  isomer  which  is  capable  of  combining  with  one  mole 
of  a  chelate  group,  or  which  is  formed  whenever  such  a  group  is  displaced, 
must  belong  to  the  cis  series.  The  application  of  this  type  of  reasoning  to 
the  geometrical  isomers  of  [Co(XH3)4Cl2]Cl  is  summarized  in  Fig.  8.21.  The 

dilute  HC1 


[Co(NH3)4CO: 


[Co(NH3)4(H20)Cir 


dilute 
H2SO4 


[Co(NH3)4(H20)2]+++ 

I   NHj  (aqueous) 

[Co(NH3)4(H20)OH]++ 
i  100° 
H 
0 

/    \ 
(NH3)4Co  Co(NH3)4 

\    / 
O 
H 


concentrated 

HC1 
(-12°) 


[Co(NH3)4Cl2]+ 

Cis  (purple) 

Fig.  8.21 


concentrated 
H2S04  +  HC1 


[Co(NH3)4Cl2]+ 
Trans  (green) 


determination  of  configuration  involves  the  reaction  of  the  binuclear  com- 
plex. 

II 

O 


(XH3)4Co  Co(NH3)4 

\    / 
O 
H 


(S04): 


42.  Werner,  Ann.,  375,  1  (1910). 


292 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


with  concentrated  hydrochloric  acid  to  give  one  mole  of  the  dichloro  com- 
plex and  one  of  the  diaquo  complex.  Assuming  that  no  rearrangement  takes 
place,  the  chloro  groups  must  occupy  adjacent  positions  and  the  salt  must 
be  cis-[Co(NH3)4Cl2]Cl.  It  is  important  to  observe,  however,  that,  in  this 
same  series  of  reactions,  a  similar  displacement  of  a  bidentate  group  (car- 
bonato)  with  hydrochloric  acid,  leads  to  a  change  of  configuration. 

Another  example  of  this  type  is  the  "proof "  of  structure  of  NH4[Co(NH3)2- 
(NCyj,  which  has  been  obtained  in  only  one  form.  Whenever  the  complex 
reacts  with  oxalic  acid,  two  nitro  groups  are  replaced  by  one  oxalate  group. 
If  the  original  complex  has  the  trans  configuration,  only  one  oxalate  com- 
plex is  to  be  expected,  but  if  the  ammonia  groups  are  adjacent  to  one  an- 
other, two  oxalat  derivatives  may  result  and  one  of  them  should  be  enan- 
tiomorphous  (Fig.  8.22). 


NO^ 


NOo 


N0< 


NH- 


H3 


,N02 


N02/ 


CO 


HgC^Qa 


c=o 


NH3 
TRANS 

NH3 


CO 


'NO?         N02 


NO- 


CIS 


CO 


NO2' 


o —  c=o 


NH: 


NH3 
NH3      N0o_l 


(SYMMETRICAL) 

NH3 
NH3  N92-_ 


NH- 


Co 


I  C=0 

(OPTICALLY    ACTIVE) 

Fig.  8.22 


OrC 


'NO2 


(SYMMETRICAL) 


Two  products  were  isolated  from  the  reaction  between  Erdman's  salt 
and  oxalic  acid;  one  of  these  was  resolved  into  optical  antimers43.  Although 
there  are  many  instances  in  which  structures  determined  by  this  method 
have  been  proven  to  be  correct,  one  cannot  disregard  the  fact  that  complex 
cobalt  compounds  are  known  to  rearrange  very  readily,  and,  therefore,  the 
assumption  that  a  molecule  retains  its  configuration  as  groups  or  atoms  are 
replaced  is  not  entirely  reliable.  This  particular  case  may  serve  as  a  good 
illustration  of  this  factor  since  the  results  obtained  by  the  oxalate  method 

43.  Shibata  and  Maruki,  J.  Coll.  Sci.  Imp.  Univ.,  Tokyo,  41,  2  (1917);  Thomas,  J. 
Chem.  Soc,  121,  2069  (1922) ;  Thomas,  ibid.,  123,  617  (1923). 


STEREOISOMERISM  OF  HEX  [COVALENT  ATOMS 


293 


do  Dot  agree  with  the  findings  oi  Riesenfeld  and  Klement44,  nor  with  x-ray 
studies  which  won4  made  on  the  silver  salt48. 

Optical  Activity,  In  certain  cases  it  is  possible  to  establish  the  configura- 
tions of  these  isomers  by  showing  thai  one  is  optically  active  and  the  other 
is  inactive.  This  procedure  offers  conclusive  proof  except  in  examples 
where  only  one  form  is  known  and  this  cannot  he  resolved;  failure  to  re- 
solve the  compound  docs  not  necessarily  mean  that  the  complex  is  sym- 
metrical. A  familiar  example  of  this  method  is  the  proof  that  the  purple 
salt,  [Co  euj  C1JC1,  which  is  optically  active,  has  a  cis  configuration;  the 
green  inactive  isomer  must  therefore  have  the  trans  configuration. 

Bailar  and  Peppardwb  used  this  method  to  determine  the  structures  of 
the  three  stereoisomeric  forms  of  dichlorodiammine(ethylenediamine)co- 
ball  (III)  ion.  (I,  III,  and  VI,  Fig.  8.23).  Salts  of  two  of  these  were  prepared 
by  Chaussy4-  who  designated  them  as  cis  and  trans  (referring  to  the  relative 
positions  of  the  chloro  groups).  Chaussy  made  no  mention  of  the  fact  that 
two  cis  ions  are  possible.  The  colors  of  these  ions  enable  one  to  determine 
the  relative  positions  of  the  chloride  groups  with  certainty,  but  do  not  dis- 
tinguish between  the  two  a's-dichloro  configurations.  The  assignment  of 
configurations,  in  this  case,  was  based  upon  the  fact  that  the  m-dichloro- 
cis-diammine  ion  (III)  is  asymmetric  while  the  as-dichloro-^ra/zs-diammine 
ion  (VI)  is  not. 

The  methods  employed  to  prepare  the  two  cis  isomers  are  of  interest. 
(Fig.  8.23).  The  preparation  of  the  a's-dichloro-cfs-diammine  salt  (III)  is 


NH3 


NH- 


44.  Riesenfield  and  Klement,  Z  anorg.  allgem.  Chem.,  124,  1  (1022). 
t:>.  Welle,  Kristallogr.,  Z.,  95A,  74  (1936). 
46.  Chaussey,  "Dissertation,"  Zurich,  1909. 


294  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

based  upon  the  fact  that  chelate  groups  can  span  only  adjacent  positions 
and,  therefore,  the  dichloro  salt  (I)  undergoes  a  rearrangement  to  produce 
the  carbonato  compound  (II).  The  preparation  of  the  m-disulfito-£rans- 
diammine  salt  (V)  is  a  good  illustration  of  a  phenomenon  known  as  the 
trans  effect  which  has  been  studied  in  some  detail  by  Chernyaev47.  Bailar 
and  Peppard19b  have  also  found  this  principle  of  trans  elimination  to  be 
useful  in  the  synthesis  of  the  as-dichloro-^rans-diammine  salt  (VI).  The 
cis-disulfitotetrammine  salt  (IV)  was  used  so  that  the  NH3  groups  trans  to 
the  sulfite  groups  would  be  labilized  and  the  ethylenediamine  would  enter 
in  the  2 , 3  positions,  to  yield  (V) . 

Chemical  Behavior.  The  possibility  of  distinguishing  between  geometric 
isomers  by  means  of  their  reactions  has  been  considered.  It  is  known,  for 
instance,  that  cis-  and  £rans-dinitrotetrammine,  and  cis-  and  £rans-dinitro- 
bis (ethylenediamine)  compounds  react  differently  toward  boiling  hydro- 
chloric acid48.  The  cis  isomer  is  dissolved  and,  upon  standing,  a  green  crys- 
talline salt  separates  from  the  purple  solution;  the  trans  isomer  forms  a 
red  precipitate  of  the  ^rans-nitrochloro  complex.  Although  this  qualitative 
test  can  be  conveniently  used  for  these  particular  dinitro  complexes,  it  does 
not  necessarily  apply  to  all  analogous  compounds.  A  typical  discrepancy  is 
found  in  the  work  of  Hurlimann49,  who  was  of  the  opinion  that  the  product, 
[Co  (Z-pn)2  (N02)2]  Br,  obtained  from  the  reaction  of  trinitrotriammine- 
cobalt(III)  and  Zezw-propylenediamine  was  the  pure  cis  isomer,  since  no 
red  precipitate  formed  when  the  complex  was  heated  with  concentrated 
hydrochloric  acid.  However,  it  has  been  shown  by  rotatory  dispersion  curves 
that  the  salt  obtained  was  a  mixture  of  the  cis  and  trans  isomers50,  and, 
furthermore,  that  trans-[Co  (Z-pn)2  (N02)2]+  does  not  give  a  red  precipitate 
when  boiled  with  concentrated  hydrochloric  acid. 

•J  Physical  Methods.  Absorption  Spectra.  In  some  cases  the  dissimilar 
spatial  arrangements  of  the  same  ligands  about  a  central  atom  results  in  a 
very  noticeable  difference  in  color.  This  difference  is  particularly  obvious 
with  the  praseo  (green)  and  violeo  (blue-violet)  series  of  isomers,  character- 
istic of  trans-  and  czs-dichlorotetrammine  compounds  of  cobalt  (III)  and 
chromium  (III).  Since  there  are  no  known  exceptions  to  this  difference  in 
color,  it  is  generally  accepted  as  conclusive  proof  of  structure  for  this  par- 
ticular type  of  compound.  Unfortunately,  dissimilarity  in  structure  is  not 
always  accompanied  by  such  a  vast  color  difference,  as  is  shown  by  the 
fact  that  the  corresponding  dinitro  complexes  differ  only  slightly  in  ap- 
pearance. 

47.  Chernyaev,  Ann.  inst.  platine,  4,  243  (1936). 

is.  Jorgensen,  Z.emorfl.  Chem.,  17, 468, 472  (1898);  Klement,Z.  anorg.  allgem.  Chem., 
150,  117  (1925). 

49.  Hurlimann,  "Dissertation,"  Zurich,  1918. 

50.  O'Brien,  McReynolds,  and  Bailar,  ./.  Am.  Chem.  Soc,  70,  749  (1948). 


STEREOISOMERISM  OF  HEXACOVALENT    {TOMS  •_,,.»:> 

In  this  same  connection  the  absorption  spectra  of  coordination  com- 
pounds have  been  thoroughly  studied  by  numerous  Investigators.  Shibata 

and  Urbain61  worked  with  cobalt  complexes  and  noticed  thai  there  were 
always  two  hands  of  maximum  absorption,  one  of  which  occurs  in  the  visi- 
ble while  the  other  is  found  in  the  near  ultraviolet.  It  was  also  observed 
that  when  two  nitro  groups  are  substituted  for  ammonia  in  the  trans  posi- 
tions, a  third  absorption  hand  occurs  in  the  short  ultraviolet62.  Shibata 
made  the  following  generalizations  from  his  studies: 

(1)  Complexes  of  analogous  constitution  absorb  similarly; 

(2)  Ligands  of  analogous  chemical  structure  absorb  similarly; 
3     ( Optical  isomers  absorb  similarly; 

(4)  Geometric  isomers  in  general  absorb  differently; 

(5)  Sign  and  magnitude  of  charge  on  the  complex  ion  do  not  affect  the 
absorption; 

(6)  The  anion  has  no  appreciable  effect. 

Generalization  (4)  is  of  interest  in  our  discussion,  because  it  may  offer  a 
possible  method  for  distinguishing  among  stereoisomers. 

Tsuchida63  formulated  some  relatively  simple  theories  to  explain  many 
of  the  complexities  of  the  spectra.  He  proposed  that  the  first  absorption 
band  (visible  zone)  is  due  to  electronic  transitions  within  the  inner  electron 
rings  of  the  transition  element  which  is  the  nucleus  of  the  complex.  He 
attributed  the  second  band  to  the  electrons  linking  the  ligands  with  the 
central  ion,  and  the  third  band  (short  ultraviolet  region)  to  a  special  type 
of  linking  of  ligands,  e.g.,  two  negative  groups  in  trans  positions.  Kuroya 
and  Tsuchida26  obtained  the  absorption  spectra  of  several  carefully  chosen 
complex  cobalt  compounds  to  show  that  the  third  absorption  band  is 
present  in  compounds  which  contain  at  least  two  negative  ligands  in  trans 
positions,  but  is  absent  if  the  negative  ligands  are  adjacent  to  each  other 
Table  8.1). 

They  say  that  the  appearance  of  the  third  band  is  independent  of  (1)  the 
nature  and  valency  of  the  central  ion,  (2)  the  ligand  in  question,  provided 
that  it  is  of  negative  character,  (3)  the  charge  of  the  complex  radical,  and 
(4)  the  configuration,  so  long  as  the  trans-pairing  condition  is  fulfilled. 
Some  question  has  recently54,  55  been  raised  as  to  whether  the  presence  or 
absence  of  this  third  absorption  band  for  a  complex  with  two  or  more  nega- 
tive ligands  can  be  taken  as  absolute  proof  of  geometric  structure.  However 
it  does  appeal-  that  in  general  the  absorption  bandfLin  the  ultraviolet  region 

51.  Shibata  and  [Jrbain,  Compt.  rend.,  157,  W.\  (1914). 

52.  Shibata.  ./.  Coll.  Sri.  Imp.  Univ.,  Tokyo,  37,   1    (1915). 

Tsuchida.  Bull.  Chem.  Soc.,  Japan,  11,  785    1936);  Tsuchida,  ibid.,  13,  388,  136, 
471   (1938). 

54.  Basolo,  •/.  Am.  Chem.  Soc.,  72,  1393  (1950  . 

55.  Shimura,  J.Am.  Chem.  Soc.,  73,  ">07'.J  (1051). 


296 


CH i:\fISTRY  OF  THE  COORDINATION  COMPOUNDS 


Table  8.1.  Absorption  Spectra  of  Some  Geometrical  Isomers 

First  Band 

Second  Band 

Third  Band 

Complex  Salt 

A 

log  € 

A 

log  € 

A 

log  « 

«s-[Co(NH3)4(N02)2]Cl 

4580 

1.99 

3250 

3.10 

//7//^-[(\>(\ir3)4(N02)2]Ci 

4490 

2.32 

3450 

3.54 

2500 

4.08 

cis-[Co  en2  (N02)2]N03 

4350 

2.10 

3250 

3.68 

trana-[Co  en2  (NOs)2]NOj 

4300 

2.20 

3380 

3.44 

2490 

4.37 

//•«//s-[Co(NH3)4CLN02]Cl 

4750 

1.87 

3380 

3.13 

2440 

4.07 

trans-[Co  en2  C1N02]C1 

4350 

2.00 

3340 

3.37 

2410 

4.35 

cis-[Co  en2  C1(NCS)]C1 

5030 

2.18 

3570 

2.75 

trans-[Co  en«  Cl(NCS)]Br 

5550 

2.10 

3460 

2.93 

2720 

3.43 

ocmir  at  the  shorter  wave  length  for  the  cisjsomer  than  for  the  analogous 
trans  compound. 

A  somewhat  different  observation  has  been  reported  by  Sueda22, 24,  who 
studied  the  characteristic  second  absorption  band  of  several  nitroammine- 
cobalt(III)  complexes  and  concluded  that  this  band  can  be  accounted  for 
by  an  additive  effect  of  groups  in  trans  positions.  The  absorption  (Fig. 
8.24)  of  cis-[Co(NH3)4(N02)2]Cl  is  assumed  as  a  sum  of  three  characteristic 


4000 


Fig.  8.24.  Absorption  spectra  of  some  cobalt  complexes. 

A.  [Co(NH3)5N02]Cl2 

B.  as-lCo(NH3)4(N02)2]Cl 

C.  <mns-[Co(NH3)4(N02)2]Cl 

D.  [Co(NH3)3(N02)3] 

E.  K[Co(NH3)2(N02)4] 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS  '_,!>7 

absorptions,  i.e.,  |  \II3— Co—  XH3)*  and  2(NH|  Co  N0a).  The  absorp- 
tion of  [Co(NH«)5NOj]C1j  canalso  be  resolved  into  2(NH*  Co  Ml)  and 
(NHy— Co-  NOj).  Since  the  absorption  of  (XII :i — Co— MI,)  can  be  ne- 
glectedf  in  comparison  with  that  of  (NH3 — Co — N02),  the  absorpt  ion  given 
by  both  Baits  shows  that,  due  to  the  Dumber  of  (NHa — Co — NO2)  groups 
contained,  the  former,  the  cis  compound,  has  double  the  absorption  in- 
tensity of  the  latter  pentammine  complex  showing  similar  curves.  With 
regard  to  the  //v///n-[(,o(NII;;)i(\()2)2]C1,  its  absorption  may  be  considered 
to  be  the  sum  produced  by  2(NH3— Co— NH3)  and  (N02— Co— N02)  and 
it  is  almost  the  same  as  that  of  (X02 — Co — N02),  since  the  absorption  of 
(XH3— Co— XH3)  is  relatively  small.  The  absorption  of  [Co(XH3)3(X02)3] 
can  be  resolved  into  (XH3— Co— XH3),  (XH3— Co— X02)  and  (X02— Co- 
XOj)  and,  as  is  expected,  the  absorption  is  represented  as  a  sum  of  those 
given  by  [Col  Ml  \'( ),]C12  and  <rans-[Co(XH3)4(X02)2]Cl.  The  absorption 
intensity  due  to  the  complex  K[Co(XH3)2(X02)4]  is  nearly  twice  that  of 
the  ^•o«s-[Co(X^H3)4(X'02)2]Cl  and  it  is  therefore  assumed  that  the  complex 
has  a  trans  configuration  and  that  its  absorption  results  from 
2  XOo— Co— X02). 

Sueda  has  applied  his  reasoning  to  a  study  of  the  structures  of  several 
aquochloroammines  of  cobalt  (III)  and  chromium(III)24,  and  also  in  estab- 
lishing the  cis  configuration  of  [Co(XH3)3(X02)3]  which  he  prepared  from 
as-[Co(XH3)3(H20)3]+++  ". 

Recent  application  of  the  crystal  field  theory  to  complex  compounds56 
permits  a  better  interpretation  of  the  absorption  spectra  of  these  com- 
pounds. This  theoretical  treatment  predicts  differences  in  the  absorption 
spectra  of  cis  and  trans  isomers  of  hexacoordinated  complexes57,  in  good 
accord  with  experimental  observations58, 26b.  However,  one  immediate  limi- 
tation is  that  for  complexes  containing  ligands  of  approximately  the  same 
crystal  field  strength  the  differences  predicted  may  be  too  small  to  observe 
experimentally. 

X-ray  Diffraction.  The  final  result  of  a  complete  x-ray  analysis  of  a  sub- 
stance is  the  determination  of  the  relative  positions  of  all  the  constituent 
atoms.  As  a  rule  this  becomes  increasingly  difficult  as  the  number  of  pa- 

*  This  represents  the  characteristic  absorption  assumed  to  be  produced  by  two 
ammonia  molecules  in  trans  positions  having  cobalt(III)  as  the  central  ion. 

t  It  Lb  convenient  to  say  that  the  absorption  capacity  due  to  the  (XH3 — Co— NH 
is  weak  compared  to  that  of  (XH3 — Co — XO2),  since  the  extinction  coefficient  of  the 
maximum  absorption  given  by  [Co(XH3)6]Cl3  is  only  ahout  40  (at  336  A),  while  that 
given  by  [Co Ml    iNOs]Cls  ,  the  weakest  absorbent  containing  the  group,  (XH3 — 
Co— X02),  is  aboul  1260  (a1  325  A). 
56.  Orgel,  J.  Chem.  Soc.,   1756  (1952). 
.">7.  Ballhausen  and  J0rgensen,  Kgl.  Danske  Videnskab.  Belskab,  Mat.  fye.  Medd., 

29,  Xo.  14  (1955). 
58.  Linhard  and  Weigel,  Z.  anorg.  Chem.,  271,  101  (1952). 


298  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

rameters  required  to  fix  these  positions  increases,  and  relatively  few  com- 
plete structure  determinations  of  hexacovalent  complex  compounds  have 
been  made.  Theoretically,  however,  it  should  be  possible  to  establish  the 
configuration  of  a  stereoisomer  by  a  careful  x-ray  study  of  the  crystalline 
compound. 

A  large  number  of  geometric  isomers  of  the  type  [Ma4YCl]X,  where  M  is 
cobalt(III)  or  chromium(III),  have  been  investigated  by  means  of  x-rays59. 
It  was  shown  that  if  Y  is  a  chloro  or  bromo  group,  the  spectra  for  the  cis 
and  trans  forms  are  different,  but  if  Y  is  a  group  coordinated  through  nitro- 
gen (NH3 ,  N02~  or  NCS-),  the  spectra  are  the  same.  The  method  was  em- 
ployed to  show  that  the  isomers  were  different,  but  not  to  establish  which 
was  cis  and  which  trans. 

A  complete  x-ray  analysis  of  the  crystal  structure  of  Ag[Co(NH3)2(N02)4] 
indicates  that  the  ammonia  groups  are  in  trans  positions45.  The  crystals  are 
tetragonal,  a  =  6.97,  c  =  10.43  A,  and  the  space  group  is  P4/nnc-(D4h). 
There  are  two  molecules  in  the  unit  cell.  This  result  differs  from  that  ob- 
tained from  chemical  evidence,  which  assigns  the  cis  configuration  to  the 
complex  ion43,  but  agrees  with  the  results  of  Sueda.22 

Rotatory  Dispersion.  The  fact  that  trans  complexes  are  not  ordinarily 
resolvable  while  those  of  the  cis  configuration  are,  is  commonly  used  to 
distinguish  between  geometrical  isomers  of  the  type  [Co(AA)2a2].  If,  how- 
ever, the  coordinating  groups  are  optically  active,  both  isomers  of  the  com- 
plex will  rotate  the  plane  of  polarized  light,  so  that  the  presence  of  optical 
activity  does  not  serve  to  distinguish  one  isomer  from  the  other.  O'Brien, 
McReynolds  and  Bailar50  have  shown  that  the  configurations  of  such  com- 
pounds can  be  conveniently  determined  by  means  of  rotatory  dispersion 
curves.  The  success  of  this  method  depends  upon  the  fact  that  complex 
compounds  containing  optically  active  donor  molecules  normally  exist  only 
in  certain  preferred  configurations  (page  313).  The  optical  activity  of  these 
compounds  is  due  largely  to  the  configurational  asymmetry  of  the  complex 
as  a  whole,  so  the  rotatory  dispersion  curves  of  complexes  having  similar 
configuration  should  exhibit  the  same  characteristics,  whether  a  certain 
type  of  ligand  is  optically  active  or  not.  Thus,  the  rotatory  dispersion 
curves  of  cis-[Co  en2  Cl2]+  and  cis-[Co  (7-pn)2  Cl2]+  should  be  quite  similar. 
It  is  also  assumed  that  in  a  compound  of  the  type  [M(AA)2X2]+  if  the  non- 
basic  constituents  (X)  are  in  trans  positions  there  can  be  no  optical  activity 
attributable  to  the  asymmetry  of  the  complex,  and  therefore  the  rotatory 
dispersion  characteristics  should  be  similar  to  those  of  the  optically  active 
base  (AA).  If,  on  the  other  hand,  the  complex  has  a  cis  configuration,  there 
should  be  an  induced  activity  and  the  rotatory  dispersion  of  the  complex 
should  resemble  that  of  a  similar  optically  active  ion  and  not  that  of  the 

59.  Stelling,  Z.  physik.  Chem.,  B33,  338  (1933). 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS  299 

Table  8.2.  Examples  oi   da  trims  Conversions 

Starting  Material  Reagent  Product 

[Co(NB       NOa)3]  propylenediamine  cm  and  trana-[Co  pna  (NOOsJNOj 

r/.s--[(\)  pn,  CljJCl  KCNS  trana-[Co  pna  (NCS)a]NCS 

tran«-[Co  pna  Cla]Cl  KCNS  «ran«-[Co  pna  (NCS)2]NCS 

cis-[Copna  Ch]C]  MI,  (aqueous)  tran8-[Co  pna  NH8Cl]Cla 

frans-[Co  pna  Cla]Cl  MI:  (aqueous)  trans-[Co  pna  NH3Cl]Cls 

cts-[Co  pn.  Cla]Cl  Ml,  (anhydrous)  franfi  [Co  pna  (NH8)a]Cla 

trans-[Co  pna  Cla]Cl  MI,  (anhydrous)  <rans-lCo  pna  (\H3),]C13 

eron«-[Co  pna  Cla]Cl  Xa,S03  «s-[Co  pna  S()3]C1 

active  base.  The  fact  that  this  is  true  was  shown  by  the  rotatory  dispersion 
curves  of  several  ethylenediamine  and  acfaVc-propylencdiamine  cobalt  fill) 
complexes  of  the  types  [Co(AA)2a2]+  and  [Co(AA)2(BB)]+  50. 

This  technique  was  applied  to  the  study  of  cis-trans  conversion50  in 
the  reactions  of  coordination  compounds  containing  optically  active  pro- 
pylenediamine (Table  8.2). 

Dipolc  Moment.  The  chemical  bond  between  two  atoms  of  the  same  or 
similar  electronegativity  is  nonpolar,  and  a  molecule  such  as  A2  has  little 
tendency  to  orient  itself  when  placed  in  an  electric  or  magnetic  field.  If,  on 
the  other  hand,  the  two  atoms  do  not  have  similar  electronegativities  (such 
as  AB)  then  the  molecule  will  orient  itself  in  such  a  field  because  it  contains 
a  permanent  dipole.  In  much  the  same  way,  it  is  possible  to  distinguish 
complex  molecules  on  the  basis  of  their  electrical  symmetry.  It  would  there- 
fore appear  that  measurements  of  dipole  moments  could  be  used  to  dis- 
tinguish between  the  cis  and  trans  isomers  of  coordination  compounds. 
Numerous  studies  of  tetracovalent  complexes  of  the  type  [Pta2X2]60  have 
been  made  by  this  method,  but  it  has  not  been  used  for  hexacovalent  com- 
pounds. This  is  due  largely  to  the  fact  that  dipole  moments  are  usually 
derived  from  measurements  of  dielectric  constants;  such  measurements  are 
difficult  to  make  in  polar  solvents.  Since  most  of  the  geometric  isomers  of 
hexacovalent  compounds  are  salts,  they  are  not  soluble  in  nonpolar  solvents. 
Perhaps  some  inner  complexes  such  as  [Co(NH3)3(N02)3]  and  [Co(gly)3] 
might  be  studied  by  this  method. 

Although  it  is  difficult  to  measure  the  dipole  moments  of  complex  salts, 
polarographic  measurements  of  the  limiting  currents  for  stereoisomers  in- 
dicate differences  which  can  be  attributed  to  a  variation  in  electrical  sym- 
metry*1. It  was  found  that  the  cations,  cis-[(\>(\II;>,' \'( ),.),!'  and  cis- 
[Co  pn2  Cl2]4  produce  largei  limiting  currents  than  the  corresponding  trans 
isomers.  This  was  attributed  to  their  nonhomogeneous  internal  electric 
fields  which  cause  the  ions  to  orient  with  respect  to  an  electrode  and  move 
toward  it  under  the  influence  of  this  force  as  well  as  by  diffusion.  Since  this 

60.  Jensen.  Z.  anorg.  Cfu  m.t  225,  97  (1935);  Jensen,  ibid.,  229,  225  (1936). 

61.  Holtzclaw,  thesis,  University  of  Illinois,  1947. 


300  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

orientation  effect  is  not  present  in  the  case  of  the  trans  isomers,  the  cis 
cation  moves  faster  and  carries  more  current. 

Recent  studies62  on  the  separation  of  cis-  and  trans-  [Co(NH3)4(N02)2]+ 
using  a  cation  exchanger  show  that  the  trans  isomer  is  more  readily  re- 
moved from  the  resin.  Since  the  charge  and  size  of  these  isomeric  complexes 
are  the  same,  it  would  appear  that  the  cis  form  is  more  firmly  held  because 
of  its  larger  dipole  moment. 

Raman  Spectra.  The  Raman  spectra,  in  principle,  should  be  applicable  to 
the  determination  of  the  configuration  of  geometric  isomers  in  coordination 
compounds.  In  actual  practice,  it  is  often  not  possible  to  obtain  sufficient 
information  by  this  method  to  make  any  structural  conclusion.  The  Raman 
spectra  of  coordination  compounds  are  also  rather  difficult  to  obtain,  be- 
cause the  solutions  of  many  of  these  compounds  are  highly  colored.  Some 
studies  have  been  made  with  tetracovalent  compounds63  but,  as  yet,  very 
little64  has  been  done  with  hexacovalent  compounds. 

Infrared  Spectra.  Recent  studies65  on  the  infrared  spectra  of  complex 
compounds  show  that  this  method  can  be  used  to  distinguish  between  cis 
and  trans  isomers.  For  example,  fewer  absorption  peaks  are  present  in  the 
spectrum  of  £rans-[Co(NH3)4(N02)2]Cl  than  in  that  of  the  cis  isomer.  This 
is  the  natural  consequence  of  the  selection  rule,  since  the  trans  complex 
has  a  center  of  symmetry  whereas  the  cis  isomer  does  not. 

Magnetic  Susceptibility.  The  magnetic  susceptibilities  of  a  large  number 
of  metallic  ammines  have  been  determined  by  Rosenbohm66.  He  observed 
that  the  diamagnetism  is  greatest  for  the  hexammines  of  cobalt  (III),  less 
for  the  pentammines,  and  still  less  for  the  tetrammines  of  this  metal.  The 
triammines  of  cobalt(III)  are  very  weakly  diamagnetic;  some  compounds 
of  this  type  exhibit  paramagnetism.  It  is  evident,  therefore,  that  the  mag- 
netism is  largely  influenced  by  the  constitution  of  the  molecule.  However, 
an  examination  of  the  geometrical  isomerides  of  cobalt(III),  chromium  (III), 
and  platinum (IV)  complexes  indicates  that  the  magnetic  susceptibilities 
of  the  cis  and  trans  forms  are  indistinguishable.  This  is  also  true  of  the  re- 
spective optical  isomers. 

Solubility.  The  difference  in  the  solubilities  of  the  stereoisomers  cannot 
be  used  to  determine  their  structures.  Perhaps,  in  most  instances,  it  can  be 
said  that  the  cis  isomer  is  more  soluble  than  the  corresponding  trans  salt, 
but  there  are  numerous  exceptions  to  this  statement  and  it  should  certainly 
not  be  taken  as  a  general  rule. 

62.  King  and  Walters,  J.  Am.  Chem.  Soc,  74,  4471  (1952). 

63.  Mathieu,  ./.  chim.  phys.,  36,  271,  308  (1939). 

64.  Mathieu,  Compt.  rend.,  204,  682  (1937). 

65.  Quagliano  and  Faust,  ./.  .1///.  Chem.  Soc,  76,  5346  (1954). 

66.  Rosenbohm,  Z.  physik.  Chem.,  93,  693  (1919). 


STEREOISOMERISM  OF  HEXACOVALENT    [TOMS 


'M)\ 


Some  Properties  of  Cis -trans  Isomers 

In tercon version  of  cis- trans  Isomers.  It  has  already  been  mentioned 
that  the  preparation  of  a  cis  compound  by  the  displacement  of  a  chelate 
group,  or  the  proof  of  structure  by  the  replacement  of  singly  bound  groups 
with  a  chelate,  is  not  reliable.  This  is  largely  because  of  t  be  ease  with  which 
some  geometric  Isomers  are  known  to  rearrange  when  in  solution.  In  many 
instances,  the  trans  isomers  can  be  obtained  by  prolonged  boiling  of  solu- 
tions of  the  cis  salts,  e.g.,  K,|Ir  ox..  CI2]18,  K  [  1  Mi  ox2  CI2]17  and  [Co  en2 
\<>    ,  V 

The  best  known  example  is  the  transformation  of  green  trans- 
[Co  en2  C  T j  1 C  T  into  violet  cis-[Co  en-..  Cl2]Cl  and  vice  versa.  Jorgensen68  dis- 
covered that  the  trans  to  cis  conversion  is  brought  about  by  evaporation  of 
the  aqueous  solution  to  dryness,  and  that  the  reverse  process  occurs  in  the 
presence  of  hydrochloric  acid.  Drew  and  Pratt69  have  suggested  a  mecha- 
oism  for  these  chanties  which  involves  the  rupture  of  a  chelate  link  between 
ethylenediamine  and  the  cobalt(III)  (Fig.  8.25). 


en 


(CIS-VTRANS) 


Pig.  8.25 

This  mechanism  was  proposed  without  any  direct  evidence  but  primarily 

on  the  analogy  that  ethylenediamine  chelate  rings  in  platinum(II)  com- 

3  have  been  opened  by  digestion  with  hydrochloric  acid7'.  There  is, 

in  fact,  little  justification  for  the  assignment  of  structures  I  and  II  to  the 

67    Werner,  Arm.,  386,  1     1912  . 

Jorgensen,  ./.  prakt.  Chem.,  39,  1     1889 

Drew  and  Pratt../.  Chem.  Soc.,  1937,  506. 
7".  Drew  and  rI'rc>s  ./.  I  1932.  2328;  1933,  1335. 


302  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

complexes  generally  represented  as  [Co  en2  C12]C1-H20  and  [Co  en2  C12]C1- 
HC1  respectively.  The  cis  hydrate  is  purple  and  the  trans  hydrochloride  is 
green ;  that  is,  the  colors  are  not  markedly  altered  by  the  presence  of  either 
water  or  hydrogen  chloride.  Structure  (I)  would  also  suggest  a  similar 
mechanism  for  the  aquation  of  the  cis  isomer,  which  leads  to  the  racemi- 
zation  of  optically  active  [Co  en2  Cl2]+  during  aquation.  However,  Mathieu71 
has  shown  that  instead  of  racemizing,  the  complex  mutorotates  to  [Co  en2 
H20  Cl]++  at  a  rate  equal  to  that  of  chloride  ion  formation,  and  with  es- 
sentially complete  retention  of  configuration.  The  mechanism  of  this 
interconversion  has  been  investigated  using  radioactive  chlorine  to  deter- 
mine the  exchange  that  takes  place  during  isomerization72.  No  evidence 
was  found  for  any  direct  exchange  of  the  coordinated  chloro  groups  with 
the  chloride  ion.  This  suggests  that  the  following  equilibria  exist  in  solution: 

cis-  and  trans-[Co  en2  Cl2]+  ^±  [Co  en2  (H20)Cl]++^±  [Co  en2  (H20)2]+++ 

The  relative  amounts  of  the  isomeric  chlorides  in  the  solid  residue  appear 
to  be  largely  controlled  by  solubility  considerations72.  The  cis  chloride  is 
less  soluble  than  the  trans  but  the  latter  forms  a  sparingly  soluble  addition 
compound  with  hydrogen  chloride.  Apart  from  its  function  as  precipitant, 
hydrochloric  acid  plays  no  essential  role  in  the  changing  of  cis  to  trans. 
This  was  shown  using  the  complex  nitrate  instead  of  the  chloride.  A  solution 
of  trans-[Co  en2  C12]N03  can  be  evaporated  to  dryness  without  isomeriza- 
tion taking  place;  conversely,  cis-[Co  en2  C12]N03  is,  by  the  same  procedure, 
converted  quantitatively  into  the  trans  salt.  In  the  case  of  the  nitrate,  the 
trans  isomeride  is  only  slightly  soluble  in  water  and  is  always  the  first  to 
precipitate. 

Ettle  and  Johnson72  have  suggested  that  the  interconversion  may  occur 
by  the  following  mechanism: 

cis-[Co  en2  Cl2]+  +  H20  ;=±  cis-[Co  en2  H20  C1J++  +  Cl~ 

11 
trans-[Co  en2  Cl2]+  +  H20  ^  trans-[Co  en2  H20  Cl]++  +  Cl~ 

However,  they  do  not  describe  how  the  rearrangement  between  the  cis-  and 
/raws-chloroaquo  complexes  takes  place.  Mathieu71  has  observed  that  the 
rate  of  racemization  of  [Co  en2  H20  Cl]+2  is  independent  of  the  rate  of 
chloride  ion  formation  and  suggests  that  this  may  occur  as  a  result  of  the 
dissociation  of  the  coordinated  water.  This  explanation  may  be  used  also 
to  account  for  the  cis-trans  interconversion  of  the  chloroaquo  complexes. 

71.  Mathieu,  Bull.  soc.  chim.,  [5]  4,  687  (1937). 

72.  Ettle  and  Johnson,  ./.  Chem.  Soc,  1939,  1490. 


en/ 

CI 

;        en/ 

5r 

i 
i 
i 

__^CI    i 
i 

.  J 

+  H20      v 

ey. 

/  C 

/  -h2o  v 

!      /Co 

1 4^--~  r 

1    en^J 

/      CO         / 

°     /      S+H2° 

N-HjC 

H20^ 

C        1 

CIS 

l^/eri 

TRANS 

activated   intermediate 
Fig.  8.26 


Ii  is  apparent  from  the  trigonal  bipyramid  structure  for  the  activated 

intermediate  that  an  approach  by  water  between  positions  4  and  5  would 
yield  the  frans-chloroaquo  complex  whereas  attack  l)etween  2  and  4  or  be- 
tween 2  and  5  would  yield  the  cis  isomer. 

There  is  some  evidence  that  the  first  steps  in  this  interconversion  (aqua- 
tion of  the  dichloro  complex)  takes  place  without  inversion  of  configuration. 
For  example  Mathieu71  has  observed  that  the  reaction 

d-[Co  en2  Cl2]+  +  H20  -+  l-[Co  en2  H20  C1J++  +  Cl" 

occurs  with  retention  of  configuration.  Direct  proof  that  the  trans  isomer 
behaves  similarly  is  not  available.  However,  since  the  rate  of  aquation  of 
and  trans-[Co  en2  NO2  Cl]+  is  rapid  as  compared  to  the  rate  of  re- 
arrangement of  the  isomers  of  [Co  en2  H20  X02]++,  it  has  been  possible  to 
show  that  both  of  the  chloronitro  complexes  aquate  with  retention  of  con- 
figuration. Furthermore,  the  suggestion  that  the  interconversion  actually 
occurs  via  the  [Co  en2  H20  CI]**  ions  instead  of  the  dichloro  complexes  is  in 
accord  with  the  numerous  observations54, 74"77  that  aquo  complexes  gener- 
ally rearrange  more  rapidly  than  the  corresponding  acido  compounds. 

Chemical  Behavior  of  Cis-Trans  Isomers.  Closely  related  to  the 
interconversion  within  an  individual  molecule  are  the  conversions  that  may 
occur  during  reactions  in  which  coordinated  groups  are  displaced.  Werner67 
made  an  extensive  study  of  such  reactions  and  some  of  the  results  obtained 
are  given  in  Table  8.3. 

It  becomes  immediately  apparent  that  no  conclusions  can  be  drawn  from 
these  results.  Reactions  1  and  2,  6  and  7,  8  and  9,  and  12  and  13  show  that 
the  configuration  of  the  product  bears  no  relation  to  the  configuration  of 
the  original  material.  Perhaps  the  most  striking  pair  is  12  and  13,  for  a 
change  of  configuration  takes  place  in  each  of  these  reactions.  A  thorough 
study  was  made  of  this  case  under  various  condition.-,  bul  the  result  wa- 
rt.  B  Stone,  and  Pearson,  J.  Am.  Chem.  Soc.,  75,  819  I  Lfl 

75.  CJspensky  and  Tschibisoff,  Z.  anorg.  Chi  m.,  164,  326    1027 

76.  Cunningham,  Buriey,  and  Friend,  Nature,  109,  1103  (1962  . 

77.  Hamm,  ./.  .1///.  Chem.  Soe.,  75,  609  (1953). 


304 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


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STEREOISOMERISM  OF  HEXACOVALENT  ATOMS 


305 


always  the  same  as  Werner  had  reported78.  Werner  isolated  the  read  ion 

product  and  separated  the  two  isomers  in  order  to  determine  the  relative 

amounts  in  which  they  were  formed.  However,  these  compounds  are  known 
to  undergo  isomeric  rearrangements,  so  the  observed  isomeric  ratio  may 

not  be  a  direct  consequence  of  the  reaction  in  question.  However,  some  <>i' 
the  reactions  studied  by  Werner  have  recently  been  reinvestigated74  using 
a  spectrophotometry'  technique  to  determine  the  ratio  of  cis  and  trans 
isomers  in  situ  immediately  following-  the  substitution  reactions.  The  re- 
sults obtained  by  this  method  were  generally  in  good  accord  with  the  earlier 
observations  reported  by  Werner. 

Werner  at  first  believed  that  substitution  normally  takes  place  with  re- 
tention of  configuration,  and  that,  whenever  this  is  not  the  case,  rearrange- 
ment takes  place  in  order  to  form  the  more  stable  isomer.  However,  it  soon 
became  apparent  to  Werner  that  this  interpretation  was  not  compatible 
with  the  experimental  facts.  For  example,  reactions  3  and  4  in  Table  8.3 
show  that  trans-[Co  en2  XCS  Cl]+  reacts  with  liquid  ammonia  to  yield  two 
parts  of  cw-  and  one  part  of  trans-[Co  en2  NHa  XCS]++;  therefore  the  cis 
isomer  is  expected  to  be  more  stable  than  the  trans  complex.  However,  the 
reaction  of  cis[Co  en2  XCS  Cl]+  with  liquid  ammonia  does  not  yield  ex- 
clusively n's-[Co  en2  X"H3  XTCS]++,  but  equimolar  quantities  of  the  cis  and 
trans  isomers. 

Werner  attempted  to  explain  these  results  by  assuming  that  the  complex 
is  surrounded  by  an  outer  sphere  of  more  loosely  held  groups.  If  the  in- 
coming group  (c)  is  oriented  in  this  outer  complex  in  a  position  adjacent  to 
the  group  that  is  to  be  replaced  (b),  there  will  be  no  change  in  configuration 
during  the  substitution  (Fig.  8.27).  However,  if  (c)  is  in  a  position  opposite 
to  (b),  the  reaction  is  accompanied  by  change  in  configuration. 


AA 


AA 


-AA    (CONFIGURATION   DOES 

NOT  CHANGE) 
AA    C 


(CONFIGURATION  D°rr  CHAf  GE) 

lie.  8.27 


"8.  Becker,  thesis,  University  of  Illinois 


306  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

The  possibility  of  predicting  the  position  of  the  incoming  group  on  the 
basis  of  electrostatic  forces  has  been  suggested79.  An  explanation  of  this 
type  might  be  used  to  interpret  the  fact  that  cis-[Co  en2  NH3  C1]C12  is  pro- 
duced by  the  reaction  of  trans- [Co  en2  C12]C1  and  aqueous  ammonia.  If  it  is 
assumed  that  the  negative  nitrogen  atom  of  ammonia  approaches  the  octa- 
hedron in  such  a  way  as  to  maintain  a  maximum  distance  from  the  nega- 
tive chloro  groups,  then  the  ammonia  would  be  in  the  plane  of  the  ethylene- 
diamine  molecules  and  it  could  be  attached  to  positions  2,  3,  4  or  5  which 
would  account  for  the  formation  of  cis-[Co  en2  NH3  C1]C12  (Fig.  8.28). 


5 

n/ 


2 

L  n Ha 


cr 

Fig.  8.28 


Although  this  explanation  appears  to  account  satisfactorily  for  the  reaction 
cited,  it  cannot  be  used  as  a  general  interpretation.  For  example,  it  would 
suggest  that  the  analogous  propylenediamine  complex,  trans- [Co  pn2  Cl2]+, 
should  react  with  ammonia  to  yield  the  a's-chloroammine  derivative; 
however,  the  product  of  this  reaction  is  the  trans  isomer50.  Furthermore,  it 
is  expected  on  the  basis  of  such  an  approach  that  cis-  [Co  en2  Cl2]+  would 
react  to  yield  trans- [Co  en2  NH3C1]++  but  the  product  is  known  to  be  the 
cis  complex.  These  results  indicate  that  the  electrostatic  effect  cannot  be 
the  sole  factor  responsible  in  determining  the  course  of  these  reactions. 

Basolo,  Stone,  and  Pearson74  have  recently  used  a  somewhat  different 
approach  to  the  problem  of  molecular  rearrangements  that  may  occur 
during  substitution  reactions  in  octahedral  complexes.  They  suggest  that 
the  reaction  involves  either  a  dissociation  process  (SN1)  or  a  displacement 
(SN2)  reaction  which  can  lead  to  different  isomeric  forms  depending  upon 
the  configuration  of  the  intermediate.  For  example  in  Fig.  8.29  the  trans 
complex  [M(AA)2ax]  is  represented  as  undergoing  a  dissociation  process 
(SN1)  by  way  of  a  tetragonal  pyramid,  to  yield  a  trans  product;  if  the 
intermediate  has  a  trigonal  bipyramid  structure,  the  product  may  be  a 
mixture  of  cis  and  trans  isomers.  However,  with  a  displacement  reaction 
(SN2)  as  shown  in  Fig.  8.30  the  product  will  have  the  cis  configuration,  if  the 
attack  of  the  incoming  group  is  from  the  "back",  but  trans  if  the  ap- 

79.  Mathieu,  Bull.  soc.  chim.,  [5]  5,  783  (1938). 


STEREOISOMERISM  OF  HEXACO}  ALENT  ATOMS 


30; 


proach  is  from  the  "front"  of  the  complex.  Molecular  rearrangements 
during  substitutions  have  been  discussed  in  terms  of  "edge"  and  "non- 
edge"  displacements79*.  It  therefore  becomes  apparenl  thai  stereochemical 
studies  alone  will  not  elucidates  detailed  mechanism  of  substitution  reac- 
tions in  octahedral  complexes.  However,  some  progress  has  already  been 
made80-88  toward  the  determination  of  the  molecularity  of  these  reactions. 
Fort  lie  reaction  [Co  en,  NO,  Up  +  1I,<>  ->  [Co  onUM)  XOJ++  +  Cl~ 
the  experimental  evidence  supports  a  dissociation  mechanism  involving  a 
tetragonal  pyramid  intermediate71,  s:;  sl. 

The  observation  that  increased  steric  hindrance  in  a  series  of  trans- 
[C\)( AAV.Ujp  compounds  is  accompanied  by  increased  rates  of  aquation 
has  been  cited  in  support  of  an  SN1  mechanism81.  Substitution  reactions  of 
cis-[Co  en-j  ( JljJ4  in  methanol  involve  either  an  SN  1  or  Sn2  process  depending 
upon  the  nucleophilic  character  of  the  reactant86. 


TETRAGONAL 

PYRAMID 


31  +9 


TRIGONAL 
BIPYRAM1D 


FlQ.  8.29.  Dissociation  process  (SnI)  for  trans-[M(AA)2ax] 

Brown,  [ngold,  and  Nyholm,  ./.  Chem.  80c.,  1953,  2071. 
BO.  Basolo,  Bergmann,  and  Pearson.  ./.  Phys.  Chi  m.t  56,  22    L952 
Bl.  Pearson,  Boston,  and  Basolo,  ./.   Am.  Chem.  Soc.,  74,    2943   (1952  ;  75,  3089 

1".- 
82.  Rutenberg  and  Taube,  •/.  Chem.  Phys.,  20,  B23  (1952 
B3.  Werner,  Ber.,46,  121  (1912). 

B4.  Pfeiffer,  Golther,  and  Angern,  Ber.,  60.  305    1927). 
B5.   Brown  and  [ngold,  •/.  <'/„/„.  Soc  .  2680    l1' 


308 


CHEMISTRY  OF  THE  cuoitDLXATION  COMPOUNDS 


Fig.  8.30.  Displacement  (Sn2)  process  for  trans- [M(AA)2ax] 


Table  8.4.  Relative  Amounts  of  Geometrical  Isomers  Anticipated  on  the 
Basis  of  Various  Reaction  Mechanisms  for  Substitutions  in- 
Octahedral  Complexes  of  the  Type  [M(AA)2ax] 


Dissociation  (SnI) 

Displacement  (Sn2) 

[M(AA)2ax] 

Tetragonal  Pyramid 

Trigonal  Bipyramid 

Rear 

Front 

cis 

trans 

cis 

trans 

cis 

trans 

CIS 

trans 

Trans 
Cis 

per  cent 

0 
100 

per  cent 

100 
0 

per  cent 

66.6 

80 

per  cent 

33.3 
20 

per  cent 

100 
66.6 

per  cent 

0 
33.3 

per  cent 
0 
100 

per  cent 

100 
0 

Optical  Isomerism 

Numerous  coordination  compounds  have  been  resolved  into  their  enan- 
tiomorphs  and  some  of  the  problems  in  this  connection  will  be  discussed. 

The  optica]  activity  found  in  coordination  compounds  is  not  always 
caused  by  the  presence  of  an  asymmetric  atom.  Experiments  have  shown 
that  molecules  or  ions  in  which  the  entire  configuration  possess  only  axial 
symmetry  may  exisl  in  enantiomorphously  related  forms.  Coordination 
compounds  are  of  this  general  type  and  many  are  known  to  have  high  op- 
tical  activity,  i.e.  [Co  en   Mr  .  [M]D  =  ±  002°.  As  is  shown  in  Fig.  8.31, 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS  309 

en 


Co 


Co 


I^Sn 


Fig.  8.31 


there  is  no  chemical  contrast  whatsoever  between  the  three  substituents 
attached  to  the  central  atom,  and  the  optical  activity  results  from  the  dis- 
symmetrical spatial  disposition  of  these  identical  substituents.  There  is  no 
"asymmetric  atom"  in  the  sense  of  the  Le  Bel-Van't  Hoff  theory,  but, 
in  contrast,  the  division  of  space  about  the  central  atom  is  a  decidedly 
symmetrical  one.  The  fact  that  the  only  prerequisite  for  optical  isomerism 
is  an  asymmetric  molecule  or  ion  can  also  be  extended  to  certain  carbon 
compounds  which  contain  no  asymmetric  carbon  atom.  A  good  example 
of  such  a  compound  is  the  dilactone,  Fig.  8.32,  which  was  resolved  by  Mills 


CO 

/ 


HOOc/ 


\ 


_/ 


) 


COOH 


O — CO 
Fig.  8.32 

and  Nodder86.  Other  compounds  of  this  spirane  type  have  also  been  re- 
solved, as  have  compounds  of  the  inositol  type87,  allenes88,  compounds  with 
restricted  rotation  about  a  single  bond89;  and,  recently90,  optical  activity 
of  the  4,5-phenanthrene  type  has  been  realized. 

Various  Types  of  Optically -Active  Isomers 

Cationic  Complex  Compounds.  Numerous  complex  cations  have 
been  resolved  into  their  optically-active  antipodes.  No  attempt  will  be 
made  to  discuss  the  preparation  and  resolution  of  all  of  these  compounds, 
but  the  general  types  which  have  been  resolved  will  be  mentioned  and  some 
examples  of  each  given.  Complex  cations  with  general  formulas  of  [M(AA)3], 
[M(AA)2(BB)],  [M(AAUi2],  [M(AA)2ab],  [M(AA)a2b2],  [M(AA)(BB)a2], 
[MfAA'Ajoj  and  [M(ABCCBA)]  have  been  separated  into  their  optically 

86.  Mills  and  Xodder, ./.  Chem.  Soc.,  117,  1407  (1920). 
B7.  Mohr,  ./.  prakt.  Chem.,  [27]  68,  369  (1903). 

88.  Pope,  Perkin,  and  Wallach,  Ann.,  371,  180  (1909). 

89.  Adams  and  Yuan,  Chem.  Revs.,  12,  262  (1933). 

90.  Newman  and  Bussey,  •/.  Am.  Chem.  Soc.,  69,  3023  (1947). 


310 


CHEMISTRY  OE  THE  COORDINATION  COMPOUNDS 


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STEREOISOMERISM  OF  HEXACOYALEXT  ATOMS 


311 


active  antipodes.  Spatial  arrangements  for  these  enantiomorphs  arc  shown 
Fig.  8.33.  For  the  last  two,  other  arrangements  arc  also  possible.  Some 


m 


[m(aa)2cbb)] 


Qmcaa)2aJ 


[MCAA")BB  Aa] 


[Waa'  a)£| 

Fig.  8.33.  Possible  forms  of  some  chelate  complexes 

specific  examples  of  these  compounds  which  have  been  resolved  are  listed 
in  Table  8.5. 

91.  Werner,  Ber.,  45,  121  (1912). 

92.  Smirnoff,  Helv.  chim.  Acta.,  3,  177  (1920). 

93.  Jaeger  and  Blumendal,  Z.  anorg.  allgem.  Chem.,  175,  161  (1928). 

94.  Jaeger  and  Bijkerk,  Proc.  Acad.  Sci.  Amsterdam,  40,  116  (1937). 

95.  Werner,  Ber.,  45,  865  (1912). 

96.  Xeogi  and  Mandal,  J.  Indian  Chem.  Soc,  13,  224  (1936). 

97.  Werner,  Ber.,  45,  433  (1912). 

98.  Jaeger,  Kristallogr.,  Z.,  58,  172  (1923). 

99.  Werner,  Ber.,  45,  1228  (1912). 

100.  Jaeger,  "Spatial  Arrangements  of  Atomic  Systems  and  Optical  Activity,"  p. 

92,  New  York,  McGraw-Hill  Book  Co.,  1930. 

101.  Werner  and  Smirnoff,  Heir.  chim.  Acta.,  3,  476,  483  (1920).  - 

102.  Xeogi  and  Mukherjee,  J.  Indian  Chem.  Soc,  11,  681  (1934). 

103.  Xeogi  and  Mandal,  ./.  Indian  Cfu  m.  Soc,  14,  653  (1937). 

104.  Werner,  Ber.,  44,  1887  (1911). 

105.  Werner  and  McCutcheon,  Ber.,  46,  3281  (1912) ;  47,  2171  (1914). 

106.  Bailar  and  Auten,  ./.  .1///.  Chem.  Soc,  56,  774  (1934). 
K)7.  Waits,  "Dissertation,"  Zurich,  1912. 

108.  Bailar,  Halsam  and  Jones,  J.  Am.  Chem.  Soc.  58,  2226    1936 

109.  Werner,  Ber.,  44,  3132  (1911). 

100.  Werner  and  Smirnoff,  Helv.  chim.  Acta.,  3,  472  (1920). 

111.  Werner,  Ber.,  44,  3272  (1911). 

112.  Mann  and  Pope,  J.  Proc.  Roy.  Soc,  London,  107A,  80  (1925). 


.-ill' 


CHEMISTIiY  OF  THE  ('OOh'l)IXATION  COMPOUNDS 


The  usual  method  employed  for  the  separation  of  these  enantiomorphous 
cations  may  be  illustrated  with  the  racemate,  [Co  en3]Cl3'3H20.  If  a  solu- 
tion containing  one  mole  of  this  salt  is  treated  with  one  mole  of  silver 
deatfro-tartrate,  there  is  formed  a  chlorotartrate,  [Co  en3]Cl(d-C4H40c). 
Slow  evaporation  of  this  solution  causes  the  gradual  deposition  of  triclinic 
crystals  of  dextro-[Co  en;{]Cl(d-C406H4)-5H20.  These  crystals  are  removed 
as  completely  as  possible;  additional  concentration  of  the  mother  liquor 
gives  a  viscous  residue.  Solutions  of  the  triclinic  crystals  and  of  the  viscous 
residue  when  treated  with  solutions  of  sodium  iodide  precipitate  the  crys- 
talline iodides  respectively:  d-[Co  en3]I3-H20  and  Z-[Co  en3]I3-H20.  Al- 
though this  procedure  gives  satisfactory  results  for  [Co  en3]+++,  the  task  of 
separal  ing  enantiomers  is  often  very  tedious  and  the  most  suitable  resolving 
agent  and  conditions  must  be  found  by  trial  and  error  for  each  particular 
complex  cation  (page  332). 

Anionic  Complex  Compounds.  The  number  of  anionic  complexes 
which  have  been  obtained  in  optically-active  form  is  considerably  less  than 
that  for  cationic  complexes.  The  spatial  arrangements  are  the  same  as 
illustrated  in  Fig.  8.33  and  specific  examples  are  given  in  Table  8.6. 


Table  8.6.  Some  Asymmetric  Anions  Which  Are  Reported  to  Have 

Been  Resolved 

[M(AA)2a2] 


[M(AA)3] 

[A1(C204)3]S113 


As< 


[Co(C204)3pa5 

[Cr(C204)3l=116 

[Cr(OOCCH2COO)3]sm 

[Fe(C204)3p118 

[Ir(C204)3p119 

[Rh(C2()4),l    « 

[Rh(OOCCH2COO» 


[Ir(C204)2Cl2]=16 
[Rh(NHS02NH)2(H20): 
[Rh(C304)2Cl2]Si " 

[M(AA)2ab] 
[Ru  py(C204)2NO]-  *« 
[M(AA)a2b2] 

[Co(NH3)2C204(N02)2]- 


113.  Burrows  and  Lauder,  ./.  .1///.  Chem.  Soc,  52,  2600  (1931);  Treadwell,  Szabados, 
and  Baimann,  Helv.  chim.  Ada.,  15,  1040  (1932);  Wahli.  Ber.,  6?,  300  (1927). 

111.  Rosenheim  and  Plato,  Ber.,  58,  2000  (1925);  Weinland  and  Heinzlei,  Ber.,  52, 
1322  (1919). 

LIS.  Jaeger,  Rec.  trav.  chim.,  38,  217  (1919). 

L16.  Jaeger,  ibid.,  38,  213  (1919);  Werner,  Ber 

117.  Jaeger,  Rec.  trav.  chim.,  38,  294  1 1019). 

lis.  Thomas,  •/    Chem.  Soc.,  119,  1140  (1921  i. 

ll'.i.  Delepine,  Compt.  rend.,  159,  239  (191  l)\ 
L917). 


45,  3061  (1012), 


Delepine,  Bull.  Soc.  chim.,  [4]  21,  161 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS  313 

In  general,  the  nun  hods  used  to  resolve  complex  anions  are  based  on 
the  same  principles  as  those  used  with  the  cations,  that  is,  the  combination 
with  an  easily  removable  optically-active  substance.  Since  the  complexes 
are  anions  in  this  ease,  the  cations  to  which  they  are  linked  must  he  re- 
placed by  optically-active  bases.  Strychnine  has  been  used  to  resolve  the 
trioxalatocobaltate(III),  chromate(III),  rhodate(III),  and  iridate(III) 
salts121.  The  strychnine  can  easily  be  removed  by  precipitation  as  the  iodide 
with  potassium  iodide,  the  potassium  salt  of  the  optically  active  anion  re- 
maining in  solution. 

Nonionic  Complex  Compounds.  Asymmetric  inner-complex  com- 
pounds are  known  to  exist  and,  theoretically  these  can  be  resolved  into 
their  optically-active  antipodes.  The  ordinary  technique  is  not  applicable 
to  the  resolution  of  these  compounds  because  they  do  not  form  salts.  Very 
few  complexes  of  this  type  have  been  obtained  in  their  optically-active 
forms.  Lifschitz27  did  obtain  some  evidence  for  the  existence  of  the  four 
possible  isomers  of  tris(rf-alanine)cobalt(III).  The  resolution  of  a  complex 
of  the  type,  [Co(DMG)2  NH3  CI],  has  been  accomplished  by  the  preferen- 
tial adsorption  of  an  antipode  on  optically-active  quartz28,  125.  Dwyer  and 
his  co-workers  have  recently  had  some  success  with  the  resolution  of  non- 
ionic  complexes  by  applying  their  method  of  "configurational  activity" 
(page  335). 

Complex  Compounds  Containing  Optically -active  Donor  Mole- 
cules. Optically-active  bidentate  molecules  or  ions  have  been  made  to 
coordinate  with  hexacovalent  metals  and  the  stereochemistry  of  some  of 
these  complex  compounds  has  been  investigated.  Complexes  of  this  type 
are  of  interest  because  they  offer  problems  for  which  there  are  no  counter- 
parts in  the  stereochemistry  of  carbon  compounds. 

Limited  Number  of  Isomers.  An  octahedral  complex  containing  three 
molecules  of  an  optically-active  bidentate  coordinating  agent  wrould  be 
expected  to  exist  in  a  large  number  of  stereoisomeric  forms.  Taking  d  and 
l  to  represent  the  signs  of  rotation  of  the  complex  as  a  wrhole;  and  d  and  /, 
the  signs  of  rotation  of  the  bidentate  molecule,  there  are  eight  possible 
combinations:  d[IU],  i>[lld],  D[ldd],  v>[ddd],  \\lll],  l[/W],  i\ldd]  and  i\ddd\. 
Moreover,  since  these  eight  cases,  when  taken  in  pairs  represent  each 
other's  mirror  images  (D[ddd]  and  l{111],  v[ldd]  and  i\dll],  etc.)  they  may  be 
combined  pair-wise  in  equimolecular  quantities  to  yield  four  racemoids  and 
twenty-four  partial  racemoids.  Experiment  has  shown,  however,  that  these 

120.  Werner,  Ber.,  47,  1954  (1914). 

121.  Jaeger,  Kec.  trav.  chim.,  38,  300  (1919). 

122.  Mann,  ./.  Chem.  Soc,  1933,  412. 

123.  Charonnat,  Coiu,,t.  rend.,  178,  1423  (1924). 

124.  Jaeger,  Rec.  trot  .  chim.,  38,  245,  251,  263,  265  (1919). 

125.  Kuroya,  Aimi,  and  Tsuchida,  J.  Chem.  Soc,  Japan,  64,  995  (1943). 


31  1  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

combinations  are  not  all  of  equal  si  ability;  in  fact,  for  octahedral  complexes 
containing  optically-active  propylenediamine92,  126,  1 ,2-cyclopentanedi- 
amineM  and  1  ,2-cyclohexanediaminew  only  the  l[///]  and  i>[ddd]  (or,  in 
other  cases,  \\rfdd\  and  i>|///|)  isomeric  ions  are  stable  enough  to  be  isolated. 
\  similar  effecl  is  observed  if  the  complex  contains  only  two  optically- 
active  coordinating  groups.  It  has  been  shown  that  ions  such  as  cis- 
[Co  i>iij  Cli]H  and  ds[Co  cptn-  (%\{  exist  in  only  two  of  the  six  possible 
torins-n|//(1lL.|  and  hlddCU]1*1*.  If  the  dichlorobis(/cro-propylenediamine)co- 
balt(III)  ion,  [Co  /-pno  Cl2]+,  is  treated  with  dea^ro-propylenediamine,  the 
ion  [Co  /-pn_-  ^/-pn]+++  apparently  forms,  but  immediately  rearranges  to  a 
mixture  of  the  more  stable  l[Co  o?-pn3]+++  and  d[Co  /-pn3]+++ 92-  128.  Analo- 
gous results  have  been  obtained  with  optically-active  cyclopentanediamine 
and  the  reactions  which  occur  are  summarized  by  Jaeger100  as: 

[ddCU]  -U  [ddl]  ->  2[ddd]  +  [III],        or  [ddd]  +  racemoid 

[UCU]  -^  [lid]  ->  2[lll]  +  [ddd],         or  [///]  +  racemoid 

These  selective  effects,  while  pronounced,  are  not  absolute,  but  relative. 
Lifschitz27  found  evidence  that  tris(d-alanine)cobalt(III)  and  chro- 
mium(III)  exist  in  B[ddd]  and  i\ddd]  forms.  It  has  likewise  been  shown  by 
Bailar  and  McReynolds129  that  the  ion  [Co  Z-pn2  C03]+  exists  inbothD[//C03] 
and  l[//C03]  forms;  they  believed  that  the  latter  is  unstable,  rearranging 
to  the  former  when  warmed  gently.  Recent  studies130  indicate  that  these 
two  forms  are  present  in  a  state  of  equilibrium  which  shifts  predominantly 
towards  d[//C03]  upon  standing  in  solution;  however,  if  this  solution  is 
evaporated  to  dryness,  the  residue  obtained  is  largely  l[//C03]. 

When  only  one  molecule  of  the  optically-active  base  is  present  in  the  co- 
ordination sphere,  there  is  some  tendency  toward  the  formation  of  pre- 
ferred orientations,  but  not  enough  to  fix  completely  the  configurations. 
Thus,  when  Jaeger  and  Blumendal93  allowed  racemic  frans- 1,2-cyclopen- 
tanediamine  to  react  with  racemic  [Co  en2  Cl2]+,  they  obtained  a  true 
racemic  mixture  of  d-[Co  en2  Z-cptn]+++  and  l-[Co  en2  d-cptn]~H'+  without 
detecting  any  of  the  other  two  possible  forms.  When,  however,  they  used 
tevo-cyclopentanediamine,  they  observed  that  the  base  entered  both  the 
D  and  l  forms  of  the  complex,  yielding  d  and  L-[Co  en2  Z-cptn]+++.  A  com- 
patible svstem,  studied  by  Jonassen,  Bailar  and  Huffman131,  reveals  that 
dextro-t&rt&nc  acid  reacts  readily  with  [Co  en2  C03]+  to  give  the  two  di- 

126.  Tschugaeff  and  Sokoloff,  Ber.,  40,  177  (1907) ; Ibid.,  42,  55  (1909). 
1  This  disregards  the  possibility  of  position  isomers  (page  286).. 

127.  Lifschitz,  Z.  physik.  Chem.,  114,  493  (1925). 

128.  Bailar,  Stiegman,  Balthis,  and  Buffman,  J.  Am.  Chon.  Soc,  61,  2402  (1939). 

129.  Bailar  and  McReynolds,  ibid.,  61.  3199  0939). 

130.  Martinette  and  Bailar,  ./.  Am.  Chem.  Soc,  74,  1054  (1952). 

131.  Jonassen,  Bailar,  and  Buffman,  •/.  .1;//.  Chem.  Soc,  70,  756  (1948). 


STEREOISOMERISM  OF  II EX ACOVALENT  ATOMS  315 

astereoisomers  d-[Co  ens  d-tart]"1  and  l-[Co  en2  d-tnrt]+,  which  differ 
strikingly  in  stability,  reactivity,  and  solubility.  It  has  recently  been  shown 
that  the  equilibrium  mixture  of  the  two  diastereoisomers  when  heated  to 
150°  changes  to  l-[Co  en,  d-tart]"*  '••-. 

These  experiments  with  salts  of  the  type  [Co  en2  C12]C1  show  that  a  mole- 
cule of  an  optically-active  base,  such  as  fevo-cyclopentanediamine,  may  be 
introduced  into  either  the  d  or  l  antipode.  Such  an  introduction  is  more 
difficult  if  two  molecules  of  the  optically-active  antipode  of  the  substitute 
are  originally  present,  instead  of  two  molecules  of  ethylenediamine.  It 
would  appear  from  this  that  there  is  a  more  pronounced  contrast  between 
a  dextro  and  levo  isomer  of  the  same  compound  than  exists  between  an 
optically-active  molecule  and  a  totally  different  substance.  The  presence  of 
such  nonrelated  molecules  in  a  coordination  sphere  appears  to  be  a  less 
serious  hindrance  to  the  entrance  of  an  optically-active  substitute  than  is 
the  presence  of  similar  molecules  having  opposed  enantiomorphous  arrange- 
ments. 

Complex  Compound  as  a  Possible  Resolving  Agent.  The  results  ob- 
tained with  optically-active  coordinating  agents  suggest  that  in  the  reaction 
between  an  optically-active  complex  and  an  excess  of  a  racemic  coordi- 
nating substance  the  complex  may  accept  one  antipode  of  the  coordinating 
agent  preferentially,  thus  effecting  a  resolution.  Investigations  of  this 
possibility  have  been  made128, 133. 

Although  the  presence  of  two  or  three  optically-active  chelate  groups 
in  an  octahedral  complex  tends  to  fix  a  definite  configuration  upon  the 
complex  as  a  whole,  and  limits  the  number  of  stereoisomers  which  can  be 
isolated  to  a  small  fraction  of  those  theoretically  possible,  this  effect  is 
considerably  less  noticeable  in  complex  ions  containing  only  one  asymmetric 
chelate  group.  As  has  already  been  indicated,  however,  while  both  the 
d  and  l  forms  of  ctoro-tartratobis(ethylenediamine)cobalt(III)  ion, 
[Co  en2  d-tart]+,  exist,  they  differ  greatly  in  reactivity131.  When  the  mixture 
of  the  two  is  shaken  with  etfrylenediamine  at  room  temperature,  part  of 
the  material  reacts  within  two  hours,  giving  d-[Co  ens]4"4"*,  and  the  re- 
mainder does  not  react  even  in  twelve  hours.  This  indicates  that  if  the 
complex  were  prepared  from  racemic  tartaric  acid,  the  active  antipodes 
would  be  displaced  at  different  rates.  This  effect  has  been  considerably 
enhanced  by  using  Zeyo-propylenediamine  in  place  of  ethylenediamine. 
Racemic  tartaric  acid  has  been  partially  resolved  by  treating  dZ-tartratobis- 
(Z-prop3denediamine)cobalt(III)  chloride  with  /-propylenediamine133a-  134. 
The  first  ion  removed  from  the  complex  was  largely  the  Z-tartrate.  Resolu- 

132.  Johnson,  thesis,  University  of  Illinois,  1948. 

133.  Jonassen,  Bailar,  and  Gott, ./.  Am.  Chein.  Soc,  74,  3131  (1952) ;  Hamilton,  thesis, 

University  of  Illinois,  1947. 

134.  Gott  and  Bailar, ./.  Am.  Chem.  Soc,  74,  4820  (1952). 


316  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

tion  of  this  acid  is  also  achieved  when  c^-tartratobis(J-propylenediamine)- 
coball  ( III)  chloride  is  made  to  react  with  racemictartrate133b.  The  Z-tartrato 
group  is  displaced  from  the  complex  ion  by  d-tartrate  and,  consequently, 
the  final  reaction  mixture  contains  largely  /-tartrate  ion  and  the  d-tartrato 
complex.  In  the  same  manner,  Z-propylenediamine  is  obtained  from  the 
reaction  of  a  mixture  of  (/-tartratobis(/-propylenediamine)cobalt(III) 
chloride  and  d-tartratobis(c?-propylenediamine)cobalt(III)  chloride  with 
i ;  1 1  ( 'inic  propylenediamine132. 

It  may  be  mentioned  in  conclusion  that  many  optically-active  complex 
salts  have  been  shown  by  Shibata  to  exhibit  a  catalytic  oxidizing  effect, 
analogous  to  the  enzymic  action  of  oxidases135.  When,  for  example,  racemic- 
3,4-dihydroxy-phenylalanine  was  oxidized  under  the  catalytic  in- 
fluence of  £m?-chloroamminebis(ethylenediamine)cobalt(III)  bromide, 
l-[Co  en2  NH3  Cl]Br2 ,  the  levo  amino  acid  was  preferentially  destroyed.  Al- 
though this  has  been  attributed  to  an  "enzyme-like  action"  by  the  inor- 
ganic complex,  Bailar136  has  suggested  as  an  additional  explanation,  that 
one  form  of  the  amino  acid  becomes  part  of  the  complex,  while  the  other 
does  not,  and  subsequent  oxidation  merely  destroys  one  or  the  other. 
Studies  of  this  type  have  likewise  been  carried  out  by  Pugh137  whose  re- 
sults are  not  entirely  in  accord  with  those  of  Shibata. 

Partial  Asymmetric  Synthesis.  The  fact  that  hexacovalent  complexes 
containing  optically-active  groups  do  not  exist  in  all  the  possible  stereo- 
chemical forms,  but  only  in  certain  preferred  configurations,  suggests  that 
these  groups  exert  a  steric  effect  on  the  coordination  sphere  of  the  central 
metal  ion  which  hinders  the  formation  of  the  other  isomers.  Thus,  existence 
of  only  d[IU]  and  h[ddd]  isomers  indicates  that  the  addition  of  I  antipodes 
to  a  complex  always  gives  rise  to  a  dextro  configuration  of  the  octahedron 
and,  likewise,  a  d  antipode  always  causes  the  formation  of  a  levo  structure. 
In  other  words,  a  preferred  configuration  is  induced  by  optically-active 
coordinating  groups,  and  reactions  which  introduce  such  groups  give  rise 
to  an  asymmetric  octahedron. 

Evidence  that  such  partial  asymmetric  syntheses  take  place  was  ob- 
tained by  a  study  of  the  molecular  rotation  of  various  platinum  complexes 
containing  different  numbers  of  coordinated  Zeyo-propylenediamine  mole- 
cules. It  was  shown126b  that  the  molecular  rotation  caused  by  each  molecule 
of  Zeyo-propylenediamine  introduced  into  various  platinum (II)  complexes 
is  about  96  degrees  (Table  8.7).  Since  the  presence  of  two  molecules  of 
active  propylenediamine  results  in  a  molecular  rotation  of  192°,  it  might 
be  expected  that  the  addition  of  a  third  active  molecule  would  give  a  com- 

l :;•"».  Shibata  and  Tsuchida,  Hull.  Chem.  Soc,  Japan,  4,  142  (1929);  Shibata,  Tonaka, 

.•in. I  Goda,  ibid.,  6,210  (1931). 
136.  Bailar,  Cfo  19,  67  (1938). 

L37.  Pugh,  Biochem.  J.,  27,  480  (1933). 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS  317 

Table  8.7.  Optical  Rotation  of  Platinum (II)  Complexes  Containing 

leVO-PBOPYLENEDl  \MINE 
Substance  MD  l-^'in 

[PW-pn  (NH,),]C1>  +25.17  +94.14 

[Pt  Z-pn  en]Cl2  +24.07  +96.28 

[Pt  Z-pn  tn]Cli  +23.60  +97.70 

[Pt  Z-pn2]Cl2  +46.37  +192.0 

pound  with  a  molecular  rotation  of  +288°.  However,  it  was  observed  by 
Smirnoff91  that  the  compounds  formed  by  addition  of  this  third  base  mole- 
cule were  L-[Pt  (/-pn3]X4  and  D-[Pt  Z-pn3]X4,  with  values  of  [M]D  equal  to 
-1027°  and  +1025°,  respectively.  If  it  is  assumed  that  only  288°  of  the 
total  is  due  to  the  three  active  propylenediamine  groups,  the  excess  must 
be  a  result  of  the  asymmetry  of  the  cation. 

A  similar  asymmetric  effect  is  observed  when  only  two  optically-active 
bidentates  are  coordinated  to  the  hexacovalent  central  ion.  This  is  clearly 
demonstrated  by  the  similar  rotatory  dispersion  curves  of  numerous  bis- 
(ethylenediamine) cobalt  (III)  ions  and  analogous  cis-bis(active-propy\ene- 
diamine)cobalt(III)  ions50.  The  rotatory  dispersion  curves  of  the  corre- 
sponding trans  isomers  resemble  that  of  active  propylenediamine  because 
the  complex  is  symmetrical  and  therefore  cannot  contribute  to  the  optical 
activity. 

Complex  Compounds  Containing  Optically -active  Unsymmetri- 
cal  Donor  Molecules.  The  most  extensively  studied  asymmetric  bidentate 
molecule  which  has  been  used  as  a  coordinating  group  is  propylenediamine. 
The  number  of  theoretically  possible  isomers  of  complexes  of  the  type 
[M  pn3]  is  greatly  increased  due  to  the  existence  of  position  isomers  as  well 
as  the  optical  isomers  (Fig.  8.34). 


.DiL.  _DtJ_  Dtj_ 


DVJ.  _DLL_ 

Fig.  8.34.  Possible  forms  of  some  complexes  containing  optically  active  ligands. 


318  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

II'  all  of  the  predicted  isomers  and  all  the  total  and  partial  racemates 
were  found,  the  chemistry  of  these  complexes  would  be  hopelessly  compli- 
cated, but  that  is  not  the  case.  For  example,  the  only  isomers  which  were 
isolated  or  identified  for  cobalt  (III)  were  d-[Co  c?-pn3]Cl3  and  l-[Co  Z-pn3]Cl3 
and  the  totally  inactive  racemic  mixture  of  these  two138.  No  effect  of  the 
position  of  the  methyl  "roups  could  he  detected.  Here  again  the  asymmetry 
of  the  coordinating  group  exerts  an  effect,  presumably  steric,  on  the  com- 
plex formed  by  cobalt  (I  II)  ion.  It  was  shown  (Fig.  8.35)  that  theoretically 
there  are  two  stereoisomers  for  each  of  the  complexes  [M-lll]  and  [M-ddd] 
(depending  on  whether  the  angular  methyl  groups  all  lie  near  the  same 
plane  or  whether  two  are  near  one  plane  and  the  third  is  further  removed 
from  it).  The  exact  nature  of  the  stereoisomeric  forms  of  the  two  stable 
isomers  are  not  known. 

The  only  conclusive  proof  of  isomerism  due  to  the  position  of  the  methyl 
group  of  the  propylenediamine  molecule  was  made  by  Werner  and  Smir- 
noff32 on  the  complex  cis-[Co  en  pn  (N02)2]X  (Fig.  8.15)  (page  286). 

A  similar  compound  containing  two  active  propylenediamine  molecules 
has  been  investigated  by  Hurlimann49  and  by  Watts107.  The  cis  modification 
of  this  ion,  [Co(d  or  /-pn)2(N02)2]+,  should  exist  in  twelve  forms  as  shown 
in  Fig.  8.35.  They  were  able  to  isolate  only  two  active  forms  and  concluded 


,N02 
'N02 

OIL. 

Fig.  8.35.  Possible  forms  of  cis-[Co  pn2  (NO«)a]+ 

that  the  position  of  the  methyl  groups  is  immaterial,  because  except  for 
these  groups,  the  three  position  isomers  for  [Co(/-pn)2(N02)2]+  or 
[Co(d-pn)2(N02)2]+  are  identical.  The  work  of  O'Brien,  McReynolds,  and 
Bailar50  casts  some  doubt  on  this  interpretation. 

Complex  Compounds  Containing  Polydentate  Donor  Molecules. 
Compounds  containing  polydentate  coordinating  groups  have  received 
only  limited  attention,  but  some  of  them  have  been  shown  to  be  optically 
active.  A  typical  example  of  a  tridentate  molecule  may  be  furnished  by 
a,j8,7-triaminopropane  which  was  investigated  by  Pope  and  Mann41b- 139. 

138.  Tschugaeff  and  Sokoloff,  Ber.}  42,  55  (1909);  Lifschitz  and  Rosenbohm,  Z.  wiss. 
Phot.,  19,  209.  211  (1920). 


STEREOISOMERISM  OF  HEXACOVALHXT  ATOMS 


319 


The  triamine  is  capable  of  displacing  the  ammonia  molecules  from  hexam- 
mine  complexes  to  yield  the  cation  containing  two  moles  of  the  organic 
amine,  [M(AA'A)2]+++.  Such  a  complex  may  possibly  exist  in  three  isomeric 
forms;  (I)  is  symmetrical  and  inactive  while  (II)  and  (III)  are  asymmetric 
and,  therefore,  optically  active  (Fig.  8.30).  Isomer(III)  may  appear  to 


I  II 

Fig.  8.36.  Possible  forms  of  [M(AA'A)2] 


III 


be  symmetrical,  but,  on  further  consideration,  it  can  be  seen  that  the 
lateral  displacement  of  the  central  atom  in  triaminopropane  destroys  the 
symmetry  of  the  complex.  Attempts  to  isolate  these  three  isomers  of  the 
cobalt  (III)  ion  were  not  successful  and  only  the  inactive  form  (I)  was  ob- 
tained. A  consideration  of  the  scale  model  of  this  complex  tends  somewhat 
to  clarify  these  results.  It  is  seen  that  it  is  sterically  impossible  for  the  tri- 
aminopropane molecule  to  occupy  three  positions  along  the  edge  of  an 
octahedron  since  the  five-membered  chain  which  includes  the  1  and  3 
amine  groups  is  by  no  means  of  sufficient  length  to  span  the  trans  positions. 
If  this  were  not  true,  trimethylenediamine  should  be  capable  of  spanning 
the  trans  positions.  The  shortest  chain  which  has  given  any  evidence  in- 
dicative of  such  behavior  contains  seven  members  (pages  259  and  277). 
This  factor  eliminates  the  possibility  of  attaining  structure  (III).  Models 
also  indicate  considerable  strain  when  the  base  behaves  in  a  tridentate 
manner  with  its  functional  groups  distributed  at  the  corners  of  an  octa- 
hedral face.  It  might  be  suspected  that  the  bonds  in  the  molecule  are  sub- 
ject to  sufficient  strain  to  allow  rapid  racemization  of  the  structure  (II), 
if  it  is  formed,  by  an  intramolecular  rearrangement  mechanism.  Pope  and 
Mann  were  able  to  obtain  slight  evidence  for  the  existence  of  the  active 
forms  by  repeated  crystallization  of  the  c?cx£ro-camphor-7r-sulfonate,  which 
gave  a  very  faintly  active  chloride.  The  activity  of  this  small  quantity  fell 
rapidly  to  zero  and  the  final  compound  was  always  homogeneous  and  in- 
active. 

The  researches  of  Morgan  and  Main-Smith110  with  ethylenediamino-bis- 
(acetylacetone), 

CH8C(OH)=CHC(CH3)=X— CH2CH2— N=C(CH3)CH=C(OH)CH3 , 

139.   Pope  and  Mann,  Proc.  Roy.  Soc,  London,  109A,  444  (1925). 
1  IC  Morgan  and  Main-Smith,  ./.  Chem.  Soc,  127,  2030  (1925). 


320  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

can  be  used  to  illustrate  the  isomerism  resulting  from  a  tetradenate  chelat- 
ing agent.  The  complex,  ICofXII^OioHisC^^lCl,  may  exist  in  five  stereo- 
chemical arrangements  (Fig.  S.37).  The  complex  ion  with  two  ammonia 


NH3  rNf^"\^lNH3 

NH, 

NH3 

OIL 

I  II  III 

Fig.  8.37 

groups  in  the  trans  positions  (I)  has  a  plane  of  symmetry  and  is,  therefore, 
inactive.  If  the  two  ammonia  groups  are  cis  to  one  another,  the  tetradentate 
molecule  can  arrange  itself  so  that  the  terminal  oxygen  groups  are  opposite 
(II),  or  adjacent  to  each  other  (III)  and,  in  addition,  each  of  these  can  exist 
in  mirror  image  forms.  Morgan  and  Main-Smith  were  able  to  obtain  all 
five  isomers  b}/  careful  fractional  crystallization  of  the  dextro-csLmphor-w- 
sulfonate.  The  optically-active  forms  slowly  changed  into  the  trans  isomer 
and  all  attempts  to  separate  a  resolvable  material  from  it  failed.  It  was 
believed  that  this  may  result  from  a  seeding  of  the  more  stable  trans  form 
but,  the  authors  were  also  unable  to  repeat  this  separation  in  a  different 
laboratory  with  new  equipment.*  Basolo141  has  studied  a  tetradentate  co- 
ordinating agent,  triethylenetetramine, 

NH2CH2CH2NHCH2CH2NHCH2CH2NH2 . 

Several  cobalt(III)  salts  containing  this  tetramine  were  isolated  but  none 
could  be  resolved  due  to  poor  solubility  relationships.  However,  Das 
Sarma141a  has  obtained  the  dichloro  complex,  [Co  trien  CyCl,  in  optically 
active  forms. 

Busch  and  Bailar143  have  resolved  [Co  enta  Br]=  and  [Co  enta]~,  in 
which  the  ethylenediaminetetraacetate  ion  is  pentadentate  and  hexaden- 
tate,  respectively.  Dwyer  and  Lions37, 39a- 39b>  141, 143  have  conclusively 
shown  that  3,6-dithia-l  ,8-bis(salicylideneamino) octane  and  its  derivatives 

:  Although  octahedral  complexes  involving  linear  tetradentate  chelating  agents 
1  heorel  ically  can  exist  in  the  five  stereochemical  tonus  shown  in  Fig.  8.37,  the  Fisher- 
Hirschfelder  models  indicate  thai  structures  II  and  III  involving  ethylenediamine- 
etylacetone)  would  be  badly  strained  as  a  result  of  the  restricted  rotation  de- 
rived from  t  he  double  bonds. 
1  II.   Basolo,  ./.  .1///.  Chi,,,.  Soc,  70,  2346  (1948 
ilia.  Das  Sarma.  and  Bailar,  ibid.,  77,  5480  (1965). 
I  12.  Dwyer  and  Gyarfas,  Nature,  168,  29  (1951). 
143.   Busch  and  Bailar,  J.Am.  Chem.  Soc,  75,  [574  (1953). 
143a.  Das  Sarma  and  Bailar,  ibid.,  76,  4051  (1954). 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS 


321 


can  function  as  hexadentate  chelating  compounds  in  one  or  another  of  two 
enantiomorphous,  strainless  configurations.  The  cobalt(III)  cation, 

[C0(C22H22N202S2)]H 

was  resolved  by  means  of  the  dex^ro-bromocamphor-x-sulfonate  and  the 

molecular  rotation  dig  green  line)  was  ±50,160°.  Solutions  of  these  salts 
can  be  boiled  for  twenty  minutes  without  racemization.  Das  Sarma  and 
Bailar148-  have  reported  the  resolution  of  the  cobalt(III),  iron(III)  and 
aluminum(III)  complexes  of 


OH 


HO 


CH=NCH,CH2NHCH2CH,NHCH2CH2N==HC 

Polynuclear  Complex  Compounds.  Most  of  the  work  with  poly- 
nuclear  complexes  was  done  by  Werner  who  isolated  the  first  optically- 
active  dinuclear  compound, 

/   \ 

en2  Co111       CoIven. 


Since  the  two  portions  of  the  ion  were  different  (Co(III)  and  Co(IY)), 
four  different  optically-active  isomers  should  be  possible;  d-[Co(III)]  and 
d-[Co(IV)];  l-[Co(III)]  and  l-[Co(IV)];  d-[Co(III)]  and  l-[Co(IV)]; 
l-[Co(III)]  and  d-[Co(IV)].  On  the  basis  of  the  modern  concept  of  resonance, 
the  last  two  combinations  are  the  same,  which  means  that  there  are  really 
only  three  possibilities.  Werner  succeeded  in  obtaining  only  two  of  these, 
one  in  which  both  the  cobalt  atoms  were  dextro  and  the  other  in  which  both 
the  cobalt  atoms  were  levo  rotatory.  The  optically-active  antipodes  have 
large  rotations,  ([a]D  =  ±815°  and  [a]E  =  ±1200°),  and  are  rather  stable 
although  the  active  cation  is  completely  racemized  after  some  weeks. 
Werner  suggests  that  the  valence  of  the  central  atom  has  a  marked  influence 
on  the  magnitude  of  optical  rotation,  basing  his  suggestion  on  the  fact  that 
the  specific  rotation  of  similar  dinuclear  complexes  containing  two  ('o(III) 
atom-  is  considerably  less.  The  data  available  are  insufficient  to  support 
his  postulate. 

It  can  readily  be  seen  that,  had  the  asymmetric  centers  in  the  above  com- 

144.  Werner,  Ann.,  375,  70  (1910);  Werner,  Ber.,  47,  1961  (1914). 


322 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


pound  been  structurally  similar,  there  should  exist  an  internally  compen- 
sated or  meso  form  as  well  as  the  dextro  and  levo  rotary  isomers.  Such  a 
binuclear  complex  would  be  analogous  to  tartaric  acid  amongst  the  active 
carbon  compounds.  The  resolution  of  a  complex  of  this  type 


Ml 


en->  Co 


Co  crif 


NO; 


Hn 


was  studied  by  Werner145.  Fractional  crystallization  of  the  dexlro-a-hromo- 
camphor-x-sulfonate  yielded  dextro  and  levo  rotary  compounds  which  gave 
a  true  racemate  when  equimolecular  quantities  of  the  enantiomorphs 
were  combined.  This  racemate  differed  from  a  third  optically-inactive 
isomeride,  which  must  have  been  the  meso  complex  (Fig.  8.38).  The  pres- 


MESQ 
Fig.  8.38.  Possible  stereochemical  forms  of  a  dinuclear  complex 

ence  of  this  meso  form  was  used  by  Werner  to  show  that  the  bridging  bonds 
between  the  two  cobalt(III)  ions  are  the  same. 

Purely  Inorganic  Complex  Compounds.  Although  Werner  success- 
fully resolved  compounds  of  the  types  [M(AA)3]  and  [M(AA)2a2]  in  which 
the  optical  activity  could  be  ascribed  to  an  octahedral  spatial  arrangement, 
some  of  his  contemporaries  objected  to  this  interpretation  on  the  basis 
that  these  compounds  contained  carbon  atoms.  It  is  now  clear  that  the 
organic  compounds  in  these  complexes  could  not  be  responsible  for  the 
observed  optical  activity,  but  at  that  time  it  was  necessary  for  Werner  to 
resolve  a  purely  inorganic  complex  in  order  to  establish  his  theory.  This  was 
successfully  accomplished  in    L914   by  the  resolution  of  the  tetranuclear 


1  16.   Werner,  Her.,  46,  3674  (1913), 


STEREOISOMERISM  OF  IIEXACOVALENT  ATOMS 


323 


complex, 


x6 


The  compound  was  prepared  b}'  the  action  of  ammonia  on  chloroaquo- 
tetramminecoball  (III)  chloride5  and  is  analogous  to  the  tris(ethylenedia- 
mine)  salts  with  the  bidentate  group  being 


(NH8)4Co 


OHN 


OHy 


(Fig.  8.39).  The  racemic  mixture  was  resolved  by  means  of  dextro-a-bromo- 


(N  H3)4Co-o 


;C0CNH3)4       (NH3)4C0( 


CNH3)4C0-0H 


A>1 

Fig. 


.39 


O  Co(NH3)4 


camphor-7r-sulfonate,  which  yielded  the  levo  rotary  ion  in  the  less  soluble 
fraction.  The  optically-active  antipodes  undergo  rapid  racemization  and 
their  rotation  is  best  studied  in  mixtures  of  water  and  acetone.  A  very 
high  molecular  rotation  ([M]56oo)  of  —  47,  610°  wras  obtained. 

Only  one  other  purely  inorganic  complex  compound  has  been  resolved 
into  its  optically-active  antipodes.  Mann122  has  successfully  resolved  cis- 
Xa[Rh(SOoX2H2)2(H20)2]  into  optical  isomerides  having  [M]578o  ±  31-34°, 
by  means  of  rf-phenylethylamine.  It  has  been  shown  that  sulphamide, 
S02(XH2)o ,  like  dimethylglyoxime146,  will  occupy  only  four  positions  in 
the  complex  of  a  hexacovalent  element. 

Optical  Activity  of  Coordinated  Atoms.  It  is  sometimes  possible  for 
an  atom  of  a  donor  molecule  to  be  rendered  optically  active  because  the 
molecule  is  coordinated  to  a  central  ion. 

Nitrogen.  Meisenheimer,  Angermann,  Holsten,  and  Kiderlen147  demon- 
strated the  tetrahedral  nature  of  the  nitrogen  atom  by  resolving  (sarco- 

146.  Tschugaeff,  Z.  anorg.  Chem.,  46,  144  (1905);  Tschugaeff,  Ber.,  39,  2692  (1906); 

Tschugaeff,  ibid.,  40,  3498  (1907);  Tschugaeff,  ibid.,  41,  2226  (1908). 

147.  Meisenheimer,  Angermann,  Holsten,  and  Kiderlen,  Ann.,  438,  217  (1924). 


324 


(  ItEUlSTRY  OF  THE  COORDINATION  COMPOUNDS 


-iiM"l)is-(ethylenediamine)cobalt(III)  chloride  into  more  than  two  opti- 
cally-active isomers. 


/ 


M 


o— c 

/     \ 

en2  Co  CH2 


NH 


Clo 


__  CH3  __ 

In  this  case,  the  complex  is  itself  optically  active,  and  the  nitrogen  atom 
acts  as  a  secondary  source  of  optical  activity,  so  that  there  should  be  four 
active  forms  of  this  complex  ([Co  +  N  +],  [Co  +  N  -],  [Co  -  N  +], 
and  [Co  —  N  —  ]).  Fractional  crystallization  of  the  ckriro-a-bromocam- 
phor-7r-sulfonate  gave  indication  that  these  forms  exist.  One  fraction, 
believed  to  be  [Co  +  N  d=],  had  a  rotation  of  [M]D  =  +2020°  and  further 
recrystallization  of  this  fraction  gave  a  slightly  soluble  portion,  [Co  +  N  +], 
with  a  rotation  of  [M]D  =  +2290°  and  a  more  soluble  portion  [Co  +  N  —  ], 
with  a  rotation  of  of  [M]D  =  +1775°.  The  rotation  of  [Co  +  N  +]  de- 
creased rapidly  to  approximately  the  orginal  value  while  that  of 
[Co  +  N  -]  increased  only  to  [M]D  =  +1825°. 

An  attempt  to  duplicate  Meisenheimer's  results  was  not  successful148. 
Mann40  attempted  unsuccessfully  to  resolve  the  complex 


Pt 


NH2— CH2 


NH — CH; 


Cls 


CH2CH2NH2HC1. 

in  which  the  only  source  of  optical  activity  is  the  asymmetric  nitrogen  atom. 
Since  the  compound  could  not  be  resolved,  it  was  suggested  that  other 
polyamines  such  as  /3/3'-diaminoethylmethylamine  and  /3-aminodiethyl- 
methylamine  be  used.  In  these  compounds  the  asymmetric  nitrogen  is 
part  of  a  tertiary  amine  group  and  should,  therefore,  possess  much  greater 
optica]  stability  than  the  secondary  amine  compounds.  At  the  same  time 
the  coordination  of  the  1  ciliary  amine  group  should  be  greatly  strengthened 
by  the  chelate  ring  of  which  this  group  is  a  part.  No  report  on  the  results 
of  this  work  seems  to  have  been  published. 


]  Is.   Baaolo,  thesis,  University  of  Illinois,  10-13. 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS  325 

Kuebler  and  Bailar148  have  prepared  and  investigated  potassium  dinitro- 
(N-methyl-N-ethylgrycine)platinate(II),  and  have  demonstrated  the  ex- 
istence of  an  asymmetric'  optically-active  nitrogen  atom  in  this  compound 
through  its  resolution  by  fractionation  with  l-quinine  and  also  by  adsorp- 
tion on  optically-active  quartz  powder.  It  should  be  noted  thai  X-methyl- 
N-ethylglycine  differs  from  sarcosine  in  having  no  hydrogen  atom  attached 
directly  to  the  nitrogen.  Part  of  the  difficulty  encountered  with  the  sarcosine 
complex  may  result  from  the  dissociation  of  the  hydrogen  atom  from  the 
nitrogen  (Chapter  12),  thus  allowing  racemization. 

Sulfur.    Tetrachloro(thiodiethylenediamine-N  ,S)platinum(IV)    hydro- 

NII, 
Cl4Pt 

s 

I 

CH2CHoNH2HCl. 

chloride,  is  an  example  of  a  complex  in  which  the  optical  activity  is  due  to 
an  element  linked  to  the  central  atom40b.  The  sulfur  atom  in  the  original  di- 
aminodiethylsulfide  molecule  has  become  asymmetric  by  the  process  of 
coordination  and  is  now  stereochemical^,  and  probably  electronically, 
identical  with  the  sulfur  atom  in  the  asymmetric  sulfoxides,  such  as 
p-amino-p-methyl-diphenyl  sulfoxide  which  has  been  resolved  by  Harrison, 
Kenyon  and  Phillips150. 

Racemic  Modifications 

Racemic  modifications  are  obtained  by  mixing  equal  amounts  of  the 
enantiomorphs,  by  chemical  syntheses,  or  by  racemization  of  an  optically- 
active  material. 

Optically-active  inorganic  complex  compounds  are  generally  optically 
unstable,  and  can  easily  be  racemized.  The  process  of  racemization  implies 
conversion  of  one  form  to  the  other  until  the  dextro  and  levo  isomers  are 
present  in  equal  amounts.  Two  theories  have  been  proposed  to  explain  the 
mechanism  of  such  a  conversion  in  coordination  compounds:  Dissociation 
and  intramolecular  rearrangement. 

Dissociation  Theory  of  Racemization.  Most  of  the  experiments  re- 
lated to  racemization  studies  have  involved  the  trisoxalato  anions.  The 
theory  of  racemization  by  dissociation118  assumes  that  an  oxalate  ion  dissoci- 
ates from  the  complex;  the  residue,  according  to  Thomas157,  undergoes  re- 
orientation to  ;i  planar  distribution  of  the  four  coordinated  groups;  and, 

140.  Kuebler  and  Bailar,  ./.  Am.  Chem.  Soc,  74,  3535  (1952). 
150.  Harrison,  Kenyon,  and  Phillips, ./.  Chem.  Soc,  1926,  2079. 


326 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


upon  recombination  of  the  third  oxalato  group,  the  original  configuration 
and  its  mirror  image  are  formed  with  equal  probability  (Fig.  8.40).  Thomas 


c2o4 


a 


+c2o< 


€ 


Fig.  8.40 


£*°4    -,  = 


czo4 


based  this  theory  on  the  fact  that  the  addition  of  silver  nitrate  to  a  solu- 
tion of  [Fe(C204)3]-  gives  an  immediate  precipitate  of  silver  oxalate,  but 
when  silver  nitrate  is  added  to  [Cr(C204)3]-  the  precipitate  forms  only  on 
long  standing151.  Other  investigators  have  shown  that  the  precipitate  so 
obtained  is  not  silver  oxalate  but  is  Ag3[M(C204)3]-6H20152  or  KnAgm- 
[M(C204)3]-:cH20153.  The  conductivity  experiments  of  Thomas  and  Fraser154 
could  not  be  checked  by  Johnson155. 

Numerous  investigations  have  been  made  to  establish  conclusively  that 
the  dissociation  theory  does  not  adequately  account  for  the  racemization 
of  the  tris(oxalato)  complexes  of  cobalt  (III)  and  chromium(III).  For 
example,  in  no  case  could  free  oxalate  ion  be  detected  in  solutions  of  tris- 
oxalatochromium(III)  or  cobalt(III)  salts,  nor  was  it  possible  to  change 
the  rate  of  racemization  of  these  active  substances  by  the  addition  of  the 
common  oxalate  ion156.  Johnson  and  Mead157  were  able  to  show  that  these 
salts  racemize  even  in  the  crystalline  state.  Finally  the  fact  that  the  dissoci- 
ation theory  is  not  correct  was  conclusively  demonstrated  by  using  oxalate 
containing  radioactive  carbon  and  determining  the  amount  of  oxalate  ex- 
change in  solutions  of  these  compounds.  If  this  theory  is  correct,  the  rate 
of  racemization  should  parallel  the  rate  of  interchange.  However,  Long158 
was  able  to  detect  no  exchange  although  the  active  complex,  K3[Cr(C204)3], 
was  slowly  being  racemized.  A  similar  study  using  inactive  [Fe(C204)3]^ 
and  [A1(C204)3]-  resulted  in  a  very  rapid  exchange,  which  implies  that  opti- 
cal activity  in  these  compounds  is  very  unlikely159.  Mathieu71  has  investi- 

151.  Thomas,  J.  Chem.  Soc,  121,  196  (1922). 

152.  Kistiakowsky,  Z.  physik.  Chem.,  6,  96  (1890). 

153.  Kranig,  Ann.  chim.,  11,  44  (1929). 

154.  Thomas  and  Frazer, ./.  Chem.  Soc,  123,  2973  (1923). 

155.  Johnson,  Trans.  Faraday  Soc,  31,  1615  (1935). 

156.  Beese  and  .Johnson,  Trans.  Faraday  Soc,  31,  1635  (1935);  Bushra  and  Johnson, 

./.  Chem.  Soc,  1939,  1911. 

157.  Johnson  and  Mead,  Trans.  Faraday  Soc,  31,  1621  (1935). 

158.  Long,  ./.  Am.  Chem.  Soc,  61,  570  (1939). 
l.v.i.  Long,  ibid.,  63,  1353  (1941). 


STEREOISOMERISM  OF  HEXACOYALEXT  ATOMS 


327 


gated  the  rate  of  change  of  optical  rotation  of  a  solution  of  dextro- 
[Co  en2  Cl2]+.  He  observed  that  the  optical  rotation  changed  to  a  fairly 
constant  value  at  the  same  rate  that  chloride  ion  was  formed.  The  resulting 
[Co  en2  H20  Cl]++  ion  then  racemized  at  a  rate  independent  of  the  rate  of 
formation  of  the  diaquo  complex.  On  the  basis  of  these  results  it  was  sug- 
gested that  the  racemization  of  [Co  en2  H2O  Cl]++  may  occur  as  a  conse- 
quence of  the  dissociation  of  the  coordinated  water  molecule  (Fig.  8.41). 


1   +-HP0 


en  1  "H2Q 


CI 


HoO, 


Co 


en 


en 


LEVO  ACTIVATED  DEXTRO 

INTERMEDIATE 

Fig.  8.41.  Racemization  of  [Co  en2  (H20)C1]++ 


Mathieu  observed  that, the  analogous  complex  [Co  en2  H20  N02]++  does 
not  racemize,  even  upon  standing  in  solution  for  several  months.  If  one 
assumes  that  the  coordinated  water  dissociates  at  a  measurable  rate82  then 
it  would  appear  that  the  intermediate  in  this  case  has  a  tetragonal  pyramid 
configuration  (Fig.  8.42)  instead  of  the  trigonal  by-pyramid  structure. 


NCfel 


Fig.  8.42 

It  has  recently  been  shown161  that  the  rate  of  racemization  of  dextro- 
[Co  en2  C1JC1  in  methanol  is  equal  to  the  rate  of  radio-chlorine  exchange. 
Therefore,  racemization  is  thought  to  occur  through  a  symmetrical  penta- 
covalent  intermediate. 

Failure  of  the  presence  of  excess  2,2'-dipyridyl  to  alter  the  rate  of 
racemization  of  [Xi(dipy)3]++  162  and  of  excess  1 ,  10-phenanthroline  to  effect 

160.  Stone,  thesis,  Northwestern  University,  1952. 

161.  Brown  and  Nyholm,  /.  Chem.  Soc,  2696  (1953). 

162.  Schweitzer  and  Lee,  ./.  Phys.  Chem.,  56,  195  (1952). 


L 


328 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Table  8.8*.  Racemization    wi>  Dissociation  Rates  of  Some  Nickel(II) 
and  Irox(II)  Complexes 


,«u  1 

Q 

Ra<  emization 

Dissociation 

k  (min->)  25° 

Ea 

K.;,l. 

AS1  E.U. 

k  (min-i)  25° 

Ea 

Krai. 

ss1  E.U. 

[Ni(o  phen),]++ 
[Ni(dipy)3]++ 
[Fe(o-phen)3]++ 
[Fe(dipy)3]+f 

18" 
I.V 
21* 
16" 

6.3  X    10    '  ,l 

1.4  X  10-1  c 

4.0  X  10  -  ■ 
3.6  X  10  a  ' 

25 
22 
31 

28 

+  1.8 
+2.7 

30 

21 

6.3  X  10~4d 

1.4  X  10"ld 

4.5  X  10-3  f 
7.3  X  lO-3^ 

25 
22 

26 

+  1.8 

+2.7 

+  10 

*  The  values  tabulated  for  the  rates  of  racemization  are  in  a  form  allowing  direct 
comparison  with  the  rates  of  dissociation  and  hence  are  twice  the  values  reported  by 
Davies  and  Dwyer. 

a  Davies  and  Dwyer,  Trans.  Faraday  Soc,  49,  180  (1953). 

b  Lee,  KolthofT,  and  Leussing,  J.  Am.  Chem.  Soc,  70,2348  (1948). 

c  Boxendale  and  George,  Nature,  162,  777  (1948). 

(1  Basolo,  Hayes  and  Neumann,  J.  Am.  Chem.  Soc.,  75,  5102  (1953). 

p  Schweitzer  and  Lee,  ./.  Phys.  ('hem.,  56,  195  (1952). 

f  Brandt  and  Gullstrom,  J.  Am.  Chem.  Soc.,  74,  3532  (1952). 

b  Baxendale  and  George,  Trans.  Faraday  Soc,  46,  55  (1950). 


the  racemization  of  [Xi(o-phen)3]++  163  has  recently  been  cited  in  support 
of  an  intramolecular  process.  However,  it  does  not  necessarily  follow  that 
an  excess  of  the  chelating  agent  should  decrease  the  rate  of  racemization. 
There  would  certainly  be  no  change  in  the  rate  of  racemization  if  the  dis- 
sociated product  were  either  symmetrical  and  thus  optically  inactive  or 
if  it  lost  its  optical  activity  very  rapidly.  In  fact,  Basolo,  Hayes  and  Neu- 
mann164 have  recently  observed  that  the  rates  of  racemization  of  these 
nickel (II)  complexes  are  the  same  as  the  rates  of  dissociation.  The  energy 
of  activation  is  identical  and,  as  is  apparent  from  the  data  summarized  in 
Table  8.8,  the  two  processes  are  the  same.  The  data  available  for  the 
analogous  iron(II)  complexes  are  included  in  Table  8.8  so  that  all  of  these 
may  be  conveniently  compared.  The  racemization  of  these  iron(II)  com- 
pounds must  involve  an  intramolecular  process  at  least  in  part164b.  It  is  in- 
teresting to  speculate  why  the1  mechanism  of  racemization  of  the  nickel(II) 
complexes  differs  from  that  of  the  iron(II)  compounds.  The  charges  on  the 
cations  arc  the  same  and  their  sizes  must  be  practically  identical.  The 
paramagnetism  of  [\i(dipy);{|++  suggests  sp3d?  type  hybridization  as  com- 
pared to  d?sp*  for  diamagnetic  |Fe(dipy)3]++.  The  more  labile  outer  orbital 
oickel(II)   complex166  may   be  expected  to  dissociate  fairly  readily  and 


L63.  I  )avies  and  Dwye 
L64.  Basolo,    Hayes,   a 

3807  '  L954). 
L66.  Taube.  Chem.  Rev, 


r,  Trans.  Faraday  Soc,  48,  244  (1952);  ibid,  49,  180  (1953). 
Qd    Neumann,   ./.    .1///.    Chem.   Soc,   75,   5102    (1953);    76, 

,50,69  (1952). 


STEREOISOMERISM  OF  HEX ACOVALEXT  ATOMS 


329 


therefore  possibly  racemize  by  such  a  mechanism.  However,  this  inter- 
pretation is  qo1  Gompatible  with  the  fad  that  [Fe(o-phen)3]++  dissociates 
faster  than  [NiCo-phen^]"*"*". 

Intramolecular  Rearrangement  Theory,  if  the  complex  does  not 
undergo  dissocial  ion,  the  racemization  must  result  from  an  intramolecular 
rearrangement.  Werner186  was  the  first  to  suggest  such  a  mechanism,  stat- 
ing that  trioxalatochromate(III)  ions  lose  their  rotatory  power  through  the 
momentary  vacation  of  one  coordination  position  by  an  oxalate  radical, 
thus  permitting  a  rearrangement  of  positions  as  it  becomes  attached  again 
Fig.  8.43).  Bushra  and  Johnson167  have  pointed  out  there  is  no  apparent 


AA, 


AA 


AA 


AA 


Yl 


AA 


AA 


DEXTRO 


LEVO 


Fig.  8.43.  Racemization  of  [M(C204)3]"  (Werner) 

racemization  of  [Co  en3]+++  whereas  [Co(C204)3]=  racemizes  at  a  measur- 
able rate,  thus  indicating  that  the  cobalt-ethylenediamine  chelate  ring  is 
not  opened  as  readily  as  the  cobalt-oxalate  ring.  They  suggest  that,  if  only 
one  chelate  ring  need  open  to  allow  racemization,  one  may  expect  the  com- 
plex [Co  en2  C204]+  to  racemize.  However,  the  loss  of  optical  activity  of  this 
compound  was  found  to  result  from  its  decomposition  rather  than  from 
inversion.  Although  the  complex  [Co  en  (C204)2]_  was  not  obtained,  the 
analogous  chromium  (III)  compound  did  racemize  and  with  an  activation 
energy  of  15.8  Kcal,  the  same  as  that  for  the  racemization  of  [Cr^OOs]35. 
On  the  basis  of  these  observations  Bushra  and  Johnson  suggest  that  the 
mechanism  of  racemization  requires  the  opening  of  two  rings  which  can 
reattach  at  the  same  positions  or  at  exchanged  positions  (Fig.  8.44). 


1  "        A  A^ 

!/- 

M 

!    -J 

AA 

\ 

AA 


DEXTRO 


M, 


U 


A  A 


LEVO 


Fig.  8.44.  Racemization  of  [M(C204)3]=  (Bushra  and  Johnson) 

166.  Werner,  Ber.,  45,  3061  (1912). 

167.  Bushra  and  Johnson, J.  Chem.  Soc,  1939,  1937. 


330 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


This  mechanism  of  intramolecular  change  by  opening  two  rings  at  two 
points  of  attachment  in  cis  positions  has  been  questioned  by  Ray  and 
Dutt168.  They  suggest  that  the  momentary  rupture  of  the  chemical  bonds 
at  these  positions  introduces  the  possibility  of  chemical  decomposition 
during  inversion  and  since  all  the  six  bonds  in  an  octahedral  complex  are 
equivalent,  (commonly  the  d2sp*  hybrid  type)  there  is  no  obvious  reason 
why  two  such  bonds  attached  to  one  and  the  same  chelate  group  will  not 
be  ruptured  at  the  same  time.  But  there  is  no  experimental  evidence  that 
chemical  decomposition  is  associated  with  inversion.  Ray  and  Dutt  have 
interpreted  their  kinetic  data  on  the  racemization  of  tris(biguanidinium)- 
cobalt(III)  chloride  in  terms  of  a  mechanism  which  does  not  necessitate 
the  opening  of  any  chelate  rings.  They  point  out  that  the  existence  of  two 
enantiomers  of  the  same  energy  content  indicates  a  potential  barrier  be- 
tween them  and  therefore  some  activation  energy  is  necessary  for  inter- 
conversion.  Addition  of  energy  to  a  molecule  leads  to  an  increase  in  trans- 
lational,  rotational  and  vibrational  motions,  and  the  molecule  is  said  to  be 
activated.  If  sufficiently  excited,  the  normal  octahedral  complex  may  lose 
its  configuration  and  assume  a  metastable  condition.  On  removal  of  the 
excess  energy,  the  molecule  returns  to  the  octahedral  form,  and,  since  the 
two  enantiomers  have  equal  energy  requirements  they  form  with  the  same 
ease. 

This  mechanism  proposed  by  Ray  and  Dutt  is  represented  in  Fig.  8.45. 
The  dextro  form  (I)  changes  to  the  activated  form  (II)  when  the  two  pairs 


A 


I  H  HI 

Fig.  8.45.  Racemization  of  [M(AA)3]  (Ray  and  Dutt) 


of  bonds  holding  y  and  z  rotate  in  opposite  directions  along  their  own  plane 
through  an  angle  of  45°  to  give  a  distorted  octahedron  with  angles  of  90° 
between  the  bonds.  The  distorted  or  activated  molecule  can  then  return  to 
its  normal  state  by  retracing  its  previous  steps  to  give  the  dextro  form  (I) 
or,  by  further  rotation  through  45°  in  the  same  direction,  it  may  degenerate 
to  produce  the  mirror  image  (III). 

Since  the  structure  of  1 ,  10-phenanthroline  does  not  allow  an  open  ring 
structure,  there  is  reason  to  feel  that  [Fe(o-phen)3]++  must  racemize  by  some 
process  of  this  type163- 164. 

168.  Ray  and  Dutt,  J.  Indian  Chem.  Soc,  20,  81  (1943). 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS  331 

Resolution  of  Racemic  Modifications.  The  problems  encountered 

and  methods  employed  in  the  resolution  of  complex  inorganic  compounds 
are  much  the  same  as  those  used  with  organic  compounds.  No  doubt  the 

biggest  difference  i>  the  tact  that  biochemical  processes,  commonly  used 
for  the  resolution  of  organic  compounds,  have  not  been  applied  to  coordi- 
nation compound.-. 

Spontaneous   Crystallization   of  the   Antipodes.    The    mechanical 

separation  of  crystals,  as  used  in  1848  by  Pasteur169  for  the  separation  of 
<l  and  /  forms  of  sodium  ammonium  tartrate,  has  been  used  for  a  few  com- 
plex compounds.  Since  most  complex  salts  form  well  defined  crystals,  it  is 
not  surprising  that  resolution  can  be  realized  by  this  method.  However, 
because  of  the  skill  and  patience  required  to  grow  suitable  crystals,  as  well 
as  the  tedious  operation  of  picking  out  the  different  types,  such  a  procedure 
is  not  practical.  It  might  be  mentioned  that  in  such  a  process  the  racemic 
crystals  must  possess  the  requisite  hemihedrism  by  which  they  may  be 
distinguished,  and  crystallization  must  yield  a  racemic  mixture  rather  than 
a  racemic  compound  or  solid  solution. 

This  method  of  spontaneous  crystallization  of  the  antipodes  from  the 
racemoid  was  first  demonstrated  with  K3[Co(C204)3]170.  A  comparison  of 
the  solubilities  of  the  racemic  compound  and  the  racemic  mixture  at  various 
temperatures  (Fig.  8.46)  demonstrates  that  the  optically-active  salts  are 
the  more  stable  phases  with  respect  to  the  racemoid  at  all  temperatures 
above  13.2°.  This  is,  therefore,  the  maximum  temperature  for  the  forma- 
tion of  the  racemate;  the  reaction  taking  place  may  be  represented  as 


-Jo 

as 


13.2°  TEMP.,°C 

Fig.  8.46.  Solubility  of  potassium  tris-oxalato  cobaltate(III) 

2K3[Co(C204)3]-3KH20  ~^=±  rf-[K,Co(C204)3]-H,0  + 

MK.3Co(C204)3]HoO  +  5H20 

The  antipodes  may  be  allowed  to  crystallize  at  temperatures  above  13.2° 
after  which  they  may  be  separated  mechanically.  Jaeger93  has  also  been 
able  to  obtain  a  racemic  mixture  of  [Rh  cptn3](C104)3-  12H20  at  tempera- 

169.  Pasteur,  Ann.  chim.  phys.,  [37]  24,  442  (1848). 

170.  Jaeger,  Rec.  trav.  chim.,  38,  250  (1919). 


332  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

lures  below  48°  and  to  .sort  the  octahedral  crystals  into  the  dextro  and  levo 
rotatory  forms. 

Preferential  Crystallization.  A  much  more  practical  way  of  accom- 
plishing a  direct  separation  of  the  enantiomorphs  in  a  racemic  mixture  is 
to  cause  one,  but  not  both,  of  the  forms  to  crystallize.  The  principle  in- 
volved is  analogous  to  that  of  causing  crystals  to  deposit  from  any  super- 
saturated  solution  by  the  addition  of  a  seed  crystal  of  the  desired  material, 
or  of  any  isomorphous  crystal.  This  procedure  was  used  successfully  by 
Werner  and  Bosshart171  in  the  resolution  of  [Co  en2  C204]+,  [Cr  en2  C204]+ 
and  [Co  en2  (N02)2]+.  They  were  able  to  show  that  if  a  crystal  of 
d-[Co  en2  C204]+  is  added  to  a  concentrated  solution  of  c?Z-[Co  en2  C204]+ 
followed  by  an  immediate  addition  of  a  small  amount  of  ethyl  alcohol  and 
ether,  a  precipitate  of  d-[Co  en2  C204]+  separates.  The  filtrate  from  this 
precipitate  is  predominantly  /-[Co  en2  C204]+.  A  similar  procedure  was 
used  to  resolve  dl-[Co  en2  (N02)2]+,  indicating  that  this  method  of  resolu- 
tion may  be  rather  general.  It  was  also  demonstrated  that  dl-[Co  en2  (N02)2]+ 
andde-[Cren2  C204]+  can  be  resolved  using  d-[Co  en2  C204]+  as  a  seed  crystal; 
this  would  indicate  that  it  is  not  necessary  to  use  an  antipode  of  the  same 
compound  but  instead  any  isomorphous  crystal  may  be  satisfactory. 

Conversion  to  Diastereoisomers.  The  most  convenient  method  avail- 
able for  the  resolution  of  optically-active  compounds  is  the  conversion  of  a 
racemic  modification  into  diastereoisomers,  which  may  then  be  separated 
by  fractional  crystallization.  The  principle  of  this  method  and  its  limitations 
need  not  be  discussed  since  they  are  analogous  to  those  encountered  with 
organic  compounds.  The  resolution  of  complex  cations  is  accomplished  by 
the  use  of  optically-active  anions  such  as  tartrate,  antimonyl  tartrate, 
o:-bromocamphor-7r-sulfonate,  camphor-7r-sulfonate,  a-camphornitronate 
and  malate;  while  for  complex  anions  one  employs  such  optically-active 
substances  as  strychnine,  brucine,  cinchonidine,  a-phenylethylamine,  mor- 
phine, quinidine  and  cinchonine.  Removal  of  the  resolving  agent  from  the 
desired  antipode  can  be  accomplished  in  various  ways  depending  upon  the 
properties  of  the  individual  complex  and  also  of  the  resolving  agent.  A 
convenient  method  is  the  separation  by  precipitation  which  is  often  in- 
stantaneous and  can  be  carried  out  at  low  temperatures,  therefore  allowing 
a  minimum  amount  of  racemization  to  take  place172.  In  other  cases,  where 
this  is  not  possible,  it  has  been  found  convenient  to  displace  the  resolving 
agenl  by  means  of  an  alcoholic  acidic  or  basic  solution  and  to  extract  the 
resulting  acid  or  base  by  washing  repeatedly  with  alcohol  to  leave  the  solid, 
insoluble  antipode178. 

171.  Werner  and  Bosshart,  Ber.,  47,  2171  (1914). 

172.  Jaeger,  Rec.  trav.  chin,  ,38,  185  (1919). 
17:;.  Bailar,  Inorganic  Synthesis,  2,  223  (1916). 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS  333 

Method  of  "Active  Racemates".  Molecules  of  inverse  configuration 
may  1  e  associated  in  a  crystal  even  though  they  may  not  have  identical 
compositions174.  This  idea  of  the  formation  of  active  racemates  (page  341  | 
has  been  extended  to  provide  a  method  of  resolul ion  of  racemic  substances171 
or  to  separate  conglomerates,  and  to  determine  the  relative  configurations 
of  homeomers  such  as  the  active  trisoxalatocobaltate(III)j  chromate(III)j 
and  rhodate(III)  in  comparison  with  active  ions  such  as  trisoxalatoiri- 
datet  [II).  Thus,  if  the  active  racemate  ar  and  b~  can  exist,  the  addition  of 
the  active  antipode  o+  to  the  racemic  compound  B  (containing  />f  +  b~) 
will  give  a  mixture  of  \>t(a+  +  lr)  +  (I  —  n)B]  where  n  represents  a  frac- 
tion o\  the  total  amount  of  racemate.  Analysis  of  the  active  racemate  would 
then  give  data  on  the  quantity  and  rotation  of  the  fraction  b  .  Since  the 
mother  liquor  from  these  racemates  contains  an  excess  of  b+,  it  too  will 
be  optically  active.  The  success  of  this  method  depends  upon  the  racemate 
separating  as  a  racemic  compound  rather  than  as  a  racemic  mixture  or 
solid  solution. 

Delepine175  verified  this  supposition  by  studying  the  following  systems: 


./-K:;[Rh(C,04);!l 

and 

tH-K.[Ir(C«04)i] 

d-K,[Ir(C204)8] 

and 

r/Z-K3[Co(C204)3 

Z-K.[Ir(C*0«)«] 

and 

^-K,[Co(C204)8] 

d-K,[Ir(C204)3] 

and 

dZ«K8[Cr(C204)8] 

d-K8[Ir(C204)a] 

and 

r//-K.,[Al(C204)3] 

r/-K3[Ir(C204)3] 

and 

.//-K,,[Fe(C,04)3] 

MCo  en3]Br3 

and 

dl-[Rh  en8]Br8 

From  the  results  obtained  it  seems  that  the  simultaneous  crystallization  of 
a  compound  B  with  an  antipode  (a+  or  ar)  of  a  homeomer  A,  should  be 
considered  as  a  sufficient  reason  for  the  existence  of  antipode  B  in  the  mixed 
crystal  and,  consequently,  of  the  occurrence  of  B  in  the  active  forms  b+  and 
b~,  each  enantiomorphic  with  a~  and  a+.  The  subsequent  separation  of 
b+  from  ar  or  of  b~  from  a+  results  in  the  resolution  of  B.  It  may  also  be 
mentioned  thai  these  experiments  did  not  lead  to  the  resolution  of 
V.  (    m;  ^or[Fe(C204)3]=1131,\ 

Preferential  Adsorption  on  Optically -active  Quartz.  Asymmetric, 
nonionic  coordination  compounds  cannol  be  converted  into  diastereoiso- 
mers,  so  this  common  method  of  resolution  is  not  applicable  to  them.  It 
has  been  demonstrated28  '-•''  thai  enantiomorphs  arc  preferentially  adsorbed 
on  optically-active  quartz;  this  technique  was  applied  to  the  resolution  of 

L74.  DeUpine,  Bull.  soc.  chim.,  [4]  29,  056  (1921   . 

175.  Delepine,  Bull.  soc.  chim.,  [57]  1,  125G  (1034). 


334  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

the  aonionic  complex,  [Co(DMG)2NH3  CI].  The  method  has  likewise  been 
used1"'  for  the  resolution  of  the  complex  ion,  cis-[Co  en  (NH3)2C03]+  and 
in  the  resolution  of  K[Pt(N02)2N(CH8)(CtH6)CH2COO]149.  The  resolutions 
in  these  cases  were  not  complete  but  the  method  is  a  useful  tool  for  deter- 
mining the  resolvability  of  certain  coordination  compounds.  It  may  also 
be  useful  in  studying  systems  which  racemize  too  rapidly  to  be  studied  by 
other  methods. 

Equilibrium  Method  of  Resolution.  Resolution  by  the  equilibrium 
method  has  been  used  successfully  for  organic  compounds176,  but  examples 
of  this  type  are  not  well  known  in  the  field  of  inorganic  complex  compounds. 
Since  the  reactions  involved  in  the  production  of  diastereoisomers  of  com- 
plex compounds  are  ionic,  the  reactions  are  instantaneous  and  shifts  in 
equilibrium  arise  from  the  relative  solubilities  of  the  diastereoisomers.  A 
typical  example  is  the  resolution  of  K3[Cr(C204)3]  by  means  of  strychnine116b. 
It  was  found  that  in  an  alcoholic  solution  the  resolution  yielded  only  the 
dextro  rotatory  complex  ion,  while  in  water  only  the  levo  rotatory  antipode 
was  obtained.  The  explanation  must  be  that  in  solution,  and  especially  at 
higher  temperatures,  there  occurs  a  very  rapid  interconversion.  Since  the 
strychnine  salt  of  the  dextro  ion  is  sparingly  soluble  in  alcohol,  it  is  pre- 
cipitated and  causes  a  shift  in  equilibrium  which  is  in  turn  established  by 
the  interconversion  of  the  levo  component.  Continued  concentration  results 
in  additional  deposition  of  the.  less  soluble  antipode  which  is  replenished  by 
interconversion  to  maintain  equilibrium  and  accounts  for  the  fact  that  only 
the  less  soluble  antipode  is  obtained.  In  this  particular  case  the  strychnine 
salt  of  d-[Cr(C204)3]=  is  less  soluble  in  alcohol  while  the  /-[Cr(C204)3]s  salt 
is  less  soluble  in  water.  Dwyer  and  Gyarfas177  have  reported  a  similar  ob- 
servation with  regard  to  the  resolution  of  [Fe(o-phen)3]++.  A  solution  of 
racemic-[Fe(o-phen)z]++  containing  an  excess  of  dextro  antimonyl  tartrate 
slowly  precipitated  the  complex  completely  in  the  form  of  Z-[Fe(o-phen)3] 
c?-(SbOC4H406)2-4H20.  This  was  attributed  to  the  lability  of  the  complex 
which  allowed  the  equilibrium  between  the  dextro  and  levo  cations  to  be 
shifted  toward  the  less  soluble  diastereoisomer  until  finally  none  of  the 
dextro  complex  remained.  The  partial  resolution  of  inorganic  complexes 
by  the  equilibrium  method  has  been  demonstrated  by  Jonassen,  Bailar 
and  Huffmann181.  It  was  found  that  while  both  the  d  and  l  forms  of  dextro- 
tartratobis(ethylenediamine)cobalt(III)  ion,  [Co  em  <7-tart]+,  form  when 
dextro-t&rt&ric  acid  reacts  with  [Co  en2  C03]+,  they  differ  greatly  in  reactiv- 
ity. When  the  mixture  of  the  two  is  shaken  with  ethylenediamine,  a  70  per 
cent  yield  of  dextro-[Co  en:;lf++  is  obtained  and  very  little  of  the  original 
material  can  be  recovered.  Evidently  the  less  reactive  form  changes  to  the 

176.  King,  Ann.  Repts.  Chem.  Soc,  London,  30,  261  (1933). 

177.  Dwyer  and  Gyarfas,/.  Proc.  Roy.  Soc.,N.S.  Wales.,  83,  263  (1950). 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS  335 

more  reactive  as  the  latter  is  used  up;  this  can  be  explained  by  assuming  thai 
the  following  reactions  take  place: 

I    dextro~[Co  en%  d  tarl  1    •  *  U  vo  |(  !o  ens  d  tai -i 

(II)  <fez<ro-[Co  on,  d-tart]+  +  en  -»  d«s*ro-[Co  cn3]*+f  +  tart 

(III)  few  [Co  (>n..  </  tnrt]f  +  en     ►  levo-[Co  (Mi;i]tM'  4-  tart 

Reaction  (II)  takes  place  more  readily  than  reaction  (III)  and,  therefore 
the  equilibrium  in  reaction  (I)  is  displaced  to  the  left  which  would  account 
for  the  fact  that  an  excess  of  dextro-[Co  en3]+++  is  obtained.  That  this  in- 
terpretation is  not  entirely  justified  has  been  recently  demonstrated132 
by  experiments  which  reveal  that  the  reaction  of  rfex/ro-tartaric  acid  with 
[Co  en»CO»]+  gives  preferentially  the  tfear/ro-cJ-tartrato  complex. 

"Configurational  Activity"  as  a  Method  of  Resolution.  Dwyer  and 
his  coworkers175  have  concluded  from  their  observations  that,  while  the 
addition  of  electrolytes,  such  as  sodium  nitrate,  to  a  pair  of  enantiomeric 
ions  in  solution  alters  the  activity  of  each  enantiomorph  to  the  same  extent, 
the  addition  of  an  electrolyte  containing  an  optically-active  anion  or  cation 
exerts  slightly  different  effects  on  the  two  enantiomeric  ions.  Consequently, 
the  possibility  of  effecting  a  resolution  exists,  and  neither  the  separation  of 
diastereoisomers  nor  the  movement  of  the  equilibrium  position  in  an  opti- 
cally labile  system  is  necessitated. 

Dwyer  has  termed  the  effect  "configurational  activity,"  and  has  dis- 
covered that  the  solubilities  of  d-  and  /-tris(l ,  10-phenanthroline)  ruthe- 
nium(II)  perchlorate  differ  by  as  much  3.5  per  cent  in  dilute  solutions  (1.0 
to  1.5  per  cent)  of  ammonium  c?-bromocamphor  sulfonate  or  sodium  potas- 
sium e?-tartrate.  At  higher  concentrations  of  the  sulfonate  or  tartrate,  the 
solubility  curves  of  the  d-  and  Z-ruthenium(II)  complexes  begin  to  converge, 
probabbr,  according  to  the  authors,  because  "the  normal  nonspecific  ac- 
tivity effect  tends  to  outweigh  the  specific  configurational  effect  at  high 
ionic  strengths." 

The  effect  has  also  been  exhibited  for  tris(2,2'-dipyridyl)nickel(II) 
iodide178,  and  for  the  tris(acetylacetone)cobalt(III)  complex142,  and  Dwyer 
and  his  associates  point  out  that,  since  the  charges  on  a  complex  ion  such 
as  [Fe(CX)G]4~  are  distributed  over  the  peripheral  atoms  of  the  ligands179, 
and  since  the  enantiomers  probably  exhibit  mirror  image  electric  fields 
about  the  antipodes,  the  "configurational  activity"  effect  may  be  due  to 
the  different  interactions  of  the  electric  fields  of  the  dextro  and  levo  forms 
of  the  enantiomeric  pair  with  the  field  of  the  added  optically-active  ion. 

Other  Probable  Methods  of  Resolution.  In  addition  to  the  methods 
of  resolution  which  have  been  used  successfully  for  separating  enantiomers 

178.  Dwyer,  Gyarfas,  and  O'Dwyer,  Nature,  167,  1036  (1951). 

179.  Pauling,  J.  Chem.  Soc,  1948,  1461. 


33G  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

of  coordination  compounds,  there  is  the  probability  that  other  techniques 
may  also  be  applicable.  In  this  connection  some  attention  has  been  devoted 
to  the  influence  of  circularly-polarized  light  on  various  asymmetric  com- 
plex compounds.  Since  it  is  known  that  circularly-polarized  light  is  ab- 
sorbed differently  by  enantiomers,  the  probability  that  the  photochemically 
sensitive  antipodes  present  in  an  optically-active  solution  will  be  decom- 
posed at  different  speeds  by  light  of  that  particular  wave-length  for  which 
absorption  is  an  optimum  has  been  considered.  In  such  a  case  the  solution 
might  be  expected  to  become  slightly  active  and  the  activity  to  be  a  func- 
tion of  time  of  exposure.  Jaeger  and  Berger180  attempted  to  show  that  this 
supposition  is  correct  by  subjecting  both  antipodes  of  K3 [Co (0204)3],  in 
separate  solutions,  to  such  a  radiation  and  in  both  cases  measure  directly 
the  decomposition  velocities.  These  experiments  were  performed  under 
various  conditions,  but  in  no  case  could  a  difference  in  speed  of  decomposi- 
tioD  of  the  dextro  and  levo  components  be  detected. 

It  is  also  possible  that  resolution  of  optically-active  complex  compounds 
can  be  accomplished  by  a  difference  in  rates  of  reaction  of  the  enantiomers. 
Such  a  kinetic  method,  unlike  the  previously  discussed  equilibrium  method, 
does  not  necessarily  involve  intercon version.  In  the  kinetic  method  it  is 
necessary  to  limit  the  amount  of  the  active  compound  used  or  to  stop  the 
reaction  at  a  given  time  before  the  reactions  are  complete.  Although  this 
type  of  resolution  is  applicable  to  relatively  slow  organic  reactions181,  it  has 
not  been  successful  with  the  ionic  reactions  encountered  in  the  production 
of  diastereoisomers  of  inorganic  complexes.  However,  reactions  which  in- 
volve the  displacement  of  groups  coordinated  to  the  central  ion  are  much 
slower,  and  there  is  a  good  probability  that  a  resolving  agent  might  dis- 
place a  particular  coordination  group  from  enantiomers  at  different  rates. 
If  we  recall,  for  example,  the  fact  that  d  and  l  forms  of  [Co  en2  d-tart]+ 
differ  greatly  in  reactivity,  it  would  be  supposed  that  these  cations  are 
formed  from  the  racemic  carbonato  salt  at  different  rates. 

Relative  Configurations  of  Analogous  Enantiomorphs 

Absolute  Configuration.  The  prefixes  dextro  and  levo  as  used  for 
optically-active  compounds  designate  the  direction  of  rotation  only  and  do 
not  supply  any  information  about  the  absolute  configuration  of  the  com- 
pounds. Some  progress  has  been  made  in  determining  absolute  configuration 
by  Kuhn  and  Bein182  with  the  simpler  complexes  of  the  type  [M(AA)3]. 
The  predictions  of  their  theory  agree  with  the  experimental  results  so  it  is 

180.  Jaeger  and  Berger,  /.Vr.  trav.  chim.,  40,  153  (1921). 

181.  Marckwald  and  Paul,  Ber.}  38,810  (1905);  39,  3654  (1906). 

L82.  Kuhn  and  Bein,  Z.  anorg.  Chem.,  216,  321  (1934);  Kuhn  and  Bein,  Z.  physik. 
Chevi.,  24B,  335  (1934). 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS  337 

concluded  that  the  model  presented  corresponds  to  the  absolute  configura- 
tion of  the  molecule.  The  determination  of  the  absolute  configuration  of 
even  the  simplest  antipode  is  extremely  difficult  and  different  theories183 
which  may  appear  logical  sometimes  end  up  assigning  opposite  configura- 
tions to  the  same  enantiomorph.  An  experimental  approach  which  makes 
use  of  x-rays  of  appropriate  wave-length  was  recently  employed  to  deter- 
mine the  absolute  configuration  of  sodium  rubidium  deatfro-tartrate184.  Al- 
though this  is  the  only  technique  reported  to  be  applicable  to  a  determina- 
tion of  absolute  configuration,  several  methods  are  available  to  determine 
relative  configurations  of  homeomers  with  considerable  certainty. 

Werner's  Solubility  Method.  Although  the  absolute  configurations  of 
a  pair  of  optical  isomers  are  generally  not  known,  the  relative  space  posi- 
tions of  analogous  compounds  may  be  found  if  the  configuration  of  a  given 
compound  be  designated.  This  has  been  realized  with  complexes  of  co- 
balt (III),  chromium(III),  rhodium(III),  and  iridium(III).  Werner185  sug- 
gested that  the  relative  configurations  of  inorganic  complex  compounds 
could  be  determined  by  comparing  the  solubilities  of  analogous  diastereoiso- 
mers.  The  resolution  of  tris(ethylenediamine)  cations  of  cobalt  (III),  rho- 
dium(III)  and  chromium(III)  by  means  of  camphornitronates  and  chloro- 
tartrates  was  used  as  an  example.  Since  the  less  soluble  diastereoisomers 
were  the  dextro  rotatory  cobalt  (III),  chromium  (III)  ions  and  the  levo 
rotatory  rhodium(III)  ion,  it  was  concluded  that  these  cations  possess  the 
same  spatial  arrangement.  Jaeger  criticized  this  theory,  stating  that,  "This 
view  is  quite  arbitrary  because,  in  general,  solubility  is  a  so  highly  compli- 
cated and  constituent  property  of  matter  that,  even  where  we  seem  to  have 
established  rules  for  homologous  series,  sometimes  most  unexpected  and 
surprising  exceptions  spring  up.  This  makes  these  rules  quite  illusory"100. 
He  suggested  that  the  crystal  form  is  a  better  criterion  for  relative  configu- 
ration and  attempted  to  demonstrate  that  the  method  suggested  by  Werner 
was  incorrect93.  Jaeger  has  since  acknowledged  that  the  method  of  solubili- 
ties is  correct  and  has  applied  it  in  studies  of  relative  configurations  of 
analogous  optically-active  antipodes94, 186. 

Rotatory  Dispersion  Curves — Circular  Diehroism.  The  fact  that 
both  the  absorption  spectra  and  the  optical  rotation  are  related  to  the 
resonance  within  a  particular  molecule  suggests  that  some  correlation  exists 
between  these  two  properties.  It  has  also  been  shown  that  certain  absorp- 
tion bands  are  directly  connected  with  the  groups  concerned  with  the  opti- 
cal rotatory  power  of  the  molecule.  Hence,  the  specific  rotation  of  a  com- 

183.  Born,  Proc.  Roy.  Soc,  London,  150A,  83  (19. 

184.  Bijvolt,  Peerdeman,  and  von  Bommel,  Nature,  168,  271  (1951). 

185.  Werner,  Bull.  soc.  chim.,  [4]  11,  1  (1912). 

186.  Jaeger,  Bull.  soc.  chim.,  [5]  4,  1201  (1937);  Jaeger,  Pro.  Acad.  Set.  Amsterdam,  40, 

2,  108,574  (1937). 


338 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


pound  is  very  different  when  the  measurements  are  made  with  light  of  a  wave 
length  which  corresponds  to  one  of  these  absorption  bands  (Fig.  8.47). 


6800     6400     6000     5600     5200  4900    4400   4000  A 


Fig.  8.47.  Absorption  spectrum  and  rotatory  dispersion  of  potassium  im-oxalato 
cobaltate(III). 

A.  Racemic-absorption  spectrum 
B-Dextro-rotatory  dispersion 
B'-Levo-rotatory  dispersion 


The  rotatory  dispersion  curves,  B  and  B',  undergo  abrupt  changes  as  the 
shaded  region  represented  by  the  absorption  curve,  A,  is  approached  and 
passed.  At  wave  lengths  of  light  remote  from  the  absorption  curve,  very 
little  change  occurs  in  the  optical  rotation  as  the  wave  length  is  changed. 
This  change  of  rotation  with  change  of  wave  length  of  light  is  called  rotatory 
dispersion. 

The  determination  of  the  optical  rotation  of  coordination  compounds, 
which  are  usually  colored  and,  therefore,  have  absorption  bands  in  the 
visible  range,  is  sometimes  difficult.  With  such  compounds  it  is  advisable 
to  determine  the  specific  rotation  at  several  different  wave  lengths  or,  at 
least,  the  wave  length  of  the  light  used  must  always  be  specified. 

Although  numerous  investigators187  have  studied  the  rotatory  dispersion 
curves  of  complex  compounds,  none  has  applied  this  technique  so  exten- 
sively or  so  successfully  as  Mathieu.  He  has  found  this  procedure  extremely 
useful  in  comparing  the  configurations  of  analogous  compounds188  and  in 
stud}nng  any  changes  in  configuration  during  displacement  reactions189. 
Mathieu  showed1880  (Fig.  8.48)  that  the  tris(ethylenediamine)  compounds 

1ST.  Bruhot,  Bull.  soc.  chim.,  [4]  17,  223  (1915);  Jaeger,  Rec.  trav.  chim.,  38,  309  (1919); 
Lifschitz,  Z.  physik.  Chem.,  105,  27  (1923);  Longchambon,  Compt.  rend.,  178, 
1828  (1924). 

188.  Mathieu,  Compt.  rend.,  119,  278  (1934) ;  Mathieu,  ibid.,  201, 1183  (1935) ;  Mathieu, 

J.  chim.  phys.,  33,  78  (1936). 

189.  Mathieu,  Bull.  soc.  chim.,  [5]  3,  463,  476  (1936) ;  Mathieu,  ibid.,  [5]  5,  105  (1938). 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS 


339 


6500   6000    5500   5000   4500   4000    3500   3000 


Fig.  8.4S.  Rotatory  dispersion  curves  of  some  tris-ethylenediamine  complexes. 
A  .  ./-[Co  en,)Br,  ;  (B),  d-[Ci  en8]I,  ;  (C),  l-[Rh  en»*]Ia  ,  (D),  l-[Ir  en,]Br3  . 

of  </-[Co(III)],  d-[Cr(III)],  Z-[Rh(III)]  and  Z-[Ir(III)]  have  the  same  config- 
uration. It  is  seen  that  these  curves  are  similar,  indicating  analogous  con- 
figurations, whereas  if  the  curves  are  different  (Fig.  8.47),  the  optically 
active  ions  have  opposite  configurations. 

This  same  technique  was  employed  by  Mathieul88a  to  corroborate  the 
conclusions  which  Werner"  made  by  means  of  his  solubility  method.  Werner 
investigated  numerous  reactions  (page  344)  involving  the  displacement  of  a 
donor  ion  or  molecule  from  the  coordination  sphere  of  an  optically-active 
complex  compound  and  showed  that  in  some  of  these  reactions,  although 
the  sign  of  rotation  may  change  when  measured  at  the  d  line  of  sodium,  the 
configuration  of  the  product  remains  the  same  as  that  of  the  original  ma- 
terial. A  typical  example  of  the  application  of  rotatory  dispersion  curves  in 
studies  of  this  type  might  be  illustrated  by  considering  the  reactions 

Zei-o-[Co  en2  Cl*]+     KCNS  ,  levo-[Co  en2  CI  NCSJ+     Na  N°2  >  dextro-[Co  en2  NCS  N021+ 

These  three  complex  cations  have  analogous  rotatory  dispersion  curves 
(Fig.  8.49)  and  must,  therefore,  possess  the  same  generic  configuration. 


+  3000, 


+  2000 

+  1000 

[M]  0 

-1000 


-2000 


, ^ 

> 

}> 

£~\ 

-f- 

y 

Hx 

^~~~— 

\ 

\  V 

i 

^ 

<±<.\ 

7000        6500 


6000 


A 


5500 


5000 


4500 


Fig.  8.49.  Rotatory  dispersion  curves  of  some  bis-ethylenediamine  complexes. 
(A),  /-[Co  en2  Cl2]+;  (B),  /-[Co  en2  CI  XCS]+;  (C),  d-[Co  en,  \CS  N02]+ 


340  CHEMISTRY  OF  Till-:  COORDINATION  COMPOUNDS 

Recently50  a  new  method  for  distinguishing  between  geometrical  isomers 
which  makes  use  of  their  rotatory  dispersion  curves  has  been  suggested 
(page  298). 

The  rotatory  dispersion  of  an  optical  isomer  is  very  closely  related  to 
another  phenomena  referred  to  as  circular  dichroism  or  "Cotton  effect." 
Although  plane-polarized  light  has  been  most  widely  used  in  the  study  of 
optical  isomerism,  some  interesting  and  fundamental  data  have  been  se- 
cured by  means  of  circularly-polarized  light.  It  was  found,  for  example, 
that  the  absorption  of  dextro  or  levo  circularly-polarized  light  is  dependent 
upon  the  wave  length.  II  the  circularly-polarized  light  is  of  a  wave  length 
in  the  neighborhood  of  the  characteristic  absorption  bands  of  groups  con- 
cerned with  the  optical  activity  of  the  molecule,  then  the  beams  of  dextro 
and  levo  circularly-polarized  light  are  absorbed  to  different  extents,  but 
at  all  other  wave  lengths  the  coefficients  of  absorption  are  equal.  This 
phenomenon  is  designated  as  the  "Cotton  effect"  because  Cotton1£0  first 
demonstrated  it  with  alkaline  solutions  of  copper  tartrates. 

The  "Cotton  effect"  and  rotatory  dispersion  of  an  optical  isomer  can  be 
related  qualitatively  by  the  fact  that  a  compound  designated  as  having  a 
positive  "Cotton  effect"  has  a  rotatory  dispersion  curve  which  changes  from 
a  maximum  rotation  to  a  minimum  rotation  in  the  direction  of  shorter 
wave  lengths.  In  the  same  manner,  a  compound  whose  rotatory  dispersion 
curve  changes  from  a  minimum  to  a  maximum  rotation  is  said  to  have  a 
negative  "Cotton  effect."  Therefore,  studies  of  rotatory  dispersions  are 
sometimes  expressed  in  terms  of  positive  or  negative  "Cotton  effect." 
Analogous  compounds  with  the  same  "Cotton  effect"  at  corresponding 
absorption  bands  have  the  same  generic  configuration;  whereas  similar 
compounds  of  different  "Cotton  effect"  have  opposite  configurations1880; 
thus,  it  is  seen  that  studies  of  the  "Cotton  effect"  may  be  used  in  determin- 
ing structures,  and,  also,  according  to  Mellor191,  in  determining  bond 
character. 

Delepine's  Active  Racemate  Method.  The  physical  characteristics  of 
a  racemic  modification  often  differ  from  those  of  the  enantiomorphs  from 
which  it  is  derived.  In  particular,  the  solid  state  of  a  racemic  modification 
may  exist  in  three  forms:  (1)  racemic  mixtures  (2)  racemic  compounds,  or 
(3)  racemic  solid  solutions.  Pvacemic  mixtures  are  produced  by  certain 
asymmetric  compounds  which  form  crystals  that  possess  hemihedral  facets 
and  are  themselves  cnantiomorphic.  A  racemic  compound  results  whenever 
a  pair  of  enantiomorphs  unite  to  form  a  molecular  compound,  all  of  the 
crystals  containing  equal  amounts  of  each  isomer  and  being  identical. 
These  crystals  have  different  physical  properties  from  those  of  the  indi- 

L90.  Cotton,  Ann.  chim.  phys.,7,  8  (1896). 

191.  Mellor,  /.  Proc.  Roy.  Soc,  N.  S.  Wales,  75, 157  (1942). 


STEREOISOMERISM  OF  HEXACO}  ALENT  ATOMS  34] 

vidua]  antipodes.  (  tftentimee  a  pair  of  enantiomorphs  arc  also  isomorphous. 
Whenever  this  situation  exists  they  may  crystallize  together  as  a  racemic 

solid  solution  without  the  formation  of  a  compound. 

ks  early  as  1921,  Delepine174  suggested  thai  similar  optically-active 
salts  which  form  isomorphous  crystals  have  the  same  relative  configuration 

regardless  of  their  optical  rotation.  This  led  to  the  method  referred  to 
Delepine's  "active  racemate"  method176  which  can  besl  be  presented  by  a 
brief  discussion.    If   two   enantiomorphs,   such   as   f/-K ;( 'o<("-_<  v.-]   and 

/-K3[Co(C204)3],  the  crystals  of  which  possess  hemihedral  facet-,  are  mixed 
in  solution  in  equimolecular  quantities  and  allowed  to  crystallize,  crystals 
of  the  racemic  mixture  are  formed.  These  crystals  represent  a  mechanical 
mixture  of  the  individual  antipodes  and,  when  put  in  solution,  they  are, 
of  course,  optically  inactive.  If  tf-K3[Cr(C204)3]  is  substituted  for  d-K  - 
[Co(C204)3],  the  crystals  which  form  will  give  an  optically-active  solution 
("active  racemate"),  because  (/-K,[Cr(C204)3]  and  /-K3[Co(C204)3]  do  not 
have  equal  rotatory  power.  Delepine  points  out  that  if  the  "active  race- 
mate''  is  a  racemic  mixture,  then  the  generic  configurations  of  the  two  anti- 
podes are  different;  however,  if  it  is  either  a  racemic  compound  or  racemic 
solid  solution,  then  the  generic  configurations  of  the  antipodes  are  the  same. 

Delepine  was  able  to  show  by  this  method  that  /-K3[Ir(C204)3]  and 
^/-K3[Rh(C204)3],  f/-K3[Ir(C204)3]  and  r/-K3[Co(C204)3],  and  d-K3[Ir(C204)3] 
and  /-K3[Cr(C204)3]  form  racemic  compounds  or  solid  solutions  of  the 
optically-active  type.  It  was,  therefore,  concluded  that  the  generic  configu- 
rations of  the  trioxalato  complexes  of  these  four  metals  are  the  same 
in  r/-K3[Co(C204)3],  /-K3[Cr(C204)3],  d-K3[Ir(C204)3]  and  /-K3[Rh(C204)3]. 
This  procedure  has  likewise  been  used  to  show  that  cobalt(III)  and  rho- 
dium(III)  complexes  of  the  same  sign  of  rotation  have  opposite  generic  con- 
figurations in  the  tris(ethylenediamine)  series192. 

The  method  of  active  racemates  is  limited  only  by  the  fact  that  the  salts 
in  question  must  form  crystals  which  have  hemihedral  facets  and  must  be 
isomorphous.  A  careful  choice  of  anions  and  cations  can  lead  to  isomor- 
phism in  quite  different  types  of  salts,  and  it  may  be  possible  to  determine 
the  generic  configurations  of  hexacovalent  metals  having  different  valem 
Thus,  the  configurations  of  analogous  zinc(II),  cobalt(III)  and  plati- 
num(TV)  complexes  might  be  related  through  the  possible  isomorphism 
such  pairs  as  [Zn  en3](X03)2-[Pt  en3](C03)2  and  [Co  en3]P04-[Zn  enj- 
S04. 

Preferential  Adsorption  on  Optically -active  Quartz.  Tsuchida, 
Kobayashi  and  Nakamura28  have  -  I  that  the  preferential  adsorption 

of  enantiomers  on  optically-active  quartz  might  furnish  a  useful  mean-  of 
enrolling  the  relative  configurations  of  ana!'  jymmetric  com- 

pounds. This  assumption  has  been  checked  experimentally118  bydetermin- 

Delepine  and  Charonnat,  Bull.  soc.  franc,  mineral,  53,  73  (1930). 


342 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


ing  the  adsorption  of  several  complex  compounds  on  finely  ground  dextro- 
quartz  powder.  The  results  of  this  investigation  confirm  the  opinion  that 
there  is  a  close  relationship  between  the  adsorption  and  the  spatial  configu- 
ration of  the  complex. 

Sonic  Reactions  of  Optically -active  Isomers 

Polynuclear  Complex  Compounds.  Werner  observed  that  groups  co- 
ordinated to  an  asymmetric  central  ion  can  be  displaced  and  a  product  ob- 
tained  which  is  still  optically  active,  although  in  many  cases  the  degree  of 
optical  rotation  or  even  the  sign  may  change.  The  optical  rotations  of  some 
of  the  products  obtained  by  the  reaction  of 

NH2 


en2  Co<m>     CodV)  en2 


X4 


with  various  reagents  are  shown  in  Table  8.9.  It  will  be  noted  that,  in  every 
case,  the  products  obtained  had  rotations  opposite  in  sign  and  smaller  than 
that  of  the  starting  material.  Mathieu189b  has  investigated  the  rotatory 
dispersions  of  some  of  these  materials  and  has  shown  that  although  the  sign 
of  rotation  changed,  the  generic  configuration  of  the  products  was  the  same 
as  that  of  the  reactant.  Thompson  and  Wilmarth161a  have  shown  that  the 
reactions  listed  in  Table  8.9  involve  a  one  electron  reduction  and  that  the 
oxidation-reduction  reaction 


NH, 


en2  Co(ni>     Co(IIX>  en2 


0; 


NH2 

/    \ 
en2  Cod11)     Co<IV>  en2 

\    / 

02 


4  + 


+  e 


is  reversible  with  an  electrode  potential  of  slightly  more  than  —1.0  volt. 
Therefore  the  structures  designated  by  Werner  and  shown  in  Table  8.9  for 
products  1  and  2  are  in  error;  there  is  good  evidence  in  support  of  the  struc- 
ture 


NH2 

/   \ 
en2  Co<m>     Co(riI>  en2 

\    / 


XrHX 


for  the  product  of  reaction  number  2181a. 

Substitution  Reactions  with  No  Change  in  Configuration.  Wer- 
ner188 postulated  thai  the  replacement  of  groups  a  and  b  in  complexes  of 

161  a.  Thompson  and  Wilmarth,  J,  Phys,  Chem.,  56,  5  (1952). 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS 


343 


MI 

/       \ 
Table  8.9.  Reactions  of  l-[en2  Co«">  Co<IV>  en2]X4 

\       / 
02 

[a]l°  =   -840°;  [M]2D°  =   -6854° 

(concentration  0.125%) 


Xo. 

Reagent 

Product 

Ul20 

Mff 

1 

MI 

NH 

/     \ 
[en>  Co"11*         Co<IV>  en2]X3  . 

\     / 

o2 

HX 

+  160 

+  1372 

2 

HX 

NH 

/       \ 

[en2  Co(ni)            Co<IV>  en2]X3 

\       / 

o2 

+  192 

+  1625 

3 

Xal 

XH2 

/       \ 
[en2  Co<ni>           Co^1")  en2]X< 

\       / 

OH 

+  110 

+990 

4 

HNO, 

XH2 

/       \ 
[en2  Co™           Co<m>  en2]X4 

\       / 
N02 

+  158 

+  1311 

5 

so2 

NH2 

/       \ 
[en2  Co"11)           Co"")  en2]X3 

\       / 
SO  4 

+200 

+  1384 

the  type  [M(AA)2ab]  takes  place  with  no  change  in  configuration.  He  sug- 
gested that  during  these  reactions  the  labile  groups,  a  and  b,  are  easily 
displaced  and  the  bidentate  groups,  AA,  remain  firmly  bound,  thus  main- 
taining the  >ame  spatial  arrangement  of  the  atoms  in  the  molecule.  This 
was  toted  by  numerous  reactions  involving  optically-active  compounds 
(Table  8.10). 

Werner  applied  his  method  of  solubilities  to  show  that  in  every  case 
the  generic  eon  figuration  of  the  product  was  the  same  as  that  of  the  react- 
ant.  This  same  conclusion  was  reached  by  Mathieu189a  who  investigated 
the  rotatory  dispersion  of  some  of  these  material-. 


344 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Table  8.10.  Reactions  of  Some  Optically-active  [M(AA)2ab]  Compounds 


Sign  of 

Sign  of 

No. 

Rota- 
tion 

Reactant 

Reagent 

Product 

Rota- 
tion 

1 

— 

[Co  en2  Cl2]+ 

K8C08 

[Co  en2  C03]+ 

+ 

2 

- 

[Co  en2  CI  NCS]+ 

NH3 

[Co  en2  NH3  XCS]+ 

+ 

3 

- 

[Co  en2  CI  NCS]+ 

NaNOa 

[Co  en2  N02  NCS]+ 

+ 

4 

- 

[Co  en2  Cl2]+ 

(NH4)2C.»04 

[Co  en2  C204]+ 

+ 

5 

- 

[Cr  en2  Cl2]+ 

(NH4)2C204 

[Cr  en2  C204]+ 

+ 

6 

— 

[Co  en2  N02  Cl]+ 

KCNS 

[Co  en2  N02  NCS]+ 

— 

The  Walden  Inversion*  in  Reactions  of  Complex  Ions  and  Inter- 
conversion  of  Enantiomorphs.  Contrary  to  Werner's  assumption  that 
labile  groups  are  always  displaced  from  the  coordination  sphere  of  a  central 
atom  without  a  change  in  configuration,  Bailar  and  Auten106  have  demon- 
strated that  certain  reactions  of  this  type  can  cause  the  interconversion  of 
enantiomorphs.  The  experiments  of  Bailar  and  Auten  (Fig.  8.50)  brought 


ci 

DEXTRO     H 


ALCOHOLIC 
HCI 


f  ALCOHOLIC 
HCI 


o-c=o 

DEXTRO-UI 


0  =  C-0 

LEVO-  321 


Fig.  8.50.  Configuration  change  in  the  reaction  of  dichloro-bis-ethylenediamine 
cobalt  (III)  ion  with  carbonate. 


to  light  the  first  example  of  a  Walden  inversion  in  the  field  of  inorganic 
complex  compounds.  It  was  shown  that  the  treatment  of  an  aqueous  solu- 
tion of  Zew-dichloro-bis(ethylenediamine)cobalt(III)  ion,  (I),  with  a  solu- 
tion of  potassium  carbonate  produces  the  eforiro-carbonato  ion,  (III),  but 
grinding  with  an  excess  of  solid  silver  carbonate  produced  the  levo  isomer, 
(IV).  This  is  converted  to  the  dextro-dichloro  ion,  (II),  by  alcoholic  hydro- 
chloric ;icid.  The  relative  configurations  of  the  complex  ions  were  assigned 
as  the  result  of  rotatory  dispersion  studies71  and  the  inversion  is  repre- 
sented as  taking  place  in  the  silver  carbonate  reaction.  Later  develop- 

*  The  use  of  the  term  "Walden  Inversion,'  in  this  connection  has  been  chal- 
lenged; however,  the  disagreement  appears  to  be  mainly  in  linguistics  and  not  of  a 
fundamental  nature  (79a). 


Temp. 

Specific  Rotation 
of  Produi  t 

-77° 

-32° 

-33° 

-22° 

+25° 

+29° 

+80° 

+43° 

+25° 

+31° 

+25° 

+29° 

STEREOISOMERISM  OF  HEXAC01  ALENT  ATOMS  345 

Table  8.11.  Effect  oi  Temperatt  re  on  Walden   [nversion 
Co  ens  Cl2]+  +  2NH,     ►  [Co  en,  (NH,),]H    <  +  2C1~ 

Reagent 
Liquid  MI 
Liquid  Nil 
Liquid  XH3 
Gaseous  Ml: 
Ml    in  CHsOH 
MI 3  in  C2H5OH 

Table  8.12.  Effect  of  Temperature  on  Walden  Inversion 
[Co  ena  Cl8]+  +  Ag2C03  -*  [Co  en2  C03]+  +  2AgCl 

Temp.  Specific  Rotation  of  Product 

0°  -10° 

15°  -100° 

25°  -106° 

50°  -78° 

75°  -28° 

90°  0° 

ments19b'  129,  195  show  that  the  reagent  is  not  the  important  factor;  instead 
the  conversions  of  /-[Co  en2  Cl2]+  to  the  cferriro-carbonato  complex  proceeds 
through  the  formation  of  an  aquated  intermediate,  while  conversion  to  the 
fevo-carbonato  compound  proceeds  directly.  The  effect  of  various  factors  on 
the  inversion  are  discussed  below: 

(1)  Effect  of  Temperature.  It  should  likewise  be  mentioned  that  experi- 
mental conditions  play  an  important  role  in  Walden  inversions.  For  ex- 
ample, the  effect  of  temperature  on  the  inversion  of  complex  inorganic 
compounds  was  first  noticed  with  the  reaction  between  l-[Co  en2  C12]C1  and 
ammonia108.  A  levo  rotatory  salt,  [Co  en2  (NH3)2]Cl3 ,  was  isolated  if  the 
reaction  took  place  at  —77°  or  —  33°C,  but  the  dextro  rotatory  product 
was  obtained  from  the  reaction  at  +25°C  (Table  8.11).  These  results  have 
been  confirmed  and  extended  by  Keeley196. 

This  effect  of  temperature  was  also  studied  for  the  reaction  of 
l-[Co  eiit  CUICI  with  silver  carbonate195.  The  data,  which  are  summarized 
in  Table  8.12,  show  that  the  chief  effect  of  low  temperatures  is  to  decrease 
the  rate  of  reaction,  and  the  effect  of  high  temperatures  is  to  cause  racemiza- 
tion. 

(2)  Effect  of  Concentration.  It  was  found  that195,  if  an  excess  of  silver  carbo- 
nate were  present,  the  levo  sail  was  obtained,  however,  if  an  excess  was 
not  present,  the  dextro  salt  was  obtained.  On  the  other  hand,  potassium 

196.  Bailar,  Jonelis,  and  Huffman,  ./.  Am.  Chem.  Soc,  58,  2224  (1936). 
196.  Keeley,  thesis,  University  of  Illinois.  1952. 


346  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Table  8.13.  Effect  of  Concentration  on  Walden  Inversion 
[Co  en2  Cl;]+  +  COr  ->  [Co  en2  C03]+  +  2C1" 

Molar  Ratio  of  Ag2C03  Specific  Rotation 

to  Complex  Present  of  Product 

0.75  +362° 

1.12  +288° 

1.50  -102° 

3.00  -160° 

4.50  -180° 

Molar  Ratio  of  K2CO3  Specific  Rotation 

to  Complex  Present  of  Product 

1.00  +240° 

1.50  +140° 

3.00  +110° 

5.00  +80° 


carbonate  produced  the  dextro  salt  at  all  times  although  the  specific  rota- 
tion decreased  with  increasing  concentration  of  potassium  carbonate  (Table 
8.13).  This  marked  racemization  was  probably  due,  however,  to  the  for- 
mation of  the  optically-inactive  trans-[Co  en2  H20  OH]++  by  the  strongly 
basic  solution. 

(3)  Nature  of  Reagent.  The  fact  that  the  particular  reagent  chosen  to 
effect  a  reaction  exerts  a  predominating  influence  on  the  configuration  of 
the  product  is  clearly  demonstrated  by  the  different  results  obtained  when 
Ag2C03  and  K2C03  react  with  l-[Co  en2  C12]C1.  There  is  no  adequate  expla- 
nation for  this.  In  an  attempt  to  determine  whether  some  correlation  exists 
between  the  type  of  reagent  and  its  influence  on  the  configuration  of  a 
particular  compound,  the  reaction  of  Hg2C03  with  Z-[Co  en2  C12]C1  was 
studied.  Mercurous  ion  and  silver  ion  both  form  insoluble  chlorides  and  car- 
bonates. They  might,  therefore,  be  expected  to  behave  similarly.  It  was 
found  however,  that  l-[Co  en2  C12]C1  reacts  with  an  excess  of  Hg2C05  to 
give  the  dextro  rotatory  carbonato  salt.  This  reaction,  which  is  much 
slower  than  that  with  silver  carbonate,  gives  results  similar  to  those  ob- 
tained with  potassium  carbonate. 

(4)  Nature  of  Solvent.  It  has  been  shown  definitely  for  carbon  compounds 
that  the  nature  of  the  solvent  plays  an  important  role  in  the  inversion  of  a 
molecule197.  Although  most  of  the  reactions  of  complex  inorganic  salts  are 
carried  out  in  water,  there  is  some  indication  that  other  solvents  or  no  sol- 
vent may  give  different  results.  For  example,  it  has  been  established  that 
the  conversion  of  /.-[Co  en2  C12]C1  to  the  dea^ro-carbonato  complex  proceeds 
through  the  formation  of  an  aquated  intermediate,  while  conversion  to  the 

197.  Senter,  J.  Chem.  Soc,  127,  1847  (1925). 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS  347 

Table  8.14.  Effect  oi  a..i\<,   \   I  Peb  Cent  Solution  of  J-[Co  en8  C1j]C1 
Before  Treating  with   \  Ten  fold  Excess  oi   Silveb  Carbonate 
t,  representa  the  time  of  :i^in^  in  minutes  and  [a]  represents  specific 
rotation  of  resultant  carbonato  sail  in  degrees 

t  [a]  t  [a] 

0  -212  7.-.  +684 

1  -183  120  +635 
3  -96  170  +587 
6  -19  186  +539 

10  +87  235  +520 

20  +250  260  +520 

40  +501  296  +462 

50  +530  360  +433 

60  +578  1080  +147 

/no-carbonato  salt  proceeds  directly198.  That  aquation  plays  an  important 
part  in  the  reaction  between  Z-[Co  en2  C12]C1  and  silver  carbonate  is  shown 
by  the  fact  that  the  rotation  of  the  carbonato  complex  obtained  depends 
upon  how  long  the  solution  of  the  dichloro  salt  is  allowed  to  stand  before 
the  silver  carbonate  is  added  (Table  8.14).  This  would  suggest  that  other 
examples  of  inversion  of  optically-active  complexes  might  be  observed,  if  it 
were  possible  to  employ  noncoordinating  solvents  in  order  to  enhance  the 
possibility  of  a  displacement  (SN2)  reaction. 

Theories  of  the  Walden  Inversion.  The  fact  that  Walden  inversions  have 
been  demonstrated  for  complex  compounds106, 108, 195  is  of  interest  in  establish- 
ing whether  the  mechanisms  proposed  for  inversions  of  the  tetrahedral 
carbon  are  sufficiently  general  to  be  applicable  to  octahedral  complex  inor- 
ganic compounds.  One  of  the  mechanisms  suggested  for  the  Walden  inver- 
sion postulates  that  every  reaction  which  involves  a  single  step  in  the  dis- 
placement of  one  group  by  another  on  a  tetrahedral  atom  should  lead  to 
inversion199.  Accordingly,  if  the  over-all  reaction  takes  place  in  an  odd  num- 
ber of  steps  the  product  will  be  the  enantiomorph  of  the  original  material, 
but  if  it  takes  place  in  an  even  number  of  steps,  the  starting  material  and 
the  product  will  have  the  same  configuration.  This  theory  was  tested  by 
Bailar,  Haslam  and  Jones108  who  studied  the  reaction  of  /-[Co  en2  C12]C1 
with  ammonia  which  yields  the  corresponding  diammine  complex.  The  two 
chloride  atoms  of  the  complex  ion  are  attached  to  the  cobalt  in  the  same 
way  and  occupy  like  positions  in  the  molecule.  It  seems  logical  to  assume, 
therefore,  that  the  same  mechanism  functions  in  their  displacement  from 
the  complex.  If  this  is  correct,  the  conversion  of  the  dichloro  salt  to  the 

198.  Bailar  and  Peppard,  /.  Am.  Chem.  Soc.,  62,  820  (1940). 

199.  Bergmann,  Polanyi  and  Szabo,  Z.  physik.  Chem.,  B20,   161    (1933);  Olson,  J. 

Chem.  Phys.,  1,  418  (1933). 


348  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

diammine  sail  must  take  place  in  an  even  number  of  steps,  and  the  theory 
mentioned  would  allow  no  inversion.  However,  it  was  shown  that  the  re- 
net  ion  docs  load  to  inversion. 

The  authors195  mention  the  possibility  that  the  displacement  of  a  negative 
chloride  group  by  a  neutral  ammonia  molecule  may  produce  such  a  pro- 
found change  in  the  complex  ion  that  the  second  step  of  the  reaction  does 
not  follow  the  same  mechanism  as  the  first.  A  more  conclusive  test  of  the 
theory  of  Bergmann,  Polanyi  and  Szabo199a  and  of  01son199b  can  be  had  if 
the  chloro  groups  were  displaced  by  other  univalent  negative  groups.  There 
has  been  no  report  made  to  date  of  a  Walden  inversion  of  this  type. 

Meisenheimer's  theory  of  the  Walden  inversion  in  reactions  of  organic 
compounds200  postulates  that  the  incoming  group  attaches  itself  to  the 
face  of  the  tetrahedron  opposite  the  group  expelled.  An  octahedron,  how- 
ever, has  four  faces  "opposite"  and  equidistant  from  each  corner.  If  it  is 
assumed  that  the  incoming  group  attaches  to  any  one  of  these  with  equal 
ease,  the  theory  of  Meisenheimer  will  predict  complete  racemization,  as  a 
study  of  the  model  will  show. 

A  consideration  of  the  models  of  these  complex  cobalt  compounds  shows 
that  the  d  isomer  may  be  transformed  into  the  I  isomer  merely  by  exchang- 
ing the  point  of  attachment  of  a  certain  two  groups.  Hence,  it  is  possible 
that  the  configuration  of  these  optically-active  cobalt  complexes  may  be 
inverted  by  the  properly  oriented  approach  of  the  incoming  group.  Such 
a  mechanism  of  inversion  does  not  necessitate  the  formation  of  a  new 
octahedron.  Basolo,  Stone  and  Pearson74  (Figs.  8.29  and  8.30)  and 
Brown  Ingold,  and  Nyholm79a  also  give  an  interpretation  of  the  Walden 
in  octahedral  structures. 

Mutarotation.  Experimental  results  show  that  in  some  instances  the 
rotatory  power  of  a  freshly  prepared  solution  of  optically-active  substances 
is  not  constant,  but  gradually  changes,  finally  reaching  a  constant  value 
(not  zero)  by  reason  of  the  establishment  of  an  equilibrium.  Such  a  change 
in  rotatory  power  is  termed  Mutarotation.  Numerous  examples  are  known 
for  organic  compounds201. 

Burgess  and  Lowry202  demonstrated  that  this  phenomenon  can  occur  in 
coordination  compounds  by  discovering  that  benzoylcamphorberyllium(II) 
mutarotates.  It  had  previously  been  reported203  that  l-hydroxy-2-benzoyl- 
camphene  exhibits  mutarotation  and  it  was  suggested  that  this  resulted 

200.  Meisenhiemer,  Ann.,  456,  126  (1927);  Meisenhiemer  and  Link,  Ann.,  479,  2.11 

(1930). 

201.  Schreiber  and  Shriner,  ./.  Am.  Chem.  Soc,  57,  1306,  1445,  1896  (1935);  Tanrent, 

Compt.  rend.,  120,  1060  (1895). 

202.  Burgess  and  Lowry,  J.  Chem.  Soc,  125,  2081  (1924). 


STEREOISOMERISM  OF  HEX  LC01  ALENT  ATOMS 


349 


Table  8.15.  Mutabotation  oi    Benzoylcamphoraltjminum(III) 


•J1  _.  pel-  cent  boIu 

lions  at   '_'( 

Chloroform 

I:. 

ozene 

Time  (mil 

.)                      [orjiioi 

Time  (min.) 

[a]  54  01 

0 

0 

(1175) 

1.-) 

7ls 

27) 

1170.5 

25 

7.V) 

is 

1K.7.7 

40 

760 

90 

1164.3 

gfi 

7()(i 

160 

L161.8 

235 

7ti!» 

265 

1158.9 

final 

772 

365 

1157.8 

1890 

1147.6 

final 

1143.8 

Ethylenebron 

ide 

Time  (min.) 

[a]sifii 

30 

570° 

45 

566 

75 

565 

195 

564 

360 

562 

1320 

558 

2820 

550 

5640 

545 

9  days 

538 

from  the  reaction 


O 


C— C— C6H. 


OH 


C=C—  C6H5 


CsHi4 


\ 


CsHi4 

\ 
C— OH  C=0 

Since  in  benzoylcamphorberyllium(II)  there  is  no  longer  a  mobile  hydrogen 


^6h5 


c8hw 


\ 


atom,  any  change  in  rotatory  power  to  a  final  constant  value  must  involve 
the  racemization  of  the  labile  asymmetric  beryllium(II)  center.  This  inter- 
pretation was  not,  at  first,  universally  accepted  because  the  tetrahedral 
configuration  of  l)cryllium(II)  had  not  yet  been  clearly  demonstrated204  and, 
therefore,  a  similar  experiment  was  carried  out  making  use  of  the  octa- 
aedral  aluminum(III)  compound205.  Some  of  the  data  obtained  with  solu- 
tions of  l)enzoylcamphoraluminum(III)  are  given  in  Table  8.15  which  shows 
tin-  rate  of  mutarotation  is  dependent  upon  the  solvent.  This  is  in  accord 
witli  the  observations162,  206«  207  that  complex  compounds  racemize  at  dif- 
ferent rates  in  various  solvents. 

Forster, ./.  Chem.  Soc,  79,  987  (1901). 

204.  Mills  and  Gotta,  •/.  Chem.  Soc,  1926,  3121. 

205.  Faulkner  and  Lowry,  ./.  Chem.  Soc.,  127,  1080  (1925). 

206.  Werner,  Ber.,  45,  3061  (1912). 

207.  Rideal  and  Thomas, ./.  Cfu  m.  Soc..  121,  1%  (1922). 


350  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

A  slightly  different  type  of  mutarotation  involving  inorganic  coordination 
compounds  is  found  in  the  experiment  reported  by  Meisenheimer,  Anger- 
mann,  Holsten  and  Kiderlen  (page  324)147. 

Asymmetric  Synthesis.  The  recent  advances  in  synthetic  organic 
chemistry  have  continually  decreased  the  apparent  gap  between  synthetic 
processes  occuring  in  the  living  cell  and  similar  reactions  in  the  laboratory; 
thus,  it  would  seem  that  even  the  most  complicated  processes  of  plant  and 
animal  metabolism  are  controlled  by  orthodox  physical  and  chemical  laws. 
Indeed,  there  is  only  one  striking  difference  between  vital  syntheses  and  their 
laboratory  counterparts.  This  is  the  fact  that  when  a  substance  whose  mole- 
cule displays  only  axial  symmetry  is  produced  by  vital  synthesis  in  a  living 
cell,  it  is  often  found  that  one  of  the  two  possible  antipodal  forms  predomi- 
nates over  the  other  in  the  resulting  product;  whereas,  the  synthesis  of 
asymmetric  molecules  in  the  laboratory  invariably  produces  the  racemic 
modification.  This  pronounced  difference  between  natural  and  laboratory 
products  has  intrigued  stereochemists  for  all  these  years. 

Absolute  Asymmetric  Synthesis.  The  preparation  of  an  optically- 
active  molecule  without  using  an  optically-active  reagent  and  without  any 
of  the  methods  of  resolution  is  called  absolute  (or  total)  asymmetric  synthesis. 
Attempts  have  been  made  to  effect  such  a  synthesis  by  employing  the  phe- 
nomenon known  as  circular  dichroism  or  "Cotton  effect"  (page  340).  One 
theory208  as  to  the  origin  of  optically-active  compounds  depends  upon  the 
fact  that  sunlight  reflected  by  the  surface  of  the  sea  is  always  in  part  ellip- 
tically  polarized209.  The  preferential  absorption  of  one  form  of  this  polarized 
light  by  a  pair  of  optical  antipodes  may  account  for  the  preferential  for- 
mation or  decomposition  of  one  enantiomorph.  Asymmetric  decomposi- 
tions, using  dextro  and  levo  circularly-polarized  light  of  a  wave  length 
comparable  to  that  of  an  absorption  band  of  the  compound  in  question, 
have  been  successfully  carried  out  for  several  organic  compounds210.  Simi- 
larly, asymmetric  formation  of  compounds  under  the  influence  of  cir- 
cularly-polarized light  has  given  positive  results  for  a  few  compounds  of 
carbon211. 

Since  coordination  compounds  are  usually  very  highly  colored  and  have 
a  pronounced  circular-dichroism  in  the  visible  region,  it  would  appear  that 
the  decomposition  or  formation  of  an  asymmetric  compound  of  this  type 
in  the  presence  of  dextro  or  levo  circularly-polarized  light  should  yield  an 

208.  Eder,  Sitzk.  Okad.  Wiss,  Wien,  Abt.  [IIA]  90,  1097  (1885);  ibid.,  94,  75  (1886). 

209.  Jamin,  Compt.  rend.,  31,  696  (1850). 

210.  Kulin  and  Braun,  Xaturwissenschaflen,  17,  227  (1928);  Kuhn  and  Knopf,  ibid., 

18,  183  (1930);  Mitchell,  J.  Chem.  Soc,  1930,  1829. 

211.  Davis  and  Heggie,  J.  Am.  Chem.  Soc.,  57,  377  (1935);  Karagunis  and  Drikos, 

Xahtririsscnschaftcn,  21,  607  (1933);  Karagunis  and  Drikos,  Nature,  132,  354 
(1933);  Karagunis  and  Drikos,  Z.  physik.  Chem.,  24B,  428  (1934). 


STEREOISOMERISM  OF  HEXACOVALENT  ATOMS  351 

optically-active  compound.  Brcdig  and  Mangold-1-  have  investigated  the 
decomposition  of  diazocamphor,  lactic  acid,  and  various  racemic  cobalt  - 
ammine  salts  by  circularly-polarized  ultraviolet  light.  In  none  of  these 
experiments  was  there  any  evidence  thai  optical  activity  was  produced.  A 
somewhat  different  approach  was  employed  by  Jaeger180  (page  336).  The 
absolute  asymmetric  synthesis  of  a  complex  inorganic  compound  has  not 
yet  been  achieved. 

Asymmetric  Synthesis.  "Asymmetric  .synthesis",  as  it  is  now  in- 
terpreted, was  first  discussed  by  Fischer213  and  later  defined  byMarckwald214 
as  that  process  which  produces  optically-active  compounds  from  symmetri- 
cally constituted  molecules  by  the  intermediate  use  of  optically-active  re- 
agents, but  without  the  use  of  any  of  the  methods  of  resolution.  Numerous 
examples215  of  asymmetric  syntheses  are  known  for  carbon  compounds. 

Coordination  compounds  containing  optically-active  donor  molecules 
have  been  found92-94 •  126  to  exist  in  only  certain  preferred  stereoisomeric 
modifications,  rather  than  in  all  the  theoretically  possible  forms.  Reactions 
leading  to  the  formation  of  this  type  of  compound  cannot  be  regarded  as 
examples  of  asymmetric  synthesis,  however,  for,  according  to  Marckwald's 
definition,  the  optically-active  reagent  is  merely  used  as  an  intermediate  in 
the  subsequent  preparation  of  an  optically-active  compound  which  no 
longer  contains  the  reagent;  this  is  not  true  of  the  numerous  examples  of 
coordination  compounds  containing  optically-active  donor  molecules,  in 
which  the  central  ion  is  rendered  optically-active  as  long  as  the  donor 
molecules  remain  coordinated. 

There  is  one  example131,  however,  in  the  field  of  inorganic  complex  com- 
pounds, which  does  fit  the  present  definition  of  asymmetric  synthesis 
(Fig.  8.51).  It  is  believed  that  these  results  are  achieved  because  of  the 


r/-i  r*s\  i+     d-H.2  tart 

'— ICo  en2  G03]    > 

racemic — [Co  en2  c/-tart]+ 


en 

>  (/-[Co  en3] 


Ca(X°2)UCoen2(NQ2)2]+ 


Fig.  8.5J  .  Asymmetric  synthesis 

difference  in  -lability  of  the  d  and  l  forms  of  dex^ro-tartratobis(ethylene- 
diamine)cobalt(III)  ion,  [Co  en.j  ^/-tart]+.  The  less  -table  n  form  reacts 
more  readily  with  ethylenediamine  or  calcium  nitrite  to  form  the  dextro 

212.  Bredig,  Mangold,  and  Williams.  Z.  Angew.  Chem.,  36,  456  (1923). 

213.  Fischer,  B<       27.  3231     1894 

214.  Marckwald,  B<    ..  37,  349    1904). 

215    Bredig  and  I   -  hem.  Z.,  46,  7  (1912);  McKenzie,  ./.  Chem.  80c. ,  85.  1249 

1904 


352  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

rotatory  tris(ethylenediamine)  and  (linitrobis(ethylenediamine)cobalt(III) 
ions  respectively. 

Asymmetric  Endue t ion.  The  phenomenon  termed  asymmetric  induc- 
tion has  been  defined  by  Kortiim216  as  the  action  of  a  force  arising  in  an 
optically-active  molecule,  which  influences  adjacent  molecules  in  such  a 
way  thai  they  become  asymmetric.  This  influence  may  be  of  two  types, 
intramolecular  and  intermolecular,  depending  upon  whether  the  systems 
involved  are  in  the  same  or  different  molecules.  The  phenomenon,  which  is 
not  entirely  understood,  has  been  well  reviewed  by  Ritchie217.  Examples 
of  asymmetric  induction  in  coordination  compounds  have  been  observed218. 
When  a  three  molar  portion  of  ortho-phenanthroline  was  added  to  a  solution 
of  zinc  f/r.r//'o-o:-bromocamphor-7r-sulfonate,  the  rotation  of  the  solution 
was  greatly  enhanced,  probably  because  of  an  asymmetric  induction. 
With  the  addition  of  strychnine  sulfate  to  [Zn(o-phen)3]++  an  abnormal 
decrease  in  the  rotation  of  the  strychnine  was  noted.  This  anomaly  was  not 
so  striking  when  c^a'-dipyridyl  was  substituted  for  the  o-phenanthroline, 
and  primary  amines  were  without  effect.  The  effect  was  attributed  to  an 
activation  caused  by  the  ortho-phenanthroline  on  coordination,  forming 
an  asymmetric  configuration  on  the  zinc  complex. 

This  phenomenon  has  been  investigated  by  Brasted219  who  concluded, 
on  the  basis  of  polarimetric,  refractometric,  conductimetric,  and  spectro- 
graphic  measurements,  that  some  type  of  compound  is  formed  between  the 
anion  and  cation  (or  complex  and  alkaloid).  This  would  indicate  that  the 
forces,  Van  der  Waals  or  ionic,  have  caused  a  distortion  in  the  configuration 
which  was  responsible  for  the  optical  activity  leading  to  a  new  observed 
rotation.  Brasted  also  showed  that  cobalt  (III)  complexes  behave  in  the 
same  manner  as  the  divalent  metal  complexes.  Dwyer178  attributes  these 
observations  to  differences  in  the  activities  of  the  labile  enantiomeric  ions 
in  the  presence  of  optically  active  cations  or  anions. 

Oxidation -Reduction.  It  has  already  been  pointed  out  that  with  com- 
plex inorganic  compounds  it  is  possible  to  achieve  conditions  which  cannot 
be  realized  with  the  carbon  compounds.  One  case  which  has  long  been  of 
interest  to  the  coordination  chemist  is  the  possibility  of  changing  the  oxida- 
tion  slate  of  the  central  metal  ion  of  an  optically-active  complex.  The 
reactions  of  the  binuclear  complexes  of  cobalt(III)  and  cobalt(IV)  which 
Werner  studied  evidently  constitute  the  first  examples  of  oxidation-reduc- 

216.   Kortiim,  Samml.  ('hem.  ('hem  Tech.  Vortage,  10  (1932). 

_M7    Ritchie,  "Asymmetric  Synthesis  and  Asymmetric  Induction,"  London,  Oxford 
University  Press,  1933. 

218.  Pfeiffer  and  Quehl,  Ber.}  64,  2667  (1931  I;  Pfeiffer  and  Baimann,  Ber.,  36,  1064 

L903). 

219.  Brasted,  thesis,  University  of  Illinois,  L942. 


STEREOISOMERISM  OF  ///.  \  iCOV  ILENT  ATOMS  353 

t ion  reactions  of  optically-active  complexes.  It  is  interesting  thai  these 
reaction.-  proceed  without  racemization. 

Dwyer*  has  recently  resolved  the  tris(o-phenanthroline)ruthenium(II) 
cation  and  has  obtained  the  optically  pure,  stable,  orange-yellow  dex- 
tro  and  levo  perchlorates.  Oxidation  with  eerie  nitrate  converts  these 
enantiomers  to  the  blue,  optically-active  [Ru(o-phen)s](C104)g ,  bul  there 
is  a  marked  drop  in  the  molecular  rotation.  However,  on  back  reduction 
with  ferrous  sulfate  the  orange-yellow  ruthenium (II)  compound  is  re- 
covered and  the  molecular  rotation  rise-  to  the  original  value.  The  observed 
rotation-  are  shown  in  Table  8.16.  It  is  of  interesi  to  note  that,  contrary 
to  the  views  of  Werner,  the  complex  of  divalent  ruthenium  has  the  larger 
rotation. 

Table  v16.  Optical  Rotation  of  Tris(o-Phe nanthroline) Ruthenium (II) 

and  (III)  Cations 

Cations 

J-[Ru  o-phen)8]++ 
RuCo-phen),]^ 
l:  .  o  phei 
;  o-phen 

Dwyer  and  Gyarfas  have  performed  similar  experiments  in  which  they 
utilized  a  different  ligand221  and  a  different  central  atom2'72.  They  also  dem- 
onstrated-'-- that  a  dynamic  electronic  equilibrium  may  exist  between  the 
oxidized  and  reduced  forms  of  a  complex  ion.  This  was  done  by  mixing  a 
solution  >8  dipy).>]'^  with  a  solution  containing  an  equivalent  quan- 

tity of  ^[Os(dipy)»]"H~f\  The  resulting  mixture  lost  its  optical  activity  very 
rapidly.  This  rapid  loss  of  optical  activity,  plus  the  fact  that  the  electron 
trai  pected  to  occur  without  inversion,  lends  support  to  one  of  the 

current  theories224  of  electron  exchange  reactions  in  aqueous  solutions. 

220.  Dwyer  and  Gyarfas,  ./.  Proc.  Roy.  Soc.  X.  8.  Wales,  83,  170  (1949). 

221.  Dwyer  and  Gyarfas,  J.P  Soc.,  A  .  S.  Wales,  83,  174  (1949). 

222.  Dwyer  and  Gyarfas,  •/.  Proc.  Roy.  Soc.  N.  S.  Wales,  83,  263  (1949). 

223.  Dwyer  and  Gyarfas,  A  166,  481  (1950). 

224.  Libby,  •/.  Phys.  Ch  n  ..  56.  863  (1952). 


Q*ff 

[Mflia 

-1818° 

-3482° 

+  1834° 

+  3494° 

-568° 

-2:>r>\ 

4-584 

+2 

7.   Stereochemistry  of  Coordination 
Number  Four 

B.  P.  Block 

The  Pennsylvania  State  University,  University  Park,  Pennsylvania 

Configurations  Encountered 

Complex  compounds  having  the  coordination  number  four  are  considered 
to  be  quite  common,  but  there  is  good  evidence  for  the  existence  of  such 
complexes  for  only  a  small  number  of  metallic  elements.  Mellor  has  sum- 
marized the  more  important  spatial  arrangements  which  have  been  sug- 
gested for  these  as  regular  tetrahedral,  pyramidal,  square  or  rectangular 
planar,  and  tetragonal  or  rhombic  bisphenoidal1.  The  first  arrangement  to 
be  established  experimentally  was  the  tetrahedral  configuration  for  the 
carbon  atom,  and  this  three-dimensional  concept  of  structure  colored  the 
thinking  of  chemists  for  many  years.  Although  Werner  explained  several 
puzzling  points  in  the  chemistry  of  some  platinum(II)  complexes  by  assum- 
ing a  planar  arrangement  of  the  four  groups  around  the  platinum  f»  the 
suggestion  was  not  accepted  by  many,  and  even  in  rather  recent  times  there 
have  been  attempts  to  explain  the  structures  of  these  compounds  on  other 
basesV27,  V28,  x1,  X18,  X40,  X41,  X42.f  Two  geometrical  configurations  for  the 
coordination  number  four  are  now  generally  accepted,  the  regular  tetra- 
hedral and  the  square  planar.  These  are  the  two  configurations  which 
Werner  recognized.  Table  9.1  shows  those  metallic  elements  for  which  a 
coordination  number  of  four  has  been  established.  In  some  cases  the  ele- 
ment has  the  configuration  in  question  only  because  the  coordinating  group 
or  groups  are  such  that  a  configuration  which  is  unnatural  to  the  element 
is  forced  upon  it. 

Tetrahedral  Configuration 

The  evidence  for  a  tetrahedral  arrangement  is  found  largely  in  complete 
structure  determinations,  either  by  x-ray  or  electron  diffraction.  Most  of 

The  references  in  this  chapter  marked  with  an  asterisk  are  of  general  interest. 
•1.  Mellor,  Chmi.  Revs.,  33,  137  (1943);  /.  Proc.  Roy.  Soc,  N.  S.  Wales,  76,  7  (1942). 
f  Reference  numbers  preceded  by  letters  refer  to  annotated  bibliography  which 
appears  al  cud  of  this  chapter. 

354 


STEREOCHEMISTRY  OF  COORDINATION  NUMHER  FOUR 


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356  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

these  structures  correspond  to  solid  compounds  and  may  have  little  rela- 
tion to  the  configuration  in  solution.  A  few  attempts  have  been  made  to 
apply  Unman  studies  to  solutions  of  species  which  have  been  studied  as 
solids  or  gases,  but  the  method  does  not  lead  to  unambiguous  results2.  For 
some  compounds  of  coordination  number  four,  there  have  been  reports  of 
resolution  into  optical  isomers,  but  in  most  cases  the  investigators  have  been 
unable  to  obtain  optically-active  fractions  free  of  the  optically-active  resolv- 
ing agenl ,  so  t  he  validity  of  the  evidence  is  doubtful.  The  Cotton  effect  has 
also  been  used  to  demonstrate  the  tetrahedral  configuration;  here,  too, 
there  is  question  as  to  the  validity  of  the  results3. 

In  addition  to  the  compounds  in  which  directed  covalent  bonds  are 
operative,  there  is  a  group  in  which  the  configuration  is  apparently  deter- 
mined by  the  principles  of  ionic  interaction.  Since  there  is  no  directed  bond- 
ing in  these  compounds,  the  configuration  is  determined  by  the  electrical 
interaction  of  the  four  ligands;  in  general,  they  have  a  like  charge,  so  mutual 
repulsion  leads  to  a  tetrahedral  arrangement.  The  possibility  for  this  con- 
figuration is  limited,  geometrically,  to  cases  in  which  the  ratio  of  the  radii 
of  the  ligand  atoms  to  that  of  the  central  atom  lies  between  0.225  and  0.4144. 
Such  compounds  have  been  said  to  contain  ionic  bonding  or  to  be  nonpene- 
tration  coordination  compounds. 

Planar  Configuration 

Mellor1  has  discussed  the  subject  of  square  planar  coordination  thor- 
oughly. The  earliest  indication  of  planar  configuration  was  Werner's  sug- 
gestion that  two  of  the  compounds  with  the  composition  Pt(NH3)2Cl2  were 
cis  and  trans  isomers.  He  further  postulated  which  was  which,  and  cor- 
related the  structures  of  the  isomers  with  their  chemical  behavior  by  means 
of  a  concept  he  called  "trans  elimination "x60.  Several  other  examples  of 
isomerism  among  platinum (II)  compounds  were  known  then,  or  were 
subsequently  discovered,  and  a  few  palladium (II)  compounds  were  known 
in  two  forms,  but  an  analogous  behavior  was  not  found  for  other  metals 
for  some  years.  As  a  result,  the  concept  was  questioned  more  and  more 
strongly,  and  the  problem  was  not  resolved  to  the  satisfaction  of  most  chem- 
ists until  the  advent  of  modern  structural  determinations. 

The  development  of  x-ray  techniques  for  structure  determinations  fur- 
nished the  additional  evidence  needed  to  satisfy  most  investigators.  Dickin- 
son demonstrated  a  square  planar  arrangement  of  the  chloride  ions  about 
the  platinum  or  palladium  atoms  in  K2PtCl4 ,  K2PdCl4  and  (NH4)2PdCl4 


vio 


2.  Mathieu,  Compt.  rend.,  204,  682  (1937). 

3.  Mellor,  J.  Proc.  Roy.  Soc,  N.  S.  Wales,  75,  157  (1942). 

1.  Wells,   "Structural  Inorganic  Chemistry,"  2nd  edition,  Oxford,  Oxford  Uni- 
versity Press,  1950. 


STEREOi  HEMISTRY  OF  COORDINATION  NUMBER  FOUR         357 

V  tew  years  later.  Pauling  explained' • €  the  planar  structure 
which  had  been  observed  for  platinum  1 1  ami  palladium (II)  compounds 
and  predicted  that  diamagnetic  compounds  of  nickel  (I  ]  .  gold  (III),  cop- 
per(III),  and  silver  III  are  also  planar.  This  theoretical  pronouncement 
created  a  renewal  of  interest  in  the  problem  and  a  large  number  of  papers 
on  the  subject  soon  appeared*"  a  < :  '•  v»-  X1"'  xir'  x  '.  For  t  he  most 

part,  these  confirmed  Pauling's  ideas,  but  some  investigators  attacked  the 
theory  of  planar  configurations  for  platinum(II),  palladium (II),  and 
nickel(II)u*0,  VM  N  M  v*  x  x:.  Others,  particularly  Jensen7,  answered 
these  objections  quite  adequately.  Jensen  made  extensive  use  of  dipole 
moment  determination-  to  show  the  existence  of  trans  planar  structures. 
While  Pauling's  prediction  that  diamagnetic  nickel(II)  and  gold(III)  com- 
pounds would  he  planar  was  verified,  it  was  also  found  that  some  silver(II) 
and  copper(II)  compounds  are  planar.  In  addition,  several  other  elements 
have  been  reported  to  exhibit  the  planar  configuration.  The  reports  are 
based  mainly  on  incomplete  x-ray  studies,  and  more  evidence  is  needed  to 
establish  the  results  conclusively.  Other  experimental  methods  which  have 
been  used  to  provide  evidence  for  planar  configuration  are  magnetic  meas- 
urements, crystal  optics,  and  resolution  into  optical  isomers. 

Theoretical  Considerations 

Isomer  Patterns  and  Configuration 

The  classical  chemical  method  for  stereochemical  investigation  involves 
preparation,  identification,  and  analysis  of  compounds,  separation  into 
isomers,  and  investigation  of  chemical  behavior.  After  a  compound  has 
been  prepared,  the  question  of  the  niimber  of  isomers  is  most  important  in 
this  method  of  attack.  Pfeiffer  elucidated  the  probable  isomer  patterns  for 
compounds  with  the  coordination  number  four,  assuming  the  regular  tetra- 
hedral,  square  planar,  and  pyramidal  configurations8.  His  result-  are  sum- 
marized in  Table  9.2.  There  is  experimental  evidence  for  the  tetrahedral 
and  planar  configurations,  but  there  is  no  case  of  isomerism  which  can  be 
explained  only  by  a  pyramidal  configuration,  although  the  isomer  pattern 
dte  distinct  for  this  configuration.  This  approach  is  illustrated  by  the 
tion  <>t'  three  geometrical  isomers  of  [Pt  |  XIb<  )II)(XH3)(py)(X02)]+ 
by  Chernyaev*9.  While  this  is  not  definitive  proof  that  the  ion  is  planar,  it 
certainly  eliminate-  the  tetrahedral  structure. 

Pauling,  •/.  .1       '  -      .  53,  1391     1931   . 

*6.  Paulii  -  " : 

7.  Jensei     '/.  .,  241,  ]  1 5 

ichemie,"  Freudenberg,  pp.  1210  ~>7.  Leipzig  and  Vienna,  Franz 

Deutic'r 


358 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Table  9.2.  Stereoisomers  of  Tetracoordinate  Structures 
(T  =  tetrahedral,  S  =  square  planar,  and  P  =  pyramidal) 


Ma* 

Ma2b2 

Ma2bc 

Mabcd 

M(AB)2 

T 

S 

P 

T 

S 

p 

T 

S 

P 

2 

1 

T 
2 

S 

3 

p 
6 

T 

2 

s 
2 

P 

Number  of  optically  active  isomers 
X umber  of  optically  inactive  isomers 

1 

1 

1 

1 

2 

2 

1 

2 

2 
1 

Configuration  and  Chemical  Reactions 

The  success  of  the  chemical  methods  of  determining  configurations  is 
dependent  upon  the  retention  of  configuration  during  reaction.  Although 
this  point  has  not  been  investigated  exhaustively,  the  coherence  of  the  facts 
and  theory  and  the  compatibility  with  configurations  assigned  by  physical 
methods  indicate  that  the  configurations  are  retained  in  reactions  of  plati- 
num complexes. 

Trans  Effect.  Some  interesting  principles  have  arisen  from  the  study  of 
the  reactions  of  platinum(II)  complexes.  These  originated  in  Werner's  ob- 
servations on  the  reactions  of  the  isomers  of  [Pt(NH3)2Cl2].  To  assign  the 
cis  and  trans  configurations,  he  assumed  that  a  phenomenon  which  he  called 
"trans  elimination"  was  operative  in  their  reactions.  This  concept  has  been 
further  developed  by  several  Russian  workers  and  is  now  one  of  the  guid- 
ing principles  in  the  assignment  of  cis  and  trans  structures  to  planar  com- 
plexes9, as  well  as  in  the  preparation  of  complexes  of  known  configuration. 
The  basic  postulate  is  that  in  a  substitution  reaction  the  group  trans  to 
the  most  electronegative  or  most  labilizing  group  will  be  replaced  (page 
146).  Fundamentally,  this  is  the  basis  for  "Kurnakov's  test"X29,  which  is 
used  frequently  by  the  Russian  investigators  to  assign  a  cis  or  trans  con- 
figuration to  a  diacido  platinum(II)  complex.  Treatment  of  the  complex 
[PtabX2]  with  thiourea  gives  [Pt(tu)4]X2  if  the  complex  is  cis,  but 
[Pt(tu)2ab]X2  if  the  complex  is  trans.  There  are  relatively  few  examples  of 
this  kind  of  isomerism  among  palladium(II)  compounds;  however,  these 
react  by  trans  elimination10.  Quagliano  and  Schubert  have  recently  dis- 
cussed the  trans  effect11.  The  phenomenon  has  been  well  documented  for 
only  a  few  classes  of  platinum(II)  compounds.  The  ultimate  resolution  of 
the  problem  awaits  the  extension  of  the  observations  to  a  broader  area. 

Reaction  with  a  Bidentate  Group.  When  a  complex  [Pta2b2]  is  treated 
with  a  bidentate  reagent  A  A,  the  cis  isomer  reacts  to  give  [Pt(AA)b2], 

9.  Chernyaev,  Ann.  inst.  platinc  (U.R.S.S.),  5,  102.,  118  (1927);  cf,  Chem.  Centr., 

1927,  II,  1557. 
id.  JonasseD  and  Cull,./.  .1///.  Chem.  Soc.,13,  274  (1951). 
•11.  Quagliano  and  Schubert,  Chem.  Revs.,  50,  201  (1052). 


1 


STEREOCHEMISTRY  OF  COORDINATION  NUMBER  FOUR        359 

whereas  the  trans  isomer  yields  (Pt(AA)ibJ,  [Pt(AA)abJ,  or  some  other 
compound  in  which  A  A  Functions  as  a  monodentate  ligand.  This  method 
for  assigning  configurations  is  also  widely  used  by  the  Russian  workers12,  u. 
Grinberg  showed  that  trans-[Pt(NO   .  Nil    _.  reacts  with  oxalic  acid  to 

form  [PI  Ilr  < ».  ,  NH,)J  whereas  the  cis  form  yields  [Pi ('(",< ),.  (NH,),]** 

Hybrid  I  >pe  and  Configuration 

In  general,  it  is  not  possible  to  predict  whether  a  configuration  will  be 

planar  or  tetrahedral.  However,  the  concepts  introduced  by  Pauling5, 6  and 
extended  by  Kimball14,  have  met  with  great  success  in  explaining  the  ob- 
served tacts  and  in  predicting  the  existence  of  diamagnetic  planar  nickel(II) 
and  gold(III)  compounds.  Pauling  approached  the  subject  from  the  con- 
sideration that  the  formation  of  covalent  bonds  between  the  central  atom 
and  the  Uganda  requires  an  overlapping  of  orbitals;  this  results  in  the  bonds 
being  so  oriented  in  space  that  maximum  overlapping  occurs.  From  a  con- 
sideration of  the  available  orbitals  in  the  atoms  of  any  element  it  is  possible 
to  predicate  what  spatial  arrangement  the  orbitals  will  take.  In  general, 
stronger  bonds  result  from  hybridization  of  the  diverse  orbitals  (angular 
strengths  are:  dsp2,  2.64;  sp*,  2.0).  Kimball  found  that  certain  combinations 
should  result  in  irregular  tetrahedral  or  pyramidal  configurations,  in  addi- 
tion to  the  regular  tetrahedral  and  square  planar  configurations  proposed 
by  Pauling,  but  these  possibilities  have  not  been  observed  for  discrete 
coordination  compounds.  Kimball's  results  are  summarized  in  Table  9.3. 

Table  9.3.  Configurations  of  Tetracoordinate  Complexes 

Configuration  Orbitals  Involved 

Regular  tetrahedral  sp3,  d3s 

Irregular  tetrahedral  d2sp,  dp3,  d3p 

Square  planar  dsp2,  d2p2 

Pyramidal  <i4 

On  the  basis  of  Pauling's  ideas,  Mellor  has  suggested  that  the  following 
species  might  also  exhibit  the  planar  configuration15:  cobalt  (I),  cobalt  (II), 
iron(II),  manganese(II),  manganese(III),  rhodium(I),  and  iridium(I). 
Although  he  has  searched  for  some  of  these,  the  planar  configuration  has 
not  yet  been  proven  for  them. 

The  Magnetic  Criterion.  It  is  possible  to  relate  the  magnetic  properties 
of  some  coordination  compounds  to  the  theory  just  discussed,  and  this  is 
one  of  its  striking  successes.  If  we  consider  only  the  3rf,  4.9,  and  4;;  levels  for 

12.  Gurin.  Doklady  Akdd.  Nauk  S.S.S.L'.,  50,  201,  205,  209  (1945);  cf.  Chem.  Abe.,  43, 

1674c,  h.  K.7.V    194 

13.  Ryabchik.  '.  rend.  acad.  set.  U.RJ5J3.,  41,  208  (191:;  . 

14.  Kimball,  .!.<■.     .  pj  ,.,.,  8,  188    L940  . 

15.  Mellor, ./.  P  oe   Ray.  .<.,,-..  N.  8.  Wales,  74,  129  (1940). 


360 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


SPECIES 


TETRAHEDRAL       Nl   ++ 


PLANAR       Nl  ++ 


ELECTRONIC       STRUCTURE 
3d  4s  4p 


•  •  • 

•  •  • 

•  [""•  r~5  I  • 

•  •  •     • 


□ 


□ 


Fig.  9.1.  Configuration  and  electronic  structure  of  nickel  complexes. 

nickel  (the  same  argument  will  apply  to  palladium  and  platinum  with  the 
substitution  of  the  proper  orbitals)  from  Hund's  rule  of  maximum  mul- 
tiplicity in  a  given  energy  level16,  the  nickel(II)  ion  should  have  two  un- 
paired electrons.  However,  if  nickel (II)  is  to  form  a  square  planar  compound 
with  dsp2  bonding,  these  electrons  will  have  to  pair,  as  indicated  in  Fig.  9.1, 
so  the  determination  of  the  magnetic  moment  of  nickel (II)  compounds 
shows  whether  dsp2  bonding  is  present.  All  the  elements  which  form  com- 
plexes with  planar  configurations,  except  copper (II)  and  silver (II),  should 
exhibit  different  magnetic  behavior  for  dsp2  bonding  than  for  the  bonding 
associated  with  the  tetrahedral  arrangement.  The  extension  of  the  magnetic 
criterion  to  elements  beyond  the  first  transition  series,  with  the  exception 
of  silver (II),  has,  however,  not  met  with  much  success17,  so  in  some  of  these 
cases  magnetic  data  will  have  to  be  supported  by  other  facts. 

Experimental  Proof  of  Configuration 

Werner's  Use  of  the  Concept  of  "Trans  Elimination" 

The  classical  arguments  of  WernerX50  are  of  historical  importance  and  are 
of  interest  in  showing  how  elegantly  a  gifted  mind  can  interpret  chemical 
data.  Two  forms  of  [Pt(NH3)2Cl2]  can  be  prepared,  one  by  the  reaction  of 
K2[PtClJ  with  aqueous  ammonia  and  the  other  by  the  reaction  of 
[Pt(NH3)4]Cl2  with  aqueous  hydrochloric  acid.  Two  forms  of  [Pt(py)2Cl2] 
are  also  known.  Treating  either  form  of  [Pt(NH3)2Cl2]  with  pyridine  yields 
[Pt(NH3)2(py)2]Cl2 ,  but  the  two  forms  of  reactant  yield  different  isomers 
of  the  product.  The  same  isomers  of  [Pt(NH3)2(py)2]Cl2  are  formed  by  the 
treatment  of  the  two  forms  of  [Pt(py)2Cl2]  with  ammonia.  When  these 
two  isomers  of  [Pt(NH3)2(py)2]Cl2  are  heated  with  hydrochloric  acid,  one 
yields  [Pt(NH3)(py)Cl2],  whereas  the  second  yields  a  mixture  of 
[Pt(NH3)2Cl2]  and  [Pt(py)2Cl2];  the  latter  two  compounds  are  identical 


16.  Hund,  Z.  Physik,  33,  345  (1925). 

17.  Mellor,  J.  Proc.  Roy.  Soc,  N.  S.  Wales,  77,  145  (1944) 


STEREOCHEMISTRY  OF  COORDINATION  NUMBER  FOUR 


361 


Cl-  a~  Cl  py4*  _         pv+  py 

CIPtCl  -^  CIPtNH,  -^»  CIPtNH,  -^»  pyPtNHa -^>  pyPtCS  -^  ClPtCl 

Cl  Cl  MI  MI  MI  MI 


or 

py 

CIPtNH 

MI 


>  same 


Cl- 


py 
.  ciptci 

MI 


MI 


NH3  + 

;PtCl 
MI 


NH,"  NH3-  NH,  NH^ci-      NHs  +  Cl-        NH' 

XH3PtXH,  ^»XH3PtCl    *=►  CIPtCl -^  pyPtpy     ^pyPtCl    -iU  CIPtCl 


XH, 


MI 


MI 

or 

Cl  + 

pyPtpy 

XH3 


cr 


MI 


Cl 

■»  pyPtpy 

Cl 


Fig.  9.2.  Trans  elimination  in  platinum(II)  complexes 


different 


with  the  isomers  from  which  the  second  form  of  [Pt(XH3)2(py)2]Cl2  was 
prepared.  An  outline  of  Werner's  explanation  based  on  trans  elimination  is 
shown  in  Fig.  9.2.  The  original  isomers  of  [PUpy^CU]  are  similar  in  con- 
figuration to  the  isomers  of  [Pt(XH3)2Cl2]  and  are  not  shown. 

Significance  of  Studies  on  Optical  Isomers 

Mills  and  his  coworkersV18,  X34  have  ingeniously  used  the  resolution  of  an 
asymmetric  substance  into  its  optical  isomers  to  gain  evidence  for  the 
planar  configuration  of  platinum(II)  and  palladium(II)  compounds.  Two 
chelating  groups,  isobutylenediamine  and  meso-stilbenediamine,  were  co- 
ordinated to  the  metal  ion.  The  ion  thus  formed  has  a  center  of  symmetry 
if  the  nitrogen  to  metal  bonds  are  tetrahedral,  but  is  asymmetric  if  the 
bonds  are  planar  (Fig.  9.3).  For  both  the  platinum  and  palladium  com- 
pounds, separation  into  optically-active  isomers  was  successful,  and  the 
cations  could  be  obtained  in  active  form,  free  of  the  material  used  for  reso- 
lution. After  destruction  of  the  complex,  the  amines  were  shown  to  be  in- 
active. Both  Mills  and  Jensen7  have  pointed  out  that  this  does  not  prove 
that  the  complexes  have  a  planar  configuration,  but  it  certainly  eliminates 
a  regular  tetrahedral  configuration. 

The  Role  of  X-Ray  Structure  Determinations 

Robertson  and  co-workers  have  carried  out  complete  x-ray  -tincture 
determinations  on  some  metal  phthalocyanines"*23,  D3,  LoJ'  x44.  Because  of 
the  large  number  of  atoms  involved,  this  is  a  particularly  interesting  ex- 
ample of  what  can  be  done  with  structure  determinations  in  favorable 


362 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


CfiH 


6n5 


C6H5 


■NH 


NH 


M 


NH 


NH; 


/ 


/ 


CH3' 
CH3 

H 


+  + 


B) 


H\ 

c6h/Sc 


C«H 


6nS 


/ 


NH2" 
NH2- 


M: 


NH 


NH; 


/CH3" 
\ 


CH- 


\ 


H 


+  + 


Fig.  9.3.  Configurations  of  tetracoordinate  complexes  containing  one  molecule  of 
isobutylenediamine  and  one  molecule  of  ???eso-stilbenediamine  bound  to  planar  (a) 
and  to  tetrahedral  (b)  central  atoms. 

cases.  The  large  organic  molecule  is  tetradentate,  with  the  four  coordinating 
nitrogen  atoms  at  the  corners  of  a  square.  The  planar  structure  (Fig.  9.4) 
does  not  vary  greatly  in  dimension  from  metal  to  metal  and  is  the  same  for 
all  of  the  metallic  ions,  irrespective  of  whether  they  ordinarily  form  planar 


Fig.  9.4.  Configuration  of  phthalocyanine  complexes  containing  divalent  metal 
ions.  M  =  Cu(II),  Be(II),  Mn(II),  Fe(II),  Co(II),  Ni(II),  or  Pt(II). 


STEREOCHEMISTRY  OF  COORDINATION  NUMBER  FOUR         363 

CH3         II  (II 

c         c         c 

/  \  /  \  /  \ 

EtOOCC  C  C  CCOOE1 

II  I  I 

CH,C N  \  CCH 

\   / 
M 

/    \ 

CllaC  N  N CCH 

I  I  I  II 

EtOOCC  C  C  CCOOEt 

\    /    \    /    \    / 
C  C  (' 

ch,      11         ch; 

Fig.  9.5.  Planar  arrangement  of  liickel(II)  and  palladium (II)  complexes  with 
pyrromethene. 

coordination  compounds  or  tetrahedral  ones.  The  stereochemistry  of  these 
compounds  is  determined  by  the  ligand  molecule. 

Although  actual  structure  investigations  have  not  been  carried  out,  a 
consideration  of  molecular  models  indicates  that  other  forced  configurations 
may  also  exist.  The  investigations  and  speculations  of  Porter  on  the  pyrro- 
mcthene derivatives18  have  been  continued  by  Mellor  and  Lockwood,  who 
measured  the  magnetic  moments  of  the  compounds  indicated  in  Fig.  9.5U2°. 
The  nickel  compound  is  paramagnetic  as  expected  for  a  tetrahedral  configu- 
ration, whereas  the  palladium  compound  is  diamagnetic.  The  stereochemical 
significance  of  this  is  not  known,  but  it  is  difficult  to  see  how  the  palladium 
complex  can  be  planar  since  the  bond  hybrid  is  most  probably  spz.  Mann 
and  Pope  have  prepared  nickel,  palladium,  and  platinum  complexes  with 
j3,/8' ,|8/,-triaminotriethylamine,  (XH^CHoCH^X;  these  should  be  tetra- 
hedral because  of  the  geometry  of  the  ligand19,  but,  again,  more  work  is 
required  to  complete  the  proof  since  octahedral  coordination  involving  sol- 
vent molecules  may  occur. 

Dipole  Moments 

A  very  complete  study  of  the  dipole  moments  of  the  compounds 
[PtX2(ER3)2]  (X  =  CI,  Br,  I,  X02 ,  or  N08  ;  E  =  P,  As,  or  Sb;  R  =  Et, 
Pr,  Bu,  or  (VJI5  ;  but  not  all  possible  combinations)  has  been  made  by 
Jensen  .  The  compounds  fail  into  two  groups,  one,  those  compounds  with 
zero  dipole  moment,  and  the  other,  those  with  an  appreciable  dipole 

18.  Porter,  J.  ■<-.,  1938,  368. 

19.  Mann  and  Pope,  J    Chem.  80c. ,  1926,  182;  Proc.  &  A109,  111 

(1925  ;  Cox  and  Webster,  Z.  Krist,  92,  is?  (19.35). 


364  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

moment.  Since  the  molecular  weights  of  some  of  these  substances  in  solu- 
tion show  them  to  be  monomeric,  the  forms  with  zero  dipole  moment  must 
be  trans  planar,  although  not  necessarily  square.  The  other  isomers  do  not 
have  to  be  cis  planar,  of  course,  but  might  have  any  of  a  variety  of  con- 
figurations. If  one  form  is  planar,  however,  it  is  reasonable  to  assume  the 
same  geometry  for  the  other  form,  expecially  since  x-ray  studies  have  shown 
the  planar  form  to  occur  in  the  solid  state.  With  the  possible  exception  of 
t,he  purely  chemical  studies  discussed  earlier,  this  study  probably  affords 
^he  best  demonstration  that  the  planar  configuration  is  not  destroyed  in 
olution  although  admittedly  the  use  of  a  nonpolar  solvent  does  not  sub- 
.ect  the  hypothesis  to  the  most  rigorous  test. 

Other  Properties 

Mellor  has  attempted  to  relate  various  properties  to  structure  so  that 
complete  x-ray  study  is  not  necessary  to  specify  a  configuration.  He  has 
used  magnetic  measurements  extensively,  particularly  in  assigning  planar 
or  tetrahedral  structures  to  nickel(II)  and  cobalt(II)  compounds.  He  has 
assumed  that  Pauling's  criteria  are  correct,  and,  on  the  basis  of  structures 
assigned  from  them,  he  has  studied  the  relationship  of  ligand  atom  to 
structure1724,  the  relationship  of  Cotton  effect  to  structure3,  and  the  rela- 
tionship of  absorption  spectra  to  structure20, 21.  In  no  case  is  there  a  clear 
pattern.  He  has  also  pointed  out  that  large  negative  or  positive  birefrin- 
gence in  the  crystal  indicates  a  planar  configuration A15.  Wells  has  amplified 
the  last  point4.  Lifschitz  has  related  the  color  of  nickel(II)  complexes  to 
their  structures1723,  and  Pauling  has  discussed  the  concept22.  More  recently 
Ray  and  Sen  investigated  the  magnetic  moments  and  colors  of  a  large 
number  of  copper(II)  complexes  and  concluded  that  the  penetration  com- 
plexes (i.e.,  dsp2  bonding)  have  magnetic  moments  of  1.66  to  1.81  Bohr 
magnetons  and  are  red,  brown,  or  violet,  whereas  the  nonpenetration  com- 
plexes have  moments  of  1.90  to  2.20  Bohr  magnetons  and  are  blue  to  green23. 
It  is  interesting  that  both  classes  are  said  to  have  planar  configurations 
although  Pauling's  considerations  would  not  predict  a  planar  configuration 
for  a  nonpenetration  type  of  complex. 

The  Relationship  of  Oxidation  State  of  Structure 

The  same  metal  in  different  oxidation  states  often  shows  different  co- 
ordinat  ion  numbers,  but  some  instances  are  known  in  which  an  element  has 
the  coordination  number  four  in  two  oxidation  states.    Copper  (I)   and 

20.  McKenzie,  Mellor,  Mills,  and  Short,  J.  Proc.  Roy.  Soc,  N.  S.  Wales,  78,  70  (1944). 

21.  Mills  and  Mellor,  J.  Am.  Chem.  Soc,  64,  181  (1942). 

*22.  Pauling,  "The  Nature  of  the  Chemical  Bond,"  2nd  edition,  pp.  81-6,  98-106, 

118-23,  Ithaca,  Cornell  University  Press,  1944. 
23.  Ray  and  Sen,  ./.  Indian  Chem.  Soc,  25,  473  (1948). 


-STEREOCHEMISTRY  OF  COORDINATION  NUMBER  FOUR         365 

silver(I)  form  tetrahedra]  compounds,  whereas  copper(II)  and  Bilver(II) 
form  planar  compounds**.  In  [Ni(CO)4],  the  nickel (0)  is  tetrahedra!    ;  in 

nickel(II)  compounds,  the  configuration  is  usually  planar  but  is  possibly 
tetrahedra]  in  gome  Cases84,  U1,  U1S.  The  oxidation  states  of  iron  and  cobalt  iii 
the  tetrahedra!  compounds  [Fe(C02)(NO)->]  and  [Co(CO)3(NO)]sl  are 
somewhat  of  a  problem  but  might  be  considered  to  be  2—  and  1  —  ,  respec- 
tively. The  only  tetracovalent  iron(II)  compound  of  which  the  configura- 
tion has  been  determined  completely  is  the  planar  phthalocyanine  . 
Cobalt (II)  is  reported  to  have  a  tetrahedral  configuration  in  some 
compounds  such  as  bis(salicyladelyde)  cobalt  (II)  and  bis(l  ,2-naphtha- 
lenediamine)cobalt(II)  acetate,  and  a  planar  configuration  in  others,  as 
exemplified  by  bis(a-benzildioxime)cobalt(II)  and  bis(thiosemicarbazide)- 
cobalt(II)T15. 

Bridged  Complexes 

The  aluminum,  gallium,  and  indium  halides,  and,  presumably,  the  cor- 
responding iron  (III)  and  gold(III)  chlorides  and  bromides  are  bimolecular 
in  the  gaseous  state.  Palmer  and  Elliott  have  shown  by  electron  diffraction 
that  the  aluminum  halides  have  a  bridged  structure  in  which  each  aluminum 
is  surrounded  by  a  tetrahedron  of  halide  ions,  the  tetrahedra  sharing  an 

j       H5 

edge     : 

X  X  X 

\  /  \  / 

Al  Al 

/     \    /    \ 
XXX 

(X  =  CI,  Br,  or  I) 

For  Au2X6  the  molecule  should  be  planar  with  the  two  square  AuX4  units 
sharing  an  edge4.  This  dimeric  structure  has  been  shown  for  [(Et2AuBr)2]C2, 
as  well  as  [(Me3AsPdCl2)2]  and  [(Me3AsPdBr2)2]V21.  It  is  interesting  that 
[(Pr2AuC\)i]  has  a  different  structure  because  of  the  rigidity  of  the  triple 
bond  between  carbon  and  nitrogen.  The  M — C  =  X — M  group  is  linear, 
and  double  cyanide  bridges  are  not  possible.  The  cyanide  group  can  serve 
as  a  bridging  unit  only  by  forming  a  large  square  molecule010: 

Pr  Pr 

I  I 

Pi-Au-CN-Au-Pr 

I  I 

N  C 

C  X 

I  I 

Pr-Au-XC-Au-l'i 

I  I 

Pr  Pr 

24.  Nyholm,  Quart.  Revs.,  3,  321  (1949). 


366 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Fig.  9.6.  Basic  unit  [Mo6Cl8]4+in  the  structures  of  [Mo6Cl8](OH)4-14H20,Mo6Cl12- 
8H20,  and  HMo3Cl7H20.  •,  Mo;  O,  CI. 


Fig.  9.7.  Structures  of  the  Nb6Cli2  ,  Ta6Bri2  and  Ta6Cli2  groups.  The  double  cir- 
cles represent  metal  atoms  and  the  single  circles  halogen  atoms. 

Extremely  interesting  structures  have  been  found  for  [Mo6Cl8](OH)4- 
14H20,  Mo6Cl12-8H20,  and  HMo3Cl7-H20Qlt  Q2,  Q3.  All  of  these  compounds 
conl  ain  the  polynuclear  unit  [Mo6Cl8]4+,  the  structure  of  which  is  shown  in 
Fig.  9.6.  Pauling  has  suggested  that  each  molybdenum  atom  forms  bonds 
with  the  four  chlorine  atoms  on  the  face  of  the  cube  nearest  it  in  a  nearly 
coplanar  configuration25.  Each  chlorine  is  shared  by  three  molybdenum  at- 
oms. A  related  structure  has  been  found  for  Nb6Cli4-7H20  and  TaeClw 
711  <  >'N1,  in  which  the  central  octahedron  of  metal  atoms  is  surrounded 
by  twelve  chloride  ions  (Fig.  9.7)  and  each  metal  ion  has  four  chlorine 
atoms  in  a  nearly  square  coplanar  relation  to  it. 

25.   Pauling,  Chem.  Eng.  News,  26,  2970  (1947). 


STEREOCHEMISTRY  OF  COORDINATION  NUMBER  F<>(  R         367 

AMHHiirni-    Ai;iM\o    FROM    SOME   OF  THE   TECHNIQUES    EMPLOYED 

to  Establish  Configi  rations 

Incomplete  X-raj    Analysis 

Unfortunately  several  of  the  conclusions  concerning  the  configurations 
of  four-coordinate  complexes  are  based  upon  incomplete  studies.  This  is 
particularly  true  of  the  x-ray  studies,  and  in  some  cases  this  has  led  to 

results  which  were  later  shown  to  be  incorrect.  A  structure  based  on  x-ray 
analysis  which  is  carried  only  to  the  unit  cell  dimensions  may  well  be  in 
error.  It  is  safer  to  include  also  symmetry  considerations  from  the  space 
group,  but  even  this  has  been  insufficient  to  yield  final  answers  in  some 


6 


Fig.  9.8.  Structure  of  Cs2Au2Cl6  .  •  =  Au;  O  =  CI. 

cases.  For  example,  on  this  basis,  Cox,  Shorter,  and  Wardlaw  reported  that 
K_Sn('l4-2H20  contains  planar  [SnCl4]=  groupings,  but  Brasseur  and  de 
Rassenfosse  showed  by  a  complete  analysis  that  the  structure  consists  of 
infinite  chains  of  [SnCl6]~4  octahedra,  sharing  edgesL1.  It  appears  that  the 
first  investigators  considered  only  discrete  coordination  units  and  were 
able  to  rule  out  the  tetrahedral  unit  but  neglected  to  consider  the  possibility 
of  condensed  structures.  Because  of  this  possibility  of  condensed  st  ructures, 
coordination  numbers  obtained  from  chemical  analysis  alone  may  not  have 
much  meaning.  For  example,  the  formula  Cd(XH3)2Cl2  appears  to  corre- 
id  to  a  compound  of  coordination  number  four,  but  actually,  the  struc- 
ture consists  of  condensed  octahedra  similar  to  those  of  the  l\JSn(,lr2H20 
structure*.  On  the  other  hand,  in  CsCuClj ,  each  copper  atom  is  Bquare 
planar,  and  the  structun  1  to  consist   of  infinite  chains  of  CuCU" 

units  joined  by  opposite  corners*29. 

26.  MacGillavry  and  Bijvoet,  Z.  Krist.,  94,  231  (1936). 


368  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Even  though  there  is  agreement  on  atomic  coordinates  there  may  still 
be  disagreement  on  the  structural  interpretation.  For  example,  Elliott 
and  Pauling  interpreted  the  structure  of  Cs2Au2Cl6  (Fig.  9.8)  as  containing 
planar  [AuCl4]~  and  linear  [AuCl2]~  units,  whereas  Ferrari  concluded  that 
the  gold  (III)  occurs  in  octahedral  [AuCle]~  units05.  Similar  disagreements 
exist  about  the  structures  of  K2CuCl4-2H20  and  CuCl2-2H20A1°- 27- 28.  The 
point  in  dispute  is  the  degree  to  which  the  different  metal  to  ligand  bonds 
can  vary  in  length  and  still  be  considered  part  of  the  coordination  sphere. 
In  Cs2Au2Cl6 ,  for  instance,  there  are  four  Au(III)-ClI  distances  of  2.42  A., 
and  two  Au(III)  —  Cln  distances  of  3.13  A.  This  problem  does  not  arise 
with  the  tetrahedral  structure,  and  with  some  of  the  platinum(II)  com- 
pounds the  structure  is  clearly  planar  since  there  are  only  four  groups  within 
a  reasonable  distance  of  the  platinum.  An  example  is  found  in  the  structure 
of  K2[PtCl4]  shown  in  Fig.  9.9V10.  Since  most  planar  structures  can  be  in- 
terpreted as  octahedral  in  the  condensed  phase,  it  has  been  suggested  that 


x> 


Fig.  9.9.  Structure  of  K2PtCl4  and  K2PdCl4  .  •  =  Pt  or  Pd;  O  =  CI. 

a  planar  structure  should  be  established  for  some  compound  in  the  gaseous 
state29.  So  far,  this  has  not  been  accomplished. 

Uncertainties  in  the  Resolution  of  Some  Optical  Isomers 

There  have  been  several  reported  resolutions  in  which  the  coordination 
compound  has  not  been  obtained  in  optically  active  form  free  of  other 
optically  active  groups.  In  these  investigations,  some  separation  into  dia- 
stereoisomers  is  accomplished,  and  the  supposed  diastereoisomers  are  shown 
to  undergo  mutarotation  in  solution.  When  the  optically-active  resolving 
component  is  removed,  however,  the  solution  of  the  coordination  compound 
is  not  optically  active.  It  is  assumed  that  the  coordination  compound  race- 
mizes  so  rapidly  that  the  active  form  cannot  be  detected.  Undoubtedly 
some  of  the  compounds  reported  to  be  tetrahedral  on  the  basis  of  such  evi- 
dence are  tetrahedral,  but,  in  view  of  the  inconclusive  nature  of  such  studies, 
it  is  desirable  to  have  additional  proof  before  considering  the  structures  to 

27.  Chrobak,  Z.  Krist.,  88,  35  (1934). 

28.  Neuhaus,  Z.  Krist.,  97,  28  (1937). 

*29.  Fernelius,  "Chemical  Architecture,"  Burk  and  Grummit,  pp.  84-90,  New  York, 
Interscience  Publishers,  1948. 


STEREOCHEMISTRY  OF  COORDINATION  NVMIiER  FOUR 


369 


be  established.  One  of  the  more  vigorous  attacks  on  the  theory  of  the  planar 
structure  of  some  platinum(II)  compounds  was  based  on  incomplete  resolu- 
tions of  this  sortX!  [  .  Reihlen  and  bis  collaborators  reported  optical 
activity  resulting  from  the  asymmetry  of  the  complex  in  bis(isobutylenedi- 

amine)platinum(II)  and  bis(isobutylenediamine)palladium(II)  ions,  and 
also  with  a  number  of  complex  species  containing  active  donor  molecules 
and  platinum(II).  However,  other  investigators  were  not  able  to  duplicate 
the  reported  partial  resolutions117,  Xu;,  so  this  work  is  generally  questioned. 

Inconsistencies  Among  Observed  Oxidation  States  and  those  Pre- 
dicted by  the  Atomic  Orbital  Theory 

The  question  of  why  eopper(II)  and  silver(II)  form  planar  complexes  and 
yet  show  no  great  tendency  to  be  oxidized  to  the  tervalent  state  is  an 
intriguing  one.  On  the  basis  of  Pauling's  theory,  the  behavior  of  gold  is 
readily  explained,  i.e.,  gold(I)  and  planar  gold(III)  compounds  exist,  but 
there  is  no  satisfactory  evidence  for  gold(II)  compounds.  The  electrons  in 
the  outermost  d,  s,  and  p  levels  and  the  bonding  possibilities  are  shown  in 
Fig.  9.10  for  the  atom  in  oxidation  states  0,  I,  II,  and  III.  The  tetrahedral 
configuration  observed  for  silver(I)  and  copper(I),  the  linear  configuration 
for  all  three  univalent  atoms,  and  the  planar  configuration  for  gold  (III) 
are  in  agreement  with  Pauling's  treatment.  Pauling22  explained  the  planar 
structure  of  the  copper(II)  compounds  by  assuming  that  the  dsp2  planar 


OXI  DATION 

STATE 

ELECTRONIC 

d 

STRUCTURE 
s                P 

0 

0       o        o       o       o   1 
0       o       o       o       o  1 

0  L 

1 

o        o       o       o       o   1 
O        0       0        o       o  1 

□  c 

1 

J 

2 

0        o       o        o 

o 
o       o       o       o 

D  C 

O        0        o        o 

O         O         0          o 

□  c 

o 

3 

o        o       o       o 
0       o       o       o 

□  c 

Pig.  9.10.  Electronic  structures  of  the  atoms  and  ions  of  copper,  silver,  and  gold. 


370  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

bonds  are  enough  stronger  than  the  spz  bonds  so  that  the  slight  difference 
in  energy  arising  from  the  promotion  of  the  unpaired  electron  from  a  3d 
orbital  to  a  Ap  orbital  is  more  than  offset.  If  this  argument  is  correct,  it  is 
difficult  to  see  why  copper(III)  and  silver(III)  compounds  are  so  hard  to 
prepare.  The  chemistry  of  gold,  on  the  other  hand,  is  what  one  would  expect. 

Inferences  Based  on  the  Atomic  Orbital  Concept  and  on  Analogy 
to  Known  Structures 

In  conclusion,  some  deductions  with  regard  to  probable  structures  will 
be  mentioned.  It  has  been  proposed  that  K4[Ni(CN)4]  and  K4[Pd(CN)4] 
should  be  tetrahedral,  since  the  central  atoms  resemble  Ni(0)  in  [Ni(CO)4] 
in  having  an  apparent  oxidation  state  of  zero30.  Linstead  and  co-workers 
have  prepared  several  phthalocyanine  derivatives  which  have  not  been 
examined  by  x-ray  methods,  but  which  almost  surely  are  planar31.  Thal- 
lium (III)  has  been  reported  to  form  both  tetrahedral  and  planar  com- 
poundsK1,  K3,  but  a  planar  configuration  is  unlikely  if  dsp2  or  d2p2  bonding 
is  required  for  its  existence,  since  only  spz  orbitals  are  available;  and  vacat- 
ing of  a  d  orbital  would  require  promotion  of  a  pair  of  electrons  from  the  d 
level  to  the  p  level  of  the  valence  shell.  The  structure  of  the  compounds  con- 
taining central  atoms  with  inert  electron  pairs  is  also  of  interest.  From 
incomplete  x-ray  work,  thallium(I)K1'  K2  and  lead (II) L1  are  reported  to 
form  planar  complexes.  Complete  structure  determinations  of  some  com- 
pounds in  this  group  should  be  made  to  determine  whether  the  coordina- 
tion number  is  really  four  or  if  a  condensed  octahedral  system  is  present. 

Annotated  Bibliography 

The  sources  cited  below  on  the  stereochemistry  of  four-covalent  com- 
pounds are  listed  by  periodic  family.  The  symbols  used  to  indicate  the  kind 
of  experimental  work  involved  are:  C,  crystal  optics;  CE,  Cotton  effect; 
D,  dipole  moment;  E,  electron  diffraction;  G,  isolation  of  geometrical 
isomers;  I,  isomorphism;  IR,  infrared  spectrum;  M,  magnetic  moment;  0, 
isolation  of  optical  isomers;  R,  Raman  spectrum;  X,  x-ray  diffraction.  If  a 
symbol  is  preceded  by  "i",  e.g.,  iX,  it  indicates  an  incomplete  study;  while 
(?)  indicates  simply  that  the  evidence  reported  supports  the  structure  listed. 

30.  Deasy,  J.  Am.  Chem.  Soc,  67,  152  (1945). 

31.  Barrett,  Dent,  and  Linstead,  J.  Chem.  Soc,  1936,  1719. 
♦32.  Cox  and  Wardlaw,  Science  Progress  32,  463  (1938). 

*33.  Hiickel,    "Anorganische   Strukturchemie,"   pp.    115-29,   Stuttgart,   Ferdinand 

Enke  Verlag,  1948. 
*34.  PfcifTer,  ./.  prakt.  Chem.  162,  279  (1943). 
•36.  Sidgwick  and  Powell,  Proc.  Roy.  Soc.  (London),  A176,  153  (1940). 


STEREOCHEMISTRY  OF  COORDINATION*   NUMBER  FOX  R         371 

Family  I: 

Copper  I  and  silver(I)  form  tetrahedraJ  complexes  and  silver(II)  and 
gold(III),  square  planar  ones.  Several  copper(II)  compounds  arc  square 
planar,  hut  at  least  one  is  tetrahedral.  Incomplete  x-ray  studies  indicate 
that  gold  (I)  forms  square  planar  bonds. 

Al.  Bezzi,  Bua,  and  Schiavianto,  Gazz.  chim.  ital.,  81,  856  (1951).  X.  In  copper  di- 
methylglyoxime  the  1  N  atoms  and  the  Cu  atom  are  coplanar. 

A2.  Barclay  and  Nyholm,  chemistry  A  Industry  1953,  378.  M.  Cul.  CH  As  (\  II.  \^ 
i'  11:     i  contains  tetrahedral  Cu  I 
Brink  and  van  Arkd.  Acta  Cryei.  5, 506(1952).  X.  (NH^iCuCli  and  (NH4)iCuBri 

contain  infinite1  chains  of  [CuX4]3"    tetrahedra. 
Al.   Brink,  Binncndijk,  and  van  de  Linde,  Acta  Cryst.  7,  170  (1954).  X.  CsCu*Cl« 

contains  infinite  double  chains  of  [CuCl4]3_  tetrahedra. 
A.V  Brink  and  MacGillavry,  Acta  Cryst.,  2,  158  (194!»  .  X.  K.CuCh  contains  infinite  , 

chains  of  [CuChl  tetrahedra.  ' 

A»'-.  Cambi  and  Coriselli,  Gazz.  chim.  ital.,  66,  779  (1936).  IM.  Some  compounds  of 

the  type  [(R»NCS3  tCu   are  tetrahedral.  (?). 
A7.  Cox,  Bharratt,  Wardlaw,  and  Webster,  /.  chem.  Soc,  1936, 129.  iX.  [Cu(py)2Cl2] 

and  [{CH  C:N  OB  C:N  I  )H)CH3}CuCl2l  have  planar  configurations. 
As.  (  Jox,  Wardlaw.  and  Webster, ./.  Chem.  Soc,  1936,  775.  iX,  C.  [(C5H4XCOO)2Cu]- 
_'11.<)   is   planar;  K,[Cu(CN)4]    is    tetrahedral;    X.[Cu  {SC(XH2)CH3}4]C1    is 
tetrahedral;  G.  UC5H4XCOO)2Cu]  is  planar. 
A1'.  Cox  and  Webster,  ./.  chem.  Soc.,  1935,  731.  iX,  C.  [{C6H4(0)(CH:XOH)}2Cu] 
ami  some  substituted  Cu  /3-diketonates  are  planar. 
A10.  Harker,  Z.  Krist.,  93,  136  (1936).  X.  [CuCl2(H20)2]  contains  planar  Cu. 
All.  Helmholz and Kruh,/.  Am.  Chem. Soc., 74, 1176  (1952).  X.  Cs2[CuCh]  contains 

tetrahedral  [CuCl4]=\ 
A12.  Koyama,  Baito,  and  Kuroya,  ./.  Inst.  Polytech.  Osaka  City  Univ.  Ser.  C,  4,  43 

(1953).  X.  Copper  acetylacetonate  is  planar. 
A13.  Lifschitz,  Z.  phys:  Chem.,  114,  491  (1925).  CE.  [Cu(d-oca)4]++  contains  tetrahe- 
dral eopper  (oca  =  oxymethylenecamphor). 
A14.   Mann.  Purdie.  and  Wells.  ./.   Chem.  Soc.  1936,   1503.   X,  I.  In   [(Et^UCuI)4l, 

Li  AsCuBr  ;!.  [Et»PCuI)4],  Cu(I)  is  tetrahedral. 
A 15.  Mellor  andQuodling,  J.  Proc  Roy.  Soc,X.S.  Wales,  70,  205  (1936).  C.  Cs2[CuCl4] 

and  [CuCl,  H20)2]  are  planar. 
Alo.  Mill.,  and  Gotts,  /.  (  hen  .  Sue.  1926,  3121.  iO.  [Cu{C6H5C(0— ):CHC(:0)- 

M  >Na}»]  is  tetrahedral. 
A17.  Muller,  Naiurwissenschaften  37,  333  (1950).  Copper  phthalocyanine  molecules 

appear  planar  in  the  field  elect  pod  microscope. 
A18.  Peyronel,  Gazz.  chim.  ital.,  73,  89    1943).  X.  [<Pr,XCS2)Cul  is  planar. 
Al'.e  Pfeiffer  and  Glaser,  ./.  prakt.    Chem.,   153,   265    (1939).    G.    [Cu{C10H6(O— ) 

(CH:NCH    [1    is  planar. 
A20.  Ray  and  Chakravarty,  ./.  Indian  Chen  .  Soc.,  18,  609  (1941  .  G.  [|C«H  NHC 

:NH  NIK'  Ml    :N}tCu]  is  planar. 
A21.  Ray  and  Dutt,  J.  Ind  -        26,51       1948     G       "  ><    1I\11:M1 

MK    Ml.  :M1    (    .    u  planar. 
A22.  R&yandGhoc  .26,  Ml    1949).  G.  [|Et«NC  :NH  MK'- 

l]  is  planar. 


372  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Aim.   Robertson,  ./.  Chem.  Soc,  1935,  615.  X.  Copper(II)  phthalocyanine  is  planar. 
A24.  Robertson,  ./.  Chem.  Soc,  1951,  1222.  X.  Copper(II)  tropolone  is  planar. 
A25.  Schlesinger,  Bar.,  58,  1877  (1925).  G.  [{OOCCRR,NH(CH2)xNHCRR"COO}Cu] 

is  planar. 
A26.  Shugam,  Doklady  Akad.  Xauk  S.S.S.R.,  81,  853  (1951);  cf,  Chem.  Abstracts,  46, 

3894d  (1952).  X.  Copper  acetylacetonate  is  planar. 
A27.  Stackelberg,  Z.  anorg.  allgem.  Chem.,  253,  136  (1947).  iX,  G.  Some  chelates 

formed  from  Cu(II)  and  aryl  aldimines  are  planar. 
A28.  Watanabe  and  Atoji,  Science  (Japan),  21,  301  (1951);  cf,  Chem.  Abstracts,  45, 

9982f  (1951).  X  (?).  [Cu(en)2]++  is  planar. 
A29.  Wells,  /.  Chem.  Soc,  1947,  1662.  X.  CsCuCl3  contains  infinite  chains  of  square 

planar  [CuCl4]=  units. 
Bl.  See  A4.  X.  CsAg2I3  contains  infinite  double  chains  of  [Agl4]3_  tetrahedra. 
B2.  Brink  and  Stenfert  Kroese,  Acta  Cryst.  5,  433  (1952).  X.  K2AgI3  ,  Rb2AgI3  ,  and 

<\H4)2AgI3  contain  infinite  chains  of  [Agl4]3~  tetrahedra. 
B3.  See  A5.  X.  Cs2AgCl3  and  Cs2AgI3  contain  infinite  chains  of  [AgCl4]s  tetrahedra. 
B4.  See  A8.  X.  [Ag{SC(NH2)CH3}4]Cl  contains  tetrahedral  Ag(I).  I,  C.  [(C5H4- 

NCOO)2Ag]  is  planar. 
B5.  Hein  and  Regler,  Ber.,  69B,  1692  (1936).  iO.  [Ag(C9H6XO)(C9H6NOH)]  and 

[Ag(C9H6XOH)2]X03  contain  tetrahedral  Ag(I). 
B6.  Mann,  Wells,  and  Purdie,  /.   Chem.  Soc,  1937,   1828.   I.  [(Pr3AsAgI)4]  (and 

[(Et3AsAgI)4]  ?)  contain  tetrahedral  Ag(I). 
CI.  Brain,  Gibson,  Jarvis,  Phillips,  Powell  and  Tyabji,  J.  Chem.  Soc.  1952,  3686.  X. 

(C7H7)2SAuCl2  contains  planar  SAuCl3  units. 
C2.  Burawoy,  Gibson,  Hampson,  and  Powell,  J.  Chem.  Soc,  1937,  1690.  X,  C,  D. 

[Et2AuBr]2  is  planar. 
C3.  Cox  and  Webster,  J.  Chem.  Soc,  1936,  1635.  X.  K[AuBr4]-2H20  contains  planar 

[AuBr4]~  ions. 
C4.  Dothie,  Llewellyn,  Wardlaw,  and  Welch,  /.  Chem.  Soc,  1939,  426.  iX.  [Au(CN)2- 

dipy]"  and  [Au(CN)2(o-phen)]~  are  planar. 
C5.  Elliott  and  Pauling,    /.    Am.    Chem.    Soc,    60,    1846    (1938).    X.    Cs2Au2Cl6 

and  Cs2AgAuCl6  contain  planar  [AuCl4]~  units.  Ferrari,  Gazz.  chim.  ital.,  67, 

94  (1937),  however,  believes  that  the  Au(III)  is  octahedrally  coordinated. 
C6.  Goulden,  Maccoll,  and  Millen,  J .  Chem.  Soc,  1950,  1635.  R.  The  Raman  spec- 
trum of  [AuCl4]~  is  consistent  with  a  planar  configuration. 
C7.  Huggins,  unpublished  work  referred  to  by  Huggins  in  /.  Chem.  Ed.,  13,  162 

(1936).  iX  (?).  [Me4N][AuCl4]  contains  planar  [AuCl4]~  ions. 
C8.  See  A15.  C.   [Me4N][AuCl4],  Na[AuCl4]-2H20,  and  K[AuBr4]  contain  planar 

Au(III). 
C9.  Perutz  and  Weisz,  J.  Chem.  Soc,  1946,  438.  iX.  [Me3PAuBr3]  is  planar. 
CIO.  Phillips  and  Powell,  Proc  Roy.  Soc.  (London),  A173,  147  (1939).  X.  [(Pr2AuCX)4] 

is  planar. 

Family  II: 

Beryllium (II),  zinc(II),  cadmium (II),  and  mercury(II)  are  tetrahedral. 
Beryllium(II)  is  planar  in  the  phthalocyanine.  The  report  that  cadmium(II) 
may  be  planar  appears  spurious. 

Dl.  Bragg  and  Morgan,  Proc.  Roy.  Soc  (London),  A104,  437  (1923).  X.  [Be40(AcO)6l 
contains  tetrahedrally  coordinated  Be(II). 


STEREOCHEMISTRY  OF  COORDINATION  NUMBER  FOUR         373 
D2.  Burgess  and  Lowry,  ./.  Chem.  Six-.,  125.  2081   (1924).  iO.   Beryllium  bensoyl 

camphor,    [CuHuOsBe],   is   tet  rahedral. 
D3.  Linstead  and  Robertson, ./.  Chem.  8oe.,  1936,  L736.  X  .  Beryllium  phthalocyanine 

is  planar. 
D4.  O'Daniel  and  Tscheischwili,  Z.  Krist.,  103.  178  (1941).  iX.  Xa,[BeF,]  contains 

tetrahedral  [BeF4J- 
D5.  O'Daniel  and  Tscheischwili,  Z.  Krist.,  104,  348  (1942).  I.    K8[BeF4]  contains 

tetrahedral  [BeF4l-. 
D6.  Hultgren,  Z.  Kriet.,  88,  233  (1934).   I.  X.  <  XII,)2[BeF4]  contains  tetrahedral 

BeF4l- 
D7.  See  AJ6.  O.  [Be{C,B  C  0     I :CHC(:0)COONa}«]  is  tetrahedral. 
D8.  Busch  and  Bailar,  ./.  Am.  Chem.  Soc.,  76,  5352  (1954).  O.  Partial  resolution  of 

bis(ben£oylacetone)beryllium    indicates    tetrahedral    configuration.    Com- 

pound  did  not  racemize  completely  in  five  hours. 
El.  Couture  and  Mathieu,  .1////.  Phys.,  [12]  3,  521   (1948).  R.  [Zn(CN)4]~  is  tetra- 
hedral in  solution. 
K2.  Danilov.  Finkelstein,  and  Levashevich,  Physik  Z.  Sowjetunion,  10,  223  (1936). 

X.  [Znl4]"  is  tetrahedral  in  solution. 
E3.  Dickinson,  ./.  Am.  Chew.  Soc,  44,  774  (1922).  iX.  K2[Zn(CN)4]  contains  tetra- 
hedral [Zn(CN)4l- 
E4.  Klug  and  Alexander,  J.  Am.  Chem.  Soc.,  66,  1056  (1944).  X.  (NH4)3ZnCl5  con- 
tains tetrahedral  [ZnCl4]=. 
E6.   I. in  and  Bailar.  J.  Am.  Chem.  Soc,  73,  5432  (1951).  O.  [(H03SC9H5NO)2Zn] 

contains  tetrahedral  Zn(II). 
E6.  MacGillavry   and  Bijovet,  Z.  Krist.,  94,  249    (1936).  X.   [Zn(NH3)2Cl2]   and 

[Zn(NH,)sBr8]  are  tetrahedral. 
E7.  Mills  and  Clark,  J.  Chem.  Soc,  1936,  175.  iO.  K2[Zn(CH3C6H3S2)2]  contains 

tetrahedral  Zn(II). 
E8.  See  A16.  iO.  [Zn{C6H5C(0— ):CHC(:0)COOXa)2]  is  tetrahedral. 
PI.  Brasseur  and  Rassenfosse,  Z.  Krist.,  95,  474  (1936).  I.  The  .Cd  in  Ba[CdCl4]- 

4H20  is  planar.  Quodling  and  Mellor,  Z.  Krist.,  97,  522  (1937),  question  the 

isomorphism  on  which  this  result  is  based. 
-    See  E3.  iX.  Ki[Cd(CN)4]  contains  tetrahedral  Cd(II). 
F3.   Evans,  Mann,  Peiser,  and  Purdie,  ./.  Chem.  Soc,  1940,  1209.  iX,  I.  [(Et3P)2- 

Cd«Br4]    and   similar  compounds  contain  bridged  tetrahedral  Cd  units.   A 

tetrahedral  structure  is  inferred  for  [(RjP)jCdX*]. 
F4.  See  E  7.  iO.  K2[Cd(CH3C6H3S2)2]  contains  tetrahedral  Cd(II). 
F5.  Pitaer,  /..  Krist.,  92,  131   (1935).  X.  [Cd(NH3)4](Re04)2  contains  tetrahedral 

[Cd(XH3)4]f+. 
CI.  See  E2.  X.  [Hgl*]"  is  tetrahedral  in  solution. 
'  r2.  Delwaulle,  Francois,  and  Weimann,  Compt.  rend.,  206,  1108  (1938).  R.  [HgBr4]- 

i>  tetrahedral. 
Bee£3    iX.  K    Bg(CN    i    Contains  tetrahedral  |Hg<<  'X 

•    iX.  I.  [(Pr»P  tHgsBr4]  and  similar  arsine  compounds  contain  l>ridged 

tetrahedral  Bg units.  A  tetrahedral  structure  is  inferred  for  |  1!  I'  iHgXj]. 
(,.">.  Jeffery,  Nature,  159,  610    l'.'17  .  X.  Co[Hg  SCN)4]  contains  tetrahedral  HgS4 

uni 
I  a;.  Ketelaar,  Z.  Krist.,  80,  L90    1931  .X.I.  kg    Bgl4]  andCu,[HgI4]  contain  tetrs 

hedral    Unh}-. 
1 .7    Bee  i:7.  iO.  K,[Hg  (11  (Ml  Bj  1]  contains  tetrahedral  Hg(II). 
G8.  Scouloudi,  A  I.  6,  051  (1953).  Same  as  G9. 


374  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

G9.  Scouloudi  and  Carlisle,  Nature,  166,  357  (1950).  X.  [Cu  (en)2][Hg(SCN)4]  con- 
tains tetrahedral  HgS4  units. 

Family  III: 

Aluminum(III),  gallium  (III),  and  indium(III)  are  tetrahedral.  The  evi- 
dence for  the  structure  of  complexes  containing  thallium  (I)  and  thal- 
lium(III)  is  incomplete  and  is  conflicting  in  the  latter  case. 

HI.  Baenziger,  Acta  Cryst.,  4,  216  (1951).  X.  Na[AlCl4]  contains  tetrahedral  [AlCh]". 
H2.  Gerdingand  Smit,  Z.  phys.  Chem.,  B50,  171  (1941).  R.  A12X6 ,  with  X  =  Cl~  Br", 

or  I~,  contains  bridged  tetrahedral  Al  units.  Kohlrausch  and  Wagner,  Z. 

Phys.  Chem.,  B52,  185  (1942),  say  the  Raman  spectrum  does  not  contradict 

such  a  structure,  but  does  not  prove  it. 
H3.  Harris,  Wood,  and  Ritter,  J.  Am.  Chem.  Soc.,  73,  3151  (1951).  X.  Fused  A1C13 

contains  paris  of  bridged  [A1C14]~  tetrahedra. 
H4.  Lippincott,  J.  Chem.  Phys.,  17,  1351   (1949).  R,  IR.  Li[AlH4]  contains  tetra- 
hedral [A1H4]". 
H5.  Palmer  and  Elliott,  J.  Am.  Chem.  Soc.,  60,  1852  (1938).  E.  A12X6  ,  with  X  = 

Cl~,  Br-,  or  I~,  contains  bridged  tetrahedral  Al  units. 
II.  Brode,  Ann.  Physik,  [5]  37,  344  (1940).  E,  Ga2Cl6  and  Ga2Br6  vapors  contain 

bridged  tetrahedral  Ga  units. 
Jl.  See  II.  E.  In2X6  ,  X  =  CI-,  Br-,  or  I-,  contains  bridged  tetrahedral  units  in  the 

vapor. 
J2.  Wood  and  Ritter,  J.  Am.  Chem.  Soc,  74,  1760  (1952).  X.  Fused  Inl3  contains 

bridged  tetrahedral  units. 
Kl.  Cox,  Shorter,  and  Wardlaw,  /.  Chem.  Soc,  1938,  1886.  iX.  [Tl(tu)4]N03  or  chlo- 
ride contains  planar  [Tl(tu)4]+,  whereas  [Me2Tl{CH3C(:0)CH:C(0— )CH3)] 

is  tetrahedral. 
K2.  Wardlaw,  unpublished,  1940,  cited  by  Sidgwick  and  Powell,  Proc  Roy.  Soc. 

{London),    A176,    153     (1940).    X(i?).     [Tl(o-phen)2]N03    contains    planar 

[Tl(o-phen)2]+. 
K3.  Watanabe,  Saito,  Shiono,  and  Atoji,  "Structure  Reports  for  1947-8,"  Vol.  11, 

pp.  393-4,  edited  by  Wilson,  N.  V.  A.  Oosthock's  Uitgevers  mij  Utrecht,  1951. 

iX.  CsTlBr4  contains  planar  [TlBr4]~. 

Family  IV: 

The  evidence  that  lead(II)  and  tin(II)  are  planar  is  incomplete.  The 
tin(II)  compound  has  been  shown  to  have  a  condensed,  not  discrete,  struc- 
ture involving  coordination  number  six. 

LI.  Cox,  Shorter,  and  Wardlaw,  Nature,  139,  71  (1937).  iX.  R2[SnX4]-2H20,  with 
R+  =  K+  or  NH4+  and  X  =  Br-  or  CI",  contains  planar  [SnX4]=.  Brasseur 
and  Rassenfosse,  Nature,  143,  332  (1939)  report  thai  K2[SnCl4]HsO  contains 
condensed  octahedral  [SnClcl=  units. 

Ml.  See  LI.  iX.  K2[Pb(C204)2],  [Pb(SC(CH3)2)2Cl2],  [Pb(OOCC6H4OH)2],  and 
[PbjC6H6C(0— ):CHC(:0)CH3)2]  contain  planar  lead  groupings. 

Family  V: 

Niobium  and  tantalum'  (in  an  indeterminate  oxidation  state)  have  four 
halogen  neighbors  in  a  displaced  planar  relationship  and  with  four  metal 


.    STEREOCHEMISTRY  OF  COORDINATION  NUMBER  FOUR        375 

atom  neighbors  form  a   pyramid.  Antimony!  1 1 1 )  exhibits  a  distorted  tct- 
rahedral  structure  in  one  compound. 

XI.  Vaughan,  Sturdivant,  and  Pauling,  ./.  .1///.  Chem.  Sac.,  72,  5477  (1960).   X. 

XI. .C!:. -711. 0  contains  [Nb6Cluf  h  units.  See  Fig.  7. 
01.  See  XI.  X.  Ta,  Br  ;-7II.<  >  and  Ta  (1     711  0  contain  [TaeXu]44  units. 
Yl.  Bystrom  and  Wilhelmi,  Arkiv  Kemi  3,  373  (1951).  X.  CsSbjF?  contains  pairs 

of  irregular  tetrahedra  of  SbF4~  sharing  a  coiner. 

I  amil>  VI: 

Chromium(VI)  is  tetrahedral;  molybdenum(II),  in  the  halogen  deriva- 
tives, has  four  halogen  neighbors  in  an  approximately  planar  relationship 
and.  with  four  more  molybdenum  atoms,  forms  a  pyramid. 

PI.  Heimlich  and  Foster,  /.  Am.  Chem.  Soc,  72,  4971  (1950).  X.  K[Cr03Cl]  con- 
tains tetrahedral  [Cr03Cl]- 

P2.  Ketelaar  and  Wegeriff,  Rec.  trav.  chim.,  57,  1269  (1938).  X.  K[Cr03F]  contains 
tetrahedral  [Cr03F]-. 

P3.  Ketelaar  and  Wegeriff,  Rec.  trav.  chim.,  58,  948  (1939).  I.  Cs[Cr03F]  contains 
tetrahedral  [Cr03F]-. 

Ql.  Brosset,  Arkiv  Kemi,  1,  353  (1949).  X.  HMo3Cl7-H20  contains  [Mo6Cl8]4+  units. 
See  Fig.  9.6. 

Q2.  Brosset,  Arkiv  Kemi,  Mineral.  Geol,  A20,  No,  7  (1945).  X.  Mo6Cl8(OH)4-14H20 
contains  [Mo6Cl8]4+  units. 

Q3.  Brosset,  Arkiv  Kemi,  Mineral.  Geol.,  A22,  Xo.  11  (1946).  X.  HMo3Cl7H20  con- 
tains [Mo6Cl8]4+  units. 

Family  VII: 

Manganese(II)  may  be  planar,  but  the  evidence  is  incomplete  except 
for  the  phthalocyanine. 

Rl.  Anspach,   Z.   Krist.,   101,    39    (1939).   X.   K2Mn(S04)2-4H20   contains  planar 

Mn:H20)4]++. 

R2.  Cox,  Shorter,  Wardlaw,  and  Way,  ./.  Chem.  Soc.,  1937,  1556.  I.  [Mn(py)2Cl2]  is 
planar.  Mellor  and  Coryell,  ./.  .1///.  Chem.  Soc,  60,  1786  (1938),  have  chal- 
lenged this  on  the  basis  of  the  magnetic  moment . 

R3.  See  D3.  X.  Manganese(II)  phthalocyanine  is  planar. 

Family  VIII: 

Xickel(II),  platinum (II),  and  palladium (II)  arc  planar.  Nickel(II), 
nickel  (0  .  cobalt(Il  .  osmium(VIII),  cobalt  in  [Co(CO)8NO]  and  [Co(C<  I 

•IIi|  and  iron  in  [Fe(CO  \'<)»,|  and  [Fe(CO)2(COH)2]  are  tetrahe- 
dral. The  evidence  thai  cobalt(II)  and  iron(Il  are  planar  in  compounds 
other  than  the  phthalocyanines  is  incomplete. 

si.  Brockway  and  Anderson,  Trans.  Faraday  Soc.,  33,  1233     1937  .  E.    Fi    CO 
NO  rahedral. 

Cambi and Cagnasso, /fend. t«f. lombardo8ci.fVf, 741  (1934  .  M.  Borne  I ■'<•  SCN] 
complexes  with  o  phenanthroline  and  at, a '-dipyridyl  are  planar. 


376  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

S3.  Ewens  and  Lister,  Trans.  Faraday  Soc,  35,  681  (1939).  E.  [Fe(CO)2(COH)2]  is 

tetrahedral. 
SI.  See  D3.  X.  Iron(II)  phthalocyanine  is  planar. 
Tl.  Biltz  and  Fetkenheuer,  Z.  anorg.  allgem.  Chem.,  89,  97  (1914).  G.  [Co(NH3)2X2], 

with  X  =  Cl~,  Br~,  or  I-,  is  planar. 
T2.  See  SI.  E.  [Co(CO)3NO]  is  tetrahedral. 
T3.  Calvin,  Bailes,  and  Wilmarth,  J.  Am.  Chem.  Soc,  68,  2254  (1946).  X  (?).  Com-' 

pounds  of  the  type  [Co(OC6H4CH:NCH2-)2]  "appear  to  be  coplanar." 
T4.  Calvin  and  Melchior,  /.  Am.  Chem.  Soc.,  70,  3270  (1948).  M.  [Co(OHCC6H40)2] 
is  planar  although  two  to  three  unpaired  electrons  are  present.  Some  cobalt 
salicylaldimines  are  planar.  See  T16. 
T5.  See  S2.  M.  Some  Co(CN)2  complexes  with  o-phenanthroline  and  «,a:'-dipyridyl 

are  planar. 
T6.  Cambi  and  Malatesta,  Gazz.  chim.  ital.,  69,547  (1939).  M.  [{C6H50(:NO— )C- 

(:NOH)C6H5}2Co]  has  one  unpaired  electron,  i.e.,  is  planar. 
T7.  Cambi  and  Szego,  Ber.,  64B,  2591  (1931).  M.  Cobalt  acetylacetonate  is  highly 

paramagnetic,  i.e.,  is  tetrahedral. 
T8.  See  R2.  iX,  G.  [Co(py)2Cl2]  is  planar.  Mellor  and  Coryell  (reference  in  R2) 

believe  one  form  is  tetrahedral,  the  other,  condensed  octahedral.  See  T10. 
T9.  See  S3.  E.  [Co(CO)3(COH)]  is  tetrahedral. 
T10.  Hantzsch,  Z.  anorg.  allgem.  Chem.,  159,  273  (1927).  G.  [Co(py)2Cl2]  is  planar. 
Rhode  and  Vogt:  Z.  phys.  Chem.,  B15,  353  (1931),  assign  different  coordination 
numbers  to  cobalt  in  the  two  forms.  See  T8. 
Til.  Jensen,  Z.  anorg.  allgem.  Chem.,  229,  282  (1936).  D.  [Co(PR3)2Cl2],  with  R  =  Et 

or  Pr,  is  either  cis  planar  or  tetrahedral. 
T12.  KrishnanandMookherji,  Phys.  Rev.,  [2]  51,  528  (1937).  M.  The  magnetic  moment 
for  Cs2[CoCl4]  corresponds  to  a  spin  only  value  for  cobalt (II).  A  tetrahe- 
dral structure  is  inferred. 
T13.  M.  Same  as  T12,  p.  774,  but  for  Cs3CoCl5  . 
T14.  See  D3.  X.  Cobalt(II)  phthalocyanine  is  planar. 

T15.  Mellor  and  Craig,  /.  Proc.  Roy.  Soc.,  N.S.  Wales,  74,  495  (1941).  M.  The  mag- 
netic moments  for  a  large  number  of  cobalt  compounds  correspond  to  either 
one  or  else  several  unpaired  electrons.  This  indicates  members  of  the  first 
group  are  probably  planar,  those  of  the  second,  tetrahedral.  No  geometrical 
isomers  could  be  found. 
T16.  Mellor  and  Goldacre,  J.  Proc.  Roy.  Soc,  N.  S.  Wales,  73,  233  (1940).  M.  Some 

cobalt (II)  compounds  are  tetrahedral. 
T17.   Powell  and  Wells,  /.  Chem.  Soc,  1935,  359.  X.  Cs3CoCl5  contains  tetrahedral 

[CoCl4]-. 
T18.  Ray  and  Ghosh,  J.  Indian  Chem.  Soc,  20,  323  (1943).  M.  Some  cobalt  (II)  com- 
pounds are  planar. 
T19.  Tyson  and  Adams,  /.  Am.  Chem.  Soc,  62,  1228  (1940).  M.  [Co(OC6H4CHO)2]  is 

tetrahedral.  See  T3. 
T20.   Varadi,  Acta  Univ.  Szeged,  Chem.  et  Phys.,  2,  175  (1949);  cf,  Chem.  Abstracts, 

44,  5661  i  (1950).  M.  [CoCl4)=  is  tetrahedral  in  solution. 

T21.   Varadi,  Acta  Univ.  Szeged,  Chem.  et  Phys.,  3,  62  (1950);  cf,  Chem.  Abstracts,  46, 

372a  (1952).  Photometer  data.  [CoCl4]=  is  tetrahedral. 
T22.  Zhdanov  and  Zvonkova,  Zhur.  Viz.  Khim.,  24,  1339  (1950);  cf,  Chem.  Abstracts, 

45,  6001e  (1951).  X.  M2[Co(NCS)4]-nH20,  in  which  M+  =  K+  or  NH4+,  con- 
tains tetrahedral  [Co(NCS)4]"  units. 


STEREOCHEMISTRY  OF  COORDINATION  NUMBER  FOl  R         .*>77 

Ul.  Baaolo  BJidMAtoush,  J.  Am.  Chem.  Soc. ,75,. 5663    1963  M.  Bifl  formylcamphor)- 
ethylenediimine-nickel(II)  although  planar  in  the  solid  is  betrahedral  in 

benzene,  toluene,  o  ,  p-,  and  ///-xylene,  and  mesitylene. 

U2.  Brasseur,  Elassenfosse  and  Pi6rard,  Compt.  rend.,  198,  1048    1934);  Brasseur  and 

R&aaenfoBse,  Bull.  80c.  franc,  mineral.,  01,  129  (1938).  \.  Ba[Ni  c\   ;;ill<> 

contains  planar  [\nCX)4]~. 

I'.'i.  Brockway  and  Cross,./.  Chem.  Phye.,  S,  828  (1935).  E.  [Ni(CO)4]  is  tetrahedral. 

I'L  Callis,  Nielsen,  and  Bailar,  ./.  .1///.  Chem.  Soc.,  74,  3461  (1952).  M.  One  nickel 

(II) -containing  dye  is  planar  and  three  are  betrahedral. 
I  •">.  See  T4.  M.  [X'uOC6H4CHO)2]  is  planar  although  paramagnetic.  Some  nickel 

saHcylaldimines  are  planar. 
I'ti.  See  A.6.  M.  Some  compounds  of  the  type  [(RjNCS    iNi]  are  planar. 
17.  See  T7.  M.   Several   nickel(II)    complexes  are  diamagnetic    (planar);   nickel 

acetylacetonate  i^  paramagnetic  (tetrahedral?). 
D8.  Cavell  and  Sugden,  ./.  Chem.  Soc,  1935,  621.  G,  M.  Several  substituted  nickel 

glyoximes  are  planar.  M.  [(R-iXCS^Xi],  with  R  =  Pr  or  Bu,  is  planar. 
D9.  Chugaev,  J.  Ruse.  Phys.  Chem.  Soc.,  42,  1466  (1910);  cf,  Chem.  Abstracts,  6,  594 
(1912).  G.  Nickel  methylglj'oxime  is  planar. 
U10.  Cox,    Pinkard,    Wardlaw,    and    Webster,    J.    Chem.    Soc,    1935,    459.    iX. 

[Xi(HOX:CHC6H40)2]  is  planar. 
I'll.   Cox,  Wardlaw,  and  Webster,  /.  Chem.  Soc,  1935,  1475.  X. 

-       /S_C=C/ 
Ka     Ni  | 

\S— 0=0, 
contains  a  planar  NiS4  unit. 
U12.  Crawford  and  Cross,  /.  Chem.  Phys.,  6,  525  (1938).  IR.  The  infrared  spectrum 
of  [Xi(CO)4]  is  compatible  with  either  tetrahedral  or  square  planar  configura- 
tion. 
U13.  Crawford  and  Horwitz, ./.  Chem.  Phys.,  16,  147  (1948).  R.  The  Raman  spectrum 

of  [Xi(CO)4]  is  compatible  with  a  tetrahedral  structure. 
014.  Curtiss,  Lyle,  andLingafelter,  Acta  Cryst.  5,  388,  (1952).  I,  iX.  [Ni(OC6H4CHO)2] 
is  tetrahedral  because  its  powder  pattern  very  closely  resembles  that  of  the 
corresponding  Zn  compound  but  not  that  of  the  Cu  compound. 
U15.  Dwyer  and  Mellor,  ./.  Am.  Chem.  Soc,  63,  81  (1941).  M.  [Xi-.(R2X3)4],  in  which 

R  —  C6H5  or  CH3C6H4 ,  contains  a  planar  X'iX4  unit. 
U16.  French  and  Corbett,  /.  Am.  Chem.  Soc,  62,  3219  (1940).  M,  CE.  Nickel  formyl 

camphor.  Xi  CuHiiOs)a'2HjO,  contains  a  tetrahedral  Xi04  unit. 
I'!7.   French,  Magee,  andSheffield,  ./.  Am.  Chem.  Soc,  64,  1924  (1942).  M,  CE.  Some 
substituted  salicylaldehyde  nickel (II)  derivatives  are  tetrahedral,  and  some 
aldimine  nickel (II)  derivatives  are  planar.  A  camphor  aldime  nickel (II) 
derivative  is  planar  in  the  solid  state,  distorted  in  an  alcohol  solution. 
U18.  Godycki  and  Rundle,  Acta  Cryst.  6,  487  (1953).  X.  Xickel  dimethylglyoximeis 

planar.  iX.  Xickel  c\  (dohexanedionedioxime  is  planar. 
D19.  Jensen,  Z.  anorg.  allgem.  Chem.,  221,  11   (1934).  G.  [(NH2CSNHNH1    Xi]S04 

contains  planar  nickel (II). 
U20.  Jensen,  Z.  anorg.  nil,,,  m.  Chem.,  229,  266  (1936).  13.  [XiX,(R3P)2],  with  X  =  CI", 
Br~,   or   I-  and  R    =    lit,   Pr,  or   Bu,   and    [NiI«(Et»As)a]  are  trans  planar. 
[Ni(N0      1.'  1'  .    is  cia  planar. 
U21.  Kleinm    and    Eladdatz,    Z.   anoTQ.    allgem.    Chem.,   250,   207    (1942).   M,   G. 
[XidIX  :( !H( ' |B  ,» »  1]  is  planar.  M.  Some  other  nickel  aldimines  arc  planar. 


378  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

U22.  Ladell,  Post,  and  Fankuchen,  Acta  Cryst.  5,  795  (1952).  X.  At  -55°  Ni(CO)4 

is  tetrahedral. 
U23.   Lifschitz,  Bos,   and  Dijkema,  Z.  anorg.  allgem.   Chem.,  242,  97   (1939).  M. 

|XilC6H5CH(NH2)CH(NH2)C6H5l2]X2    and    [Ni(C6H5CHNH2CH2NH2)2]X2 

contain  planar  nickel  in  some  eases  and  possibly  tetrahedral  nickel  in  others 
(deductions  In   Pauling,  ref.  8). 

IJI  Mellor  and  Craig,  J.  Proc.  Roy.  Soc,  N.  S.  Wales,  74,  475  (1941).  AI.  Planar  or 
tetrahedral  configurations  are  assigned  to  many  nickel  compounds,  all  che- 
lates. An  attempt  is  made  to  relate  configuration  to  kind  of  atoms  directly 
bonded  to  nickel. 

U25.  Mellor  and  Lockwood,  J.  Proc.  Roy.  Soc,  N.  S.  Wales,  74,  141  (1940).  M.  A  sub- 
stituted nickel  pyrromethene  is  tetrahedral. 

U26.  See  A15.  C.  K2[Ni(CN)4]-H20  contains  planar  LNi(CN)4]=. 

U27.  Milone  and  Tappi,  Atti  accad.  sci.  Torino,  Classe  sci.fis.,  mat.  nat.,  75,  445  (1940). 
X.  Nickel  dimethylglyoxime  and  nickel  methylethylglyoxime  are  planar. 

U28.  Peyronel,  Z.  Krist.,  103,  157  (1941).  X.  [(Pr2NCS2)2Ni]  contains  a  planar  NiS4 
grouping. 

U29.  Rayner  and  Powell,  J.  Chem.  Soc.  1952,  319.  X.  One  half  the  Ni  atoms  in  Ni- 
(CN)2(NH3)  C6H6  are  tetrahedrally  surrounded  by  four  C  atoms. 

U30.  Reihlen  and  Htihn,  Ann.,  499, 144  (1932).  iO.  [(CH3C9H5NCH2NH2)2Ni]  contains 
nonplanar  nickel. 

U31.  See  A  23.  X.  Nickel(II)  phthaiocyanine  contains  a  planar  NiN4  group. 

U32.  Robertson  and  Woodward,  J.  Chem.  Soc,  1937,  219.  X.  Nickel(II)  phthaio- 
cyanine is  planar. 

U33.  Speakman,  Acta  Cryst.  6,  784  (1953).  X.  NiC26Hi4N8  contains  a  planar  NiX4 
unit. 

U34.  Sugden,  J.  Chem.  Soc,  1932,  246.  G,  M.  Nickel  methylbenzylglyoxime  is  planar. 

U35.  See  T19.  M.    [Ni(OC6H4CHO)2]  is  tetrahedral  and   [Ni(OC6H4CH:NH)2]  is 
planar. 
VI.  Brasseur  and  Rassenfosse,  Mem.  acad.  roy.  Belg.,  Classe  sci.,  16,  No.  7  (1937). 

I.  Several  complex  cyanides  contain  planar  [Pd(CN)4]=. 
V2.  Brasseur,  Rassenfosse,  and  Pierard,  Z.  Krist.,  88,  210  (1934).  I.  Ba[Pd(CN)4]- 

4H20  contains  planar  [Pd(CN)4]=. 
V3.  Cahours  and  Gal,  Compt.  rend.,  71,  208  (1870).  G.  [(Et3As)2PdCl2]  exists  in  two 

forms. 
V4.  Chatt,  Mann,   and  Wells,   /.    Chem.   Soc,   1938,   2086.   iX.    [Bu3PClPdC204- 

PdClPBu3]  contains  bridged  planar  palladium  units. 
V5.  See  U  10.  iX,  I.  [Pd(OC6H4CH:NOH)2]  is  planar. 
V6.  Cox  and  Preston,  /.  Chem.  Soc,  1933,  1089.  iX.  [Pd(en)2]Cl2,  [Pd(NH3)4]Cl2  , 

and  (NH4)2[PdCl4]  contain  planar  groupings. 
V7.  Cox,  Saenger,  and  Wardlaw,  J.  Chem.  Soc,  1934,  182.  I.  [(Me2S)2PdCl2]  contains 
a  planar  unit. 

r    /s-c=o\ 

V8.  See  U  11.  X.  K2    Pdl  contains  a  planar  PdS4  unit. 

L    \s-c=o/ 

V9.    Dickinson,  Z.  Krist.,  88,  281    (1934).  X.   [Pd(NH3)4]Cl2-H20  contains  planar 

[Pd(NH3)4]++. 
\  10.    Dickinson,  ./.  ,1///.  Chem.  Soc,  44,  2404  (1922).  X.  K2[PdCl4]  and  (NH4)2[PdCl4] 

contain  planar  [PdCl4]". 
Vll.   Dwyerand  Mellor,./.  Am.  Chem.  Soc,  56,  1551  (1934).  G.  Bis(antibenzylmethyl- 

glyoxime)palladium(II)  is  planar. 


STEREOCHEMISTRY  OF  COORDINATION  NUMBER  FOUR         379 

V12.  See  U  18.  iX.  Palladium  dimethylglyoxime  La  planar. 

V13.  Grinberg  and  Shul'man,  Compt.  rend.  acad.  eci.  (U.R.8.S.)  [N.  S.],  1933,  215. 

G.  [Pd  NH,)»X8]  and  [Pd(pj   A],  X  =  Cl-orBr-,  are  planar. 
V14.  Janes,  ./.  .1//;.  (In m.  Soc.,  57,  171  (1935).  M.  Several  palladium  complexes  are 

diamagnetic. 
V15.  Jensen,  Z.  anorg.  allgem.  Chem.,  226, 97    1935  .  1).  [PdCh(SEtj) a]  is  trans  planar. 
V16.  Jensen,  Z.  //m»/-«/.  allgem.  Chem.,  229,  225  (1936).  D.  [1MC1    Ki  Bb)2]  is  trans 

planar. 
Yl 7.  Erauss  and  Brodkorb,  Z.  anorg.  allgem.  Chan.,  165,  73  (1927).  G.  [Pd(py)2Cl2] 

and    [(EtNHj)jPdCli]   are  planar.   Drew,    Pinkard,   Preston,   and    Wardlaw, 

./.  ('hem.  Soc,  1932,  1895,  believe  the  isomerism  is  not  geometric  but  is  poly- 

merism.  i.e.,  [Pd(py),Cl,]  and  [Pd(py)4][PdCl4]. 
V18.    l.idstone   and  Mills,  /.   Chem.  Soc,  1939,   1754.  O.    [{XH2C(CH3)2CH2NH2( - 

Pd(XH2CHC6H5CHC6H5XH2)]-H-  is  planar. 
V19.  Mann.    Crowfoot,    Gattiker,    and   Wooster,   /.    Chem.   Soc.,   1935,    1642.    iX. 

[(NH,),PdC204]  is  planar.  iX,  G.  [(XH3)2Pd(X02)2]  is  planar. 
V20.   Mann  and  Purdie,  /.  Chem.  Soc.,  1935,  1549.  C.  [PdX2Cl>],  in  which  X  =  Et2S, 

Et3P,  or  Et3As,  is  planar. 
V21.  Mann  and  Wells,  J.  Chem.  Soc,  1938,  702.  X.  [Me3AsPdBr2]2  contains  bridged 

planar  units. 
V22.   See  U25.  M.  A  substituted  pyrromethene  of  palladium(II)  is  diamagnetic  but 

cannot  be  planar. 
V23.   See  A  15.  C.  [Pd(XH3)4]Cl2-H20  and  K2[PdCl4]  contain  planar  palladium  com- 
plexes. 
V24.  Pinkard,  Sharratt,  Wardlaw,  and  Cox,  /.  Chem.  Soc,  1934,  1012.  G.  Palla- 

dium(II)  glycinate  is  planar. 
V25.   Poral-Koshits,  Doklady  Akad.  Nauk  S.S.S.R.,  58,  603  (1947);  cf,  Chem.  Ab- 
stracts, 46,  4313d  (1952).  X.  K2[Pd(X02)4]  structure  determined.  Abstract  does 

not  give  full  details. 
V26.  Reihlen  and  Hiihn,  Ann.,  489,  42  (1931).  iO.   [{NH2C(CH3)2CH2NH2}2Pd]++  is 

not  planar. 
-"     See  U  30.  iO.  [(CH3C9H5XCH2XH2)2Pd]-H-  is  not  planar. 
V28.  Rosenheim    and    Gerb,    Z.    anorg.    allgem.    Chem.,    210,    289     (1933).    iO. 

[Pd(OC6H4COO)2]    is  not  planar. 
V29.  Theilacker,  Z.  anorg.  allgem.  Chem.,  234,  161   (1937).  X.  K2[PdCl4]  contains 

planar  [PdCl4]=. 
V30.  Wells,  Proc.  Roy.  Soc.  (London),  A167,  169  (1938).  X.  [(CH3)3AsPdCl2]2  contains 

bridged  planar  groupings. 
WI.  Jaeger  and  Zanstra,  Rec  trav.  chim.,  51,  1013  (1932),  also  appeared  in  Proc 

Koninkl.    Nederland.    Akad.    Wetenschap.,    35,    610,    779,    787,    (1932).    X. 

M[Os03X],  in  which  M+  =  K+,  XH4+,  Rb*,  Tl+,  or  Cs+,  contains  tetrahedral 

[Os03X]-. 
XI.  Angell,  Drew,  and  Wardlaw,  ./.  Chem.  Soc,  1930,  349.  G.  The  two  forms  of 

[(Et2S)2PtCl2]  are  structural,  not  cis-trans,  isomers  (the  formulation  pro- 
posed is  much  less  likely  than  cis-trans  isomerism  when  considered  from  the 

standpoint  of  modern  concepts). 
X2.  Bokil,  Valnshteln,  and  Babareko,  Izvest.  Akad.  Nauk  S.S.S.R.,  Otdel.  Khim. 
•'•  1951,  0tl7;  cf,   Chem.   Abstracts  46,  5927d    (1952).  Electronographic 

KPtCl«NH|  and  KPi  Br  XII,  contain  planar  Pi , 
X3.  Bozorth  and  Pauling,  Phys.  Rev.,  [2]  39,  537  (1932).  X.  The  data  of  Bozorth  and 


380  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Haworth  (Phys.  Rev.,  [2]  29,  223  (1927))  show  that  Mg[Pt(CN)4]-7H20  con- 
tains planar  [Pt(CN)4]=. 

X  }.  See  V  1.  I.  Several  complex  platinum(II)  cyanides  contain  planar  [Pt(CN)4]=. 

X5.  See  V  2.  I.  Ba[Pt(CN)4]-4H20  contains  planar  [Pt(CN)4]". 

X6.  Brosset,  Arkiv  Kemi,  Mineral.  Geol  ,  A25,  Xo.  19  (1948).  X.  [Pt(NH3)2Br2][Pt- 
(XH3)2Br4]  contains  planar  [Pt(NH3)2Br2]. 

X7.  Cahours   and  Gal,   Compt.   rend.,   70,  897    (1870).   G.   There   are   two  forms 
of  [(Et3P)2PtCl2]. 

X8.  See  V  3.  G.  [(Me3P)2PtCl2]  and  [(Et3As)2PtCl2]  both  exist  in  two  forms. 

X9.  Chernyaev,  Ann.  inst.  platine  (U.S.S.R.)  4,  243  (1926);  cf,  Chem.  Abstracts,  21, 
2620    (1927).    G.    [Pt(NH2OH)(NH3)(py)(N02)]2[PtCl4]    contains    a    planar 
cation  (three  isomers  found).  Several  other  compounds  contain  planar  plati- 
num (II)  because  they  exist  as  cis-trans  isomers. 
X10.  Cox,    J.    Chem.    Soc.,    1932,    1912.    X.    [Pt(NH3)4]Cl2-H20    contains    planar 

[Pt(NH3)4]++. 
Xll.  See  U  10.  G,  iX.  [Pt(OC6H4CH:NOH)2]  is  planar. 
X12.  See  V  6.  iX.  [Pt(en)2]Cl2  is  planar. 


X13.  See  V  7.  iX,  G.  [{ (CH3)2S|2PtCl 

/s-c=o 

X14.  SeeU  11.  X.  Ks 
X15.  See  V  10.  X.  K 


is  planar. 

contains  a  planar  PtS4  unit. 


Pt[ 

^s— c=o, 

PtCl4]  contains  planar  [PtCl4]=. 
X16.  Drew  and  Head,  J.  Chem.  Soc.,  1934,   221.   G.  [Pt{NH2C(CH3)2CH2NH2}2]Cl 

and     [Pt(NH3)(EtNH2){NH2C(CH3)2CH2NH2!]Cl2    contain     planar     plati- 

num(II). 
X17.  Drew,  Head,  and  Tress,  J.  Chem.  Soc,  1937,  1549.  Attempted  O.  [Pt{NH2C- 

(CH^CH^Ho}^       and       [PtJNH2C(CH3)2CH2NH2}{NH2CH2CH(CH3)- 

CH2NH2}]+4  could  not  be  resolved. 
X18.  Drew,  Pinkard,  Wardlaw,  and  Cox,  J.  Chem.  Soc,  1932,  988,  1004.  G.  A  third 

isomer  reported  for  [Pt(NH3)2Cl2].  Structural  isomerism  proposed.  The  third 

isomer  proved  to  be  a  mixture  of  the  first  two.  See  V  17. 
X19.  Drew  and  Wyatt,  J.  Chem.  Soc,  1934,  56.  G.  [PtCl2(Et2S)2]  is  planar. 
X20.  Grinberg,  Helv,  Chim.  Acta,  14,  455  (1931).  G.  [Pt(NH3)2Cl2]  reactions  related  to 

planar  structure. 
X21.  Grinberg,  Z.  anorg.  allgem.  Chem.,  157,  299  (1926) ;  Ann.  inst.  platine  (U.R.S.S.), 

5,  365  (1927).  G.  [Pt(NH3)2(SCN)2]  is  planar. 
X22.  Grinberg  and  Ptitzuin,  J.  prakt.  Chem.,  [2]  136,  143  (1933);  Ann.  inst.  platine 

(U.R.S.S.),  9,  55  (1932).  G.  [Pt(NH2CH2COO)2]  is  planar. 
X23.  Grinberg  and  Razumova,  Zhur.  Priklad.  Khim.  27,  105  (1954);  cf.  Chem.  Ab- 
stracts 48,  6308a  (1954).  The  reaction  of  [Pt{  (C6H5)3P}2C12]  with  ethylene- 
diamine  shows  it  to  be  the  cis  isomer. 
X24.  Hantzsch,  Ber.,  59,  2761  (1926).  G.  [Pt(py)2Cl2]  is  planar. 
X25.  Hel'man,  Karandashova,  and  Essen,  Doklady  Akad.  Nauk  S.S.S.R.,  63,  37 

(1948);   cf,    Chem.    Abstracts,   43,    1678i    (1949).    G.    [Pt(py)(NH3)ClBr]    is 

planar  (three  isomers). 
X26.  See  V15.  D.  [PtX2(R2S)2],  in  which  X  =  CI",  Bi-,  I",  or  N02~  and  R  =  Et, 

Pr,  i-Pr,  Bu,  s-Bu,  i-Bu,  or  C6H5  ,  is  planar. 
X27.  See  V  16.  D.  [PtX2(R3E)2],  in  which  X  =  Cl~,  Br",  I~,  NOr,  or  NO,"",  R  =  Et, 

Pr,  Bu,  or  C6H5  ,  and  E  =  P,  As,  or  Sb,  is  planar. 
X28.  Klason,  Ber.,  28,  1493  (1895).  G.  [PtCl2{  (CH3)2S}2]  is  planar. 


STEREOCHEMISTRY  OF  COORDINATION  NUMBER  FOUR         381 

X29.  Kuraakov,  ./.  Ritas.  Phya.  CKem.  Sac,  25,  565  (1803);  cf,  Chem.  Centr.,  65,  I, 

460  (1894),  G.  Thiourea  reacts  with  cis  |IVM1,  ,Cltjor   [Pt(py)8Cl«]  to  yield 
I'i  til) 4] Ch  and  with  the  trans  compounds  to  yield  [Pt(tu)2Cl2]. 
X30.  Lambot,  Rail.  soc.  roy.  set.  Litge,  12,  541  (1943);  cf,  Chem.  Abstracts,  40,  5656" 

(1946).  X.  Ki[Pl   \< »    ;!  contains  a  planar  PtN<  unit. 
X:>1.  Lifschitz    and    Froentjes,    Z.    anorg.    allgem.    Chem.,    233,     1     (l!)37j.    (1. 

[PtXi  (11  CHSEtCOOH),],  in  whirh  X  =  C1-,  Br~,  etc.,  is  planar. 
X32.  NfatMeu, /.cairn,  pays.,  36, 308  (1939).  R.[Pt(N^^ 

[Pt(en)t]Cli  ,  and  [Pt(py)«Clt]  contain  planar  or  octahedral  platinum (II)  in 

solution. 
X33.  See  Ai:».  C.  :l't  \"H3)4]C12-H20,  K2[PtCl4],  Ba[Pt(CX)4]-4H2(>,  Mg[Pt(CN)«]- 

7HiO,  and  LiK[Pt(CN)4]-3HjO  contain  planar  platinum(II). 
X34.  Mills  ami  Quibell,  J.   Chem.  Soc,  1935,  839.  O.   [Pt{NH2CH2C(CH3)  >XH2)- 

jXH.CHCeHsCHCeHsXHoj]-^  is  planar. 
X35.  Monfort,  Rull.  soc.  roy.  sci.  Liege,  11,  567  (1942);  cf,  Chem.  Abstracts,  38,  41743 

(1944).  X.  KXa[Pt(CX)4]-3H20  contains  planar  [Pt(CN)4]-. 
Petren,  Z.  anorg.  allgem.  Chem.,  20,  62  (1899).  G.  Two  forms  of  [Pt(SEt2)2Cl2] 

arc  reported. 
See  V24.  G.  [Pt(XH2CH2COO)2]  is  planar. 
X38.  Ramberg,  Ber.,  43,  580  (1910);  46,  3886  (1913).  G.  [Pt(OOCCH2SEt)2]  is  planar. 
X39.  Sec  V26.  iO.    [PtjXH2C(CH3)2CH2NH2}2]++  and   [Pt{CH3C9H5XCH2NH2J2]++ 

are  not  planar. 
X40.  Reihlen  and  Hiihn,  Ann.,  519,  80  (1935).  iO.  [Pt(NH2CH2CHC6H5XH2){CH3- 

(    II.-OCgH^XCHoXH-.j]^  is  not  planar  or  tetrahedral. 
X41.  Reihlen  and  Xestle,  Ann.,  447,  211  (1926).  G.  "Trans"  [Pt(XH3)2Cl2]  is  a  dimer 

in  liquid  ammonia  and  the  planar  nature  of  platinum (II)  is  therefore  suspect. 
X  12.  Reihlen,  Seipel,  and  Weinbrenner,  Ann.,  520,  256  (1935).  iO.  [Pt(dipy){NH2CH- 

(C6H5)CH2XH2j]++  is  not  planar. 
X43.  See  A23.  Platinum(II)  phthalocyanine  contains  a  planar  PtX^4  grouping. 
X44.  Robertson  and  Woodward,  J.  Chem.  Soc,  1940,  36.  X.  Platinum(II)  phthalo- 
cyanine is  planar. 
X45.  See"V28.  iO.  [Pt{  (XH^oCe^CH,),]^  is  not  planar. 
X46.  Rosenheim  and  Handler,  Ber.,  59,  1387  (1926).  Attempted  O.  [Pt{  (NH2)2C6H3- 

CH3)2]++  could  not  be  resolved. 
X47.  Roy,  Indian  J.  Phys.,  13,  13  (1939).  R.  The  Raman  spectrum  of  [Pt(en)2]Cl2  is 

compatible  with  square  planar  [Pt(en)2]++. 
X48.  Ryabchikov,  Compt.  rend.  acad.  sci.  U.R.S.S.,  27,  349  (1940).  G.  K2[Pt(S203)2] 

contains  a  planar  Pt02S2  grouping. 
X49.  See  V29.  X.  K2[PtCl4]  contains  planar  [PtCl4]=. 
X50.  Werner,  Z.  anorg.  allgem.  Chem.,  3,  267  (1893).  G.  [Pt(XH3)2Cl2]  and  [Pt(py)2Cl2] 

are  planar. 
X51.  Wunderlich  and  Mellor,  Acta  Cryst.  7,  130  (1954).  iX.  In  K[PtCl3C2H2]H20  the 

Pt  and  3  CI  atoms  are  coplanar.  The  fourth  planar  position  is  occupied  by  the 

C2H2  double  bond. 


IU.   Stereochemistry  and  Occurrence  of 

Compounds  Involving  the  Less  Common 

Coordination  Numbers 

Thomas  D.  O'Brien* 

University  of  Minnesota,  Minneapolis,  Minnesota 

The  term  "coordination  number"  in  the  chemical  sense  refers  to  the 
number  of  groups  attached  to  a  central  atom  and  may  depend  upon  the 
nature  of  the  central  atom,  the  valence  of  the  central  atom,  the  nature  of 
the  coordinating  group  and  the  nature  of  the  anion.  "Coordination  num- 
ber" in  a  crystallographic  sense,  however,  is  quite  different.  It  refers  to 
the  number  of  nearest  neighbors  of  an  atom  in  the  crystal,  and  is  dependent 
only  on  the  radius  ratio.  In  many  cases  the  two  coordination  numbers  are 
identical,  so  there  is  no  ambiguity,  but  this  cannot  always  be  assumed. 

Coordination  Number  Two 

Only  those  elements  in  Group  I  of  the  Periodic  Table,  including  hydrogen, 
seem  to  have  a  consistent  tendency  to  exhibit  a  coordination  number  of 
two.  In  a  few  cases,  elements  in  other  periodic  groups,  which  can  exist  with 
a  valence  of  one,  may  also  be  two-coordinate.  There  are  only  two  possible 
geometrical   configurations,   linear,   O — M — O,   and   angular,   O — M 

\ 

o, 

and  no  cases  of  stereoisomerism  are  known. 

It  has  been  shown1  that  in  the  compounds  KHF2  and  NH4HF2  the  two 
fluorine  atoms  are  linked  linearly  through  the  hydrogen,  (F — H — F)~, 
giving  hydrogen  a  coordination  number  of  two.  There  are  many  similar 
examples  in  compounds  exhibiting  hydrogen  bonding,  of  which  dimeric 
acetic  acid, 

*  Now  at  Kansas  State  College,  Manhattan,  Kansas. 

I.  Belmholz  and  Rogers,  /.  Am.  Ch  em.  Soc,  61,  2590  (1939);  ibid.,  62,  1533  (1940). 

382 


COMPOUNDS  INVOLVING  LESS  COMMON  COORDINATION  NUMBERS    383 

O— H O 

/  \ 

t  II— C  CH,f 

\  / 

0-    H— O 

is  typical.  The  bonding  in  these  cases  is  doubtless  due  to  dipole  attractions, 
and  is  not  truely  covalent. 

The  Group  IB  elements  in  their  univalent  state  all  exhibit  the  coordina- 
tion number  t>\  two.  although  the  copper!  1 1  compounds  are  not  -<>  common 
and  are  often  less  stable  than  those  of  silver!  I  i  and  gold(I).  Rosenheim 
and  Loewenstamm*  reported  the  preparation  of  bis(thiourea  copper(I) 
chloride.  [Cu{SC  Ml.  »}jCl,  in  which  they  believi  the  thiourea  is  coordi- 
nated to  the  copper  atom  through  the  sulfur*.  Spacu  and  Murgulescu4 
report  a  number  of  compounds  in  which  anionic  copper(I)  has  a  coordination 
number  of  two.  assuming  thiosulfate  ion  is  a  bidentate  group,  as  in 
Na[CuSsOs].  This  aecessitates  an  improbably  small  angle  for  the  covalences 
of  the  copper. 

Silver(I)  forms  the  well-known,  linear  diamminesilver(I)5,  [Ag(XH3)2]+, 
and  dicyanosih  er(I)6,  [Ag(CX)2]~~,  ions.  Fyfe7  prepared  silver(I)  diammines 
with  acridine,  quinoline,  isoquinoline,  and  pyridine  and  found  that  the 
order  of  stability  of  the  complexes,  acridine  >  quinoline  >  isoquinoline  = 
pyridine,  is  the  same  as  the  order  of  the  electron  densities  on  the  nitrogen 
atoms  in  the  amines.  It  has  also  been  shown  that  silver(I)  forms  only  mono 
and  bis  benzoate  complexes  in  solution8.  With  ethylenethiourea 


XHCH: 

/ 

S=C 

\ 

XHCH2/  2J 


X 


is  tunned,  where  M  is  silver(I)  or  gold(I).  The  silver  salt  in  which  X  is  a 
halide  is  unaffected  by  light. 

A  dimethyldithioethylene  go)d(I)  complex  salt, 

2.  Rosenheim  and  Lowenstamm,  Z.  anorg.  Chem.,  34,  62  (1903). 

l:    tl         '         17.  297    L884  . 
\    Spacu  and  Murgulescu,  hull.  Sue.  stiinte  cluj.,  5,  344    L934 
ind  Wyckoff,  '/.    KrUt.t  87,  264    I  I 

6    li  /    8  84.  _  , 

7.   I  ■     169.  I 

v  3        •..  3,  L318    194 


384 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


('II:: 


Aii 


CH2 


S— CH2 


CH3 


CI, 


is  also  known.  Two  coordinate  complexes  of  gold(I)  have  been  prepared 
with  tertiary  arsines9.  The  compounds  are  characterized  by  their  solubility 
in  nonpolar  solvents,  insolubility  in  water  and  sharp  melting  points. 

Many  alkali  metal  salts  of  metal  amides  have  been  reported  by  Franklin10. 
Among  them  are  compounds  of  the  type  K[M(NH2)2],  where  M  is  silver  (I) 
or  thallium  (I). 

The   rather   curious  halogen   compounds   [Br(py)2]C104 ,    [I(py)2]N03 , 

py 


and 


I 


NO: 


have  been  prepared11.  On  the  basis  of  solubilities,  Yatsi- 


mirskii12  has  formulated  a  series  of  complexes  wThich  contain  anionic  central 
atoms  and  cationic  ligands.  These  formulations  are  exemplified  by  the 
species,  [Ag2Cl]N03 ,  [Ag2Br]N03 ,  and  [Ag2I]N03 .  The  stability  increases 
in  the  order,  chloride  <  bromide  <  iodide.  The  conditions  favorable  to 
the  formation  of  such  complexes  are  low  electron  affinity  of  the  anion,  high 
electron  affinity  of  the  cation,  and  large  radius  of  the  cation. 


Coordination  Number  Three 

On  the  basis  of  theoretical  considerations,  Kimball13  offers  the  trigonal 
plane  (I),  unsymmetrical  plane  (II),  and  trigonal  pyramid  (III)  as  possible 
structures  for  three  coordinate  complexes  (Fig.  10.1).  The  unsymmetrical 
plane  would  give  rise  to  geometric  isomerism,  and  the  trigonal  pyramid  would 
show  optical  isomerism  in  complexes  of  the  type  [MXYZ).  The  other  struc- 
ture, being  completely  symmetrical,  would  give  no  stereoisomerism.  Mann14 


9. 

K). 


12. 
13. 
14. 


Dwyer  and  Stewart,  J.  Proc.  Roy.  Soc,  N.  S.  Wales,  83,  177  (1949). 

Franklin,  "Nitrogen  System  of  Compounds,"  New  York,  Reinhold  Publishing 

Corp.,  1935. 
Carlsohn,  "Uber  eine  Neue  Klasse  von  Verbindungendes  positive  einwertigen 

.Jods,"  Leipzig,  1932;  Ber.,  68B,  2209  (1935). 
Ynisin.irskii,  Doklady  Akad.  Nauk  S.S.S.R.,  77,  819  (1951). 
Kimball,  ./.  Chem.  Phys.,  8,  188  (1940). 
Mann,  ./.  Chem.  Soc,  1930,  1745. 


COMPOUNDS  INVOLVING  LESS  COMMON  COOHDIXATIOX  XCMIiKliS    :*N."> 


CI) 


(E) 
O   =   cent  ral  atom 

Fig.  10.1 


(HE) 


proved  that  the  sulfur  atom  in  tel rachlorol ^^'-diaminodiethylsulfide Iplal  - 
inum  (IV)  has  the  trigonal  pyramid  configuration  by  resolving  the  com- 
plex into  its  optical  antipodes.  The  complex  has  the  structure 


/CHZ-CHZ-NH2 


,1    XCH* 

CH? 


Silverl  1 1  and  copper(I),  in  addition  to  being  two-coordinate,  also  form  a 
number  of  compounds  in  which  they  are  apparently  three-coordinate. 
Compounds15  containing  ethylenethiourea,  like  [Ag{SC(XH)2(CH2)2!:i)(,l 
and  [Cu{SCl  XH)2(CH2)2}3]2S04  are  known,  as  are  the  corresponding 
thiourea  salts2.  The  corresponding  nitrates  contain  four  molecules  of  the 
ethylene  thiourea  per  metal  atom,  so  that  it  might  be  suspected  that  the 
anions  in  the  chloride  and  sulfate  are  coordinated. 

The  reddish  chlorocuprates,  the  chlorocadmates,  and  the  chloromercu- 
rates,  [CuCl3]~,  [CdCl3]~,  and  [HgCl3]~,  are  all  well-known,  but  it  has  been 
shown  that  the  metals  in  these  do  not  have  a  coordination  number  of  three 
in  the  solid  state.  The  cadmium  compound  consists  of  chains  of  CdCU 
octahedra  joined  laterally16  as  shown  in  Fig.  10.2.  The  mercury  compound 
is  of  a  different  crystalline  structure17. 

The  red  color  obtained  when  potassium  tetracyanonickelate(II)  is  re- 


Fig.  10.2 


15.  Morgan  and  Burstall,/.  Chem.  So,-.,  1928,  143. 

16.  Braaseui  and  Pauling,  ./.  Am.  ('hem.  8oc.,  60,  2886  (1938). 

17.  Harmsen,  Z.  Krist.,  100,  208  (1939;. 


380  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

duced18  is  believed  to  be  due  to  the  formation  of  potassium  tricyanonick- 
elate(I),  K2[Ni(CN)3].  Dark  red  solutions  of  potassium  tricyanonick- 
elate(I),  when  exposed  to  the  air,  lose  their  color  and  precipitate  part  of 
I  heir  nickel  as  nickel(II)  hydroxide  and  the  remainder  as  potassium  tetra- 
cyanonickelate(II).  From  polarographic  studies,  Caglioti,  Sartori,  and 
Silvestroni19  estimate  the  potential  of  the  couple  [Ni(CN)4]=-[Ni(CN)]3= 
lo  be  —0.(3844  volts.  The  validity  of  the  measurement  is  disputed  by  Kol- 
thoff  and  Hume20,  who  found  that  the  tetracyanonickelate(II)  ion  undergoes 
an  irreversible  two-electron  reduction  at  the  dropping  mercury  electrode. 
They  have  also  shown  that  the  tricyanonickelate(I)  ion  is  subject  to  anodic 
oxidation  but  not  to  further  polarographic  reduction.  Recent  x-ray  studies20a 
indicate  that  the  tricyanonickelate(I)  ion  is  dimeric,  [Ni2(CN)6]4~. 

Other  compounds  in  which  copper  is  reported  to  have  a  coordination 
number  of  three  are  the  blue-black  [CuNOCl2],  [CuNOBr2],  and 
[CuNOS04]21,  the  dark  green  triamminecopper(I)  octacyanomolybdate 
(VI)22,  and  triamminecopper(I)  halides23.  Although  Biltz  and  Stollenwerk23 
write  the  formulas  of  the  halides  as  [Cu(NH3)3]X,  it  is  quite  possible  that 
the  halogen  is  also  coordinated,  giving  the  copper  a  coordination  number  of 
four. 

Franklin  has  reported  amides  of  the  general  formula,  K[M(NH2)3],  in 
which  M  is  lead(II),  beryllium,  calcium,  strontium,  barium,  or  tin(II). 

It  is  believed  that  the  solubility  of  silver  chloride  in  a  concentrated  solu- 
tion of  cesium  chloride  is  due  to  the  formation  of  the  trichloroargentate(I) 
ion,  [AgCl3]=,  in  wrhich  the  silver  is  three-coordinate24.  The  simple  ammino 
compound  [Ag(NH3)3]X  has  also  been  reported25. 

It  is  believed  that  the  iodine  is  the  central  atom  in  a  cationic  complex 

Ag" 


with  three  silver  atoms  attached  as  ligands, 


Ag-I 


(N03)226.  This 


Ag. 
complex  ion  was  shown  to  migrate  to  the  cathode  during  electrolysis. 

Thallium  alcoholates  w'hen  dissolved  in  polar  solvents  are  typically  salt- 
like in  their  behavior.  They  are,  however,  also  soluble  in  nonpolar  sol- 

18.  Belluci  and  Corelli,  Atti.  accad.  Lincei,  22,  II,  579  (1913). 

19.  Caglioti,  Sartori,  and  Silverstroni,  Ricera  Sci.,  17,  624  (1947). 

20.  Kolthoff  and  Hume,  J.  Am.  Chem.  Soc,  72,  4423  (1950). 
20a.  Mast  and  Pfab,  Nalurwissenschaften,  39,  300  (1952). 

21.  Manchot,  Ann.,  376,  308  (1910);  Gall  and  Mengdahl:  Ber.,  60B,  86  (1927). 

22.  Bucknall  and  Wardlaw,  /.  Chem.  Soc,  1927,  2981. 

23.  Biltz  and  Stollenwerk,  Z.  anorg.  Chem.,  119,  97  (1921). 

24.  Wells  and  Wheeler,  Am.  J.  Sci.,  [3]  44,  155  (1892). 

25.  Biltz  and  Stollenwerk,  Z.  anorg.  Chem.,  114,  1176  (1920);  ibid.,  119,  97  (1921). 

26.  Helhvig,  Z.  anorg.  Chem.,  25,  157  '1900). 


COMPOUNDS  INVOLVING  LESS  COMMON  COOL'I)/ \  ATION  NUMBERS    387 


II) 


OCjiis 

CH3 

H-C              0 -Tt             C  — CH3 

III                  1 

1 

HC 

H3C~CYT^°\//C"H 

1 

OC2H5 


(H) 
1  i«;.  10.3 


\ 


Pb-OH 


CH, 


(m) 


vents  such  as  benzene,  and  they  have  been  shown27  to  be  tetrameric  in 
that  solvent,  possibly  with  a  three-coordinate  structure  as  in  (I)  (Fig.  10.3). 
In  similar  solvents,  thallium(I)  ethyl  acetoacetate  is  dimeric  and  three- 
coordinate27  (II). 

Menzies28  lias  reported  a  nonionic  basic  lead  acetonylacetonate  with  the 
formula  shown  in  (III)  (Fig.  10.3).  There  is,  however,  no  evidence  to  indi- 
cate that  the  substance  is  not  dimeric,  the  lead  atoms  being  linked  together 

OH 

\    /      \    / 
through  the  hvdroxvl  groups,       Pb  Pb      ,  giving  the  metal  a  coor- 

/     \       /     \ 
OH 

dinatioD  number  of  four. 

Coordination  Number  Five 

From  theoretical  considerations,  a  coordination  number  of  five  should  be 
the  least  likely  to  exist,  although  there  are  many  examples  in  which  atoms 
are  apparently  five-coordinate.  Kimball13  gives  the  following  as  geometrical 
possibilities: 


TRIGONAL 
B  I  PYRAMID 


TETRAGONAL 
PYRAMID 


PENTAGONAL 
PLANE 


PENTAGONAL 
PYRAMID 


Fig.  10.4.  Some  possible  configurations  for  coordination  number  five 


Duli't'Y  has  <\t<-ii<l<'(l  the  study  of  the  bipyramidal  structure,  calculating 
the  extent  to  which  d  electrons  are  involved  in  the  hybridization29. 
On  the  basis  of  electron  diffraction  studies,  iodine(V)  fluoride  was  first 

27.  Sidguick  and  Sutton,  /.  Chem.  Soc,  1930,  1461. 

28.  Menzies,./.  «.,  1934,  1756. 

20.  Duffey,  ./.  Chi  m.  Phys.t  17,  106  H049) ;  Proc.  S.  Dakota  Acml .  Sri.,  28,  07  (1949). 


388  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

reported  to  have  the  trigonal  bipyramidal  structure30,  but  subsequent  x-ray 
examination  showed  that  the  I-F  distance  was  much  less  than  would  be 
expected31.  As  a  result  of  studies  on  the  infrared  and  Raman  spectra  it  has 
been  postulated32  that  the  molecule  has  the  tetragonal  pyramidal  structure, 
with  an  unshared  pair  of  electrons  occupying  a  position  equivalent  to  the 
unique  position  of  the  fifth  fluorine  atom,  but  below  the  base  of  the  pyra- 
mid on  the  perpendicular  to  the  plane,  (Fig.  10.5).  De  Heer33  states  that 
the  structure  is  still  uncertain  but  that  dipole  moment  studies  could  provide 
final  proof  of  the  structure.  From  the  Raman  spectrum34,  it  is  believed  that 
bromine (V)  fluoride  also  has  the  tetragonal  pyramidal  configuration. 

For  many  years  the  structures  of  the  pentahalides  of  phosphorus,  arse- 
nic, and  antimony  were  debated,  but  it  is  now  accepted  that  phospho- 
rus (V)  chloride  in  the  vapor  state  is  made  up  of  trigonal  bipyramidal 
molecules35.  However,  in  the  crystalline  state  it  consists  of  PC14+  and  PC16~~ 
ions36, 37.  Measurements  of  the  electrical  moment38,  dielectric  constant,  and 


Fig.  10.5.  The  structure  of  iodine  pentafluoride 

conductivity39  in  inert  solvents  indicate  ionic  character,  so  it  is  assumed  that 
the  same  ions  exist  in  solution  as  exist  in  the  crystalline  state.  Phospho- 
rus (V)  bromide  is  composed  of  PBr4+  and  Br~  ions40. 

Compounds  of  the  type  R2[MX5]  have  been  prepared,  where  R  is  an 
alkali  metal  ion,  thallium(I),  or  an  ammonium  ion;  X  is  a  halide,  and  M  is 
antimony  or  bismuth.  In  addition,  bismuth  forms  a  corresponding  nitrate40 
and  the  trichlorodiamminebismuth(III)  complex41.  On  the  basis  of  color 

30.  Braune  and  Pinnow,  Z.  Physik,  B35,  239  (1937). 

31.  Rogers,  Wahrhaftig,  and  Schomaker,  Abstracts,  111th  Meeting  of  Am.  Chem. 

Soc,  April,  1947. 

32.  Lord,  Lynch,  Schumb,  and  Slowinski,  /.  Am.  Chem.  Soc,  72,  522  (1950). 

33.  De  Heer,  Phys.  Rev.,  83,  741  (1951). 

34.  Burke  and  Jones,  J.  Chem.  Phys.,  19,  1611  (1951). 

35.  Brockway  and  Beach,  J.  Am.  Chem.  Soc.,  60,  1836  (1938). 

36.  Clark,  Powell,  and  Wells,  J.  Chem.  Soc.,  1942,  642. 

37.  Moureu,  Magat,  and  Wetroff,  Compt.  rend.,  205,  545  (1937);  Clark,  Powell,  and 

Wells:  /.  Chem.  Soc,  1942,  642. 
,v  Trunel,  Compt.  rend.,  202,  37  (1936). 

39.  Holroyd,  Chadwick,  and  Mitchell,  /.  Chem.  Soc,  127,  2492  (1925). 

40.  Powell  and  Clark,  Nature,  145,  971  (1940). 

41.  Schwarz  and  Striebach,  Z.  anorg.  Chem.,  223,  399  (1935). 


COMPOUNDS  INVOLVING  LESS  COMMOh  COORDINATION  NUMBERS    389 

and  vapor  pressure  of  ammonia,  Schwarz  and  Strieback  postulate  that 
throe  chloride  ions  and  two  ammonia  molecules  are  attached  to  each  bis- 
muth atom.  However,  an  alternative  structure  could  be 

CI 

/  \ 

(Cl)o(NH3),Hi  Bi(Cl)2(XH,), 

\     / 

giving  the  bismuth  a  coordination  number  of  six.  A  dark  violet  antimony 
salt  of  the  formula  Tl[SbCl5]  is  known,  in  which  the  antimony  is  apparently 
tetravalent4*.  A  deep  color  of  this  kind  is  often  attributed  to  the  presence 
of  two  valence  states  of  an  element  in  one  compound,  so  the  compound  may 
well  be  TLJSb^^Sb^Clio].  The  same  applies  to  the  dark  violet  K2[TiF5]. 
This  may  be  a  mixed  titanium(II)  and  titanium(IV)  dinuclear  complex. 
However,  discrete  [SbF5]=  groups  exist  in  K2SbF5  (page  8). 

The  metal-organic  compound  (CH3)3SbCl2  in  the  crystalline  form  has 
been  shown  to  have  the  three  methyl  groups  in  the  plane  of  the  metal  atom 
with  the  two  chlorine  atoms  at  the  two  apices  of  a  trigonal  bipyramid43. 
The  compound  is  not  dissociated  in  inert  solvents.  It  slowly  undergoes 
stepwise  hydrolysis  in  water,  first  to  (CH3)3SbC10H  and  finally  to 
(CH3)3Sb(OH)2  .  The  first  product  is  a  very  strong  base  while  the  latter  is  a 
very  weak  base,  suggesting  that  the  first  may  be  a  substituted  stibonium 
hydroxide  (coordination  number,  four),  while  the  final  dihydroxide  is  simi- 
lar in  structure  to  the  original  dihalide  (coordination  number,  five). 

Many  compounds  are  known  in  which  the  central  atoms  appear  to  be 
five-coordinate  in  the  solid  state,  but  since  dissociation  takes  place  in  solu- 
tion, crystal  structure  studies  are  necessary  to  establish  the  true  coordina- 
tion number.  Cs3CoCl5  has  been  shown44  to  be  made  up  of  tetrahedral 
tctiachlorocobaltate(II)  ions  and  odd  cesium  and  chloride  ions,  so  the 
cobalt  is  actually  four-coordinate.  Klug  and  Alexander45  showed  that 
Ml^ZnCls  is  composed  of  tetrachlorozincate(II)  tetrahedra  and  am- 
monium and  chloride  ions  as  addenda.  Perhaps  diethylenetriamine  penta- 
chlorocuprate(II)49,  [dien-H3]  [CuCl5],  is  also  composed  of  planar  or  tetra- 
hedral tetrachlorocuprate(II)  ions  with  odd  chloride  ions  in  the  lattice. 

It  has  been  proved  that  the  compound  T12A1F5  is  composed  of  infinite 
chains  of  hexafluoroaluminate(III)  octahedra  in  which  the  two  opposite 
corners  are  shar*ed46  (Fig.  10.6). 

42.  Wells:  "Structural  Inorganic  Chemistry,"  p.  232,  London,  Oxford  University 

Press,  1945. 

43.  Wells,  Z.  Krist.,  99,  367  (1938). 

44.  Powell  and  Wells,  ./.  Chem.  Soc,  1935,  360. 

45.  Klug  and  Alexander,  J.  Am.  Chem.  Soc,  66,  1056  (1944). 

46.  Brosset,  Z.  anorg.  Chem.,  235,  139  (1937). 


390  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

F 
F    F/ 


Fig.  10.6.  The  structure  of  T12A1F5 

A  number  of  fluoride  and  oxyfluoride  compounds  which  apparently  have 
the  coordination  number  of  five  have  been  reported47.  Of  these,  there  is  some 
evidence  that  tetrafluorooxychomate(V)  and  pentafluoromanganate(IV) 
ions  are  actually  five-coordinate48.  Potassium  pentafluoromanganate(IV) 
is  only  slightly  colored,  and  its  x-ray  powder  patterns  show  that  no  im- 
purities such  as  potassium  fluoride,  manganese  (III)  fluoride  or  potassium 
hexafluoromanganate(IV)  are  present.  There  is  no  proof  of  structure  for 
these  compounds. 

Copper  is  also  reported  to  be  five-coordinate  in  the  black  crystalline 
compounds,  K3[Cu(N02)5],  Rb3[Cu(N02)5]50,  and  Tl3[Cu(N02)5]51.  Combes52 
prepared  the  ethylenediaminebisacetylacetone  (enac)  copper  salt  shown  in 

H 
CH3-C-C=C- CH3 
3      II  I  3 

CH2-N  O 

I         xo/ 

CH2  — N  O 

II  I 

CH3-C-C-C  —  CH3 
H 

Fig.  10.7.  The  structure  of  ethylenediamineacetylacetone  copper (II) 

Fig.  10.7,  which  is  violet  in  color  and  nonionic.  Morgan  and  Main-Smith53 
showed  that  this  complex  adds  one  molecule  of  ethylenediamine  and  one 
molecule  of  water  and  turns  dark  green.  When  placed  in  a  vacuum  desic- 
cator over  sodium  hydroxide  or  calcium  chloride,  two  molecules  of  water 
and  one  of  ethylenediamine  are  lost  from  two  molecules  of  the  salt,  pro- 

47.  Huss  and  Klemm,  Z.  anorg.  Chem.,  262,  25  (1950);  Zachariasen,  J.  Am.  Chem. 

Soc,  70,  2147  (1948);  Cefola  and  Smith,  Natl.  Nuclear  Energy  Ser.,  Div.  IV, 
14,  Transuranium  Elements,  Pt.  I,  822  (1949). 

48.  Sharpe  and  Woolfe,  J.  Chem.  Soc,  1951,  798. 

49.  Jonassen,  Crumpler,  and  O'Brien, J.  Am.  Chem.  Soc.,  67,  1709  (1945). 

50.  Kurtenacker,  Z.  anorg.  Chem.,  82,  204  (1913). 

51.  Cuttica  and  Paciello,  Gazzetta,  52,  141  (1922). 

52.  Combes,  Compt.  rend.,  108,  1252  (1889). 

53.  Morgan  and  Main-Smith,  J.  Chew.  Soc.,  1925,  2030;  ibid.,  1926,  913. 


COMPOUNDS  INVOLVING  LESS  COMMON  COORDINATION  NUMBERS    391 

ducing  the  bridged  dinuclear  compound 

[(enac)CuNH,CH  (II  Ml  ,Cu(enac)], 

in  which  the  copper  seems  to  have  a  coordination  number  of  five. 

Thorium  forms  the  aonelectrolyte  |Th  IV '( 'li( '.-.I I.-,N ]  and  the  complex 
salt,  NaJTh  ,N  ..('<>.  •,!•  L2H20,  the  latter  being  isomorphous  with 
Na€[Ce<nr)(CO,)s]-12HiOM.   Lortie  showed  thai  teo  of  the  twelve  water 

molecules  are  removed  very  easily,  while  the  other  two  are  removed  only 
with  difficulty. 

Kay  and  Dutt55  carefully  dehydrated  the  yellow  diamagnetic  silver  penta- 
cyanoaquocobaltate(III)  complex  and  obtained  a  compound  with  the 
formula  Agj[Co(CN)s],  This  compound  is  deep  blue  in  color  and  paramag- 
netic, both  properties  indicating  unpaired  electrons.  Similarly,  Adamson56 
has  prepared  potassium  pentacyanocobaltate(II),  K3[Co(CX)5],  and  postu- 
lated that  the  cobalt  has  a  coordination  number  of  five  in  solution;  however, 
the  electronic  configuration  and  molecular  structure  of  the  complex  are 
still  open  to  question.  It  is  possible  that  the  true  ionic  species  in  solution 
is  pentacyanoaquocobalate(II)  ion,  [H2OCo(CX)5]^,  as  has  been  shown  to 
be  the  case  with  pentachloroindate(III)  ion,  which  is  actually  pentachloro- 
aquoindate(III)57. 

Cobalt  is  apparently  five-coordinate  in  the  bis-salicylaldehyde-7,Y'-di- 
aminodipropylamine  salt  (I)5S.  The  crystalline  compound  shown  in  (II) 

I 


9- 


O      0^\?  I  Co^- O  -^coC 

Co'  I  H2C"NC         "9  °^    ^N"CH2 

(CH^-N-fcH^a 

CD  (n) 

Fig.  10.8 

(Fig.  10.8)  was  prepared  byDiehl59,  who  assumes  a  coordination  number  of 
five  for  the  cobalt  because  of  a  water  bridge  in  the  dinuclear  molecule.  This 
seems  to  be  the  first  case  reported  in  which  a  water  molecule  acts  as  a 

54.  Lortie,  Compt.  raid.,  188,  915  (1929). 

55.  K;.v  and  Dutt,  Current  Science,  5,  476     \{.vtf). 

56.  Adamson, ./.  Am.  Chem.  Soc,  73,  5710  (1951). 

57.  Klut..  Kummer,  and  Alexander,  ./.  Am.  Chem.  8oc.,  70,  3064     1948). 

58.  Calvin,  et  al.,  ./.  Am.  Chem.  Soc,  68,  2254,  2012    194 

59.  Diehl,  et  al.,  Iowa  StaU  College  J.  of  Sri.,  21,  No.  3,  27s  [1947  . 


392 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


bridging  group.  It  is  possible  that  this  water  is  not  actually  coordinated 
but  is  lattice  water. 

Both  iron  and  ruthenium  form  pentacarbonyls  of  the  general  formula 
M(CO)5  .  It  has  been  shown  by  electron  diffraction  studies  that  in  iron 
pentacarbonyl  the  carbonyl  groups  are  distributed  around  the  iron  at  the 
apices  of  a  trigonal  bipyramid  (Chapter  16). 

Tribromobis(triethyl  phosphine)nickel(III)  is  an  unusual  compound  in 
two  respects:  it  contains  nickel  (III)  and  it  exhibits  a  coordination  number 
of  five.  Molecular  weight  determinations  in  benzene  solution  indicate  that 
it  is  monomeric  and  not  dissociated.  The  magnetic  moment  is  consistent 
with  the  presence  of  one  unpaired  electron.  On  the  basis  of  dipole  moment 
measurements,  Jensen  and  Nygaard60  have  assumed  that  the  molecule 
exists  in  the  form  of  a  tetragonal  pyramid. 

In  postulating  mechanisms  for  the  reactions  of  complex  compounds, 
especially  aquation,  some  investigators  propose  the  formation  of  inter- 
mediates, in  which  a  normally  6-coordinate  central  atom  has  a  coordina- 
tion number  of  5  or  7.  The  number  5  is  indicated  when  the  reaction  seems 
to  be  a  SN1  type,  and  7  when  the  reaction  appears  to  be  the  SN2  type.  In 
view  of  the  transient  nature  of  such  complexes  they  will  not  be  discussed 
further  here.  (See  pp.  327  and  329). 

Coordination  Number  Seven 

The  coordination  number  of  seven  is  quite  rare,  and  the  fact  that  it 
appears  generally  in  the  heavier  atoms,  such  as  zirconium,  niobium,  tanta- 
lum, and  iodine,  leads  one  to  suspect  that  J  electrons  are  significant  in 
bonding,  although  structures  have  been  deduced  which  require  only  s,  p, 
and  d  orbitals.  The  halogens  in  general  (especially  fluorine)  seem  to  favor 


(I)  (II) 

Fig.  10.9.  Coordination  number  seven 


cm) 


this  coordination  number.  Three  structures  have  been  proposed  for  mole- 
cules and  ions  exhibiting  the  coordination  number  of  seven  (Fig.  10.9).  They 
are  (I)  the  trigonal  prism18  in  which  a  seventh  coordination  position  exists 
beyond  one  lace,  (II)  the  octahedron  with  a  seventh  bond  beyond  the 
center  of  one  face18,  and   (III)  the  pentagonal  bipyramid61.  The  hybrid 

GO.  Jensen  and  Nygaard,  Acta  Chan.  Scand.,  3,  474  (1949). 
61.  Duffey,  ./.  ('hem.  Phys.,  18,  943  (1950). 


(  DM  POUNDS  INVOLVING  LESS  <  VMM  ON  COORDINATION  NUMBERS    393 

states  proposed  for  these  configurations  are  (I)  '/'.sp2,  d4p3,  dbp2;  (II) 
dtsp,  (/;.s'/;:5;  (III)  sp'ut'K  and  other  hybrid  configurations  requiring/  elec- 
trons81, 

Compounds  of  the  general  formula  R*wM(IV)Fi  arc  known,  in  which  K  is  a 
sodium,  potassium  or  ammonium  ion,  and  M  is  silicon,  I  ilanium,  zirconium, 
hafnium,  or  lead.  The  ammonium  "heptafluorosilicate"  has  been  reported 
to  be  made  up  of  discrete  hexafluorosilicate(IV),  ammonium,  and  fluoride 
ions68,  so  the  authors  propose  to  write  the  formula  (NH^SiFel'NHJT  to 
emphasize  that  the  central  atom  is  six-  rather  than  seven-coordinate.  On 
the  other  hand,  the  analogous  compound,  potassium  heptafluorozirconate 
(IY\  K3[ZrF7],  has  been  shown  to  consist  of  finite  heptafluorozirconate (IV) 
ions  in  the  crysalline  state64,  the  zirconium  atom  being  at  the  center  of  an 
octahedron  of  fluorine  atoms  with  the  seventh  or  odd  fluorine  above  the 
cent ta-  of  one  face.  The  octahedron  is  somewhat  distorted  by  a  forced  sepa- 
ration of  the  atoms  at  the  corners  of  this  face.  Hassell  and  Mark65  have 
shown  that  the  hafnium  and  zirconium  compounds  are  isomorphous,  so 
hafnium  probably  has  a  coordination  number  of  seven  in  its  analogous 
compound.  Another  fourth  group  element,  tin,  is  apparently  seven-co- 
ordinate66 in  the  compound  Na(C5H5NH)2[Sn(NCS)7]. 

Klemm  and  Huss  prepared  potassium  heptafluorocobaltate(IV)  by  the 
action  of  gaseous  fluorine  on  mixtures  of  potassium  chloride  and  cobalt(II) 
chloride67.  X-ray  studies  indicate  that  it  probably  has  a  structure  similar 
to  that  associated  with  the  salts  of  the  heptafluorozirconate (IV)  ion  (Struc- 
ture II,  Fig.  10.9). 

The  elements  of  the  fifth  Periodic  Group  form  compounds  of  the  general 
formula  R2(I)[M(V)F7]  where  R  is  potassium,  hydrogen  or  ammonium  ion  and 
Z\I  is  antimony,  niobium,  or  tantalum.  Neither  arsenic  nor  vanadium  seems 
to  form  this  type  of  compound.  Both  the  niobium  and  tantalum  compounds 
are  truly  seven-coordinate  since  their  finite  heptafluoro  ions  have  been 
proved  to  exist.  Hoard  and  coworkers68  have  shown  that  in  the  solid  state 
the  seventh  fluorine  atom  is  added  beyond  the  center  of  one  of  the  rectangu- 
lar faces  of  a  trigonal  prism.  A  number  of  hydroxy  organic  derivatives  of 
niobium  and  tantalum,  such  as  those  with  catechol,  (NH^NbCKCeKUC^s], 
and  with  acetylacetone,  (NH^INbO^HeC^],  are  reported  to  be  seven- 
coordinate69. 

62.  Shirmazan  and  Dyatkina,  Doklady  Akad.  Nauk  S.S.S.R.,  77,  75  (1951). 

63.  Hoard  and  Williams,  J.  Am.  Chem.  Soc,  64,  633  (1942). 
54.  Hampson  and  Pauling,  ./.  Am.  Chem.  Soc,  60,  2702  (1938). 

65.  Hassel  and  Mark,  Z.  Phys.,  27,  89  (1924). 

66.  Weinland  and  Barnes,  Z.  anorg.  Chem.,  62,  250  (1909). 

67.  Klemm  and  Huss,  Z.  anorq .  Chem.,  258,  221  (1949). 

68.  Hoard,  J.  Am.  Chem.  Soc,  61,  1252  (1939) ;  ibid.,  63,  1 1  (1941). 

69.  Rosenheim  and  Roehrich,  Z.  anorg.  Chem.,  204,  342  (1932). 


394 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Other  compounds  reported  to  contain  seven-coordinate  atoms  are  the 
black  (CN3NH2.H)3[Pt<IV>l7]7(),  dark  red-brown  (CH3NH2H)4[RuCl7]71,  and 
K3[U02F5]. 

Iron  enneacarbonyl,  Fe2(CO)9  ,  is  postulated  to  contain  seven-coordinate 
iron  (Chapter  16). 

On  the  basis  of  Raman  and  infrared  spectra32,  iodine (VII)  fluoride  has 
been  assigned  the  pentagonal  bipyramidal  structure  ((III),  Fig.  10.9). 

Coordination  Number  Eight 

In  general,  substances  containing  eight-coordinate  central  atoms  can 
give  rise  to  so  many  stereoisomers  that  a  chemical  determination  of  their 
structures  is  almost  impossible.  The  configurations  of  only  a  few  eight-co- 
ordinate groups  have  been  studied. 

The  cube  (I)  was  the  first  structure  proposed  for  eight-coordinate  com- 
plexes72; this  configuration  was  shown  by  Penny  and  Anderson73  to  be  con- 
sistent with  the  theory  of  molecular  orbitals.  The  Archimidean  tetragonal 
antiprism13, 74  (II),  a  trigonal  prism  with  two  extra  bonds  at  the  extremities 
of  the  unique  axis75  (III),  the  dodecahedron13- 76  (IV),  and  a  trigonal  prism 
in  which  the  two  extra  bonds  extend  beyond  the  centers  of  two  of  the  rec- 
tangular faces13  (V)  have  also  been  considered  to  be  feasible  configurations 
(Fig.  10.10). 


t 


M 


<^ 


(D 


(n)  cnr)  (m) 

Fig.  10.10.  Coordination  number  eight 


Csn 


Calculations  made  by  Duffey77  indicate  that  either  the  dodecahedron 
or  the  tetragonal  antiprism77, 78  may  be  attained  through  a  hybrid  of  the 
type  d4spz,  while  the  trigonal  prism13  in  which  the  extra  bonds  appear  in 
rectangular  faces  may  assume  dbsp2  hybridization.  However,  the  trigonal 
prismatic  structure  in  which  the  last  two  ligands  are  added  above  the  cen- 

70.  Anon.,  Chem.  Centr.,  II,  143  (1914). 

71.  Gutbier,  Ber.,  56,  1008  (1923). 

72.  Pfeiffer,  Z.  anorg.  Chem.,  105,  26  (1919). 

73.  Penny  and  Anderson,  Trans.  Faraday  Soc,  33,  1363  (1937). 

74.  Huttig,  Z.  anorg.  allgem.  Chem.,  114,  25  (1920). 

75.  Marchi  and  McReynolds,  J.  Am.  Chem.  Soc,  65,  333  (1943). 

76.  Hoard  and  Nordsieck,  J.  Am.  Chem.  Soc.,  61,  2853  (1939). 

77.  Duffey,  J.  Chem.  Phys.,  18,  1444  (1950). 

78.  Duffey,  J.  Chem.  Phys.,  18,  746  (1950). 


COMPOUNDS  INVOLVING  LESS  i  OMMOh  COORDINATION  NUMBERS    395 

ters  of  the  triangular  faces  cannot  be  realized  in  the  absence  of /orbital* 
It  is  also  reported  thai  /  orbitala  are  required  in  the  cubic  structure18,  s". 
Definite  evidence  for  the  presence  of  /electrons  in  eight-coordinate  mole- 
cules has  been  reported  by  Sacconi81,  who  studied  the  magnetic  properties 
of  uranium(IV)  complexes  with  a  series  of  /8-diketones.  The  results  indicate 
thai  two  5/ electrons  are  involved  in  the  bonding. 

Probably  the  most  widely  studied  compounds  are  the  octacyanides  of 
molybdenum  and  tungsten,  which  have  the  formulas  M4(I)[M(IV)(CX  )s| 
and  M,;1  [M^(CX)8].  Potassium  octacyanomolyhdate(IV)  is  yellow  and 
can  be  prepared  by  air  oxidation  of  potassium  hexachloromolybdate(III) 
in  the  presence  of  excess  potassium  cyanide,  or  by  the  reduction  of  molybde- 
num(V)  compounds  with  potassium  cyanide.  Hoard  and  Xordsieck76  have 
shown  the  existence  of  individual  octacyanomolybdate(IV)  ions,  with  the 
eight  cyanide  groups  arranged  at  the  apices  of  a  dodecahedron.  The  carbon- 
nitrogen  bonds  are  colinear  with  the  molybdenum-carbon  bonds.  It  is 
presumed  that  the  orbitals  used  are  four  4c?,  one  bs  and  three  5p,  although 
Van  YleclO2  has  predicted,  on  theoretical  grounds,  that  s,  p,  d,  and/ orbitals 
must  all  be  available  for  bonding  in  order  to  attain  symmetrical  distribution 
of  eight  coordinated  groups.  It  is  interesting  to  note  that  /  electrons  do  not 
appear  in  neutral  atoms  until  element  58,  cerium.  One  must  assume,  then, 
that  in  the  octacyanomolybdate(IV)  ion,  where  there  are  several  more 
electrons  than  there  would  be  if  the  system  were  electrically  neutral,  the 
4/  orbitals  are  comparable  in  stability  to  other  orbitals  in  the  4  shell.  On 
the  basis  of  effective  atomic  number,  one  would  expect  a  greater  stability 
for  octacyanomolybdate(IV)  (E.A.X.,  54)  than  for  octacyanomolybdate(V) 
(E.A.X.,  53),  and  the  former  is  actually  more  stable. 

Some  of  the  substituted  octacyanides  which  have  been  reported  are 
W  OHMCN),]*-88,  [Mo(CX)7H2Op- 8*,  [Mo(OH)4(CX)4]4- 85,  and 

[Mo(OH)3(CX)4H20]s  86. 

Fluorine  also  seems  to  favor  eight-coordination  as  exhibited  in  the  com- 
pounds   (XH4)3H[PbF8]87,   H^SbFs]88,    Xa3[TaF8]89,    and   the   well-known. 

79.  Shirmazan  and  Dyatkina,  Doklady  Akad.  Sauk  S.S.S.R.,  82,  755  (1952). 

80.  Racah,  J.  Chem.  Phys.,  11,  214  (1943). 

81.  Sacconi,  Atti  accad.  uazl.  Lincei,  Rend.  Classe  sci.fiz.,  mat.  e  nat.,  6,  639  (1949). 

82.  Van  Vleck,  J.  Chem.  Phys.,  3,  803  (1935). 

83.  Collenberg,  Z.  anorg.  Chem.,  136,  249  (1924). 

84.  Young,  J.  Am.  Chem.  Soc,  54,  1402  (1932). 

85.  YonderHeide  and  Hofman,  Z.  anorg.  Chem.,  12,  285  (1896). 
Bucknall  and  Wardlaw,  ./.  Chem.  Soc,  1927,  2989. 

B7.  RufT;  Z.  anorg.  Cht  n   .  98,  27  (1916). 

yv  Morgan  and  Buratall,  "Inorganic  Chemistry,"  p.  145,  New  York,  Chemical 

Publishing  Co.,  1937. 
89.  de  Marignac,  Compt.  rend.,  63,  86  (1866). 


396  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

highly  volatile  osmium (VIII)  fluoride.  Hoard90  has  shown  by  x-ray  crystal 
analysis  that  the  octafluorotantalate(V)  ion  forms  a  tetragonal  antiprism. 
Kimball13  predicts  that  osmium (VIII)  fluoride  will  be  found  to  have  a 
face-centered  prismatic  structure. 

An  attempt  by  Marchi  and  McReynolds91  to  determine  the  structure  of 
K^U^OJJ  by  chemical  means  was  only  partially  successful.  They 
assumed  four  possible  structures;  the  cube  (I),  the  Archimidean  anti- 
prism(II),  the  trigonal  prism  with  two  extra  bonds  along  the  unique 
axis(III),  and  the  dodecahedron  with  triangular  faces(IV).  The  trigonal 
prism  with  two  extra  bonds  along  the  normal  to  two  of  the  rectangular 
faces(V)  was  also  mentioned  as  an  alternative  structure.  Of  these,  (I)  and 
(III)  would  not  show  optical  isomerism  for  an  ion  of  the  type  of  [U^O^J4- 
while  (II),  (IV),  and  (V)  would.  Structure  (II)  would  have  six  optical 
isomers,  while  (IV)  and  (V)  would  each  have  ten.  The  authors  succeeded 
in  isolating  four  optical  isomers  by  fractional  precipitation  of  the  strych- 
nine salt.  One  pair  of  optical  isomers  racemized  rapidly,  and  the  other  pair 
was  stable.  These  results  eliminate  structures  (I)  and  (III)  but  do  not 
distinguish  between  (II),  (IV),  and  (V). 

Other  compounds  reported  in  which  the  central  atom  apparently  has  a 
coordination  number  of  eight  are  the  octammines,  MX2-8NH3 ,  where  M 
is  calcium,  strontium,  barium92,  or  lead93;  metal  acetylacetonates, 
M(C5H702)4 ,  where  M  is  zirconium94,  hafnium95,  thorium95,  uranium96, 
polonium97,  or  cerium98;  tetrakis(ethylenediamine)  chromium  (III)  chlo- 
ride99, and  tetrakis(ethylenediamine)cadmium(II)  iodide100;  other  oxalate 
complexes  similar  to  the  uranium  compound  [M^O^]4-  discussed  above, 
where  M  is  zirconium101,  hafnium101,  thorium102,  or  tin103;  tin (IV)  phthalocya- 
nine104;  and  tetrakis(8-hydroxyquinoline)plutonium(IV)105. 

90.  Hoard,  Paper  presented  at  the  6th  annual  symposium,  Div.  Phys.,  and  Inorg. 

Chem.,  Columbus,  Ohio,  December,  1941. 

91.  Marchi  and  McReynolds,  J.  Am.  Chem.  Soc,  65,  333  (1943). 

92.  Huttig,  Z.  anorg.  Chem.,  123,  31   (1922);  ibid.,  124,  322  (1922);  ibid.,  125,  269 

(1922). 

93.  Biltz  and  Fischer,  Z.  anorg.  Chem.,  124,  230  (1922). 

94.  Von  Hevesy  and  Logstrup,  Ber.,  59B,  1890  (1926). 

95.  Young,  Goodman,  and  Kovitz,  J.Am.  Chem.  Soc,  61,  876  (1939). 

96.  Biltz,  Z.  anorg.  Chem.,  40,  220  (1904). 

97.  Servigni,  Compt.  rend.,  196,  264  (1933). 

98.  Scagliarini,  Atti  accad.  Lincei,  [6]  4,  204  (1926). 

99.  Lang  and  Carson,  J.  Am.  Chem.  Soc,  26,  759  (1904). 

100.  Barbier,  Compt.  rend.,  136,  688  (1903). 

101.  Tchakirian,  Compt.  rend.,  204,  356  (1937). 

102.  Brauner,  ./.  Chem.  Soc,  73,  956  (1898). 

103.  Rosenheim  and  Platsch,  Z.  anorg.  Chem.,  20,  309  (1899). 
101.  Barret,  Dent,  and  Linstead,  ,/.  Chem.  Soc,  1936,  1733. 

Hi.").  Pat  ton,  Natl.  Nuclear  Energy  Ser.  Div.  IV,  14B,  Transuranium  Elements,  Pt.  I, 
853  (1949). 


COMPOUNDS  INVOLVING  LESS  COMMON  COORDINATION  NUMBERS    397 

Coordination  Numbeh  Greatbb  than  Eight 

Coordination  aumbers  greater  than  eight  have  been  postulated  for  such 
compounds  as  \a.,ZrF<>  and  many  hydrates  and  ammoniates.  In  some  of 
these,  such  as  [Nd(HjO)J(BrOa)  .  the  central  atom  has  a  coordination 
aumber  of  nine  in  the  crystallographic  sense,  hut  it  is  doubtful  whether 

these  coordination  aumbers  exist  in  the  original  Werner  sense. 

1  Mit'tey  '"  has  predicted  that  compounds  of  the  type  M  '  <  >sF9  should  have 
a  structure  consisting  of  a  trigonal  prism  with  one  atom  added  to  each  of 
the  four-sided  faces.  He  refers  to  this  st  net  lire  as  an  irregular  t  ripyrannd. 
Shirmazan  and  Dyatkina6-  offer  several  hybrid  configurations  as  con- 
sistent with  this  structure.  Of  these,  only  sp*db  does  not  require/  electrons. 

106.  Duffey,  J.  Chem.  Phys.,  19,  553  (1951). 


I.   Stabilization  of  Valence  States 
Through  Coordination 

James  V.  Quagliano 

Notre  Dame  University,  Notre  Dame,  Indiana 

and 

R.  L  Rebertus 

Shell  Development  Co.,  Emeryville,  California 

One  of  the  most  familiar  and  useful  chemical  concepts  is  that  of  relative 
stability  of  chemical  compounds,  and  the  coordination  theory  accounts  for 
the  existence  and  relative  stabilities  of  many  complex  compounds.  Mul- 
liken1  has  pointed  out  that  by  sharing  or  transferring  electrons  a  nucleus  in 
a  molecule  tends  to  be  surrounded  by  a  stable  electronic  configuration  with 
a  total  charge  approximately  equal  to  that  of  the  nucleus.  However,  the 
term  "stability"  is  vague  and  is  used  in  many  different  ways.  Reference  is 
made  to  stability  toward  aquation,  thermal  decomposition,  oxidation, 
reduction,  and  other  types  of  reactions.  Hydrogen  peroxide,  for  example, 
is  unstable  toward  decomposition  into  water  and  oxygen  but  is  very  stable 
toward  decomposition  into  hydrogen  and  oxygen2. 

In  this  chapter  stability  toward  oxidation  and  reduction  is  emphasized, 
and  of  especial  interest  are  those  valence  states  which  cannot  exist  unless 
stabilized  through  coordination. 

Quantitative  Measurement  of  the  Degree  of  Stabilization 

Oxidation  Potentials 

The  concept  of  electron  loss  or  gain  has  long  been  associated  with  oxida- 
tion or  reduction.  As  applied  to  the  formation  of  an  essentially  ionic  com- 
pound, as  by  the  reaction  of  chlorine  with  sodium,  this  concept  is  nearly 
correct.  Ambiguity  arises,  however,  when  an  attempt  is  made  to  apply  elec- 
tron loss  or  gain  to  covalent  compounds.  Moeller3  suggests  that  it  is  more 

I     Mulliken,  Phys.  Rev.,  41,  60  (1932). 

2.  Hildebrand,  Chem.  Revs.,  2,  395  (1926). 

3.  Moeller,  "Inorganic  Chemistry,"  New  York,  John  Wiley  &  Sons,  Inc.,  1952. 

398 


STABILIZATION  OF  VALENCE  STATES  399 

nearly  correct  to  consider  oxidation-reduction  as  an  increase  or  decrease  in 
oxidation  state;  this  may  be  brought  about  with  no  change  iii  the  number 
of  electrons  associated  with  a  particular  nucleus.  This  tendency  toward  an 
increase  or  decrease  in  oxidation  state  can,  in  many  instances,  he  measured 
quantitatively  and  expressed  as  the  oxidation  potential  of  a  half  cell  re- 
action.* Potential  data  have  been  published  by  Latimer4. 

In  general,  the  oxidation  potential  of  any  half-reaction  is  altered  when 
the  activities  of  the  reactants  or  products  are  changed.  The  potential  of  the 
half-cell  reaction, 

Fe++  -*  Fe+++  +  e", 

can  he  described  in  terms  of  the  Nernst  equation, 

E  =  E°  -  RT/nF  In  aFeWaFe^, 

where  E  is  the  potential  at  any  activity  of  product  or  reactant,  E°  is  the 
standard  potential  taken  at  unit  activities,  n  is  the  number  of  electrons  in- 
volved in  the  reaction,  F  is  the  Faraday  constant,  T  is  the  absolute  tempera- 
ture, R  is  the  gas  constant,  and  a  is  the  activity  of  product  or  reactant. 

One  method  of  changing  the  activity  of  a  product  or  a  reactant  is  to  co- 
ordinate the  ion  in  question  with  a  complexing  agent.  The  resulting  change 
in  oxidation  potential  is  a  quantitative  measure  of  the  degree  to  which  the 
particular  valence  is  stablized  relative  to  the  couple  consisting  of  aquated 
ions.  (It  is  customary  in  writing  equations  for  half-cell  reactions  in  aqueous 
solutions  not  to  describe  aquated  ions,  though  this  would  be  more  nearly 
correct.)  A  few  examples  of  this  phenomenon  follow. 

Iron (Il)-Iron (III)  Couple.  It  was  shown  in  1898  by  Peters5  that  the 
oxidation  potential  of  mixtures  of  iron  (I I)  and  iron  (III)  chlorides  in  hydro- 
chloric acid  depends  upon  the  concentration  of  the  acid.  The  system  was 

'  ( '<  mfusion  sometimes  arises  in  the  literature  with  regard  to  convention  of  sign  of 
potentials  for  oxidation-reduction  couples.  See,  for  example,  Latimer,  /.  Am.  Chem. 
Soc,  76,  1200  (1954).  If  the  number  of  electrons  required  to  balance  the  equation  is 
written  on  the  right  hand  side,  any  half-cell  reaction  expressed  as  'reduced  state  = 
oxidized  state  +  n  electrons'  may  be  described  with  an  oxidal  ion  potential.  A  positive 
value  indicates  that  the  reduced  form  of  the  couple  is  a  better  reducing  agenl  than 
Hj  .  This  is  based  on  the  selection  of  thermodynamic  conventions  by  (i.  X.  Lewis 
but  is  commonly  referred  to  as  Latimer's  system.  This  convention  will  lie  adhered  to 
in  this  chapter  except  in  the  discussion  of  polarography.  Polarographers,  in  general, 
choose  to  write  the  requisite  number  of  electrons  on  the  left  in  the  general  form: 
oxidised  Btate  -f  n  electrons  =  reduced  state,  and  the  sign  of  potential  is  the  opposite 
of  Latimer's  sign  for  any  half -cell  reaction. 

4.  Latimer,  "Oxidation  Potentials/'  2nd  Edition,  New  York,  l'rentice-Hall,  Inc., 

1952. 

5.  Peters,  Z.  physik.,  26,  193  (1898). 


400  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Table  11.1.  Effect  of  Coordination  on  the  Iron  (II)-Iron  (III)  Couple 

Equation  Potential  (£°) 

Fe^  ^5  Fe^  +  e"  -0.771 

[Fe(CN)6]4-  ±5  [Fe(CN),]-  +  e~  -0.36 

Fe++  +  6F-  ±5  [FeF6]s  +  e~  -0.40 

Fe^  +  2P043  ±=>  [Fe(P04)«]"  +  e~  -0.61 

Table  11.2.  Stabilization  of  Iron  (II)  by  Coordination 

Equation  Potential  (£°) 

[Fe  (dipy)  3]++  *±  [Fe  (dipy )  3]+++  +  e~  -1.10 

[Fe  (o-phen)  ^  <=±  [Fe  (o-phen)  3]+++  +  e~  - 1 .  14 

[Fe(nitro-o-phen)3]++  +±  [Fe(nitro-o-phen)3]+++  +  e~  -1.25 

studied  in  more  detail  by  Carter  and  Clews6,  who  found  that  the  oxidation 
potential  decreases  as  the  concentration  of  the  acid  is  increased.  The  change 
in  potential  was  explained  by  a  change  in  the  ratio  of  th<  iron  (I I)  to  iron- 
fill)  ions  as  a  result  of  the  complexing  of  the  iron(III)  ion  with  chloride 
ions.  PopofT  and  Kunz7  confirmed  the  report  of  Carter  and  Clews.  Similar 
investigations  were  made  in  sulfuric  acid  medium  by  Glover8,  and,  again, 
evidence  for  complex  formation  was  reported. 

In  Table  11.1  standard  potentials  are  listed  for  the  iron(II)-iron(III) 
couple  in  the  presence  of  different  complexing  agents.  The  hexacyanofer- 
rate(II)  ion  is  thermodynamically  less  stable  toward  oxidation  than  is  the 
aquated  iron(II)  ion,  and  the  apparent  chemical  stability  of  the  hexacyano- 
ferrate(II)  ion  is  attributed  to  the  slowness  of  the  rate  of  oxidation  under 
usual  experimental  conditions.  Rate  of  oxidatior  or  reduction  should  not  be 
confused  with  thermodynamic  stability.  The  data  in  Table  11.1  indicate 
that  cyanide,  fluoride,  and  phosphate  stabilize  iron  (III)  against  reduction 
to  a  greater  degree  than  does  water. 

Many  complexing  agents  stabilize  the  dipositive  state  of  iron.  Of  these, 
the  ones  listed  in  Table  11.2  also  possess  properties  desirable  in  indicators 
for  oxidimetry. 

Cerium (Ill)-Cerium (IV)  Couple.  A  study  of  the  influence  of  complex 
formation  on  the  oxidation  potentials  of  cerium(III)-cerium(IV)  ni- 
trates in  nitric  acid  by  Noyes  and  Garner9  revealed  the  lack  of  dependence 
of  the  oxidation  potential  upon  the  acid  concentration  over  a  relatively 
short  range.  Kunz10  found  little  change  in  the  oxidation  potential  of  cerium- 
(III)  and  cerium(IV)  sulfates  in  solutions  of  sulfuric  acid.  G.  F.  Smith  and 

6.  Carter  and  Clews, ./.  Chem.  Soc,  125,  1880  (1924). 

7.  Popoff  and  Kunz, ./.  Am.  Chem.  Soc,  51,  382  (1929). 

8.  Glover,  ./.  Chem.  Soc,  1933,  10. 

9.  Noyes  and  Garner,  J.  Am.  Chew.  Soc,  58,  1265  (1936). 
10.  Kunz,  J.  Am.  Chem.  Soc,  53,  98  (1931). 


STABILIZATION  OF  VALENCE  STATES  401 

his  co-workers  extended  the  potential  measurements  to  acid  concentrations 

as  high  as  S  normal11.  They  found  that  the  potential  of  the  system  in  mix- 
tures o(  nitrate  and  sulfate"  at  lower  acid  concentrations  exhibited  the  con- 
stancy reported  by  the  previous  investigators  but  that  at  higher  acid  con- 
centrations the  oxidation  potential  decreased  markedly.  However,  the 
results  of  experiments  in  perchloric  acid  solution  showed  an  opposite 
effect.  The  formation  and  stability  of  complex  ions  are  undoubtedly  re- 
sponsible for  the  potential  changes  in  nitrate  and  sulfate  media  hut  not  in 
perchloric  acid  solution.  An  extensive  study  of  the  system  in  perchloric 
acid  solution  was  made  by  Sherrill,  King  and  Spooner12  to  determine  the 
effect  of  perchlorate  ion  concentration  and  hydrogen  ion  concentration.  The 
potential  was  found  to  vary  with  hydrogen  ion  concentration  and  was  de- 
pendent upon  the  hydrolysis  of  cerium (IV)  perchlorate  to  form  the  ions 
Ce(OH)+++  and  Ce(OH)2++.  Postulating  that  these  complex  ions  exist  in 
solution,  Heidt  and  Smith13  presented  evidence  for  the  formation  of  dimers 
resulting  from  the  splitting  out  of  water  from  the  hydroxyl  groups  of  these 
ions. 

Thallium(I)-Thallium(III)  Couple.  Investigations  of  the  thallium(I)- 
thallium(III)  couple  show  that  the  oxidation  potential  depends  to  a  large 
extent  on  the  nature  of  various  complex  ions  present14.  Thallium(I)  chloride 
in  hydrochloric  acid  is  more  easily  oxidized  to  thallium  (III)  than  is  thal- 
lium^) sulfate  or  nitrate  in  solutions  of  sulfuric  or  nitric  acid,  resulting  from 
the  formation  of  stable  chlorothallate(III)  complexes.  Since  nitric  acid 
and  perchloric  acid  do  not  appreciably  alter  the  oxidation  potential  of  the 
thallium(I)-thallium(III)  couple,  it  was  assumed  that  no  complex  forma- 
tion occurs  with  the  anions  of  these  acids. 

Zinc(O)-Zincdl)  Couple.  The  complexes  formed  by  zinc  ion  with 
hydroxyl  ion  are  among  the  most  stable  and,  from  the  standpoint  of  theo- 
retical significance,  the  most  interesting  of  the  numerous  zinc  coordination 
compounds.  The  data  of  Table  11.3  indicate  that  amphoterism  may  lead 
to  the  stabilization  of  a  valence  state  through  coordination. 

Cobalt(H)-Cobalt(III)  Couple.  The  aquated  cobalt(III)  ion  reacts 
with  water  to  liberate  oxygen.  On  the  other  hand,  the  hexacyanocobaltate- 
(II)  ion  is  a  powerful  reducing  agent  and  is  oxidized  by  water  with  the 

11.  Smith,  Sullivan,  and  Frank,  hid.  Eng.  Chem.,  Anal.  Ed.,  8,  449  (1936);  Smith  and 

Getz,  Ind.  Eng.  Chem.,  Anal.  Ed.,  10,  191  (1938);  ibid.,  10,  304  (1938). 
!_\  Sherrill,  King,  and  Spooner,  ./.  Am.  Chem.  Soc,  65,  170  (1943). 

13.  Heidt  and  Smith,  J.  Am.  Chem.  Soc,  70,  2476  (1948). 

14.  Spencer  and  Ahegg,  Z.  anorg.  Chem.,  44,  379  (1905);  Gruhe  and  Hermann,  Z. 

Elektrochem.,  33,  112  (1927);  Partington  and  Stonehill,  Trans.  Faraday  Soc,  31, 
1357  (1935);  Sherrill  and  Haas,  J.  Am.  Chem.  Soc,  65,  170  (1943);  Noyes  and 
Garner,  ./.  Am.  Chem.  Soc,  58,  1268  (1936). 


402  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Table  11.3.  Stabilization  of  Zinc (II)  Through  Hydroxyl  Ion  Coordination 

Equation  Potential  (£«) 

Zn  <=>  Zn++  +  2  e"  0.762 

Zn  +  20H"  ^  Zn(OH)2  +  2e-  1.245 

Zn  +  40H-  *±  Zn02-  +  2  H20  +  2e~  1.216 

Table  11.4.  Stabilization  of  Cobalt(III)  Through  Coordination 

Equation  Potential  (£°) 

Co++^±Co+++  +  e~  -1.82 

[Co(NH3)6]"H"  <=*  [Co(NH3)6]+++  +  e-  -0.1 

[Co(CN)6]4-  <=±  [Co(CN)6p  +  e-  +0.83 

evolution  of  hydrogen15.  Table  11.4  shows  the  wide  variations  in  the  oxida- 
tion potentials  of  the  cobalt(II)-cobalt(III)  couple  in  the  presence  of  co- 
ordinating groups.  The  hexamminecobalt(III)  ion,  a  slightly  better  oxidiz- 
ing agent  than  the  hydrogen  ion,  is  a  much  weaker  oxidizing  agent  than 
aquated  cobalt  (III)  ion,  but  a  more  powerful  oxidizing  agent  than  the 
hexacyanocobaltate(III)  ion.  Stabilization  of  cobalt(III)  against  reduction 
to  cobalt(ll)  is  favored  by  coordination  with  cyanide  ion  as  compared  with 
ammonia  and  water. 

Half -Wave  Potentials — Polarography 

Ease  of  reduction  or  oxidation  of  a  complex  ion  at  the  dropping  mercury 
electrode  is  different  from  that  of  the  aquated  metal  ion,  and  half-wave 
potentials  obtained  under  such  conditions  that  the  reactions  are  reversible 
have  the  great  advantage  of  thermodynamic  significance  and  may  be  re- 
lated to  ordinary  standard  potentials.*  If  the  reduction  of  the  complex 
proceeds  reversibly,  the  values  of  dissociation  constants  of  the  complex 
and  the  number  of  coordinated  groups  can  be  calculated16  from  the  change 
in  half -wave  potential.  Irreversibility  of  a  process  can  easily  be  determined 
by  this  method,  and  many  processes  reported  in  the  literature  as  reversible 
by  classic  methods  have  been  found  to  be  irreversible  at  the  dropping 
mercury  electrode.  Application  of  the  polarographic  technique  has  brought 
forth  many  examples  of  stabilization  of  oxidation  states  through  coordina- 
tion. 

*  The  electropositive  metals  exhibit  high  energies  of  formation  when  they  proceed 
from  the  pure  metal  to  the  amalgam,  and,  consequently,  the  half -wave  potential  is 
more  positive  than  the  standard  potential.  The  less  electropositive  metals  that  read- 
ily form  amalgams,  zinc,  lead,  cadmium,  bismuth,  thallium,  and  silver,  have  re- 
versible amalgam  electrodes,  and  in  certain  instances  the  half-wave  potentials  of  these 
metal  ions  may  be  nearly  equal  to  the  standard  oxidation  potentials. 

15.  Bigelow,  Inorganic  Syntheses,  2,  225  (1946). 

16.  Kolthoff  and  Lingane,  "Polarography,"  1st  ed.,  p.  164,  New  York,  Interscience 

Publishers,  Inc.,  1941. 


STABILIZATION*   OF  VALENCE  STATES  403 

Polarography  of  Copper  Complexes.  Equated  copper(II)  ions  are 

reduced  directly  to  the  amalgam  at  the  dropping  mercury  electrode,  and 
only  a  single  polarographic  wave  can  l>c  obtained  in  the  absence  of  complex- 
ing agents.  The  potential  of  the  CuH  «=*  Cu(Hg)  system  is  more  positive 

than  that  of  the  CuM  *=*  Cu(IIg)  system,  and  COpper(I)  ions  cannot  exist 

at  the  potential  at  which  copper(II)  ions  are  reduced. 

The  stability  and  composition  of  the  complex  ions  formed  by  copper(Il) 
ions  (5  X  10-4  molar)  with  glycinate  and  alaninate  ions  were  determined 
by  Keefer17.  The  complexes  formed  are  mainly  [Cu(gly)o]  or  [Cu(alan)2] 
when  the  concentration  of  the  complexing  agent  is  from  0.08  to  0.1  molar, 
and  the  stable  glycinate  complex  is  [Cu(gly)3]~  at  higher  concentrations. 
Under  the  conditions  of  pH  and  concentration  studied,  two  electrons  are 
involved  in  the  electrode  reduction,  indicating  tne  instability  of  the  cop- 
per(I)  glycinate  or  alaninate  complexes.  Two-electron  reductions  were  also 
observed  by  Onstott18  for  the  bis(ethylenediamine),  bis(propylenediamine), 
and  bis(diethylenetriamine)  complexes  of  copper (II). 

Table  11.5.  Potentials  for  the  Polarographic  Reduction  of  Copper 

Am. mines 

Equation 

[Cu(NH,),]+  +  Hg  +  e-  «=>  Cu(Hg)  +  2NH3 
[CuCNH,)*]-"-  +  Hg  +  2e~  <=t  Cu(Hg)  +  4NH, 

[Cu(NH3)4l++  +  e~  +±  [Cu(NH,)2]+  +  2NH3 

Certain  complexing  agents  that  form  stable  copper(I)  complexes  shift 
the  half -wave  potential  of  the  Cu+  «=*  Cu(Hg)  system  in  the  negative  di- 
rection more  than  that  of  the  Cu~H~  «=*  Cu+  system,  so  two  distinct  polaro- 
graphic waves  result.  Table  11.5  lists  potential  values  for  the  ammines  of 
copper19.  Two  waves  result  when  copper(II)  ion  is  reduced  in  ammoniacal 
solution.  Thiocyanate,  chloride,  and  pyridine  complexes  behave  simi- 
larly19- 20. 

Iron  Oxalato  Complexes.  In  the  presence  of  oxalate  ions,  the  half- 
wave  potential  of  the  aquated  iron(III)  ion  shifts  to  a  more  negative  value 
because  of  the  formation  of  [Fe(C204)3]-  19a.  Consideration  of  the  half-wave 
potential  of  the  tris(oxalato)ferrate(III)  ion  as  a  function  of  oxalate  ion  con- 
cent ration  revealed  that  the  formula  of  the  iron(II)  complex  produced  in  a 

17.  Keefer,  ./.  Am.  Chem.  Soc,  68,  2329  (1946). 

18.  Onstott,   thesis,  University  of  Illinois,   1948;   Laitinen,   Onstott,    Bailar,   and 

-     ,iin.  ./.  Am.  Chem.  Soc,  71,  1550  (1949). 

19.  Stackelberg  and  Freyhold,  Z.  Elektrochem.,4&,  120  (1940);  Lingane,  Chem.  Revs., 

29,  1    1941   :  Bchaap,  Laitinen  and  Bailar,  ./.  .1//'.  Chem.  80c.,  16,  5868    1954). 

20.  Lingane  and  Iverlinger,  Ind.  Eng.  Chem.,  Anal.  Ed.,  13,  77  (1941 >;  Korshunov  and 

Malvugina.  ./.  Gen.  Chem.,  U.S.S.R.,  20,  425  (1950). 


vs.  N.C.E 

-0.522 

-0.397 

-0.262 

404  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

0.001  to  0.002  molar  solution  of  iron(III)  ion  in  the  presence  of  0.15  molar 
oxalate  ion  concentration  is  [Fe(C204)2]=,  but  when  the  concentration  of 
oxalate  ion  is  in  greater  excess,  the  species  is  the  complex  [Fe^O^]4-  19b. 
These  results  were  essentially  confirmed  by  Schaap19c. 

Tin  Complexes.  Although  the  standard  potential  of  the  tin(II)-tin(IV) 
couple  is  about  0.15  volts,  the  tin (IV)  ion  is  not  reduced  at  the  dropping 
mercury  electrode21.  Solutions  of  tin(IV)  ion  in  1  to  2  molar  perchloric  acid 
solution  give  no  indication  of  a  reduction  wave  before  the  discharge  of 
hydrogen.  The  predominant  species  in  solution  is  believed  to  be  the  hexa- 
quotin(IV)  ion,  and  apparently  the  failure  of  this  ion  to  be  reduced  can 
be  attributed  to  its  slow  rate  of  reduction.  Furthermore,  no  reduction  of 
tin  (IV)  ion  at  the  dropping  mercury  electrode  takes  place  in  sodium  hy- 
droxide, tartrate,  or  acidic  oxalate  media22.  Either  the  complexes  formed 
are  too  stable  to  be  reduced,  or  they  are  reduced  at  such  slow  rates  that  no 
appreciable  reduction  can  take  place  during  the  short  life  of  each  mercury 
drop. 

The  hexachlorostannate(IV)  ion  is  reduced,  however,  when  the  chloride 
ion  concentration  is  greater  than  4  molar.  The  two  well-defined  waves  which 
result  are  attributed  to  the  reduction  of  the  hexachlorostannate(IV)  ion  to 
the  tetrachlorostannate(II)  ion,  followed  by  the  reduction  of  the  latter 
complex  to  the  metal.  A  fairly  well-defined  doublet  wave  is  also  obtained 
in  the  reduction  of  the  hexabromostannate(IV)  ion  in  the  presence  of  a 
large  excess  of  bromide  ion21.  In  these  cases  the  activation  energy  has  been 
greatly  diminished  by  converting  the  hexaquotin(IV)  complex  to  the  chloro- 
or  bromostannate(IV)  complex. 

Antimony  Complexes.  Pentapositive  antimony  is  a  fairly  strong  but 
slow  oxidant.  The  failure  of  the  reduction  of  antimony (V)  in  perchloric  acid 
or  dilute  hydrochloric  acid  media  indicates  a  situation  analogous  to  that 
encountered  with  tin.  In  solutions  containing  large  concentrations  of 
chloride  ion,  antimony (V)  shows  reduction  first  to  the  tripositive  state  and 
then  to  the  amalgam23.  The  failure  of  the  reduction  to  take  place  in  the 
presence  of  a  small  concentration  of  chloride  is  attributed  to  the  presence  of 
ions  of  the  type  [Sb02Cl2]_  and  [SbOClJ-.  Presumably,  these  species  are 
converted  to  the  hexachlorostibnate(V)  ion  as  the  chloride  ion  concentration 
is  increased. 

Uranium  (V).  Kolthoff  and  Harris  have  studied  the  polarographic  be- 
havior of  uranium (VI)  in  acidic24  and  basic25  solutions.  In  moderately  con- 

21 .  Lingane,  ./.  Am.  Chem.  Soc,  67,  919  (1945). 

22.  Lingane,  Ind.  Eng.  Chem.,  Anal.  Ed.,  15,  583  (1943). 

23.  Lingane  and  Nishida, ./.  Am.  Chem.  Soc.,  69,  530  (1947). 

24.  Harris  and  Kolthoff,  ./.  Am.  Chem.  Soc,  67,  1484  (1945);  Kolthoff  and  Harris, 

J .  An,.  Chem.  Soc.,  68,  1175  (1946). 

25.  Harris  and  Kolthoff,  ./.  Am.  Chem.  Soc,  69,  446  (1947). 


STABILIZATION  OF  VALENCE  STATES  405 

centrated  acid  (0.01  toO.'J.U  HC1)  iir:mium(VI)  oxychloride  gives  two  well- 
defined  reduction  waves,  the  first  being  one-half  the  height  of  the  second. 
Consideration  of  current-voltage  data  revealed  the  first  to  correspond  to  B 
reversible  reaction.  Since  the  half-wave  potential  of  this  wave  did  not  shift 
with  changing  hydrogen  ion  concentration,  the  following  one-electron  re- 
duction was  postulated. 

[U02]++  +  e~  <=±  [U02]+ 

The  second  wave,  a  two-electron  irreversible  reduction,  corresponds  to  the 
reduction  of  pentapositive  uranium  to  the  tripositive  state. 

Complexes  of  Cadmium — Successive  Formation  Constants.  The 
chloro-,  bromo-,  and  iodocadminm  complexes  were  investigated  polaro- 
graphically  by  Strocchi26.  Jt  was  reported  that  such  species  as  CdX+, 
CdXj  ,  (MX.  .  and  (\L\4=  exist  in  solution,  the  species  present  depending 
upon  the  relative  concentrations  of  the  ions,  and  all  are  reduced  to  the  amal- 
gam reversibly.  If  only  one  complex  species  exists  over  a  considerable  range 
of  concentration  of  complexing  agent,  and  if  this  species  is  reduced  reversibly, 
the  formula  and  dissociation  constant  may  be  calculated  according  to  the 
method  described  by  Kolthoff  and  Lingane16.  However,  the  method  has 
not  been  applied  to  systems  involving  mixtures  of  consecutively  formed 
complex  ions.  Bjerrum27  and  Leden28  have  developed  methods  for  determin- 
ing successive  formation  constants,  and  subsequently  De  Ford  and  Hume29 
have  described  a  mathematical  treatment  of  half-wave  potential  data  which 
makes  possible  the  identification  of  successively  formed  complex  species 
and  the  calculation  of  their  dissociation  constants.  These  investigators 
successfully  applied  this  mathematical  analysis  to  the  study  of  the  com- 
plexes of  cadmium,  CdSCN+  Cd(SCN)2 ,  Cd(SCN)r,  and  Cd(SCN)4=; 
the  calculated  formation  constants  are  11,  56,  6,  and  60,  respectively30. 

Vanadium  Complexes.  The  polarographic  characteristics  of  vanadium 
in  noncomplexing  media  have  been  studied  by  Lingane31.  In  both  acid  and 
ammoniacal  solution,  vanadium (V)  undergoes  stepwise  reduction,  first 
to  the  tetrapositive  state,  and  then  to  the  dipositive  state.  Evidence  was 
presented  for  the  existence  of  complexes  in  which  vanadium  displays  valence 
states  of  2+,  3+,  4+,  and  5+  in  the  presence  of  some  other  complexing  agents32. 
The  formation  of  complexes  is  greatly  influenced  by  the  presence  of  hy- 

26.  Strocchi,  Gazz.  chim.  ital.,  80,  234  (1950). 

lijerrum,    "Metal   Ammine  Formal  ion   in   Aqueous  Solution,"  Copenhagen,   P. 
Maaae  and  Son,  1941 . 

28.  Leden,  Z.  physik.  Chem.,  188A,  160  (1941   . 

29.  DeFord  and  Hume,  J.Am.  Chi  m.  Soc.,  73,  5321  (1951). 

30.  Hume,  DeFord,  and  Cave,  •/.  .1///.  Chi  m.  Soc.,  73,  5323    1951). 

31.  Lingane, ./.  Am.  Chem.  Soc,  67,  182  (1945 

32.  Lingane  and  Meites,  J.  Am.  Chem.  Soc,  69,  1021     1947 


400  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

droxyl  groups  in  the  complexing  agent,  for  vanadium  tends  to  coordinate 
preferentially  wit h  oxygen.  The  tartrate  ion  with  its  two  adjacent  hydroxyl 
groups  forms  more  stable  complexes  than  does  the  citrate  ion,  which  con- 
tains only  one  hydroxyl  group.  In  alkaline  solution  the  hydrogen  of  the 
hydroxyl  group  is  replaced  by  an  equivalent  of  the  coordinating  metal 
ion.  Half -wave  potentials  show  that  the  oxalate  ion,  which  contains  no 
hydroxyl  group,  forms  the  least  stable  complex  of  the  series33. 

The  Significance  of  Standard  Potential  Values  for  Irreversible  Sys- 
tems 

It  has  been  pointed  out  that  oxidation  potentials  become  altered  when 
the  activity  quotient  term  of  the  Nernst  equation  is  varied.  This  may  arise 
when  the  equilibrium  conditions  of  a  system  become  changed  through 
complex  formation.  Many  oxidation  potentials  cannot  be  measured  directly 
and  must  be  calculated  from  thermal  data,  or  estimated,  for  the  Nernst 
equation  applies  without  reservation  only  to  reversible  systems.  Conse- 
quently, the  significance  of  the  standard  potential,  E°,  is  limited  in  some 
cases. 

On  the  basis  of  isotopic  exchange  studies,  Taube34  has  observed  that  ex- 
change between  an  oxidized  form  and  a  reduced  form  of  the  same  complex, 
one  of  which  contains  a  radioactive  central  atom,  proceeds  most  easily 
when  the  electronic  bonding  orbitals  of  the  two  forms  are  identical.  Such 
exchange  could  proceed  by  the  electron  transfer  mechanism.  For  example, 
an  exchange  of  electrons  between  [Fe(CX)6]4_  and  [Fe(CN)6]3_  ions  in 
neutral  solution  and  in  0.05  molar  sodium  hydroxide  was  observed  to  take 
place  within  one  minute35.  Each  of  these  ions  has  the  d2sps  octahedral  con- 
figuration. Some  investigators36  have  suggested  that  these  conditions  also 
favor  electrode  reversibility.  Conversely,  where  a  difference  in  electronic 
bonding  orbitals  exists  between  the  oxidized  form  and  the  reduced  form 
of  a  particular  complex,  slowness  or  lack  of  exchange  is  observed  in  most 
cases,  and  it  is  believed  that  electrode  irreversibility  should  also  exist.  In 
many  instances  the  interrelationship  between  the  ligand  and  the  central 
ion  imposes  a  new  electronic  configuration  upon  either  the  oxidized  or  the 
reduced  form  of  a  complex,  and  oxidation  states  may  be  stabilized  to  a 

33.  Lingane  and  Meites, ./.  Am.  Chem.  Soc,  69,  1882  (1947). 

34.  Taube,  Chew.  Revs.,  50,  69  (1952). 

35.  Thompson.  ./.  .1///.  Chem.  Soc,  70,  1045  (1948). 

36.  Lyons,./.  Electrochem. Soc,  101,363, 376  (1954) ;  Lyons,  Bailar  and  Laitinen, ibid., 

101,  410  (1954).  Libby,  "Theory  of  Electron  Exchange  Reactions  in  Aqueous 
Solutions,"  p.  39,  preprint,  Symposium  on  Electron  Transfer  and  Isotopic 
Reactions,  Division  of  Physical  and  Inorganic  Chemistry,  American  Chemical 
rciety,  and  Division  of  Chemical  Physics,  American  Physical  Society,  Notre 
Dame,  June  11-13,  1952. 


STABILIZATION  OF  VALENCE  STATES  407 

marked  extent.  When  this  happens,  the  bond  between  the  central  atom  and 
the  ligand  of  the  stabilized  form  seems  to  lose  all  lability,  and  exchange 
studies  indicate  that  an  equilibrium  no  longer  exists  between  the  complex 
and  its  constituents. 

Stabilization  of  Unusual  Oxidation  States  Through 
Coordination 

An  interesting  and  important  aspect  of  stabilization  through  coordina- 
tion is  the  stabilization  of  unusual  valence  states.  The  methods  for  charac- 
terizing unusual  oxidation  states  include  the  use  of  analytical  data,  chem- 
ical properties,  magnetic  susceptibility  measurements,  and  x-ray  studies37. 

Copper(I)  and  Copper(III) 

The  unipositive  state  of  copper  is  stabilized  by  coordination  with  thiourea 
to  such  an  extent  that  the  copper(I)  complex  is  formed  even  when  cop- 
per (I  I)  ion  is  used  as  a  reactant38.  Similarly,  ethylenethiourea  reacts  with 
copper(II)  ion  to  form  the  stable  copper(I)  complex, 

[Cu(ethylenethiourea)4]N0339. 

(\>pper(I)  complexes  with  the  cyanide  ion  are  among  the  most  stable  cya- 
nides, and  hydrogen  sulfide  fails  to  precipitate  any  sulfide  of  copper  when 
added  to  solutions  of  potassium  tetracyanocuprate(I).  In  most  of  these 
complexes  the  copper  achieves  the  coordination  number  of  four.  Some 
alkyl-substituted  phosphines  and  arsines  combine  with  equimolar  quantities 
of  copper(I)40,  but  these  complexes  are  polymeric. 

The  complex  K3[CuF6],  prepared  by  allowing  a  mixture  of  potassium 
chloride  and  copper(II)  chloride  to  react  with  fluorine  at  250° 41,  is  decom- 
posed by  water.  More  stable  copper(III)  complexes  have  been  prepared 
by  the  peroxysulfate  oxidation  of  copper(II)  with  the  periodate  and 
tcllurate  complexing  groups42.  Some  interesting  analytical  applications  of 
copper(III)  complexes  are  described  by  Kleinberg43. 

37.  Kleinberg,    "Unfamiliar   Oxidation   States,"   Lawrence,   University   of    Kansas 

Press,  1950;  Kleinberg,  J.  Chen,.  Ed.,  29,  324  (1952);  Mellor,  "Some  Recent 
Developments  in  the  Chemistry  of  Metal  Complexes."  Report  of  the  Bris- 
bane Meeting  of  tlu'  Australian  and  New  Zealand  Association  for  the  Ad- 
vancement of  Science,  Vol.  XXVIII,  131,  (1951). 

38.  Rosenheim  and  Loewenstamm,  Z.  anorg.  Chem.,  34,  62    1903 

39.  Morgan  and  Burstall,  •/.  Ckem.  Soc.,  1928,  143. 

in.  Mann.  Purdie,  :m(l  Wells,  ./ .  Chem.  Soc.,  1926,  2018;  Kabesh  and  Nyholm,  ./. 

1951,  38. 
U.  rHemm  and  Hubs,  Z.  anorg.  Chem.,  258,  221  (1949  . 

42.  Mai:.-  n.  itnl..  71,  If,;.  580  [1941 

43.  Kleinberg,  J.Cht  29,  326  (1952). 


408 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Silver(II)  and  Silver(III) 

The  existence  of  higher  oxidation  states  of  silver  is  well  established44. 
Silver  has  been  found  to  be  dipositive  in  the  complex  formed  with  quinolinic 
acid45.  A.  A.  Noyes  and  co-workers  established  the  presence  of  silver(II) 
and  silver(III)  in  oxidizing  solutions.  When  ozone  was  passed  into  a  solu- 
tion of  silver(I)  nitrate  in  nitric  acid,  it  was  shown  that  the  metal  oxidized 
to  a  nitrato  silver(II)  complex.  This  conclusion,  drawn  from  a  considera- 
tion of  color,  oxidizing  potential  of  the  solution,  and  increased  solubility  of 
the  compound  in  solutions  with  higher  nitrate  concentration45a,  agrees  with 
that  given  by  Weber46  on  the  basis  of  transference  experiments. 

A  cooled  aqueous  solution  of  silver  sulfate  and  ethylenedibiguanide  reacts 
with  potassium  peroxy sulfate  to  form  a  silky,  red,  crystalline  precipitate 
of  a  silver(III)  salt.  It  is  stable  at  ordinary  temperatures  and  can  be  re- 
crystallized  from  warm,  dilute  nitric  acid.  The  tripositive  silver  ion  in  this 
diamagnetic  complex  has  the  same  electronic  configuration  as  the  nickel (II) 
ion47.  A  solution  of  the  complex,  acidified  in  the  presence  of  potassium 
iodide,  liberates  two  equivalents  of  iodine  for  every  atom  of  silver,  and  the 
molar  conductivity  of  the  nitrate  indicates  the  presence  of  a  tripositive 
complex  cation48.  The  constitution  of  this  cation  is  represented  by 

NH 


CH2— NH— C— NH— C— NH 


NH 

II 
CH2— NH— C— NH— C— NH2 

II 
NH 

and  the  quadridentate  nature  of  the  ligand  explains  the  stability  of  the  tri- 
positive state  of  silver.  The  pK  values  for  the  dissociation  of  the  complex 
and  for  the  displacement  of  the  silver(III)  ion  by  hydrogen  ion  are  52  and 
29,  respectively48b. 

McClelland49  has  found  that  pyridine  forms  two  complex  ions  with  sil- 

44.  Bailar,  J.  Chem.  Ed.,  21,  523  (1944). 

45.  Berbieri,  Atti.  Acad.  Lincei,  17,  1078  (1933);  Noyes,  DeVault,  Coryell,  and  Deahl, 

J.  Am.  Chem.  Soc,  59,  1326  (1937). 

46.  Weber,  Trans.  Am.  Electrochem.  Soc,  32,  391  (1917). 

17.  Manchot  and  Gall,  Ber.,  60,  191  (1927). 

18.  Ray  and  Chakravarty,  ./.  Indian  Chem.  Soc.,  21,47  (1944);  Sen,  Ray,  and  Ghose, 

ibid.,  27,619  (1950). 
49.  McClelland,  thesis,  University  of  Illinois,  1950. 


STABILIZATION  OF  VALENCE  STATES  409 

ver(II),tris(pyridine)silver(II)  ion  and  tetrakis(pyridine)silver(]  1 1  ion.  Bis- 
(dipyridyl)silver(II)  Ls  formed  by  oxidizing  silver(I)  with  eerie  ammonium 

nitrate  in  nitric  acid,  and  its  dissociation  constant  is  2.5  X  10~19.  The  stand- 
ard potential  oi  the  dipyridyl  complexes  of  silver(I)  and  silver(II)  is  0.814 
volts  vs.  the  hydrogen  electrode  at  25°. 

Manganese  (I) 

Manganese  in  the  unipositive  state  was  reported  to  have  been  prepared 
by  the  reduction  of  the  cyano  complex  of  divalent  manganese  with  granu- 
lated aluminum17  and  by  electrolytic  reduction50.  The  crystalline  product, 
KolMmCN"^],  was  said  to  be  a  powerful  reducing  agent.  Klemm51  ques- 
tioned the  identity  of  this  compound  because  it  was  found  to  be  para- 
magnetic, whereas  the  formula  indicates  it  should  be  diamagnetic.  However, 
Tread  well  and  Raths52  have  prepared  the  compound  electrolytically  and 
report  it  to  be  diamagnetic.  Christensen,  Kleinberg,  and  Davidson53  have 
obtained  excellent  evidence  for  manganese  in  the  zero  and  unipositive  oxi- 
dation states  by  treatment  of  a  liquid  ammonia  solution  of  potassium 
hexacyanomanganate(III)  with  potassium  metal.  The  yellow  product  so 
obtained  has  the  composition  K5Mn(CN)6-K6Mn(CN)6-2NH3 .  Their 
conclusions  are  based  on  studies  of  reacting  ratios,  chemical  analysis,  re- 
ducing power,  and  magnetic  measurements  (the  effective  magnetic  moment 
is  1.25  Bohr  magnetons  as  compared  to  a  calculated  value  of  1.73  for  a  sin- 
gle unpaired   electron). 

Nickel(O),  Nickel(I),  and  Nickel(IV) 

In  a  study  of  the  reduction  of  nickel  salts  in  anhydrous  liquid  ammonia, 
Eastes  and  Burgess54  isolated  a  unipositive  nickel  compound  K2[Ni(CN)3]. 
The  reaction  of  this  compound  with  an  excess  of  the  alkali  metal  produces 
K4[Xi(CX)4],  in  which  nickel  has  an  apparent  valence  state  of  zero.  The 
negative  radical  [Xi(CX)4]4~  is  isoelectronic  with  nickel  carbonyl,  and  based 
upon  the  electronic  configuration  of  the  latter  molecule  as  postulated  by 
Pauling55,  an  explanation  of  the  zero  valence  of  nickel  is  offered  by  Deasy55. 

Many  complexes  of  nickel(IV)  have  been  reported.  Klemm57  reports  the 
fluoro  complex  K2[XiF6],  and  the  tetrapositive  state  of  nickel  is  confirmed 

50.  Grube  and  Brause,  Ber.,  60,  2273  (1927). 

51.  Klemm,  Angew.  CTiem.,  63,  396  (1951). 

52.  Treadwell  and  Raths,  Heir.  chim.  Acta,  35,  2259  (1952);  ibid,  35,  2275  (1952). 

53.  Christensen,  Kleinberg,  and  Davidson,  J.  Am.  Chem.  Soc,  75,  2495  - 1953). 
.54.  Eastes  and  Burgess,  J.  Am.  Chem.  Soc.,  64,  1187  (1942). 

55.  Pauling,  "The  Nature  of  the  Chemical  Bond,"  p.  252,  Ithaca,  Cornell  University 

Press,  1944. 

56.  Deasy,  ./.  Am.  Chem.  Soc.,  67,  152    1945 

57.  Klemm  and  Huss,  Z.  anonj.  Cfu  m.t  25,  221  (1949). 


410 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


by  magnetic  evidence.  Hieber  and  Bruck58  describe  nickel (IV)  complexes 
of  the  types : 


'  ./V^1 


Cobalt(I),  Cobalt(III),  and  Cobalt(IV) 

A  number  of  stable  polynuclear  compounds  were  prepared  by  Werner59 
in  which  the  peroxide  ion  02=  functions  as  a  bridging  group,  and  the  analy- 
ses indicated  the  presence  of  both  tripositive  and  tetrapositive  cobalt.  The 
compound 

NH2 


(NH3)4Co 


Co(NH3)4 


X4 


was  among  those  prepared,  and  the  presence  of  both  cobalt  (III)  and  co- 
balt (IV)  is  supported  by  chemical  and  physical  evidence.  These  /x-peroxo 
type  compounds  are  decomposed  by  heating  with  sulfuric  acid  to  produce 
mononuclear  ammines  with  the  liberation  of  oxygen.  The  presence  of 
tetrapositive  cobalt  is  supported  by  titration  with  arsenite60  and  by  mag- 
netic susceptibility  measurements60- 61. 

When  aqueous  solutions  of  potassium  hexacyanocobaltate(III)  are  re- 
duced electrolytically,  a  deep  brown  solution  of  a  unipositive  cobalt  com- 
plex results62.  The  existence  of  cobalt  (I)  was  confirmed  polarographically 
by  Hume  and  Kolthoff63.  According  to  Malatesta64,  most  cobalt (II)  salts 
react  with  aromatic  isonitriles  in  alcoholic  solution,  undergoing  reduction 
and  forming  complex  salts  of  cobalt(I)  with  the  formula  [Co(CNR)5]X. 
The  salts  in  which  X-  is  perchlorate,  chlorate,  iodide,  and  nitrate  were 

58.  Hieber  and  Bruck,  Naturwissenschaften,  36,  312  (1949). 

59.  Werner,  Ann.,  375,  1  (1910). 

60.  Gleu  and  Rehm,  Z.  anorg.  allgem.  Chem.,  237,  79  (1938). 

61.  Malatesta,  Gazz.  chim.  ital,  72,  287  (1942). 

62.  Grube,  Z.  Elektrochem.,  32,  561  (1926). 

63.  Hume  and  Kolthoff,  J.  Am.  Chem.  Soc,  71,  867  (1949). 

64.  Malatesta,  Angew.  Chem.,  65,  266  (1953). 


STABILIZATION  OF  VALENCE  STATES  411 

isolated  and  found  to  be  yellow  or  brown  crystalline  solids.  They  are  soluble 
in  polar  solvents  and  arc  reported  to  be  diamagnetic  and  of  unlimited  sta- 
bility in  air.  The  preparation  of  some  of  these  salts  requires  the  presence  of 
mild  reducing  agents,  while  others  form  merely  upon  warming  an  alcoholic 
solution  of  the  constituents. 

Platinum(ni),  Platinum(V),  Platinum(VI),  and  Platiiiuiii(VIII) 

A  number  of  compounds  formed  by  the  reaction  of  chloroplatinie  acid 
with  various  thio  compounds,  such  as  sulfides,  mercaptans,  and  disulfide-, 
in  which  the  platinum  exhibits  the  unusual  valence  states  of  three,  five,  six, 
and  eight  have  been  described  by  Ray  and  his  co-workers65.  The  evidence 
for  the  variations  in  the  valency  of  platinum  was  obtained  by  the  reaction 
of  platinum(IY)  chloride  and  the  organic  ligand  given  by  the  following 
equation"*. 

x  (HSC0H4SK)  +  PtCl4  -»  [Pt(S  C2H4  SH),]  x  -  3,  4,  5,  6,  or  8 

Molecular  weight  determinations650  and  chemical  reactions65d  were  of  much 
value  in  elucidating  the  constitution  of  the  platinum  complexes.  Some  of 
the  compounds  do  not  correspond  to  the  empirical  formulas  but  are  poly- 
mers. The  unusual  valence  states  of  platinum  are  explained  by  the  great 
coordinating  power  of  the  sulfur  atom  in  the  organic  ligand,  and  the  particu- 
lar valence  state  that  platinum  assumes  is  a  function  of  the  two  variables, 
concentration  and  temperature.  At  low  temperatures  platinum  exhibits  its 
maximum  valency,  and  at  approximately  100°  only  trivalent  platinum  com- 
pounds are  obtained.  The  relative  ease  with  which  the  ligands  are  liberated 
might  indicate  that  some  of  the  organic  groups  are  not  truly  bound  to  the 
platinum,  and  all  of  the  valences  mentioned  above  may  not  exist. 

Chromium (II),  Chromium (IV),  and  Chromium(V) 

Chromium  is  stabilized  in  the  dipositive,  tetrapositive,  and  pentapositive 
oxidation  states.  Some  chromium(II)  complexes  most  stable  toward  oxi- 
dation contain  hydrazine  as  a  complexing  agent66.  The  reducing  properties 
of  hydrazine  account  in  part  for  this  stability.  The  dihydrazine  complexes 
of  the  chloride,  bromide,  and  iodide  of  dipositive  chromium  have  been  pre- 
pared. ( 'hromium(II)  complexes  of  a,a'-dipyridyl,  hexamethylenetetra- 
mine.  0-phenanthroline,  and  8-hydroxyquinoline  have  also  been  reported67. 

66.  Raj  and  Ghoee, ./.  Indian  Chi  m.  Soc.}  11,  737  (1034);  Ray,/.  Chem.  80c.,  123,  133 

Soc.,  2,  178    1926  :  I:  :■  .  Bose  Raj  .  and 
I:  :     <  haudhury,  ./.  Indian  Chem.  80c.,  5,  139  (1928). 
_••,  Ber.t  46,  L505    L913). 

67.  Beriberi  and  Tettamanzi,  Atti.  Acad.  Lincei,  15,  ^77    L932);  Hammett,  Walden, 

and  Edmonds,  ■/.    I  Joe.,  56,  1002  'VX-W);  Hume  and  Stone,  •/.  .1///. 

Chem.  Soc,  63,  1200  (1941). 


412  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

A  tetrapositive  chromium  compound  was  reported  by  Klemm  and  Huss68; 
the  complex  K2[CrF6]  is  formed  when  a  mixture  of  potassium  chloride  and 
chromium(III)  chloride  is  fluorinated. 

Chromium  (V)  was  first  reported  by  Weinland69,  who  succeeded  in  isolat- 
ing the  complexes  K2[CrOCl5]  and  (pyH)  [CrOClJ. 

Some  Factors  Which  Contribute  Toward  Stabilization  of 
Oxidation  States  Through  Coordination 

The  factors  contributing  to  the  stabilization  of  valence  are  numerous 
and  interdependent.  Some  conclusions,  however,  can  be  drawn  from  con- 
sideration of  the  nature  of  the  coordinating  group,  the  central  metal  ion, 
and  the  bond  between  them.  Douglas70  has  reviewed  several  contributing 
factors,  and  his  criteria  are  included  in  these  considerations. 

Nature  of  the  Coordinating  Group 

Reducing  Tendencies.  Complex  compounds  formed  by  metallic  ions 
with  unsaturated  compounds,  such  as  the  metal-olefin  complexes,  tend  to 
stabilize  the  lower  valence  states  of  the  central  metal  ion.  Stable  compounds 
have  been  prepared  by  the  reaction  of  potassium  tetrachloroplatinate(II) 
with  unsaturated  alcohols,  acids,  aldehydes,  and  ketones71  (Chapter  15). 
The  extremely  stable  dihydrazine  complexes  of  chromium(II)  are  accounted 
for  by  the  reducing  character  of  the  complexing  agent. 

Steric  Factors.  a,a:'-Dipyridyl  reacts  with  iron (II)  to  form  the  stable, 
intensely  colored  complex,  [Fe(Ci0H8N2)3]++,  but  the  introduction  of  certain 
substituents  into  the  ring  produces  a  marked  decrease  in  the  coordinating 
ability  of  the  base.  This  shielding  effect  is  shown  by  the  failure  of  a-(a'- 
pyridyl)-quinoline  to  complex  with  iron (II)72.  Large  groups  often  prevent 
an  ion  from  exhibiting  its  maximum  coordination  number,  and  forced  con- 
figurations may  result.  Mann  and  Pope73  investigated  complexes  of  nickel 
(II),  palladium (II),  and  platinum(II)  with  tris(2-aminoethyl) amine  and 
established  the  formula  [Mtren]++.  Such  an  ion  must  be  an  irregular  tetra- 
hedron. 

The  steric  effects  associated  with  the  replacement  of  hydrogen  atoms  of  a 
coordinated  amine  by  alkyl  groups  have  been  studied  by  Basolo  and 
Murmann74.  With  the  groups,  methyl,  ethyl,  and  n-propyl,  the  stabilities 
of  the  complexes  formed  by  N-alkylethylenediamine  with  copper(II)  and 

68.  Huss  and  Klemm,  Z.  anorg.  allgem.  Chem.,  262,  25  (1950). 

69.  Weinland  and  Mitarb,  Ber.,  38,  3784  (1905). 

70.  Douglas,  /.  Chem.  Ed.,  29,  119  (1952). 

71.  Pfeiffer  and  Hoyer,  Z.  anorg.  allgem.  Chem.,  211,  241  (1933). 

72.  Smirnoff,  Helv.  chim.  Acta,  4,  802  (1921). 

73.  Mann  and  Pope,  /.  Chem.  Soc.,  1926,  482. 

74.  Basolo  and  Murmann,  J.  Am.  Chem.  Soc.,  74,  5243  (1952). 


STABILIZATION  OF  VALENCE  STATES  413 

Table  11.6.  The  Effbct  of  Chelation  on  ran  Stability  of  Cad  mm  m  Ammi  n 

Complex  Dissociation  Constant 

[CcKNHs)*]-^  3.3  X  10-7 

[CdCen),]-""  6.7  X  10~13 

[Cd(pn)3]-H-  5.4  X  10"13 

[Cd(dien)2]++  7.6  X  10~16 

nickel(II)  decrease  as  the  size  of  the  alkyl  group  increases.  The  n-butyl 
derivative  is  more  stable  than  anticipated;  this  might  arise  from  a  shielding 
effect  as  a  result  of  entwining  of  the  butyl  group  about  the  metal  ion.  As 
might  be  expected,  N-iso-propylethylenediamine  forms  complex  ions  of 
lesser  stability  than  those  formed  by  N-normal-propylethylenediamine. 
Steric  effects  are  greater  with  hexacovalent  than  with  tetracovalent 
nickel(II). 

Chelation.  Some  complexing  agents  have  a  greater  tendency  to  occupy 
two  coordination  positions  than  one.  These  so-called  chelate  groups  form 
complexes  of  enhanced  stability  (Chapter  5),  the  most  stable  complexes 
resulting  from  the  formation  of  five  and  six-membered  rings.  The  effect  of 
chelation  is  illustrated  in  Table  11.6  by  the  comparison  of  the  dissociation 
constants  of  cadmium  chelate  complexes  with  that  of  the  ammine  complex. 

The  most  probable  explanation  for  the  increased  stability  is  the  simple 
one  that  if  one  of  the  two  coordinating  linkages  is  broken,  the  other  can 
keep  the  coordinating  group  near  the  central  ion  until  the  broken  bond  is 
reformed.  This  explanation  is  supported  by  experiments  using  radioactive 
"tracers."  The  study  of  the  racemization  of  optically-active  tris(oxalato)- 
chromate(III)  ion  revealed  that  the  mechanism  of  the  transformation  does 
not  involve  an  ionization  of  oxalate  groups75.  A  suggested  mechanism  in- 
volves an  intramolecular  rearrangement  (Chapter  8). 

Nature  of  the  Central  Metal  Ion 

Electronegativity.  Coordination,  in  general,  is  favored  by  a  small  ion 
of  high  charge.  Preferential  coordination  of  a  metal  ion  with  a  given  ele- 
ment is  a  function  of  the  electronegativity  of  the  metal  ion.  Thus,  alumi- 
num, beryllium,  and  zinc  coordinate  tightly  with  oxygen  in  a  ligand;  zinc, 
chromium,  cadmium,  cobalt,  and  nickel  coordinate  preferentially  with 
nitrogen-containing  ligands;  and  tin,  lead,  antimony,  silver,  mercury,  and 
the  platinum  metals  prefer  either  halogen  or  sulfur-containing  ligands 
( Chapter  1). 

Coordination  Number.  If  a  metal  achieves  its  maximum  coordination 
number  in  the  formation  of  a  complex  compound,  the  resulting  compound 
is  generally  more  stable  than  compounds  in  which  fewer  groups  are  co- 

75.  Long,  J.  Am.  Chem.  Soc,  61,  570  (1939). 


414  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

ordinated76.  The  exact  reason  why  a  metal  fails  to  fill  all  available  coordina- 
lion  positions  is  undoubtedly  a  combination  of  many  factors,  but  certainly 
the  size  of  the  ligand  relative  to  the  metal  is  one  such  factor. 

Availability  of  Bond  Orbitals.  Pauling77-78  points  out  that  the  d  orbitals 
of  the  penultimate  shell  are  of  great  significance  in  bond  formation.  The 
transition  elements  have  inner  d  orbitals  of  about  the  same  energy  as  the 
s  and  p  orbitals  of  the  valence  shell,  and  it  is  with  these  elements  that  com- 
plex formation  occurs  most  extensively  if  their  d  orbitals  are  not  completely 
occupied  by  unshared  electron  pairs. 

Nature  of  the  Bond 

Effective  Atomic  Number.  The  stability  of  coordination  compounds 
is  sometimes  related  to  the  attainment  or  near  attainment  of  the  number  of 
electrons  of  the  next  rare  gas  in  the  period.  Sidgwick76  has  described  this 
number  of  electrons  as  the  effective  atomic  number  (E.A.N.)  of  the  central 
metal  ion.  Thus  the  ammines,  [Pt(NH3)6]Cl4  and  [Co(NH3)6]Cl3 ,  and  the 
metal  carbonyls,  such  as  Mo(CO)6  and  Ni(CO)4 ,  appear  to  owe  their  sta- 
bility to  the  rare  gas  configuration  of  the  central  atom.  In  the  hexacyano- 
ferrate(II)  ion  the  coordinated  metal  has  36  electrons,  but  in  the  hexa- 
cyanoferrate(III)  ion  the  metal  has  only  35  electrons.  The  proponents  of 
the  effective  atomic  number  concept  would  explain  the  instability  of  the 
latter  on  the  basis  of  its  electron  deficiency.  In  like  manner,  the  great 
stability  of  tris(a,a/-dipyridyl)iron(II)  bromide,  which  was  resolved  by 
Werner79,  and  the  instability  of  tris(ethylenediamine)iron(III)  chloride 
may  be  related  to  the  effective  atomic  number  concept.  Similarly,  Gil- 
christ80 has  offered  explanations  of  the  stabilities  of  some  of  the  platinum 
group  complexes.  The  compounds,  K3[RuCl6]  (E.A.N.  =  53)  and  K3[OsCl6] 
(E.A.N.  =  85),  are  unstable,  but  if  a  nitrosyl  group  replaces  a  chloro  group, 
the  effective  atomic  number  of  each  is  increased  to  that  of  the  next  rare 
gas.  The  resulting  compounds,  K2[RuCl5NO]  and  K2[OsCl5NO],  are  ex- 
tremely stable. 

Although  the  above  explanations  on  the  basis  of  the  effective  atomic 
number  concept  seem  plausible,  it  must  be  pointed  out  that  not  only  is 
this  highly  formalistic,  but  direct  application  of  the  principle  is  possible 
only  with  a  minority  of  complexes,  and  it  is  not  possible  to  predict  the  stabil- 

76.  Sidgwick,  "The  Electronic  Theory  of  Valency,"  p.  163.  Oxford  University  Press, 

London,  1946. 

77.  Pauling,  "The  Nature  of  the  Chemical  Bond,"  p.  92,  Ithaca,  New  York,  Cornell 

University  Press,  1944. 

78.  Pauling,  J.  Am.  Chem.  Soc,  63,  1367  (1931). 

79.  Werner,  Ber.,  45,  433  (1912). 

80.  Gilchrist,  Chem.  Revs.,  32,  321  (1942). 


STABILIZATION  OF  VALENCE  STATES  415 

it v  of  any  complex  on  the  basis  of  this  concept  alone.  However,  it  holds  for 

all  volatile  carbonyls  and  oitrosyls. 

Hybridization  of  Orbitals*  On  the  basis  of  quantum  mechanics,  Paul- 
ing18 developed  a  theory  which  satisfactorily  accounts  for  the  relative 
strengths  r»i  bonds  formed  by  the  different  atoms,  the  molecular  configura- 
tion, and  the  magnetic  behavior  of  complex  compounds.  Postulating  thai 
the  stronger  bond  between  two  atoms  will  be  formed  by  the  two  orbitals 
which  can  overlap  more  with  each  other  and  that  the  bond  so  formed  will 
be  in  the  direction  in  which  the  orbital  has  its  greatest  density,  Pauling 
derived  a  number  of  results  of  chemical  and  stereochemical  significance 
(.Chapter  9). 


\A.    Theories   of  Acids,   Bases,  Amphoteric 

Hydroxides  and  Basic  Salts  as  Applied  to 

The  Chemistry  of  Complex  Compounds 

Fred  Basolo* 
Northwestern  University,  Evonston,  Illinois 

The  fact  that  bases  are  electron  pair  donors  and  acids  are  electron  pair 
acceptors  was  first  pointed  out  by  Lewis.  It  follows  that  the  interaction  of 
an  acid  and  a  base  results  in  the  formation  of  a  coordination  compound 
which  subsequently  may  or  may  not  yield  ions.  Excellent  accounts  of  the 
early  concepts  of  acids  and  bases  have  been  written  by  Walden1,  by  Luder 
and  Zuffanti2a,  and  by  Audrieth2b. 

The  oxonium  theory  of  acids  and  bases,  proposed  by  Werner3  shortly 
after  the  advent  of  the  water  theory,  was  the  first  attempt  to  indicate  the 
importance  of  the  solvent  in  acid-base  relationships  (the  Arrhenius  theory 
disregarded  the  solvent).  Although  Werner's  interpretations  were  only 
partially  correct,  he  succeeded  in  showing  that  the  solvent  is  a  principal 
agent  in  electrolytic  dissociation,  instead  of  being  merely  a  passive  medium 
in  which  solutes  are  dispersed.  In  his  studies  of  the  hydroxoamminecobalt- 
(III)  complexes,  Werner  discovered  that  they  react  with  water  in  the 
following  manner: 

[Co(NH3)5OH]++    +  HOH  ;=±  [Co(NH3)5OH2]+++  +  OH~ 

nonionized  hydroxyl  ionized  hydroxyl 

'  Mr.  Stephen  J.  Bodnar  helped  in  the  preparation  of  this  chapter.  His  help  is 
gratefully  acknowledged. 

1.  Walden,  "Salts,  Acids  and  Bases,"  New  York,  McGraw-Hill  Book  Co.,  Inc., 

1929. 

2.  Luder  and  Zuffanti,  "The  Electronic  Theory  of  Acids  and  liases,"  New  York, 

John  Wiley  &  Sons,  Inc.,  1946;  Andreth,  "Twenty  third  Annual  Priestley  Lec- 
tures: Acids,  Bases,  and  Nonaqueous  Systems"  Ypsilanti,  Michigan,  Uni- 
versity Lit  Imprinters,  1949. 

3.  Werner, Z.anorg.Chem.,  3,  267  (1893);  16, 1  (1897);  Werner,  Ber.,  40,  4133  (1907); 

Werner,  "New  Ideas  on  Inorganic  ( Ihemisl  ry,"  1  ranslated  by  Hedley,  London, 
Longmans,  Green  and  Company,  1911. 

416 


ICIDS,  BASES,  AND  AMPHOTERIC  HYDROXIDES  417 

By  analogy  he  postulated  that  no  metal  hydroxide  dissociates  until  it  is 
hydrated,  indicating  this  by  the  reaction 

MOB    •    HOB  _-  [MOHaJOB  ^  [MOH,]4  +  OB 

lie  called  the  hydroxide,  M(  )I  I,  an  anhydro  base  and  the  compound  which 
actually  dissociates,  [MOHJOH,  an  aquo  base.  Similarly,  Werner  postu- 
lated that  the  ordinary  "hydrogen"  acids,  in  analogy  to  the  complex 
plat inic  acids,  form  hydrates  in  aqueous  solution,  and  that  the  acid  hydro- 
gen comes  from  the  water  and  not   the  anhydro  acid;  viz: 

[PtClj(OB),]  +  2HOH  ^±  H-,[PtCl2(OH)d  ^±  2H+  +  [PtCl2(OH)4]=. 

Thus,  in  effect,  an  anhydro  acid  is  a  compound  which  combines  with  the 
hydroxyl  group  of  water,  liberating  an  excess  of  hydrogen  ions, 

A  +  HOH  ^±  H[AOH]  ^±  H+  +  [AOH]~ 

anhydro  acid  aquo  acid 

while  an  anhydro  base  is  a  compound  which  combines  with  the  hydrogen 
ion  of  water  to  produce  an  excess  of  hydroxyl  ions, 

B  +  HOH  ^±  [BH]OH  ^±  [BH]+  +  OH" 

anhydro  base  aquo  base 

The  reaction  between  an  aquo  base  and  an  aquo  acid  results  in  the  forma- 
tion of  an  aquo  salt, 

H[AOH]     +  [BH]OH  ;=±  [BH][AOH]  +  H20 
aquo   acid       aquo  base         aquo  salt 

Therefore,  the  reaction  between  potassium  hydroxide  and  hydrochloric 

acid  was  written: 

KOHJOH      +     H[HC10H]    ->  [KOH2][HC10H]  +  H20 

aquopotassium         aquohydrogen  .iquopotassium 

hydroxide  chloride  chloride 

According  to  this  theory,  it  is  to  be  expected  that  basic  metallic  hydroxides 
and  analogous  compounds  would  always  form  aquo  salts  when  neutralize*  1 
with  acids.  Werner  states  that  the  instability  of  the  free  aquo  salts  in  no 
way  contradicts  the  assumption  of  the  existence  of  aquo  bases  and  aquo 
salts  in  solution,  but  shows  rather  that  a  relationship  exists  between  the 
strength  of  the  base  and  the  stability  of  the  aquo  salts;  the  stability  de- 
creases  as  the  strength  of  the  base  increase-.  Consequently,  the  phenome- 
non that  the  strongesl  metallic  hydroxide  bases  (those  of  the  alkali  metals) 
preferably  yield  anhydrous  .-alts  is  to  be  explained  by  the  assumption 
that  the  aquo  -ait-,  which  are  originally  formed,  are  too  unstable  to  be 
isolated. 
These  ideas  -hocked  the  followers  of  Arrhenius  and  gave  rise  to  severe 


418  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

criticism  from  numerous  investigators  in  the  field;  others  simply  passed 
over  Werner's  oxonium  theory  as  being  of  no  importance4.  A  criticism 
raised  by  Walden1  mentions  the  difficulty  encountered  if  ethyl  alcohol  is  used 
instead  of  water.  He  suggests  that  it  would  be  necessary  for  the  alcohol  to 
dissociate  in  two  different  ways,  allowing  the  formation  of  [HC10C2H5]H 
and  [KC2H5]OH.  It  would  appear  that  this  may  not  be  a  justifiable  objection 
because  of  the  analogy  of  OH~  and  OC2H5~  which  allows  a  designation  of 
[KH]OC2H5  for  the  alcoholobase.  Although  some  of  the  ideas  of  the  theory 
are  wholly  consistent  with  present  views,  it  did  not  achieve  wide  acceptance. 
Solvents  other  than  water  were  seldom  considered  as  media  for  acid- 
base  reactions  prior  to  1905;  in  that  year,  Franklin5  demonstrated  the 
striking  similarity  between  reactions  carried  out  in  liquid  ammonia  and 
those  known  to  occur  in  aqueous  solutions6.  Liquid  ammonia  ionizes  into 
ammonium  and  amide  ions,  just  as  water  ionizes  into  hydronium  and 
hydroxide  ions. 

2NH3  ;=±  NH4+  +  NH2" 

2H20  ^±  H30+  +  OH- 

In  liquid  ammonia,  substances  like  ammonium  chloride  are  acids  and  sub- 
stances like  sodium  amide  are  bases.  Acids  and  bases  in  ammonia  solution 
neutralize  each  other  just  as  they  do  in  aqueous  solutions: 

NH4CI  +  NaNH2  -»  NaCl  +  2NH3 

H3OCI  +  NaOH    ->  NaCl  +  2H20 

acid  base  salt  solvent 

It  was  also  observed  that  hydrogen  was  liberated  by  the  reaction  of  an 
active  metal  and  ammonium  ions  in  liquid  ammonia,  a  reaction  which  is 
exactly  analogous  to  that  which  takes  place  in  aqueous  medium.  Additional 
experimental  evidence  in  support  of  the  close  similarity  between  wrater  and 
liquid  ammonia  wTas  furnished  by  the  fact  that  zinc  amide,  insoluble  in 
liquid  ammonia,  is  dissolved  upon  the  addition  of  either  ammonium  chloride 
or  sodium  amide,  just  as  zinc  hydroxide  is  soluble  in  either  an  excess  of 
hydronium  chloride  or  sodium  hydroxide : 

/-VTT—  OTT~ 

[Zn(H20)4]++  r-^-*  [Zn(H20)2(OH)2]  ,  H|Q+  »  [Zn(OH)4]= 
[Zn(NH3)4]++  ^==^  [Zn(NH3)2(NH2)2l  *==^  [Zn(NH2)4]~ 

4.  Lamb  and  Yngve,  J.  Am.  Chem.  Soc,  43,  2352  (1921). 
:..   Franklin,  ./.  Am.  Chem.  Soc,  27,  820  (1905). 

6.  Franklin,  "The  Nitrogen  System  of  Compounds,"  New  York,  Reinhold  Publish- 
ing Corp.,  1935. 


ACIDS,  BASES,  AND  AMPHOTERIC  HYDROXIDES  419 

This  analogy  between  the  hydronium  ion  and  ammonium  ion  suggested 

that  the  acid  properties  result  from  the  solvated  proton  in  each  instance. 
Some  of  the  more  extensively  studied  protonic  solvents  are  acetic  acid7, 8, 
hydrogen  sulfide9*  l0,  n,  hydrogen  fluoride12,  sulfuric  acid  13,  14?  and  hydroxyl- 
amine1-'.  Experiments  carried  out  in  nonprotonic  solvents  such  as  phosgene16, 
sulfur  dioxide17,  selenium  oxychloride18,  and  bromine  trifluoride19  revealed 
that  certain  generalizations  can  be  made  for  any  solvent  system  (Table 
12.1).  G.  B.  L.  Smithls,  in  an  excellent  review  of  the  subject,  defines  an 
acid  as  an  electron-pair  acceptor  toward  the  solvent,  and  a  base  as  an  elec- 
tron-pair donor  toward  the  solvent. 

One  of  the  more  recent  concepts  of  acid-base  phenomena20  (often  referred 
to  as  the  "Positive-negative"  Theory)  defines  an  acid  as  any  substance 
capable  of  giving  up  a  cation  or  combining  with  an  anion  or  electron,  and 
a  base  as  any  substance  capable  of  giving  up  an  anion  or  electron,  or  of 
combining  with  a  cation.  Usanovich  suggests  that  neutralization  reactions 
be  considered  as  shown  in  Table  12.2.  Sodium  oxide  is  a  base  because  it  is 
capable  of  giving  up  the  anion  0=  and  silicon  dioxide  is  an  acid  because  it 
combines  with  this  anion.  In  the  reaction  of  sodium  with  chlorine,  sodium 
is  the  base  because  it  gives  up  an  electron  and  chlorine  is  the  acid  since  it 
combines  with  the  electron.  This  implies  that  oxidation  and  reduction  are 
nothing  more  than  special  cases  of  acid-base  phenomena.  Partly  because  of 
this2a  and  also  because  of  the  stress  placed  upon  salt  formation,  and  the 
reasoning  involved  in  making  ions  so  important,  the  theory  has  been  widely 
criticized. 

7.  Davidson,  J.  Am.  Chem.  Soc,  50,  1890  (1928);  Davidson,  Chem.  Rev.,  8,  175 

(1931). 

8.  Davidson  and  McAllister,  J.  Am.  Chem.  Soc,  52,  519  (1930). 

9.  Quam,  J.  Am.  Chem.  Soc.,  47,  103  (1925). 

10.  Quam  and  Wilkinson,  /.  Am.  Chem.  Soc,  47,  989  (1925). 

11.  Wilkinson,  Chem.  Rev.,  8,  237  (1931). 

12.  Weiser,  "Inorganic  Colloid  Chemistry,"  Vol.  II,  New  York,  John  Wiley  &  Sons, 

Inc.,  1935;  Simons,  J.  Am.  Chem.  Soc,  54,  129  (1932). 

13.  Kendall  and  Davidson,  J.  Am.  Chem.  Soc,  43,  979  (1921). 

14.  Kendall  and  Landon,  /.  Am.  Chem.  Soc,  42,  2131  (1920). 

15.  Audrieth,  ./.  Phys.  Chem.,  34,  538  (1930);  Audrieth,  Trans.  III.  StaU  Acad.  Sci., 

22,  385  (1930) ;  Audrieth,  Z.  physik.  Chem.,  A165,  323  (1933). 

16.  Germann,  ./.  .1//    ('hem.  Soc,  47,  2461  (1925);  Germans  and  Timparry,  ibid.,  47, 

2275  (1925). 

17.  Jander  and  Wickert,  /.  physik.  Chem.  A178,  57  (1936);  Jander  and  [mmig,  Z. 

org.  allgem.  Chem.,  233,  295  (1937);  Jander  and  Ullmann,  ibid.,  233,    105 
(1937);  Jander  and  Schmidt,  Wien.  Chem.  Ztg.,  46,  49  (1943]  . 

18.  Smith,  Chem.  Rev.,  23,  165  (1938). 

19.  Sharpe  and  Emeleus,  J.  Chem.  Soc,  1948,  2135;  Banks,  Emeleus,  and  Wool!', 

ibid.,  1949,  2861;  Woolf  and  Emeleus,  ibid.,  1949,  2865;  Sharpe,  Quart.  R 
Chem.  Soc,  London,  IV,  No.  2  (1950). 

20.  Usanovich,  J.  Gen.  Chem.,  U.S.S.R.,  9,  182  (1939). 


420  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Table  12.1.  Different  Solvent  Systems 

A.  Ionization  of  Various  Solvents: 
solvent  — *  acid  +  base 
2H20^±  [HH20]+  +  OH- 
2NH3^  [HNH3]+  +  NH2- 
2HC2H3()2;=±  [HHC2H302]+  +  C2H30r 
2H2S^  [H-H2S]+  +  HS- 
2H2S04^  [HH2S04]+  +  HSOr 
2C0C12^±  [C0C1-C0C12]+  +  ci- 
4S02^±  [SO-2S02]++  +  SOr 

2BrF3  ^  BrF2+  +  BrFr 

B.  Neutralization  reactions  in  Various  Solvents: 
acid  +  base  — >  salt  +  solvent 

[H-H20]+,  X-  +  M+,  OH-  ->  MX  +  2H20 
[H-HC2H302]+,  X-  +  M+,  C2H30r  ->  MX  +  2HC2H302 
[C0C1-C0C12]+,  [AlCU]-  +  M+,  CI-  ->  M[A1C14]  +  2C0C12 
*[SO-2S02]^,  Xr  +  M2+,  SOr  -»  2MX  +  4S02 

C.  Reaction  of  a  Metal  with  an  Acid  in  Various  Solvents: 
metal  +  acid  — >  metal  ion  +  reduction  product  +  solvent 
2M  +  2[HH20]+  ->  2M+  +  H2  +  2H20 

2M  +  2[H-NH3]+  -*  2M+  +  H2  +  2NH3 

2M  +  2[HNH2OH]  -*  2M+  +  H2  +  2NH2OH 

2M  +  2[C0C1-C0C12]+  ->  2M+  +  CO  +  3C0C12 

D.  Electrolysis  of  Various  Solvents: 

Cathode  Reaction  Anode  Reaction 

base  — >  oxidation  product  + 
acid  +  e~  — ■>  reduction  product  +  solvent  solvent  +  e~ 

2[HH20]+  +  2e~  ->  H2  +  2H20  40H~  -»  02  +  2H20  +  4e" 

2[H-NH3]+  +  2e  ->  H2  +  2NH3  6NH2~  ->  N2  +  4NH3  +  6e~ 

2[C0C1-C0C12]+  +  2e-  ->  CO  +  3C0C12  2C1"  ->  Cl2  +  2e~ 

*[SO-2S02]++  +  2e~  ->  SO  +  2S02  SOr  -»  S03  +  2e~ 

E.  Amphoterism  in  Various  Solvents: 

base  m  t     t  base 

cation     ^  amphoteric  precipitate     v  anion 

acid  acid 

OH"  OH- 

[M(H20)x]+   ,       /   MOH  /   [M(OH)x]<*-» 

H3U1"  H3U"'" 

|M(NH3),]+  *==*  MNH*  *=±  [M(NH2)J<»-T 

[M(Hc,H,o,y  ■g.^o,^  mch,o,  ^^^  ww^- 

ua-  110- 

Cl_  CI- 

IM(COCl,)„]*  ■lcocl.coci!r'MCl-lcoci.coci,r''MCl^-r 

*  The  rate  of  exchange  of  sulfur  in  solutions  of  thionyl  halide  in  sulfur  dioxide  is 
extremely  slow.  These  results  indicate  that  there  is  a  negligible  amount  of  thionyl 
ion  in  these  solutions  so  that  the  simple  ionization  picture  represented  here  is  in  need 
of  some  modification.  Johnson,  Norris,  and  Huston,  J.  Am.  Chem.  Soc.}  73,  3052 
(1951). 


ACIDS,  BASKS,  AND  AMPHOTERIC  HYDROXIDES  421 

Table  12.2.  Some  Neutralization  Reactions  According  to  the 
Positive  Negative  Theory 


Acid 

+ 

Base 

-> 

Salt 

Si02 

+ 

NaiO 

-> 

NaSiO, 

BnSi 

+ 

(NH4)2S 

-♦ 

(NH4)s[SnS, 

AgCN  ■+ 

N;,('N 

-> 

Na[Ag(CN)2 

SnCl; 

+  2KC1 

-» 

Ki[SnCl$] 

CI, 

+  2Na 

-» 

2NaCl 

The  Proton  Theory 

The  one-element  theory  of  acids  and  bases  has  been  very  successfully 
modernized  into  what  is  known  as  the  proton  theory21, 22,  which  defines 
an  acid  as  a  substance  that  gives  up  a  hydrogen  ion  and  a  base  as  a  sub- 
stance that  accepts  a  hydrogen  ion: 

A    ^=±   B-   +  H+ 
acid         base 

However,  this  equation  is  purely  hypothetical,  for  an  acid  will  not  give  up 
a  proton  unless  a  base  is  present  to  accept  it,  so  that  an  exchange  of  a  pro- 
ton from  an  acid  to  a  base  produces  an  acid  conjugate  to  the  original  base 
and  a  base  conjugate  to  the  original  acid.  The  ionization  of  hydrogen 
chloride  is  written: 

HC1  +  H20  ;=±  H30+  +  Cl~ 
acid        base  acid         base 

The  reaction  toward  the  right  takes  place  because  of  the  tendency  of  hydro- 
gen to  form  the  coordinated  [H(OH2)]+  ion. 

The  fact  that  this  theory  is  both  general  and  useful  has  been  extensively 
discussed2113, 23.  Its  greatest  shortcoming  lies  in  the  fact  that  it  is  not  adapt- 
able to  nonprotonic  systems  and  does  not  include  as  acids  substances  which 
contain  no  hydrogen. 

The  Electronic  Theory 

Lewis24  suggested  that  the  behavior  of  acidic  and  basic  substances  might 
be  described  entirely  in  terms  of  electrons.  In  his  own  words,  "It  seems  to 
me  that  with  complete  generality  we  may  say  that  a  basic  substance  is  one 
which  has  a  lone  pair  of  electrons  which  may  be  used  to  complete  the  stable 

21.  Brpustcl,  Ree.  iron,  chim.,  42,  718  (1923);  Br0nsted,  Chem.  Rev.,  5,  231  (1923). 

22.  Lowry,  CI  A  Industry,  42,  1048  (1923). 

/     PI         Chem.,  30,  777     1926) ;  Hall,  Briscoe,  Hammett,  Johnson, 
Alyea,  McReynolds,  Hazlehurst,  and  Luder,  ''Add-  and  Bases,"  Journal  of 
Chemical  Education,  Easton,  Pennsylvania,  1941. 
24.  Lewis,  •/.  Franklin  Inst.,  226,  293  (1938). 


422  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

group  of  another  atom,  and  that  an  acid  substance  is  one  which  can  employ 
a  lone  pair  from  another  molecule  in  completing  the  stable  group  of  one  of 
its  own  atoms.  In  other  words,  the  basic  substance  furnishes  a  pair  of  elec- 
trons for  a  chemical  bond,  the  acid  substance  accepts  such  a  pair." 

The  electronic  theory  of  acids  and  bases  has  been  reviewed  by  Luder2*'  2,\ 
Since  the  theory  defines  an  acid  as  a  substance  capable  of  accepting  a  pair 
of  electrons,  and  a  base  as  a  substance  capable  of  donating  a  pair  of  elec- 
trons, it  requires  that  the  first  step  in  a  neutralization  reaction  be  the  for- 
mation of  a  coordinate  covalent  bond ;  this  appears  to  be  extremely  general : 

A  +    B  ->  A:B 

acid  base  coordination  compound 

[H(OH2)]+  +   :0:H-   ->  2H20 

F  H  F  H 

II  II 

F— B       +  :N— H  -»  F— B:N— H 

II  II 

F  H  F  H 

The  theory  makes  no  mention  of  the  solvent  (not  even  the  necessity  of  a 
solvent),  nor  is  anything  said  about  protons. 

The  Acid-Base  Properties  of  Some  Coordination  Compounds 

The  effect  of  coordination  on  acid-base  properties  may  be  considered, 
qualitatively,  on  the  basis  of  ionic  size  and  charge.  The  maximum  amount 
of  distortion  is  exerted  by  small  cations  of  high  ionic  charge26,  acting  on 
large,  polarizable  anions.  This  polarization  effect  explains  why  oxides  of 
large  metal  ions  with  small  positive  charge  react  with  water  to  form  bases, 
e.g.,  Na20  +  H20  ->  2NaOH,  CaO  +  H20  ->  Ca(OH)2 ,  while  oxides  of 
nonmetals  or  of  small  metals  in  the  higher  oxidation  states  react  with  water 
to  form  acids,  e.g.,  C120  +  H20  ->  2HC10,  Cr03  +  H20  ->  H2Cr04 .  In 
all  of  these  compounds  an  atom  of  oxygen  is  interspersed  between  the  hy- 
drogen atom  and  the  remainder  of  the  molecule ;  the  basic  or  acidic  charac- 
ter seems  to  depend  largely  upon  the  relative  attractive  forces  between  the 
oxide  ion  and  the  hydrogen  ion,  on  the  one  hand,  and  the  remainder  of  the 
molecule  on  the  other,  modified  by  the  energy  of  hydration  of  the  resulting 
ions.  This  being  the  case,  hydroxides  of  sodium  and  chlorine  behave  differ- 
ently because  of  the  difference  in  the  sizes  of  the  respective  ions.  Since 
t  he  sodium  atom  is  large,  the  bond  between  it  and  oxygen  is  weak  and  cleav- 

25.  Luder,  Chem.  Rev.,  27,  547  (1940);  Luder  and  Zuffanti,  ibid.,  34,  345  (1944). 

26.  Fajana  and  Joos,  Z.  physik.,  23,  (1924). 


ACIDS,  BASES,  AND  AMPHOTERIC  HYDROXIDES  123 

Table  12.3.  Ionic  PorBNTiAM  ro»  Cations  01   the  First  Two  Short  Pbriodb 


Cations 

i. 

Be 

B^+ 

c* 

N 

0^ 

i 

Hydroxide 

1.29 
base 

2.64 
amphoteric 

3.87 

acid 

5.16 

acid 

6.71 
acid 

(8.19) 

acid 

(10) 

acid 

Cations 

Vi 

\i, 

Al 

- 

ps+ 

S«+ 

Cl«* 

Hydroxide 

1.02 
base 

1.76 
base 

2.45 
amphoteric 

3.13 
amphoteric 

3.83 
acid 

4.55 
acid 

5.20 
acid 

age  occurs  at    (1),  while  the    chlorine   atom  is   small  and  forms  a  rela- 
tively strong  bond  with  oxj'gen  so  that  cleavage  occurs  at  (4). 


N*  i    •<*    I    h 

(I)         (2) 


CI    I    9/    I  h 

13)        U) 


The  same  conclusions  were  reached  by  Cartledge27  in  his  paper  on  ionic 

potential.  He  defines  the  ionic  potential,  </>,  as  4>  =  — ,  in  which  Z  is  the  oxi- 

r 

dation  state  of  the  ion  and  r  is  the  radius  of  the  ion.  Since,  in  any  comparison 
of  the  properties  of  two  different  ions,  the  increasing  ionic  charge  and  in- 
creasing ionic  radius  act  in  opposite  directions,  it  is  apparent  that  the  ratio 
of  charge  to  radius  (0)  must  be  considered  in  any  predictions  of  relative 
properties.  Cartledge2*  has  pointed  out  that  ions  in  which  V^  <  2.2  are 
basic,  those  with  3.2  >  y/$  >  2.2  are  amphoteric,  and  those  with  v^  >  3.2 
are  acidic  (Table  12.3). 

These  observations  on  the  relation  between  polarization  and  ionic  po- 
tential can  be  used  to  explain  the  fact  that  although  cobalt  (III)  hydroxide 
is  a  very  weak  base,  hexamminecobalt(III)  hydroxide  is  as  strong  a  base  as 
the  alkali  hydroxides4.  This  results  from  an  increase  in  the  effective  radius 
of  the  cation,  and  a  consequent  decrease  in  the  ionic  potential,  since  the 
oxidation  state  is  not  changed.  The  unavailability  of  orbitals  to  form  co- 
valent  bonds  must  also  be  considered.  Boric  acid29  is  an  extremely  weak 
monobasic  acid  (K  =  6  X  10-10);  the  phenolphthalein  end  point  (Fig.  12.1) 
is  reached  when  only  10  to  20%  of  the  acid  has  been  neutralized.  Hilde- 
brandM  followed  the  change  in  pH  when  varying  amounts  of  mannitol 
were  added  to  boric  acid  (Fig.  12.1).  Curve  E  corresponds  roughly  to 
A'  =  10~5  and  shows  that  the  excess  mannitol  magnifies  K  by  about  104, 

27.  Cartledge,  /.  Am.  Chem.  Soc.,  50,  2855,  2863  (1928). 

28.  Cartledge,  ibid.,  52,  3076  (1930). 

29.  Jorgensen.  Z   angew.  Chem.,  ot!>  (1896). 

30.  Hildebrand,  •/.  Am.  Chem.  Soc.,  35,  860  (1913 


424 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


PH 


tt 

y/> 

BO^S- 

^G\£ 

^  / 

A, 

yS 

B 

^ 

^ 

C 

0 

0 

^ 

1^^ 

"E 

xT 

0      0.1      0.2   0.3    0.4    0.5    0.6    0.7    0.8     0.9     10      I.I 
EQUIVALENTS   NAOH    PER    MOLE    H3BO3 

Fig.  12.1.  Titration  curves  of  mixtures  of  boric  acid  and  mannitol. 

Curve  A    0.8  g  mannitol  per  100  ml.  O.liV  H3BO3  . 

Curve  B     2.4  g  mannitol  per  100  ml.  O.liV  H3BO3  . 

Curve  C     4.0  g  mannitol  per  100  ml.  O.liV  H3BO3  . 

Curve  D     5.6  g  mannitol  per  100  ml.  O.liV  H3BO3  . 

Curve  E    7.2  g  mannitol  per  100  ml.  O.liV  H3BO3  . 

making  it  possible  to  titrate  boric  acid  conveniently  using  phenophthalien 
as  the  indicator.  Although  the  exact  structure  of  these  complex  acids  has 
not  been  conclusively  established,  it  is  known  that  the  hydroxy  groups  are 
attached  to  the  boron  in  such  a  way  as  to  displace  a  proton,  and  thus  in- 
crease the  acid  strength.  Lowry31  proposes  the  quadricovalent  structure  for 
the  mannito-boric  acid  complex: 


11 


HO 


0 


B  C6H1204 

/   \    / 
HO  O 


Cationic  Complexes 

Bases.  Werner  has  called  attention  to  the  variation  in  basicity  of  a 
series  of  hydroxo  complexes3b  (Table  12.4).  His  qualitative  studies  showed 
thai  :  (1)  will  precipitate  silver  oxide  from  silver  nitrate;  (1)  through  (3) 
liberate  ammonia  from  NH4+  in  the  cold;  (1)  through  (5)  absorb  carbon 
dioxide;  (1)  through  (8)  react  alkaline  to  litmus  while  (9)  and  (10)  are 
neutral;  (1  )  through  (8)  are  more  soluble  in  acetic  acid  than  in  water;  from 
acetic  acid  solutions  of  (1)  through  (3)  the  salts  precipitate  as  aquo  salts, 
while  (It  hrough  (8)  yield  hydroxo  salts;  all  of  these  cations  appear  to  form 


31.  Lowry,  J.  Chcm.  Soc,  1929,  2853. 


ACIDS,  BASES,  AND  AMPHOTERIC  HYDROXIDES 
Table  12.4,  Werner's  Series  op  Basic  Cations 


125 


No. 

Cations 

M 

Cation- 

1 

2 
3 

4 
5 

[Co(NH,)«(NO,)OH]+ 
[Co(NH      OHJ++ 
[Co(NH,)4(H^))OHJ++ 

[Co  en,    11  0  OH         1,2) 

[Co  en,    H,0  <>1I)++  (1,6) 

6 
7 
8 
9 
10 

[Cr(MI        B,0    <)H]++ 
[Cr(XH3)2(H20)2(OH)2]+ 
[CoiMI      |.v     Ho  <)II]++ 
[Ru(NH     .  \0)OH]++ 
[Pt(XH3)4(OH)2]++ 

Table  12.5.  Conductance  Ratio  of  Some  Ammixecobalt(III)   Btdroxides 


N 
1 

2 
3 
4 
5 
6 
7 
8 
9 


Cation 

a  (%)  (1.33  X 

lCo(XH,)4CO,]+ 

97.6 

tran*-[Co(NH,)4(NO,)8]+ 

95.0 

[Co(XH3)6]+++ 

89.5 

[Co  en3]+++ 

88.6 

cis  [Co(XH3)4(X02)2]+ 

81.2 

[Co(XH3)5H20]+++ 

53.5 

[Co(XH3)3H20(X02)2]+ 

36.0 

[Co  en2  (H20)2]+++ 

27.3 

[Co(XH3)4(H20)2]^+ 

24.6 

10-3  m; 


t'  (%) 


82.9 

84.8 
74.0 


aquo  salts  with  strong  mineral  acids  but  even  from  solutions  of  this  type 
(9)  and  (10)  are  still  isolated  as  the  hydroxo  complexes. 

Werner  ascribed  this  decrease  in  basic  strength  from  the  moderately 
si  rong  base  (1)  to  the  nonbasic  ion  (10)  to  a  difference  in  affinity  for  the  hy- 
drogen ion.  Werner's  observations  have  been  reviewed  by  Br0nsted23a  and 
the  results  interpreted  in  terms  of  more  modern  concepts  (page  421). 

Coordination  of  the  metal  of  a  weak  base,  MOH,  results  in  the  formation 
of  a  stronger  base,  [MAJOH,  due  to  the  increase  in  cationic  size.  Lamb  and 

Yngve4  determined  the  conductance  ratio  (  a  =  -^  J  for  a  series  of  ammine- 

cobalt(III)  hydroxides  at  0°,  and  found  that  many  of  them  are  as  highly 
ionized  as  the  hydroxides  of  the  alkalis  (Table  12.5).  Hall34  points  out  that 
if  the  more  probable  assumption  (rejected  by  Lamb  and  Yngve)  is  made, 
that  the  aquo  cations  are  transformed  to  hydroxo  compounds,  in  Werner's 
sense,  the  more  useful  figures  (a)  are  obtained. 

Acids.  The  acidity  of  aqueous  solutions  of  salts  can  be  accounted  for  by 
the  loss  of  protons  from  the  hydrated  cations. 

[M(H20),]++  +  H20  ^±  [M(H20)I_,OH]+  +  H30+ 

For  instance,  as  early  as  190G  Bjerrum35  reported  a  value  of  0.89  X  10~4 
as  the  dissociation  constant  at  2.5°  for  the  reaction 

34.  Hall,  Chem.  Rev.,  19,  89  (1936). 

35.  Bjerrum.  Kgl.  Dm, she  Videnskab.  SeUkabi  Skeifter,  [7]  4,  1  (1906). 


426  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

[Cr(H20)6]+++  +  H20  ;=±  [Cr(H20)5OH]++  +  H30+ 

and  a  few  years  later  Denham36  assigned  it  a  value  about  twice  as  great. 
Lamb  and  Fonda87  arrived  at  an  average  value  of  1.58  X  10~4  at  25°  which 
is  comparable  to  a  more  recent  determination  by  Br0nsted  and  Volqvartz38. 
The  acidity  of  aquoammines  is  due  to  loss  of  protons  from  the  coordinated 
water  molecules,  although  with  the  ammines  of  heavier  metals,  the  acidity 
of  the  coordinated  ammonia  is  noticeable.  Tschugaev39  and  Griin- 
berg40*' 41a- 41b  have  demonstrated  this  by  the  conversion  of  platinum  am- 
mines to  the  corresponding  amido  or  basic  salts: 

[Pt(NH3)5Cl]+++  +  OH-  ^±  [Pt(NH3)4NH2Cl]++  +  H2() 

Corresponding  amido  compounds  of  cobaltammines  are  not  known,  but 
evidence  for  this  type  of  reaction  has  been  obtained  from  exchange  reactions 
with  heavy  water423  • 42b. 

[Co(NH3)6]+++  ^  [Co(NH3)5NH2]++  +  H+ 

[Co(NH3)5NH2]++  +  HDO  ^±  [Co(NH3)5NH2D]+++  +  OH- 

H+  +  OH-  ^±  H20 

Ionization  of  a  hydrogen  ion  from  one  of  the  coordinated  ammine  groups  in 
the  bis(ethylenediamine)gold(III)  ion  has  been  demonstrated  by  Bailar 
and  Block42c.  This  phenomenon  has  also  been  reported  by  Dwyer  and  Ho- 
garth, who  studied  the  ethylenediamine  complexes  of  osmium42d.  The  study 
of  metal  ammine  complexes  furnishes  some  insight  into  the  properties 
of  aquo  ions.  The  dissociation  constants  for  some  of  these  ions  are  known 
fairly  accurately  (Table  12.6).  The  equilibrium  constants  are  calculated 

36.  Denham,  ./.  Chem.  Soc,  93,  53  (1908). 

37.  Lamb  and  Fonda,  J.  Am.  Chem.  Soc,  43,  1154  (1921). 

38.  Br0nsted  and  Volqvartz,  Z.  physik  Chem.,  134,  97  (1928). 

39.  Tschugajeff,  Z.  anorg.  allgem.  Chem.,  137,  1,  401   (1924);  Tschugajeff,  Compt. 

rend.,  160,  840  (1915);  161,  699  (1915). 

40.  Griinberg  and  Faermann,  Z.  anorg.  allgem.  Chem.,  193,  193  (1930);  Griinberg  and 

Gildengershel,  Izvest.  Akad.  Nauk  S.S.S.R.,  Otel.  Khim.  Nank,  479  (1948). 

41.  Griinberg  and  Rvabchikov,  Acta.  Physiocochim .  U.S.S.R.,  3,  555  (1935);  Griin- 

berg, ibid.,  3,  573  (1935);  Griinberg  and  Rvabchikov,  Compt.  rend.  acad.  set. 

U.S.S.R.,  4,  259  (1936);  Griinberg,  Bull.  acad.  set.  U.S.S.R.,  Classe  sci.  chin,., 

350  (1943). 
L2a.  Anderson,  Spoor,  and  Briscoe,  Nature,  139,  508  (1937). 
L2b.  Anderson,  Spoor,  and  Brisco,  Nature,  139,  508  (1937);  Anderson,  Briscoe,  and 

Spoor,  J.  Chem.  Soc,  1943,  36] ;  Garrick, Nature,  139,  507  (1937) ;  James,  Ander- 

BOn,  and  Briscoe,  Nature,  139,  109  (1937). 
lie.  Block  and  Bailar, ./.  Am.  Chem.  Soc,  73,  4722  (1951). 
12.1    Dwyer  and  Hogarth, ./.  Am.  Chem.  Soc,  76,  1008  (1953). 


ACIDS,  BASES,  AND  AMPHOTERIC  HYDROXIDES 


427 


Table  12.6*.  Acid  Strength  <>r  Somk  Comim.kx  Cations 


Acid 


pKa 


[Co  en,  (OH)a]+ 

L,  (13) 

[Co(NH,)4(OH),r 

L,  (12) 

[Co(NH       NO,),(H,0)]+ 

L,  (11) 

Pi   Ml      NH2C1]++ 

G,  10.9 

[Pt  en,  Cl2]++ 

(i.  10.4 

[Pt(NH    -Cl,]++ 

G,  9.8 

[Pt(NH      OIIJ+++ 

G,  pKa,  ,  9.5;  pKa, 

,  10.7 

IPUXH3)5Br]+++ 

G,  pKa,  ,  8.2;  pKas 

,  10,1 

[Pt(NH,)6Cl]-^ 

G,  pKa,  ,  8.1;  pKaa 

,  10.5 

[Pt(NH,)6r 

G,  pKa,  ,  7.9;  pKaa 

,  10.1 

[Ru(XH3)4(XO)OHr 

W,  7 

[Pt  en  (XH3)4p+ 

G,  pKai  ,  6.2;  pKa2 

,  10.0 

[C0(M1          ;.\(),HoO]  +  + 

W,  6 

[Rh  \H:.),H,0]+++ 

B,  5.86 

[Cr(H20)4Cl2]+ 

L,  5.72;  Bj,  5.42 

[Co(XH3)5H,0]+++ 

B,  5.69;  W,  (5-6) 

[Co(XH3),(H20)2]+++ 

B,  5.22;  W,  (5-6) 

[P1    (>n3]4+ 

G,  pKai  ,  5.5;  pKa2 

,  9.8 

*cis-[Co  en2  (HoO),]4^4 

W,  (3-4) 

*<rans-[Co  en2  (H20)2]+++ 

W,  (3-4) 

[A1(H20)6]+++ 

B,  4.95 

[Co(XH3)3(H20)3]+++ 

B,  4.73 

[Cr(H,0),]+++ 

B,  3.90,  L,  3.80,  Bj, 

4.05,  ] 

[Co(XH3)2(H20)4]+++ 

B,  3.40 

[Co(XH3)2(H20)3OH]++ 

W,  (2-3) 

[Cr(XH3)2(H20)4]+++ 

W,  (2-3) 

[Co(XH3),(H20)4]+++ 

W,  (2-3) 

[Ru(XH3)4XOH20]+++ 

W,  (2) 

[Pt(XH,)4(H20)2]4+ 

W,  (2) 

[Fe(H20)6]+4+ 

B,  2.20 

D,  3.75 


*  In  this  table,  B  refers  to  Br0nsted,  Bj  to  Bjerrum,  D  to  Denham,  G  to  Grunberg, 
L  to  Lamb,  and  W  to  Werner.  This  table  is  taken  from  a  review  article  by  Hall34  to 
which  the  data  of  Griinberg40  are  added.  Xote  that  in  a  few  cases  Griinberg40b  has 
demonstrated  the  polybasicity  of  the  complex  platinum(IV)  ion  acids.  The  third 
dissociation  constant  was  evaluated  with  difficulty  in  only  a  few  cases,  and  it  was 
demonstrated  that  the  ratio  K2/K3  is  much  smaller  than  K]/K-i  . 

**  Bjerrum  and  Rasmussen,  Acta  Chem.  Stand.  6,  1265  (1952)  report  the  following 
pK«  values:  cisiCoeno^O),]4^-,  pKal  =  6.06,  pKa2  =  8.19;  trans  [Co  en2  (H20)2]+++, 
pK.,  =  4.45,  pKa,  =  7.94. 

as  shown  below : 

[Co(XH3)5H20]+++  ^±  [Co(XH3)5OH]++  4-  H+ 

„       [Co(XH3)5OH]++[H+] 

A   = — — =    1    V    10-6  43 

[Co(XH3)5H20]+++  X 

In  the  case  where  the  proton  is  liberated  from  a  coordinated  ammine  group, 
43.  Br0nsted  and  King,  Z.  physik  Chem.,  130,  699  (1927). 


128  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Table  12.7.  Relative  Stabilities  of  Amminecobalt(III)  Ions44 

(1)  */Yms-[Co(NH3)4(N(),),r  (4)  aV[Co(NH3)4(N02)2]+ 

(2)  |(\»u\II,)6]+++  (5)  [Co(NH3)4(H20)2]+++ 

Co(NH8)  NOJ++  (6)  [Co(NH,)6H20]+++ 

Table  12.8.  Dissociation  Constants  of  Some  Metal  Ammixes 

Ammine  Kc 

[Ag(NH3)2]+  6.8    X  10-8 

[Cu(NH,)8]+  1.5    X  10-9 

[Cd(NH3)4]++  1.0    X  10-7 

[Zn(NH3)4]++  2.6    X  lO"10 

[Co(NH3)6]++  1.75  X  10-5 

the  expression  for  K  is  as  illustrated  below : 

[Pt(NH3)5Cl]+++  ^±  [Pt(NH3)4NH2Cl]++  +  H+ 

[Pt(NH3)4NH2Cl]++[H+] 
K  =  - — — — - — =   7.9  X  10"9  40 

[Pt(NH3)5Cl]+++ 

In  addition  to  dissociation  constants,  the  relative  stabilities  of  a  series 
of  amminecobalt(III)  ions  were  determined44  and  it  was  found  that  the 
stabilities  decrease  in  the  order  shown  in  Table  12.7.  The  concentration 
dissociation  constants  are  very  small  (Kc  =  2.2  X  10~34  for  [Co(NH3)6]+++) 
as  is  expected  from  the  well  known  chemical  stability  of  these  cations.  It 
can  be  supposed  that  the  greater  the  dissociation  constant  (greater  the 
tendency  to  liberate  ammonia)  of  these  ions,  the  weaker  their  acid  strength; 
that  this  is  usually  true  can  be  seen  by  a  comparison  of  the  relative  stabili- 
ties of  the  complexes  in  Table  12.7  with  their  relative  acid  strengths  given 
in  Table  12.6.  This  is  further  illustrated  by  the  very  small  acid  strengths 
of  the  more  highly  dissociated  metal  ammines  listed  in  Table  12.8. 

Br0nsted43  deduced  that  in  the  homologous  series  of  aquoammine- 
cobalt(III)  ions,  the  acid  strength  is  a  statistical  factor  based  upon  the 
number  of  coordinated  aquo  groups.  This  requires  that  a  hexaaquo  ion  be 
six  times  as  strong  an  acid  as  a  monoaquo  ion.  Br0nsted  and  Yolqvartz38 
found  that  although  the  calculated  influence  of  the  statistical  factor  is  in 
qualitative  agreement  with  the  values  found  for  the  dissociation  constants 
of  aquoamminecobalt(III)  ions  (Table  12.9),  it  is  insufficient  to  account 
quantitatively  for  the  differences  found. 

Br0nsted48  has  called  attention  to  the  relation  between  acid  strength 
;iii(l  the  charge  on  an  aquo  cation;  Werner  found  [Co(NH3)50H]++  to  be 
less  basic  than  |(  !o(Nl  I:i  ),X()2OH]+  which  means  that  [Co(NH3)5H20]+++ 
is  more  acidic  than  ((,o(XII:;)i  NO2  I1L>()|++.  Br0nsted  deduced  from  such 
examples  that  the  higher  the  positive  charge  on  the  complex,  the  stronger 

II.  Lamb  and  Larson,  ./.  Am.  Chem.  Sac,  42,  2024  (1920). 


ACIDS,  BASES,  AND  AMPHOTERIC  HYDROXIDES  429 

Table  12.9.  Dissociation  Constants  of  Somk  TkiposiTIVB   Acid- 


\ 

Cation 

A'„  X  10" 

No. 

5 
6 

7 
8 

Cation 

K,i  X  10" 

1 

2 

3 
4 

[Co(NH3)5H,>0]+++ 
[Co,  Ml     .  11.. <»,]+++ 
[Co(NH3)3(H20)3]+++ 

[Co(XH3)2(H2())4]+++ 

2.04 
6.03 
18.8 
400. 

[Rh(NH8)5H,0]+++ 

[A1(H,0)6]+++ 
[Cr(H20)6]+++ 
[Fe(H20)6]+++ 

i.:;s 
11.2 
126. 
6300. 

the  acid.  This  is  a  Logical  consequence  of  the  greater  repulsion  oi  a  proton 
by  the  more  posit ive  cation.  Lamb  and  Yngve4  found  that  the  substitution 
of  an  additional  nitro  group  decreased  the  acid  strength  still  further.  Like- 
wise, Tschugajeff39  has  prepared  a  series  of  hydroxoammineplatinum(IV) 
ions  and  noticed  that  [Pt(XH3)5OH]+++  is  a  much  weaker  base  than  the 
corresponding  cobalt(III)  complex,  [Co(NH3)50H]++,  which  has  a  smaller 
positive  charge.  There  is  also  a  considerable  difference  in  the  acidic  strength 
of  hexammineplatinum(IV)  and  hexamminecobalt(III)  ions;  the  latter  has 
little  tendency  to  behave  as  an  acid42a  while  the  former  is  readily  soluble  in 
alkaline  solution,  from  which  the  amido  complex  can  be  isolated39a. 

[Pt(NH8)6]4+  +  H20^  [Pt(NH8)6NH2]+++  +  H30+ 

It  should  be  mentioned,  however,  that  this  difference  in  acidity  between 
[Pt(XH3)6]4+  and  [Co(XH3)6]+++  is  greater  than  anticipated  merely  on  the 
difference  in  cationic  charge. 

The  influence  of  the  oxidation  state  of  the  central  atom  on  the  acid 
strength  of  complex  ions  has  been  demonstrated  by  comparing  the  proper- 
ties of  [Co(XH3)6]+++  and  [Pt(XH3)5Cl]+++.  The  net  charge  on  the  cations 
is  the  same,  but  the  cobalt  (III)  ion  is  almost  neutral  while  the  platinum(IV) 
is  strongly  acid. 

A  careful  consideration  of  the  relative  acid  strengths  shown  in  Table  12.6 
reveals  the  fact  that  no  definite  predictions  can  be  made  from  the  structure 
of  the  cation  alone.  However,  it  is  apparent  that  the  charge  and  size  of  the 
complex,  the  charge  of  the  central  atom  and  the  statistical  factor  must  all 
exert  considerable  influence.  Likewise,  the  ammine  cations  are  in  general 
far  less  acidic  than  the  corresponding  aquo  cations. 

( rriinberg  has  published  a  series  of  interesting  papers4011'  41  concerned  with 
the  effect  of  geometrical  isomerism  on  acid  strength.  In  investigating  the 
acid-base  properties  of  cis-  and  /ra/i.s-diacmodiammineplatinum(II),  lie 
found  that  the  first  ionization  of  the  trans  isomer  is  greater  than  that  of 
the  cis  form,  and  that  the  two  ionization  constants  oi  the  cis  isomer  are 
nearly  alike,  while  those  of  the  trans  isomer  are  quite  different  from  each 
other.  The  explanation  of  this  observation  is  given  in  terms  of  the  trans 
effect45  (see  Chapter-  3  and  8). 


45.  Chernyaev,  ann.  inst.  pl/itinc,  4,  243  (1936). 


430 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


This    was   illustrated  by  Grunberg41  with  the  geometrical  isomers  of 
diaquodiammineplatinum(II) : 


II,  ()              NH3 

\     / 
P1 

/     \ 
_H3N               OH2_ 

++ 

+  H20^± 

H20               NH3 

\     / 
Pt 

/     \ 
_H3N              OH  _ 

+ 
+  H30+  (1) 

"H20              NH3" 

\     / 
Pt 

/     \ 
_H3N               OH  _ 

+ 

+  H>0^± 

"HO                NH3~ 

\     / 
Pt 

/     \ 

_H3N              OH  _ 

+  H,0+  (la) 

"H20              NH3" 

\     / 

Pt 

/     \ 
_H20              NH3_ 

++ 

+  Ho()  ^± 

"  HO              NH8~ 

\     / 
Pt 

/     \ 
_H20              NHg_ 

+ 
+  H30+  (2) 

"  HO              NH3~ 

V 

/  \ 

_H.O              NH,_ 

+ 

+  H2O  ^ 

"  HO              NHf 

\     / 

Pt 

/     \ 
_  HO              NH3_ 

+  H30+  (2a) 

Fig.  12.2.  The  trans-effect  principle  as  applied  to  the  first  and  second  acid  dissoci- 
ation constants  of  a  Werner  complex. 

Since  it  is  the  group  trans  to  the  aquo  group  that  affects  its  ionization, 
(Fig.  12.2),  the  first  ionization  (1)  of  the  trans  isomer  is  greater  than  that 
(2)  of  the  corresponding  cis  form,  because  the  polarizability  of  water  is  less 
than  that  of  ammonia  (RH2o  =  3.76;  RNh3  =  5.61).  The  cis  isomer  should 
show  very  little  difference  in  the  two  ionization  constants,  K\  or  (2)  and 
K2  or  (2a),  because  the  group  opposite  the  ionizing  group  is  XH3  in  both 
cases;  while  the  two  ionization  constants  of  the  trans  isomer  should  differ 
markedly  since  K\  or  (1)  is  a  measure  of  ionization  with  water  opposite  the 
ionizing  group  and  K2  or  (la)  is  the  same  measurement  with  a  much  more 
highly  polarizing  group  (ROH  =  5.1)  trans  to  the  aquo  group.  In  this  case 
the  stronger  trans  effect  of  the  hydroxo  group  should  result  in  a  value  of 
/v2  smaller  than  that  of  Ki  .  Although  the  same  conclusions  are  reached  on 
the  basis  of  a  smaller  charge  on  the  cation,  this  is  not  justified  in  that  it 
also  predicts  different  ionization  constants  for  the  cis  isomer.  Ryabchikov468 
carried  out  potentiometric  titrations  with  the  cis  and  trans  isomers  of 
diaquodiammineplatinum(II)  ion  and  found  that  the  cis  isomer  behaves 
as  a  monobasic  acid,  while  the  trans  isomer  gives  the  type  of  curve  charac- 
beristic  of  dibasic  acids.  The  observation  that  the  cis  isomer  is  monobasic 


46a.  Ryabchikov,  Ann.  aecteri  platine,  Inst,  chim.,  gen.  (U.S.S.R.)  16,  35  (1938). 


ACIDS,  BASES,  AND  AMPHOTERIC  HYDROXIDES  L31 

is  indeed   unexpected   in   view   of   the   fact    that    the   monovalent    cation, 
tPt(NHi)iHjO(OH)]+  should  certainly  be  a  weaker  acid  than  [Pt(NH 
1 1  <  0»]+1-  Therefore,  the  acid  constants  of  these  two  isomers  were  carefully 

redetermined  by  Jensen1''  and  the  pl\„  values  obtained  were:  Cis  [Pt- 
Ml  11:<)',];-  pKmi  =  5.56,  pK.s  =  7.32;  trans  [Pt(XH3).»(II,()),]i2 
pK :.  =  4.32,  pKa:  =  7.38.  These  results  are  not  inconsistent  with  Grun- 
berg's  interpretations  of  relative  acid  strength  on  the  basis  of  the  polariz- 
ability  of  the  trans  ligand.  In  the  first  place  the  trans  isomer  is  the  stronger 
acid  as  explained  previously.  Secondly  the  ApKa  =  1.70  observed  for  the 
cis  isomer  may  be  attributed  to  the  difference  in  charge  on  the  cation.  The 
greater  difference,  ApKa  =  3.00,  for  the  trans  isomer  can  be  said  to  result 
from  the  larger  polarizing  effect  of  the  trans  hydroxo  group  compared  to 
the  original  aquo  group  in  the  first  dissociation  step.  It  is  of  interest  that 
this  same  polarization  treatment  can  account  for  the  acid  dissociation 
constants  of  cis  and  trans  isomers  of  [Co  en2(H20)2]+3  46d  and  [Co  en2X02- 
H20]+2  46b. 

Anionic  Complexes 

Werner  first  called  attention  to  the  almost  complete  analogy  between  the 
union  of  anhydrides  with  water  to  give  oxyacids,  and  the  union  of  metal 
halides  with  hydrogen  halides  to  form  the  halo  acids. 

H>0  +  S03  -»  HS04 

HF  -f  BF3  -»  HBF4 

2HC1  +  PtCl4  -*  HoPtCh 

The  various  factors  known  to  effect  the  acid-base  strengths  of  complex 
cations  can  be  expected  to  have  similar  effects  on  complex  anions.  For 
example,  it  was  pointed  out  (page  429)  that  the  larger  the  charge  on  a 
cation,  the  greater  its  repulsion  of  a  proton  and  consequently  the  stronger 
its  acid  properties;  in  much  the  same  way  it  has  been  shown47  that  while 
[Fe(CX)6]s  is  a  very  weak  base,  [Fe(CX)6]4_  is  about  as  strong  a  base  as 
benzoate  ion.  This  would  indicate  that  the  more  negative  a  complex  anion, 
the  greater  the  proton  attraction  and  therefore  the  stronger  its  basic  proper- 
ties. 

Mention  has  also  been  made  of  the  increased  basic  strength  of  [Co  a6] 

(OH)3  overCo(OH)3  due  to  the  coordination  of  six  "a"  groups  to  the  cobalt- 

(III)  ion.  In  much  the  same  way,  certain  weak  acids  are  greatly  strengthened 

by  coordination  (page  423).  This  is  illustrated  by  the  weak  acid  IK'X 

K    =  7.2  X  10-lu)  as  compared  to  the  relatively  strong  acid  H4[Fe(CX)6] 

46B.  Stone,  thesis,  Northwestern  University,  1952. 

46C.  Jensen,  Z.  anorg.  Chem.  24_\  s?     l'.)39). 

46D.  Bjerrurn  and  Rasmusaen,  Acta.  ('firm.  Stand.,  6,  1265  (1952). 

47.  Kolthoff  and  Tomsicek,  J.  Phya.  Chem.,  39,  945  (1935). 


432  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

(K\  =  6.8  X   10~5)4S.  A  similar  explanation  might  be  given  for  the  fact 
thai  water  is  neutral  while  complexes  in  which  oxygen  is  the  donor  atom 

( IIj|S()4),  II[C104],  etc.)  are  often  strong  acids. 

Relative  Acid -Base  Strength 

In  the  preceding  discussion  an  attempt  has  been  made  to  account  for 
increasing  or  decreasing  strengths  of  acids  and  bases.  The  generalizations 
made  are  concerned  with  the  acid  strength  toward  a  reference  base,  OH-, 
or  the  basic  strength  towards  the  acid,  H30+,  in  the  solvent,  water.  The 
tact  that  it  is  impossible  to  arrange  acids  or  bases  in  a  single  monotonic 
order  of  strength  has  been  clearly  stated  by  Lewis24.  He  points  out  that  the 
relative  acid-base  strengths  depend  upon  the  solvent  chosen  as  wrell  as 
upon  the  particular  base  or  acid  used  for  reference. 

It  has,  however,  been  suggested25a  that  on  the  basis  of  the  electronic 
theory  of  acids  and  bases,  the  relative  strengths  of  acids  correspond  to 
the  tendency  to  accept  pairs  of  electrons  while  the  strengths  of  bases  de- 
pend on  their  tendency  to  donate  pairs  of  electrons.  If  this  wrere  all  that 
need  be  considered  it  should  certainly  be  possible  to  construct  a  monotonic 
series  of  acids  and  bases.  Perhaps  a  more  correct  interpretation  of  acid-base 
strength  is  that  suggested  by  Lingaf elter49 :  (a)  the  strength  of  an  acid  cor- 
responds to  the  strength  of  the  bond  it  can  form  with  a  base,  or  (b)  the 
strength  of  an  acid  corresponds  to  the  decrease  in  free  energy  upon  forma- 
tion of  a  bond  with  a  base.  Certainly  the  interatomic  forces  of  a  coordina- 
tion compound  (neutralization  product)  involve  not  only  the  bonding  forces 
of  the  covalent  bond,  but  also  electrostatic  forces  which  depend  upon  the 
magnitude  and  separation  of  charges  and  the  presence  or  absence  of  dipole 
moments  in  either  acid  or  base  and  steric  effects. 

Pauling50  has  pointed  out  the  variation  in  the  strength  of  bonding  orbitals 
of  different  types.  Since  the  factors  contributing  to  bond  strength  can  vary 
more  or  less  independently,  the  relative  strengths  of  a  series  of  bases  may 
depend  on  the  particular  acid  used  in  making  the  comparison.  That  this  is 
the  case  has  been  shown49  by  a  consideration  of  some  equilibrium  constants 

(K  =  —r — ^rrr, —  t^    ,  as  a  measure  of  the  strength  of  an  acid  or  base, 
[acid]  [base]   / 

The  equilibrium  constants  at  25°  for  the  reactions 

H+  +  B  ^  HB+,        Ag+  +  2B  ;=±  [AgB>l+,        Cu+  +  2B  ^±  [CuB2]+, 

and        Hg++  +  4B  ^  [HgB4]++ 

are  given  in  Table  12.10. 

is.  BrittoD  and  Dodd,/.  Chem. Soc. ,  1988,  154:};  LanfordandKiehl,/.Pfcys.  Chem., 

45,  300  (1941   . 
19.  Lingafelter,  •/.  .1///.  Chem.  Soc.,  63,  1999  (1941). 
50.  Pauling,   "The   Nature  of  the  Chemical  Bond,"  Ithaca,  New   York,  Cornell 

University  Press,  L945. 


ACIDS,  BASES,  AND  AMPHOTERIC  HYDROXIDES  433 

Table  12.10.  Equilibrium  Constants  fob  Some  Acids  with  Different 


R 

I  i  BBENCE    B  ISES 
Acid 

Base 

ii 

\ 

i'u 

n« 

CN 

2.5  X  10" 

2.6  X  1018 

1  X  10" 

2.5  X  10" 

Ml 

1.8  X  L09 

1.7  X   K)7 

so,- 

1  X   10- 

3.5  X  108 

Cl 

Weak 

3.4  X  L0« 

9  X  1015 

Br- 

Weaker 

8.3  X  105 

4.3  X  I0-'1 

I 

Weakest 

7.1  X  108 

1.9  X  1030 

In  each  series  there  is  a  reversal  of  relative  strengths  of  some  of  the  bases 
upon  changing  the  reference  acid,  showing  thai  no  single  arrangement  of 
basic  strength  can  be  made  which  will  be  applicable  to  all  cases.  These 
peculiarities  in  relative  acid  strengths  seem  to  be  connected  with  the  fact 
that  different  metals  have  different  coordinating  power  toward  various 
ligands. 

The  difference  in  acid-base  strengths  depending  on  the  reference  base 
or  acid  can  sometimes  be  explained  on  the  basis  of  molecular  structure;  this 
possibility  has  been  more  carefully  investigated  with  organic  compounds 
than  in  the  field  of  inorganic  chemistry.  A  good  example  is  the  reversal  of  the 
relative  strengths  of  triethylamine  and  ammonia;  ammonia  is  the  weaker 
base  toward  the  proton,  but  much  stronger  toward  ra-dinitrobenzene. 
Lewis  and  Seaborg61  explain  this  behavior  as  being  a  result  of  the  double 
chelation  which  is  possible  through  hydrogen  bonding  in  the  case  of  am- 
monia but  not  with  triethylamine: 


The  researches  of  Brown  and  co-workers52  demonstrated  a  complete  re- 
versal in  the  basic  strength  of  ammonia  and  primary,  secondary  and  terti- 
ary amines.  They  collected  data  on  the  dissociation  constants 

R  \:BR3'^R;iX:  +  BR3' 

[R  =  CH3  and/or  H;        C2H5  and/or  H.        R'  =  CH3  ,  C2H3  ,  CH(CH3)2 

or  C(CH3)3] 

and  equilibrium  constants  for  the  displacements 

i:  \:    •    IT    \:BR'3  —  R3N:BR'3  +  R"3N: 

51.  Lewis  and  Seaborg,/.  .1///.  Chem.  Soc.,  62,  2122  (1940). 

52.  Brown,  Moddie,  and  Gerstein,  •/.  .1///.  Chem.  Soc..  66,  431  (1944);  Brown,  Bar 

tholomay,  and  Taylor,  foid.,68,  I3fi  (1944);  Brown,  ibid.,  67, 374  I L945);  Brown. 
ibid.,  67,  378  (1946) ; Brown,  ibid.,  67,  503  (1945);  Brown,  ibid.,  67,  1452  (1945). 


434  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Table  12.11.  Relative  Base  Strength  of  Some  Amines  Compared  to  Different 

Reference  Acids 

Amine  H+  B(CH3)3  B(CH2CH3)3    B(CH(CH3)2)3    B(C(CHa)3)3 

Ml  4  4  1  1  2 

CH,NHj  2  2  1 

(CH3)2NH  1  1  3 

(CH3)3N  3  3  2  2  4 

(B) 

MI  4  3  1 

C2H5NH2  2  1  2 

(C2H5)2NH  1  2  3 

(C2H5)3N  3  4  4 

*  Relative  basic  strengths,  1  >  2  >  3  >  4. 

Some  of  the  results  obtained  are  tabulated  (Table  12.11)  to  show  the  rela- 
tive base  strength  of  different  amines  as  compared  to  various  reference  acids. 
The  steric  effects  arising  from  the  substitution  of  organic  groups  on  co- 
ordinated ethylenediamines  have  also  been  studied53  (see  Chapter  8). 

Amphoterism 

An  amphoteric  substance  is  one  which  is  capable  of  behaving  either  as 
an  acid  or  a  base.  Kraus57  considers  all  elements  of  the  4th,  5th,  6th  and  7th 
groups,  having  a  deficiency  of  electrons  with  respect  to  the  rare  gas  con- 
figuration, to  be  amphoteric,  and  Cartlege27  states  that  all  substances  of 
which  the  square  roots  of  the  ionic  potentials  lie  between  2.2  and  3.2  are 
amphoteric.  Contrary  to  such  generalizations,  even  lithium58  and  barium59  are 
amphoteric  under  some  conditions.  Again,  the  solvent  is  found  to  play  an 
important  role;  iron  (III)  hydroxide  is  not  amphoteric  in  water  but  iron  (II  I) 
cyanide  is  definitely  amphoteric  in  liquid  hydrogen  cyanide. 

The  mechanism  of  amphoterism  is  still  obscure  and  there  are  several 
theories  concerning  the  processes  of  dissolution  of  metallic  hydroxides  in 
an  excess  of  alkali.  The  discussion  which  follows  gives  a  brief  account  of 
three  of  these  theories  and  some  of  the  experimental  evidence  supporting 
each  of  them.  A  more  general  interpretation  of  amphoterism  is  also  pro- 
posed and  the  mechanism  of  these  reactions  is  related  to  the  behavior  of 
the  more  stable  Werner  complexes. 

53.  Basolo  and  Murmann, ./.  Am.  Chem.  Soc,  74,  5243  (1952);  Irving,  "A  Discussion 
on  Coordination  Chemistry,"  Paper  No.  4,  Butterwick  Research  Lab.,  I.C.I. , 
Sept.  21-22,  1950. 

57    Krause,  ■/.  Chem.  AW.,  8,  2126  (1931). 

58.  Krause  and  Krzyzanski,  Ber.,  70,  1975  (1937). 

59.  Beholder  and  Patsch,  Z.  anorg.  allgem.  Chem.,  222,  135  (1935). 


ACIDS,  BASKS.  AM)  AMPHOTERIC  HYDROXIDES  435 

Theories  on  the  Mechanism  of  Amphoterism 

The  Theory  ol"  Peptization.  The  fact  thai  in  most  cases  a  Large  in- 
definite excess  of  hydroxide  beyond  that  required  for  the  formation  of  a 
compound  such  as  NasZnOs  must  be  used  to  dissolve  an  amphoteric  hy- 
droxide has  suggested  the  possibility  that  no  true  compound  is  formed, 
but  that  the  insoluble  hydroxide  is  merely  peptized.  Many  experiments 
have  failed  to  establish  definitely  the  formation  of  a  true  compound. 
The  studies  of  Britton60  suggest  that  only  in  the  case  of  aluminum  is  a 
true  compound  formed,  while  Mahin61  consider.-  that  even  aluminum  forms 
mainly  colloidal  suspensions.  Weiser12a  believes  it  likely  that  the  first  step 
in  the  dissolution  of  some  or  all  hydroxides  is  peptization,  followed  in  most 
by  compound  formation. 

The  concentrations  of  the  hydroxide  ion  in  alkaline  solutions  of  ampho- 
teric hydroxides  have  been  determined62  (by  measurements  of  electrical 
conductivity  and  of  the  velocity  of  esterification)  to  be  larger  than  would  be 
expected  if  neutralization  of  the  metal  hydroxide  has  taken  place.  Accord- 
ing to  this  view,  hydroxide  ions  are  preferentially  adsorbed  by  the  particles 
of  insoluble  metal  hydroxide,  forming  negatively  charged  colloids.63  Evi- 
dence for  this  theory  is  given  by  the  fact  that  in  many  cases  (e.g.,  Cu(OH)2 
and  Cr(OH)a)  precipitates  of  the  metal  hydroxide  appear  on  standing,  or 
precipitation  may  be  brought  about  by  the  addition  of  an  electrolyte. 

Although  colloidal  suspensions  are  markedly  different  from  most  crystal- 
loid solutions,  it  is  well  known  that  true  solutions  and  colloidal  dispersions 
of  the  same  material  are  different  in  degree  only.  The  gradual  transition 
in  properties  from  true  solution  to  colloidal  dispersion  has  been  shown  for 
hydrophilic  colloids64  m  that  the  properties  of  true  solutions  of  low  molecular 
weight  amino  acids  are  similar  to  colloidal  dispersions  of  high  molecular 
weight  amino  acids  and  proteins.  A  similar  observation  has  been  made65 
for  the  transition  in  properties  from  a  true  solution  of  aluminum  chloride, 
through  the  more  basic  salts,  to  the  aluminum  oxychloride  hydrosol.  In 
fact,  some  colloid  chemists66,  concerned  primarily  with  the  structure  of  the 
micelle  rather  than  the  stability  of  the  suspension,  visualize  the  formation 
of  certain  colloids  as  a  continual  increase  in  the  molecular  weight  of  polynu- 
clear  complexes  until  colloidal  dimensions  are  reached  (see  page  457). 

60.  Britton.  "Hydrogen  Ions,"  3rd  Ed.,  Vol.  II,  London,  Chapman  and  Hall.  1942. 

61.  Mahin.  Engraham,  and  Stewart.  ./ .  .1///.  ('hem.  Soc,  35,  30  (1913). 

62.  Hantzsch,  7. .  anorg.  allgem.  rhem.,  30,  289  (1902). 

63.  Davis  and  Farnham.  ./.  Phijs.  Chem.,  36,  1056  (1932). 

64.  Loeb,  "Proteins  and  the  Theory  of  Colloidal  Behavior,"  1st   Ed.,  New  York, 

McGraw-Hill  Book  Company.  Inc.,  1922. 

65.  Whitehead  and  Clay,  /.  Am.  Chem.  Soc.,  66,  1844  (1934). 

66.  Whitehead,  CJu  -    R<     .  21,  113    Ifl 


136  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

The  Oxy-acid  Theory .  The  mechanism  proposed  in  1899  by  Bredig67  is 
often  referred  to  as  the  oxy-acid  theory  and  can  be  illustrated  by  the  equi- 
libria 

M     +  OH-  ^  MOH  ^  MO"  +  H+ 

A1+++  +  30H-  ^  Al(OH)3  ^  A103=  +  3H+. 

Studies  of  the  solubility  of  amphoteric  hydroxides  in  excess  of  alkali 
have  Led  to  the  conclusion  that  insoluble  hydroxides  react  with  excess 
alkali  to  form  definite  stoichiometric  compounds  instead  of  merely  being 
peptized  (page  435).  For  example,  Hildebrand68  followed  the  read  ion 
between  zinc  hydroxide  and  sodium  hydroxide  by  means  of  the  hydrogen 
elect  rode.  He  came  to  the  conclusion  that  the  hydrogen  zincate  ion, 
IlZn<  >2  ,  exists  in  the  presence  of  excess  sodium  hydroxide.  Mellor69  men- 
tions the  formation  of  sodium  meta-  and  ortho-chromite,  XaCrC>2  and 
\a;Cr03,  and  Grube  and  Feucht70  claim  that  dissolution  of  cobalt (II) 
hydroxide  in  potassium  hydroxide  is  due  to  the  formation  of  the  compound 
K2C0O2  .  Copper(II)  hydroxide  dissolves  appreciably  in  concentrated 
alkali  solutions,  giving  a  deep  blue  color,  and  the  bulk  of  the  evidence 
supports  the  view  that  the  coloration  is  due  to  the  cuprate  ion,  Cu02=,  and 
not  to  colloidal  copper(II)  oxide71. 

The  most  extensively  studied  hydroxide,  by  far,  is  that  of  aluminum; 
some  of  the  observations  made  on  this  amphoteric  hydroxide  support  the 
oxy-acid  theory.  Prescott72  states  that  since  one  mole  of  sodium  hydroxide 
is  needed  to  dissolve  one  mole  of  aluminum  hydroxide,  the  solution  must 
contain  the  meta-aluminate  ion,  A102~;  while  Herz73  points  out  that,  if  the 
aluminum  hydroxide  is  dried  before  treatment  with  the  excess  of  alkali, 
the  ortho-aluminate,  A103^,  is  formed.  Studies  with  the  hydrogen  electrode 
have  indicated  to  Blum74  and  to  Britton60  that  the  meta-aluminate  is 
formed. 

The  type  of  information  which  has  been  collected  by  these  investigators, 
and  by  many  others,  can  be  illustrated  by  a  brief  review  of  some  hydrogen 
elect  rode  studies  made  by  Britton  and  his  co-workers  (Fig.  12.3).  The  curve 
represent-  the  titration  of  a  solution  of  aluminum  sulfate  with  a  solution 
of  sodium  hydroxide.  The  solution  becomes  neutral  when  the  sodium  hy- 
droxide is  added  in  an  amount  slightly  less  than  is  required  for  the  forma- 

67.  Bredig,  Z.  Electrochem.,  6,  6  (1S09). 

68    Hildebrand  and  Bowers,  ./.  .1///.  Chem.  Nor.,  38,  785  (1916). 
(*>9.  Mellor.  "A  Comprehensive  Treatise  on  Enorganic  and  Theoretical  Chemistry," 
Vol.  EII,  p.  191,  New  York,  Longmans  Green  and  Company,  1928. 

70.  Grube  and  Feucht,  Z.  Electrochem.,  28,  568  (1922). 

71.  Jirsa,  Kolloid  Z..  40,  28  (1926). 

72.  Prescott,  ./.  .1///.  Chen.  Soc.,  2,  27    1880 

73.  Hers,  Z   anorg.  allgem.  Chem.,  23,  222  (1900). 

74.  Blum../.  .U/.  r/„,„.  >•„,-..  35,  1499  (1913). 


ACIDS.  BASES  AND  AMPHOTERIC  HYDROXIDES 


i:J7 


13 
I2J 
If 

10 
pH   9 

8 

7 
6 
5 
4 
3 


80         90 


0  10        20         30        40         50         60         70 

ml.  NdOH   (~  0.09N) 
Fig.  12.3.  Titration  of  aluminum  ion  with  sodium  hydroxide  (100  ml  of  0.00667  M 


tioD  of  aluminum  hydroxide,  owing  to  the  retention  by  the  precipitate  of 
some  of  the  acid  radical  present  in  the  original  salt.  This  precipitate  dis- 
solves  completely  when  approximately  one  more  equivalent  of  sodium 
hydroxide  is  added,  the  dissolution  being  reflected  by  the  characteristic 
aluminate  inflexion  of  the  titration  curve,  extending  over  a  pH  range  from 
8  to  10.5.  The  precipitate  dissolved  completely  when  4.13  equivalents  of 
sodium  hydroxide  had  been  added.  Hence,  it  is  concluded  that  the  formula. 
XaAlOj  ,  probably  represents  the  condition  in  which  aluminum  hydroxide 
exist-  in  solutions  of  alkali.  However,  information  of  this  type  does  not 
rule  out  the  possibility  that  the  formula  is  either  Xa[Al(OH)4]  or 
Xa[Al(H2()).2(OH)4]. 

< MJier  so-called  amphoteric  ions,  such  as  those  of  beryllium,  zinc,  chrom- 
Lum(III),  tin(II),  and  zirconium,  exhibit  similar  behavior,  but  according 
to  Britton60,  none  of  them  show  such  sharp  inflexions  as  does  aluminum. 
Britton  also  states  thai  only  in  the  ease  of  aluminum  hydroxide  is  the 
amount  of  alkali  required  for  the  solution  of  the  hydroxide  approximately 
equal  to  thai  suggested  by  the  formula  and  also  independent  of  the  concen- 
tration of  the  sodium  hydroxide  used76.  Britton  suggests  thai  this  is  possibly 
due  to  the  fad  thai  although  other  hydroxides  may  be  acidic  in  their  be- 
havior toward  alkali,  they  are  such  weak  acids  thai  the  hydrogen  ion  con- 
centration of  the  alkali  solution  is  scarcely  affected  by  their  presence. 

A  consideration  of  the  tremendous  amount  of  information  which  has 


5.  Britton,  Analyst,  46,  363    1921   . 


438  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

been  collected  reveals  that  there  is  no  conclusive  evidence  for  the  existence 
of  ions  such  as  A102~,  Pb044~,  Zn02=,  in  solution,  as  is  proposed  by  the 
oxy-acid  theory.  No  doubt  the  strongest  support  for  the  existence  of  these 
ions  in  solution  comes  from  the  fact  that  mixed  oxides  such  as  NaA102 , 
Iv.ZnOo ,  and  Ca2Pb04 ,  are  known  to  exist  in  the  solid  state.  Most  of  these 
compounds  can  be  made  only  by  fusion  of  a  mixture  of  the  constituent 
oxides,  and  x-ray  analyses76  show  that  they  are  essentially  ionic  crystals, 
often  with  structures  closely  related  to  those  of  the  simple  oxides.  Although 
it  is  customary  to  refer  to  substances  in  solution  as  having  the  same  formu- 
las as  in  the  solid  state,  it  is  well  known  that  this  is  not  always  necessarily 
the  case. 

The  Hydroxo -Complex  Theory.  A  somewhat  different  explanation  of 
the  dissolving  of  metallic  hydroxides  (almost  intermediate  between  the 
oxy-acid  theory  and  the  theory  of  peptization)  was  first  proposed  by 
Pfeiffer77  in  1908.  According  to  this  view  amphoterism  is  represented  by  the 
equilibria 

QTT—  OTT~ 

[M(H,0)»]*+  ^^  [M(H20)„-x(OH),]  *==*  [M(OH)n]C»-*>- 

[A1(H20)6]+++  ^=±  [Al(H20)8(OH)8]  ^=±  [Al(OH),]- 

The  maximum  number  of  hydroxo  groups  which  may  combine  with  the 
metal  ion  is  determined  by  the  coordination  number  of  the  metallic  ion, 
but  the  actual  number  varies  with  the  concentration  of  hydroxide  ion.  This 
concept,  which  is  referred  to  as  the  hydroxy -complex  theory,  is  mentioned  in 
only  a  few  textbooks78;  in  fact,  Wells76  states,  "...  there  is  no  evidence  for 
the  existence  of  complex  ions  in  these  solutions."  However,  several  pieces 
of  evidence  can  be  marshalled  to  support  the  theory.  The  oxy-acids  may  be 
divided  roughly  into  three  classes78a : 

(1)  Simple  oxy-acids,  formed  by  the  lighter,  strongly  electronegative 
elements.  The  composition  of  these  oxy-acids  is  governed  primarily  by 
direct  considerations  of  the  valency  of  the  central  atom,  and  there  is  little 
tendency  to  form  true  ortho-acids.  (H2S04  rather  than  S(OH)6  and  H3PO4 
rather  than  P(OH)5). 

(2)  Complex  oxy-acids,  formed  by  the  heavier,  weakly  electronegative 
or  amphoteric  elements.  The  composition  of  these  is  determined  by  the 
necessity  of  completing  the  coordination  sphere  of  the  central  atom 
(H[Sb(OH)6]  and  H6[IOfl]). 

76.  Wells,  "Structural  Inorganic  Chemistry,"  Oxford.  Clarendon  Press,  1945. 

77.  Pfeiffer,  Ber.,  40,  W36  (1908  . 

:-  Emeleus  and  Anderson,  "Modern  Aspects  of  Inorganic  Chemistry,"  New  York, 
I).  Van  Nostrarid  Co.,  1945;  Pauling,  "General  Chemistry."  San  Francisco, 
W.  11.  Freeman  and  Company,  (1947);  Sneed  and  Maynard,  "General  Inor- 
ganic Chemistry,"  p.  396,  New  York,  D.  Van  Nostrand  Co.,  1942. 


ACIDS,  BASES.  AND  AMPHOTERIC  HYDROXIDES  YM) 


B,PtCle-6H,0 


>  BilPtCUOH] 

UK)  nun 
Ba(OH)i 


Bunlight 

NaOH 


Ba[PtCl(OH 


->  Xa,[Pt(OH)6] 


NaOH 


PtCl«-5H,0 


■»  Na,[PtCl4(OH),l 


MI 


-  .MI,)2[PtCl2(OH)J 


Fig.  12.4.  Format  ion    of    the    chloro-hydroxo   complexes   of   platinum. 

(3)  Poly-acids,  formed  by  the  elements  of  groups  VB  and  VIB.  These 
are  discussed  in  Chapter  14. 

The  second  group,  termed  here  complex  oxy-acids,  include  the  metal 
hydroxides  capable  of  behaving  as  acids,  that  is,  the  amphoteric  hydroxides. 
Reactions  between  these  acids  and  varying  amounts  of  alkali  produce  solu- 
tions which  in  some  cases  are  known  to  yield  crystalline  compounds  of 
definite  composition  not  dependent  on  that  of  the  original  solution79.  Thus, 
the  alkali  stannates  and  plumbates  all  contain  three  molecules  of  water 
(Xa20-Sn02-3H20)  which  are  lost  only  at  temperatures  considerably  above 
100°,  when  complete  decomposition  of  the  salt  takes  place80;  the  more  highly 
hydrated  salts  (BaOSn02-7H20)  lose  water  readily,  down  to  the  last  three 
molecules.  The  salts  may  therefore  be  derived  from  an  anion  [Sn(OH)6]=,  in 
which  the  coordination  number  of  the  central  atom  is  satisfied;  removal  of 
the  constitutional  water  breaks  up  the  complex  anion  completely. 

The  fact  that  Pfeiffer,  who  worked  with  Werner,  looked  upon  alkaline 
solutions  of  amphoteric  hydroxides  as  coordination  compounds  with  hy- 
droxo  groups  attached  to  the  central  metal  ion  is  not  at  all  surprising.  A 
considerable  number  of  well  defined  complexes  are  known  in  which  the 
hydroxy]  ion  is  coordinated  to  the  central  atom,  i.e.,  [Co(XH3)50H]++. 
In  many  instances  the  metal  acceptor  also  forms  an  amphoteric  hydroxide 
and  it  therefore  seems  reasonable  to  suppose  that  the  metal  could  be  com- 
pletely surrounded  by  hydroxo  groups  instead  of  being  attached  to  only 
one  or  two  such  groups.  The  analogy  between  Werner's  complexes  and 
hydroxo  anion-  is  particularly  emphasized  by  the  nearly  complete  series 
Ol  compounds  between  H2[PtCU]  and  H2[Pt(OH)6],  worked  out  by  Miolati81 
(Fig.  12. 1 >.  Numerous  investigators  have  demonstrated  thai  the  amphoteric 

79.  Footer.  Z.  Electrockem.,  6,  30]  (1899  ;  Beholder,  Angeto  Chem.,  46,  5090    19 

Muller.Z.  Electrockem.,  33,  134    I 
B0.  Belucci  and  Parravano,  Z.  anorg.  Chem.,  45,  142  (1905). 
Bl.  Miolati, Z. amarg.  Chem., 22,  145    1900  ; 26,  209    1901);88,251  (1903 


440 


CHEMISTRY  OF  THE  COORD1 X ATION  COMPOUNDS 


10         20        30 
MOLE  °/o  N6.0H 


Fig.  12 
tic  acid. 


20        30       40 
Yo  NOlC2H302 

5.  Solubility  of  Zn++  in  NaOH  in  water  and  NaC2H302  in  glacial  ace- 


10 

MOLE 


behavior  observed  in  the  water  system  is  found  in  other  solvent  systems, 
and  that  reactions  in  different  solvents  support  the  hydroxy-complex 
theory  of  amphoterism. 

The  fact  that  certain  amides  which  are  insoluble  in  liquid  ammonia,  are 
dissolved  either  by  acid,  NH4+,  or  by  base,  NH2~,  was  reported  independ- 
ent ly  by  Franklin82  and  Fitzgerald83  (see  page  418).  Franklin6  has  given 
an  excellent  summary  of  some  other  examples  of  salts  of  amphoteric  amides 
and  imides. 

Similar  observations  have  been  made  with  glacial  acetic  acid  as  a  solvent. 
Davidson8  points  out  that  when  a  small  amount  of  sodium  acetate  solution 
is  added  to  a  solution  of  zinc  chloride  in  acetic  acid,  a  precipitate  of  zinc 
acetate  is  formed;  this  dissolves  when  additional  sodium  acetate  is  added. 
A  detailed  study  of  this  phenomenon  showed  that  the  analogy  between 
this  reaction  and  that  of  zinc  hydroxide  and  sodium  hydroxide84  in  water 
is  far  from  being  a  superficial  one.  The  solubility  curve  of  zinc  acetate  in 
acetic  acid  containing  varying  amounts  of  sodium  acetate  at  constant  tem- 
perature is  strikingly  similar  to  the  curve  for  zinc  hydroxide  in  aqueous 
sodium  hydroxide  solutions  (Fig.  12.5).  The  solid  phase  which  appears  at 
high  concentrations  of  the  sodium  compound  may  be  formulated,  in  each 
case,  as  a  ternary  addition  compound.  The  composition  of  these  two  ternary 
compounds  is  very  similar,  as  is  evident  from  the  following  comparison: 
Zn(OH)2-2NaOH-2H20  or  Xa2[Zn(OH)4]  -2H20  in  water  and  Zn(C2H302)2- 
2NaC2H802-4HC2Hs02  or  Na2[Zn(C2H302)4]-4HC2H;A  in  acetic  acid.  The 
same  sort  of  results  have  been  obtained  with  copper(II)7b' 85,  lead  (II)86, 
and  silver(I)87. 

Nbnprotonic  systems  have  likewise  been  investigated  in  connection  with 
amphoterism17*-  v\    It    has   been   observed   that   the   addition   of  a  base 

82.   Franklin,  ./.  .1///.  ('hem.  Soc.,  29,  1274  (1907). 
83    Fitzgerald,  ibid.,  29,  056  (1907). 

84.  Gourdioon,  Rec.  trav.  chim.,  39,  505  (1920). 

85.  Muller,  /.  physik.  Chem.,  105,  73  (1924);  114,  129  (1925). 
Tehrman  and  Leifer,  ./.  .1///    Ch  m.  Soc,  60,  1  12  (1938). 

s7    Peterson  and  Dienea,  •/.  Phys.  and  Colloid  Chen,.,  55,  1299  (1951). 
58.  Janderand  Hecht,  Z.  anorg.  allgem.  Chem.,  240,  287  (1943) 


ACIDS,  BASES,  AND  AMPHOTERIC  HYDROXIDES  111 

[(CII:;lj\]-jS( )..  ,   to  a  sulfur  dioxide  solution  of  aluminum  chloride  results 

in  the  precipitation  of  the  amphoteric  sulfite,  A1.(S(  n , ,  which  can  lie  dis- 
solved by  adding  more  of  the  base  to  yield  the  salt,  | ( C ' 1 1.; » ,N |;.| Al < S<  ):i);i]. 
Acid-base  reactions  in  different  solvents  were  discussed  on  page  lis  and 
the  close  analogy  of  amphoteric  behavior  in  various  systems  was  sum- 
marized in  Table  12.1.  It  may  be  mentioned  in  addition  that  iroiulll 
cyanide89  and  silver  cyanide'"'  are  amphoteric  in  liquid  hydrogen  cyanide 
and  that  several  alcoholates,  when  dissolved  in  alcohol,  increase t  hehydrogen 
ion  concentration  of  the  alcohol1". 

The  existence  of  hydroxo  complex  compounds  in  solutions  of  amphoteric 
hydroxides  in  strong  bases  has  found  support  in  a  determination  of  ionic 
weights  by  the  dialysis  method  of  Brintzinger92!  Using  chromate  ion  as  a 
standard,  it  was  found  that  the  following  ions  exist  in  solution: 

[Sb(OH)a]-  [Gaa(OH)8]- 

[Sb(OH)4]-  [Zn,(OH)8r 

[Ge(OH).]-  [Be10(OH),o]2°- 

[Al,(OH)8]- 

Although  these  values,  except  for  antimony  and  germanium,  appear  to 
differ  markedly  from  what  might  be  expected,  they  merely  represent  poly- 
nuclear  forms  of  the  mononuclear  complex  structures;  aluminum,  gallium, 
and  zinc  are  present  as  binuclear  complexes  while  beryllium  is  present  in 
the  decanuclear  form  of  [Be(OH)4]=. 

Much  more  convincing  proof79b  for  the  existence  of  these  hydroxo  complex 
compounds  is  furnished  by  the  successful  crystallization  of  well  defined 
salts  of  definite  composition  from  strongly  alkaline  solutions  of  amphoteric 
hydroxides  (Table  12.12).  Attempts  to  produce  nickelates100,  bismuthates1"1, 
mercurates",  and  borates59  failed  to  yield  well  defined  crystalline  com- 
pounds, probably  because  the  corresponding  hydroxides  are  extremely 
weak  acids.  In  general,  the  salts  were  made  by  adding  a  cold  solution  of  a 
Ball  of  the  metal  to  a  cold  concentrated  solution  of  sodium  hydroxide. 

vi.  Jander  and  Scholz,  Z.  pkyaik.  ('hem.,  192,  163  (1943  . 
Jander  and  Gruttner,  £er.,81,  114  (1948). 

91.  Meerwein,  Ann.,  455,  227  (1927 

92.  Brintzinger  and  Osswold,  Z.  angew.  Chem.,  47,  61  (1934  . 

93.  Scholder  and  Weber,  Z.  anorg.  allgem.  Chem.,  215,  365  (1933  ;  Scholder  and 

Bendrich,  ibid.,  241,  76    L939 

94.  Beholder,  Felsenstein,  and  Apel,  Z.  anorg.  allgem.  Chem.,  216,  138  (1938 
-(■holder  and  Weber,  Z   anorg.  allgem  Chem.,  216,  L50  (1933). 
Beholder  and  Patsch,  Z.  anorg.  allgi  m  Chi  m.,  216,  176    193 
Scholderand  Patsch.  ,/„</..  220,  til    1934 

98.  Scholder  and  Patech,  ibid.,  217,  21  1    1934 

-(■holder  and  Btaufenbiel,  Z.  anorg.  allgem  Chi  m.,  247,  259  (1941). 
-■•holder.  Z.  anorg  allgem  Chem.,  230,  209    1934  . 
101.  Scholder  and  Stobbe   Z  anorg.  allgem  Chem.t  847,  392    1941 


142  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Table  12.12.  Some  Hydroxo  Salts  Prepared  by  Scholder  and  His  Co-workers 

Zincates*3 

Na[Zn(OH)8]-3H20  Ba[Zn(OH)4]H20 

Na[Zn(OH)8]  Sr[Zn(OH)4]-H20 

Na2[Zn(OH)4]-2H20  Ba2[Zn(OH)6] 

Na2[Zn(OH)4]  Sr,[Zn(OH)6] 

Cuprates  (II)94 

Na2[Cu(OH)4]  Ba2[Cu(OH)6]-H20 

Sr[Cu(OH)4]H20  Sr2[Cu(OH)6]-H20 

Cobaltates  (II)95 

Na2[Co(OH)4]  Sr2[Co(OH)6] 

Ba2[Co(OH)6] 

Stannates  (II)96 

\a[Sn(OH)3]  Sr[Sn(OH)3]2.H20 

Ba[Sn(OH)3]2-2H20  Ba[Sn(OH)3]2 

Ba[  (HO)  2Sn-0-Sn  (OH)  2]  Sr  [  (HO)  2Sn-0-Sn  (OH) ,] 

Chromates  (III)97 

Na3Cr(OH)6  Ba3[Cr(OH)6]2 

Na4Cr(OH)7H20-2-3H20  Sr3[Cr(OH)6]2 

\a,Or(OH)8-4H20 

Plambates  (II)98 

Na2[Pb(OH)4](?)  Ba[Pb(OH)4] 

Na[Pb  (OH)  3]  Ba[Pb  (OH)  3X] 

Na2[Pb(OH)3X]  BaNa2[Pb(OH)6] 

(X  =  C1-,  CNS-,  Br-,  or  I~) 

Cadmates" 

Na2[Cd(OH)4]  Ba2[Cd(OH)6] 

Xa[Cd(OH)3H20]H20(?)  Sr2[Cd(OH)6] 

Na2[Cd(OH)3Br] 

The  compound,  Na2Cu02-2H20  or  Na2[Cu(OH)4],  loses  one  mole  of  water 
at  approximately  200°C,  at  which  temperature  the  color  changes  from  blue 
to  black.  Additional  heating  to  a  temperature  of  500°C  results  in  a  gradual 
loss  of  water  amounting  to  less  than  0.05  moles.  However,  if  the  black 
residue  is  intimately  mixed  with  potassium  dichromate,  the  second  molecule 
of  water  is  readily  lost  at  approximately  500°C.  If  it  is  assumed  that  the 
structure  of  the  compound  is  Na2Cu02-2H20  (oxy-acid  theory)  it  would  be 
expected  that  the  two  moles  of  water  would  be  liberated  under  approxi- 
mately the  same  conditions,  and  probably  below  200°.  According  to 
Scholder,  removal  of  the  first  mole  of  water  is  not  possible  until  the  complex 


[CMS,  BASES,  AND  AMPHOTERIC  HYDROXIDES  443 

has  been  decomposed,  which  accounts  tor  the  high  temperature  required, 
Na,[Cu(OH)4]  — >  2NaOH  +  Cu(OH)., . 
him  blue 

Following  this  decomposition  the  amphoteric  hydroxide  readily  loses  one 
mole  of  water, 

Cu(OH)2— >CuO  +  H20 
blue  black 

The  second  mole  of  water  is  not  easily  liberated  because  of  the  extreme 
stability  of  sodium  hydroxide;  however,  at  much  higher  temperatures  a 
small  portion  of  this  water  is  gradually  lost  due  to  the  reaction 

CuO  +  2XaOH     >200°  >  Na2Cu02  +  H20 

This  is  supported  by  the  fact  that,  if  potassium  dichromate  is  mixed  with 
the  black  residue,  the  second  mole  of  water  is  readily  lost  at  approximately 
500°C, 

K2Cr207  +  2XaOH  -^1>  Na2O04  +  K2Cr04  +  H20  | 

Similar  dehydration  experiments  have  been  carried  out  with  other  hydroxo 
salts  (Table  12.12)  and  analogous  results  obtained.  Although  most  of  the 
hydroxoplumbates  are  unstable,  replacement  of  one  or  more  of  the  hydroxo 
groups  by  halide  or  thiocyanate  ions  increases  the  stability  of  the  complex, 
particularly  if  the  halogen  is  iodine.  The  fact  that  partial  replacement  of 
the  hydroxo  groups  by  other  anions  is  possible  is  further  evidence  that  the 
amphoteric  property  depends  upon  the  formation  of  complexes. 

Experiments  of  this  type  have  likewise  been  performed  in  the  presence 
of  pyrocatechol  and  crystalline  compounds  containing  both  hydroxo 
groups  and  pyrocatechol  groups  have  been  isolated102. 

It  may  be  that  the  dissolution  of  some  amphoteric  metallic  hydroxides  is 
a  result  of  peptization  and  that  in  other  cases  it  involves  the  formation  of 
true  compounds.  Seward103  has  pointed  out  that  in  many  cases  the  hydroxy- 
complex  theory  is  easily  reconciled  with  the  formation  of  colloidal  solutions. 
In  a  slightly  alkaline  solution  of  a  weak  metallic  hydroxide,  the  complexes 
formed  may  contain  a  large  number  of  metal  hydroxide  molecules,  a  few  of 
which  may  be  coordinated  to  additional  hydroxyl  ions.  A  complex  contain- 
ing, for  example,  one  hundred  molecules  of  metal  hydroxide  and  one  extra 
hydroxyl  ion  which  is  coordinated  to  a  metal  ion  would  constitute  a  colloidal 

102.  SchoklerandSchletz,Z.  anorg.  aUgem  Chem.,*ll,  161  (183 

103.  Seward.  ./.  Chem.  Ed.,  11,  567  (1934). 


444  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

particle.  In  Mich  a  case  the  amount  of  base  used  to  dissolve  the  precipitate 
is  small.  When  the  concentration  of  base  is  increased,  the  relative  number 
of  molecules  of  metal  hydroxide  containing  extra  hydroxyl  ions  will  in- 
crease until  the  coordination  number  of  each  of  the  metal  ions  is  approxi- 
mately satisfied.  Thus,  a  true  solution  will  form,  and  from  it  definite  com- 
pounds may  crystallize. 

A  consideration  of  the  available  data  indicates  that  a  much  more  gen- 
eralized definition  of  amphoteric  substances  is  required.  On  the  basis  of 
G.  X.  Lewis'  extended  acid-base  concept,  it  can  be  said  that  an  amphoteric 
substance  is  one  which  is  capable  of  either  donating  or  accepting  a  share  in 
a  pair  of  electrons.  An  application  of  this  principle  to  inorganic  amphoteric 
substances  suggests  that  they  are  complexes  which  are  capable  of  under- 
going both  of  the  following  reactions  to  such  an  extent  that  the  sign  of  the 
charge  on  the  complex  changes:  (1)  negative  or  neutral  ligands  may  replace 
neutral  or  positive  ligands  of  the  complex,  and  (2)  positive  or  neutral  ligands 
may  replace  neutral  or  negative  ligands  of  the  complex.  With  this  general 
interpretation  of  amphoterism  the  analogy  between  the  reactions  of  certain 
metallic  hydroxides  and  Werner  complexes  is  immediately  apparent  (Fig. 
12.0). 

The  chief  difference  between  the  reactions  of  the  zinc  complexes  and  those 
of  the  cobalt  complexes  is  that  the  equilibria  in  the  former  are  easily  re- 
versible, while  those  of  the  cobalt  complexes  can  be  made  to  go  in  either 
direction,  but  with  some  difficulty.  The  existence  of  easily  reversible  reac- 
tions in  the  case  of  the  zinc  complexes  makes  it  difficult  to  isolate  definite 
intermediate  compounds,  but  the  chemistry  of  the  more  stable  Werner 
complexes  is  well  defined  and  in  many  instances  it  has  been  possible  to 
isolate  all  of  the  intermediates  in  a  series  of  reactions  similar  to  that  given 
in  Fig.  12.6.  This  general  viewpoint  allows  a  better  understanding  of  ampho- 
terism in  any  system  than  the  older  concepts,  which  are  often  limited  to 
metallic  hydroxides  in  aqueous  medium. 


lZn(H20)2(OH)2l— 


[Co(NH3)3(NO, 


k  [Zn(H20)(OH)3]-  *=±  [Zn(OH)4]= 


(2)  (2) 


VL±  r^xx  ^   ^XT1+  J®-*  r^/TT   ^   1++ 


[Zn(H20)3OH]+  *==  [Zn(H20)4r 


(l)     l      *    -    "       '       (l) 
=*  [Co(NH3)2(N02)4]-  J=±  [Co(NH3)(N02)5]=  J= 


[Co(N02) 


-  [Co(NH3)4(N02)2]+  ^=^  [Co(NH3)5(N02)]++  <F^ 


(1)       l v • '"«-  *'*'         (1)       ■     ~-'      •'"      ™  (!) 

[Co(NH3)6]+++ 
Fig.  L2.6.  Equilibria  illustrating  the  general  principle  of  amphoterism. 


ACIDS,  BASES,  AND  AMPHOTERIC  HYDROXIDES  145 

Basic  Salts 
Structures  Based  on  the  Coordination  Theory 

Any  Bait  which  contains  an  oxide  or  hydroxide  group,  either  in  the  ionic 
or  coordinated  state,  is  referred  to  as  a  "basic  salt."  Many  basic  salts,  such 

as  white  lead  and  antinionyl  chloride,  are  of  90mewha1  indefinite  composi- 
tion, and  are  often  considered  to  be  simple  mixtures  of  the  "normal"  sail 
and  the  metallic  hydroxide.  Some  of  them,  however,  are  polynuclear  com- 
plexes, held  together  by  oxide  or  hydroxide  "bridges"  in  which  an  oxygen 
atom  is  coordinated  to  two  metal  atoms  (see  Chapter  13).  These  substances 
are  insoluble  in  water,  but  tend  to  be  hydrolyzed  by  it  ;  acid-  convert  them 
to  normal  Baits;  and  bases,  to  hydroxides.  These  reaction-  account  for  the 
variable  composition  of  precipitates  obtained  from  their  solutions.  The 
hydroxoammines,  however,  are  readily  obtained  as  crystalline,  water  solu- 
ble basic  salts  of  definite  composition.  They  are  formed  by  the  action  of 
3  on  aquoammines,  into  which  they  can  be  readily  reconverted. 

[Co(NH      IL<'\    .-   Co(NH,)6OH]X,  +  HX*.  « 

[PtNII      H.<>  0H)]C1,^  [PtiXII     .  <>H),]C1,  +  HCl39a 

There  is  little  possibility  of  bridging  in  these  cases  as  the  coordination 
sphere  of  the  metal  ion  is  completely  filled  by  groups  which  are  tightly 
held. 

Werner*   pointed  out  that  many  basic  salts  contain  three  moles  of  hy- 
droxide for  each  mole  of  normal  salt, 

atacamite  CuCl>-3Cu(OH)-> 

langite  CuS<  >;  :;(  u  OH)»HiO 

basic  zinc  nitrate  Zn  NO     r3Zn  0H)S 

basic  cobalt  carbonate  CoCO  -3Co(OH)j 

-   Its  PbX*-2Pb  0H)S  . 

The  amount  ot  water  present  in  any  basic  salt  is  almost  without  exception 
sufficient  to  permit  the  existence  of  the  hydrated  oxide  or  hydroxide  groups. 
In  cases  where  the  water  is  in  excess  of  thai   needed  to  form  hydroxide 
ipe  Werner  suggested  that  this  is  attached  to  the  "outer"  metal  atoms. 
The  structures  of  solid  basic  Baits  proposed  by  Werner  are  in  harmony 
with  experimental  -tudies  of  partially  hydrolyzed  .-alt-  in  solution  but  not 
in  complete  accord  with  the  results  of  x-ray  studies  of  these  salts76.  How- 
lidity  of  Werner'.-  views  has  been  justified  iii  certain  cases  by 
Weinland,  Stroh,  and  Paul104.  Their  measurements  of  the  electrical  conduc- 
tivity of  solutions  of  basic  lead  salts,  Pb(OH  X.  showed  the  presence  of  a 

int.  Weinland,  Stroh,  and  Paul,  Ber.,  55,  2706    1922  ;  X   mom.  aUgem.  Chem.,  129, 


4  Hi 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


bivalent  cation.  They  therefore  wrote  the  formula 


H 

o 

/   \ 
PI)  Pb 

\   / 
O 
H 


X: 


just  as  Werner  had  written 


Cu 


:cu 


ci2l 


co3, 


Potentiometric  studies  on  the  hydrolysis  of  uranium (VI)105,  bismuth(III)106, 
copper(II)107,  and  scandium(II)108  indicate  that  the  formation  of  bridged 
polymeric  cations  in  basic  solutions  of  metallic  ions  may  be  a  general  phe- 
nomenon (Chapter  13). 

The  hydrolysis  of  covalent  halides  probably  proceeds  through  the  addi- 
tion of  water  to  form  aquo  complexes  which  then  lose  protons  to  the  solvent. 
This  may  be  illustrated  by  the  hydrolysis  of  stannic  chloride.  Step  (1)  in- 
volves the  addition  of  water  to  satisfy  the  coordination  number  of  tin,  and 
step  (2)  is  a  hydrolysis  reaction  as  already  described;  in  step  (3)  the  stronger 
acid,  H30+,  displaces  the  weaker  acid,  [Sn  H20  Cl3  OH].  This  displaced  acid 
is  in  turn  capable  of  accepting  another  pair  of  electrons  which  are  donated 
by  a  molecule  of  water ;  the  process  is  repeated  until  the  final  hydrolysis 
product,  H2[Sn(OH)6],  is  obtained. 

(1)  [SnCl4]  +  2H20     hydration >  [Sn(H20)2Cl4] 


(2)  [Sn(H20)2Cl4]  +  H20 


hydrolysis 


[Sn  (H20)  Cl4  OH]-  +  H30+ 


(3)  [Sn(H20)Cl4  OH]-    dissociation  >  [sn  (H20)  CIa  OH]  +  Cl" 

(4)  tSn(H20)Cl3  OH]  +  H20    hydmtion  >  [Sn(H20)2Cl3  OH] 

In  the  particular  case  of  tin (IV)  chloride,  the  initial  addition  compound, 
[Sn  (1120)204] -3H20  and  the  product  of  the  second  stage  of  hydration, 

105.  Ahrland,  Acta  Chem.  Scand.,  3,  374  (1949). 

inc.  Graner  and  Silten,  Nature,  160,  715  (1947) ;  Graner  and  SilhSn,  Acta  Chem.  Scand., 
1,  631  (1947). 

107.  Pederaen,  Kgl.  Danske  Videnskab.  Selskab.  math.-fys.  medd.,  20,  No.  7,  24  pp. 

(1943)  ;cf.  Chem.Ab8.,38,  48545  (1944). 

108.  Kilpatrick  and  Pokrae,  J .' Electrochem .  Soc,  100,85  (1953). 


acids.  BASES,  AND  AMPHOTERIC  HYDROXIDES  447 

[Sn(HiO)iCUOH]-HA  are  both  known.  The  acid  nature  of  the  latter  com- 
pound, as  required  by  the  above  mechanism,  is  shown  by  the  formation  of 
a  salt  with  cineole10*.  The  postulated  mechanism  for  the  hydrolysis  of 
covalenl  halides  is  in  accord  with  the  fact  that  some  compounds  of  this 
type  in  which  the  coordination  number  of  the  central  atom  is  satisfied 
(CC1<  and  SFt)  are  very  difficultly  hydrolyzed. 

Structures  Rased  on  X-ray  Studies 

The  x-ray  studies  of  Feitknecht110  show  that  the  actual  structure  of  basic 
Baits  is  one  in  which  the  metals  are  equivalent  in  that  they  are  surrounded 
by  the  same  number  of  oxygen  groups.  In  the  structure  proposed  by 
Werner  only  the  central  metal  atom  is  coordinated  to  six  oxygens  while  the 
other  metal  atoms  form  a  part  of  the  chelate  donor  molecules  and  are 
attached  to  only  two  oxygens.  The  x-ray  data,  therefore,  need  not  be  re- 
garded as  a  contradiction  of  Werner's  views,  but  merely  as  a  modification 
to  the  more  logical  structure  in  which  the  metal  atoms  are  so  arranged  that 
they  all  tend  to  be  coordinately  saturated.  This  results  in  a  type  of  polymer- 
ization similar  to  that  found  in  silica,  each  macromolecular  sheet  of  the 
crystal  lattice  representing  a  polynuclear  complex  of  infinite  size. 

The  structure  suggested  by  Feitknecht  is  one  in  which  there  are  layers 
of  hydroxide  interleaved  with  layers  containing  the  metal  ions  and  acid 
anions.  For  bivalent  metals  the  layer  lattice  is  of  the  cadmium  iodide  type. 
The  spacing  between  the  layers  may  be  variable,  and  the  intermediate 
layers  may  be  almost  unordered  in  structure.  This  gives  rise  to  the  possi- 
bility of  nonstoichiometric  compounds,  which  are  formed  by  inserting 
different  amounts  of  metal  salt  into  the  intermediate  layers.  It  seems,  how- 
ever, that  these  double  layer  lattice  structures  are  metastable,  and  tend  to 
give  compounds  of  the  formula  MX-2-3M(OH)2  as  the  limiting  type.  In  such 
a  structure  the  hydroxide  layer  is  a  giant  molecule  which  permits  varying 
amounts  of  water  as  well  as  normal  salt  and  which  is  insoluble;  these  are  all 
characteristic  properties  of  basic  salts. 

109.  Pfeiffer  and  Angera,  Z.  anorg.  Chem.,  183,  189  (1929). 
IK).  Feitknecht,  Helv.  Chim.  Acta,  16.  427  (1933). 


IvJ.   Olation  and  Related  Chemical 
Processes 

Carl  L  Rollinson* 

University  of  Maryland,  College  Park,  Maryland 


Olated  compounds  are  complexes  in  which  metal  atoms  are  linked  through 
bridging  OH  groups.  Such  a  group  is  designated  an  ol  group1  to  distinguish 
it  from  the  hydroxo  group;  i.e.,  a  coordinated  OH  linked  to  only  one  metal 
atom.  The  process  of  formation  of  ol  compounds  from  hydroxo  compounds 
is  called  olation.  In  a  review  of  the  theory  of  olated  compounds,  Basset2 
gives  the  following  examples: 


r0H\     \ 
Cu 

OhT        L 


OH 


CI; 


/0H\ 

(NH3)4Cc/       ^Co(NH3)4 
OH 


4  + 


(NH3)3Co^OH-^Co(NH3)3 
OH^ 


+++       i— 


/0H\ 


(en)2Co  Xo(en)2 

OH 


4  + 


Chromium  complexes  analogous  to  the  above  cobalt  complexes  display 
remarkably  similar  properties. 

Olation  is  often  followed  or  accompanied  by  oxolation  or  anion  penetra- 
tion or  both.  Oxolation  is  conversion  of  ol  groupstobridgingo.ro  groups; 
each  ol  group  loses  a  proton.  Anion  penetration  consists  of  replacement  of  a 
coordinated  group,  such  as  an  anion  or  a  hydroxo,  aquo  or  ol  group  by 
another   anion. 

Mr.  Harold  .).  Matsuguma  helped  in  the  preparation  of  this  chapter.  His  assist- 
ance is  gratefully  acknowledged. 
1     WCrncr,  Ber.,  40, 2113  (1907). 
_'.  Basset,  Quart.  Rev.,  1,  246  (1947). 

448 


OLATIOh    AND  RELATED  CHEMICAL  PROCESSES 


I  I'.i 


The  N  \ii  re  \\i)  Significance  of  Olation 

PfeillVr  observed  that  a  blue-violet  compound  is  formed  when  the  red 
hydroxo-aquo-bis(ethylenediamine)  chromium  (III)  chloride  is  heated  at 
L20  C.  He  had  previously  suggested4  that  one  coordinate  bond  of  cadi  of 
two  metal  atoms  might  be  shared  by  one  OH  group,  and  therefore  formu- 
lated the  reaction  as  follows: 


(en)  j  Cr 


/ 


OH 


II, o 


CI,  + 


H»0 


IK) 


Cr  (en), 


CI, 


120 


2  moles  of  the  red  sail 


OH 

/    \ 
(en),  Cr  Cr  (en) o 

\    / 

OH 

blue-violet  salt 


Cl<  +  2H«() 


It  is  evident  that  a  reaction  of  this  type  could  occur  readily  only  with  a  cis 
salt,  since  the  trans  salt  would  have  to  rearrange.  The  red  salt  is  thus  as- 
signed the  cis  configuration5.  Werner6  prepared  octammine-/x-diol-dico- 
balt(III)  sulfate  by  a  similar  reaction: 


Ml 


OH 


OH. 


II,  o 


HO 


Co<XH.,)4 


S04 


OH 

/    \ 
XII    4C0  Co(NH8)4 

\    / 
OH 


Si  >4),  +  2H,0 


The  following  hexol1,7  is  of  interest   because  it  is  a  completely  inorganic, 

3.  Pfeiffer,  Z.  anorg.  Chem.,  56,  261  (1907  . 
I.  Pfeiffer,  Z.  anorg.  Chem.,  29,  107  (1901 

.").  Emele'iu  and  Anderson,  "Modern  Aspects  of  Inorganic  Chemistry,"  p.  89,  Wu 
Y*ork,  I).  Van  Nostrand  Company,  Inc.  Lfl 

6.  Werner,  Bt  r.,  40,  1437    L907 

7.  Jorgensen,  Z.  anorg.  Chem.,  16,  184    1897 


450  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

optically-active  compound;  it  was  resolved  by  Werner8: 


Col         ^Co(NH3)4 
\0\< 


In  this  compound  each  of  three  tetrammine-dihydroxo-cobalt(III)  com- 
plexes acts  as  a  bidentate  chelating  group. 

The  examples  given  may  be  represented  conventionally  as  follows: 


NH3  NH3 

H 


H3N 
H3N 


ZZZTD 

L , — Ln-L ■ — lh 


O 

H 

NH3 


NH3 
NH3 


NH: 


J\         \zi  [Co(NH3)4(OH)2 


NH2CH2CH2NH2 


The  other  possible  bridging  arrangements  for  two  octahedral  atoms  are 
linkage  by  one  and  by  three  bridging  groups;  three  is  the  maximum  number 
of  bridges  attainable  since  two  octahedra  can  have  only  one  face  in  com- 
mon: 


Moreover,  more  than  two  atoms  can  be  linked  chain-wise: 


h      h  r 


Gmelin's  uHandbuch"9  contains  an  excellent  summary  of  bridged  cobalt 

S.  Werner,  Ber.,  47,  3087  (1914);  Compt.  rend.,  159,  426  (1914). 

9.  Gmelin,  "Handbuch  der  anorganischen  Chemie,"  Teil  B,  S.N.  58,  pp.  332-374, 
Berlin,  Verlag  Chemie,  G.m.b.h.,  1930. 


OLATION  AND  RELATED  CHEMICAL  PROCESSES  451 

compounds,  including  those  in  which  the  ol  group  is  the  only  bridging  group, 

and  those  in  which  the  ol  group  and  sonic  other  group,  such  as  peroxo  or 
oitro,  act  as  bridges.  The  summary  includes  complexes  containing  as  many 

as  four  cobalt  atoms.  Similarly,  Mellor1"  lists  polynuclear  chromium  com- 
pounds of  from  two  to  four  chromium  atoms. 

the  "Continued"  Process  of  Olation 

Instead  of  reaching  a  definite  termination,  as  in  the  reactions  just  men- 
tioned, the  olation  process  may  continue,  with  the  formation  of  polymers. 
This  may  occur  if  the  product  of  each  successive  step  contains  aquo  or 
hydroxo  groups.  Although  much  of  the  evidence  regarding  such  polymers 
is  indirect,  Werner's  theory,  as  extended  by  Pfeiffer,  Bjerrum,  Stiasny,  and 
others,  has  been  consistently  successful  in  accounting  for  the  experimental 
observations  and  predicting  the  behavior  of  these  compounds. 

The  continued  process  of  olation  starts  with  the  hydrolysis  of  salts  of  such 
metals  as  aluminum  and  chromium.  Pfeiffer11  suggested  that  the  acidity  of 
solutions  of  such  salts  is  due  to  conversion  of  aquo  to  hydroxo  groups,  e.g.: 

[Cr(H20)6]+++^  [Cr(H20)5(OH)]++  +  H+ 
[Cr(H20)5(OH)r+;=±  [Cr(H20)4(OH)2]+  +  H+ 

The  degree  of  hydrolysis  increases  as  the  temperature  is  raised12.  It  is  also 
dependent  on  the  nature  of  the  anion12,  and  especially  on  the  pH  of  the  solu- 
tion. 

If  alkali  is  added  to  a  warm  solution  of  such  a  salt,  but  not  enough  for 
complete  neutralization,  polymerization  occurs  instead  of  the  precipitation 
of  the  basic  salt  or  hydroxide.  For  example,  Bjerrum14  showed  that  ag- 
gregates up  to  colloidal  dimensions  are  formed  when  basic  chromic  chloride 
solutions  are  heated,  and  similar  results  have  been  obtained  by  other  in- 
vestigate] 

These  results  may  be  explained  on  the  basis  of  a  series  of  hydrolytic  and 

in.  Mellor,  "A  Comprehensive  Treatise  of  Inorganic  and  Theoretical  Chemistry," 

Vol.  11.  j)]).  407-0.  London,  Longmans  Green  and  Co.,  Ltd..  1931. 
11.  Pfeiffer,  £«r.,  40, 4036  (1907). 

illgren,  Z.  pkys.  Chen-  .85,  406  (1913) . 

13.  Cupr,  Collection  Czechoslov.  Chem.  Communs.,  1,  167  (1929);  cf.,  Chem.  Ah*.,  24, 

1013  (1930). 

14.  Bjerrum,  Z.  phys.  Chem.,  59,  336  (1907);  73,  724  (1910);  110,6.56  (1921 

15.  Riess  and  Barth.  Collegium,  778,  62  (1935). 


152 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


olation  reactions18.  The  first  steps  might  be  formulated  as  follows: 


/ 


H20 


(H,0)4Cr 


(H20)4Cr 


OH 


H20 


+ 


H20. 

II, () 


HO 


OH 


(H40)4Cr 


\ 


H20 


+  H+ 


Cr(H20), 


OH 

/    \ 
(H20)4Cr  Cr(H20)4 


OH      H20 


+  H20 


If  the  reacting  groups  in  each  ion  are  in  the  cis  positions,  a  completely  olated 
ion  may  be  formed: 


OH 

/    \ 
(HoO)4Cr  Cr(H20)4 

I  I 

OH      H20 


OH 

/    \ 

(H20)4Cr  Cr(H20)4  f-  H20 

\    / 
OH 

Further  hydrolysis  and  olation  might  result  in  such  polymers  as  the  tetra- 
hydroxy-dodecaquo-)U-decaol-hexachromium(III)   ion : 

-i  4+ 
H20  H20  H20  Hp  H20  H20 

H0\  J   ^0H\  I      ^0I_K  ^°H^  /°H^  /OH>^  /OH 


.Cr 


Cr 


Cr 

HCT         XOHX       ^OH^       ^OH 

HP  H20  H20  H20 


Cr  Cr'  /CrV 

^OH^  I  ^OH^  I   ^OH 


H20 


H,0 


Stiasny17  postulated  the  existence  of  such  polymers  in  Bjerrum's  solutions14 
and  aggregates  of  ionic  weight  400 — 1000  have  been  detected  in  such  so- 
lutions15. The  possibility  of  formation  of  the  following  types  of  olated  com- 
pounds must  also  be  admitted  (A  =  a  coordinated  molecule  or  ion) : 


•OH 


a4 —  Cr 

HO 

\ 

A4 Cr 


OH- 


n  + 

Cr A4 

OH 

/ 

Cr A4 

A. 


OHv  .OH. 

XC^       N^       XCr 


-im+ 

>H\I>H\ 

r  Cr 


These  processes  involve  the  aquo  groups  attached  to  the  metal  atoms 

1(>.  Stiasny  and  Balanyi,  Collegium,  682,  86  (1927). 

17.  Stiasny,  "Gerbereichemie",  p.  348,  Dresden  und  Leipzig, TheodorSteinkopff,  1931. 


OLATIOA   AND  RELATED  CHEMICAL  PROCESSES  153 

not  at  the  ends  of  the  chain,  as  well  as  those  at   the  ends,  so  cross-linked 

polymers  may  he  formed,  as  shown  in  the  diagrammatic  formula: 

HO     OH         HO     OH  HO     OH         HO     OH 

,        N         V^OH^/       OH^\/^OH     \/ 

(H20)2 C<"  ^Cr^  ^Cr^  J>  (H20)2 

_  X)H^        ^OH^         ^OhT  _ 

Because  y^i  the  octahedral  configuration  of  complexes  of  metals  such  as 
chromium,  the  bonds  of  a  given  metal  atom  occur  in  pairs,  each  of  which 
lies  in  a  plane  perpendicular  to  the  planes  of  the  other  two  pairs.  Thus  such 
-linked  polymers  are  three-dimensional. 

These  processes  account  for  the  results  obtained  when  a  warm  solution 
of  a  chromium  salt  is  titrated  with  a  base. With  the  addition  of  an  incre- 
ment of  base,  the  pll  rises  immediately,  but  falls  slowly  if  the  solution  is 
allowed  to  stand  before  more  base  is  added.  This  continues  with  successive 
increments  of  base  until  enough  base  has  been  added  to  precipitate  the 
hydrated  oxide.  As  base  is  added  to  the  solution,  the  hydrogen  ions  are 
removed.  The  equilibrium  then  shifts  in  the  direction  of  further  hydrolysis 
and  elation,  with  the  formation  of  more  hydrogen  ions.  In  this  way  an 
amount  of  base  can  be  added,  without  precipitation,  which  would  cause 
precipitation  if  it  were  added  all  at  once. 

The  changes  in  pH  accompanying  the  titration  of  scandium  perchlorate 
with  base  cannot  be  explained  by  the  formation  of  hydroxo  or  dihydroxo 
monomers  alone18.  The  data  obtained  are  consistent  with  the  assumptions 
that  a  monomeric  monohydroxo  compound  is  formed  and  is  in  equilibrium 
with  dimeric,  trimeric,  and  more  highly  aggregated  species.  Kilpatrick  and 
Pokras18  obtained  the  equilibrium  constants  for  the  reactions 

[Sell  (0    <>II,]+++^  [Sc(H,()),()Iir+  +  H+ 

2[Sc(H  0    <»II]++^  [Sc(H20)4OH]2^ 

They  found  that  these  two  reactions  predominate  during  the  addition  of  the 
first   0.3  equivalent   of  base,  but    that   the  addition  of  more  base  leads  t<> 
higher  aggregation.  Gran6r  and  Sillen1-'  found  similar  behavior  in  the  case 
of  bismuth  perchlorate. 
Gustavson20  conducted  a  Beries  of  studies  on  the  chromium  complexes 

In.   Kilpatrick  and  Pokras,  /.  Electrochem.  So,-..  100,  85  (19 

Jran^r  and  Silten,  Acta  Chem.  Scand.,  1,  631     1947  ;  Nature,  160,  715    1947  . 
20.  Gustavson,  •/.  Am.  Leather  Chem.  Assoc.,  47,  151     1952);  44,  388     I'M'-  ;  ./    Coll. 

i.     •: 


454  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

involved  in  leather  tanning.  By  ion  exchange  methods  he  found  that,  in 
strongly  basic  solutions,  30  per  cent  of  the  chromium  complexes  were 
cationic  and  70  per  cent  noncationic.  Since  electrophoresis  showed  no 
negative  complexes,  he  concluded  that  neutral  complexes  predominate  in 
such  solutions.  He  established  the  empirical  formula  of  these  to  be 
[Cr2(OH)5Cl]°.  It  was  also  found  that  hydroxo  and  ol  compounds  lead  to 
cationic  complexes  in  highly  acid  solutions.  Electrophoresis  of  such  chro- 
mium solutions  showed  the  presence  of  complexes  of  very  low  or  negligible 
ionic  mobility.  Extremely  basic  chromium(III)  chlorides  also  contain 
components  having  little  or  no  ionic  mobility.  Gustavson  subjected  these 
basic  chromium  solutions  to  dialysis  for  four  weeks,  and  upon  analysis  of 
the  dialysate  he  found  that  91  per  cent  of  the  chromium  had  been  removed. 
The  remainder  of  the  chromium  was  present  in  the  form  of  compounds 
having  the  approximate  formula  [Cr405(H20)i2Cl2].  The  average  molecular 
weight  was  found  to  be  600. 

In  another  investigation  of  chromium  complexes,  Gustavson20c  carried 
out  the  separation  and  quantitative  determination  of  cationic,  anionic,  and 
neutral  complexes  in  solutions  of  basic  chromium  chlorides  and  sulfates  by 
filtering  them  through  layers  of  cation  and  anion  exchange  organolites. 
He  reports  the  existence  of  the  following  species:* 

Cr2(OH)2Cl4  Cr2(OH)2Cl4-NaCl        Cr2(OH)3Cl3 

Cr4(OH)7Cl5  Cr4(OH)9Cl3  Cr4(OH)2(S04)5 


Cr2(OH)2(S04)2        Cr4(OH)6(S04)3 

ft 

In  preparing  his  solutions,  he  boiled  the  appropriate  chromium  salts  with 

sodium  carbonate  to  effect  a  gradual  change  in  the  pH  and  the  gradual 

olation  of  the  various  complexes.  It  was  found  that  most  of  the  complexes 

present  in  basic  chromium  sulfates  or  ordinary  sulfates  are  of  the  form20b: 

[Cr2(OH)2(S04)3]=        or        [Cr2(OH)2(S04)]++ 

Castor  and  Basolo21  have  applied  a  kinetic  technique  to  the  study  of 
heterogeneous  dehydration  of  hydrated  salts,  and  were  able  to  identify 
hydrates  intermediate  between  those  found  by  thermodynamic  methods. 
Thus,  in  addition  to  the  4-,  2-,  and  1 -hydrates  of  manganese(II)  chloride, 
they  identified  a  3.5-  and  a  3-hydrate.  Complete  dehydration  yields  the 
anhydrous  metal  chloride.  However,  in  the  case  of  zirconyl  choride  8-hy- 
drate,  dehydration  was  shown  to  proceed  through  7.75-,  7.5-,  7-,  and  6.5- 

*  In  these  formulas,  and  others  in  this  chapter,  the  possible  presence  of  coordi- 
nated water  molecules  is  disregarded.  It  is  probable  that  all  of  the  complexes  dis- 
cussed contain  at  least  enough  water  to  fill  the  coordination  spheres  of  the  metal  ions. 
21.  Castor  and  Basolo,  J.  Am.  Chem.  Soc,  75,  4804,  4807  (1953). 


OLATION  AND  RELATED  CHEMICAL  PROCESSES 


455 


hydrates  to  the  6-hydrate;  complete  dehydration  involves  hydrolysis  and 

produces  zirconium  dioxide.  Fractional  hydrate  formation  was  explained 
on  the  basis  of  the  reactions: 


2[RM  oil     II,<))K  ^ 


II 
(  I 

/    \ 
_B  H»0  M  M  OB  R_ 


•-•»• 


•Mo 


H 
O 

/    \ 
R  H  «»  M  M  OB  R 


R— M 


M— R 


-     Ho 


2    R(OH)M 


M(H>0)R 


(c) 


M— R 


+  H20 


H  H 

O      R'    H     R'     0 

/   \l        I        1/   ' 
R— M  M — 0 — M 

\    /  \   . 

O  O 

H  H 

Denk  and  Bauer22  found  that  when  aluminum  reacts  with  a  deficiency 
of  dilute  hydrochloric  acid,  six  times  as  much  aluminum  is  dissolved  as  is 
required  for  the  formation  of  simple  aluminum  chloride,  A1C13  ,  and,  from 
the  resulting  solutions,  they  isolated  the  "%  basic"  aluminum  chloride  in 
-table  form.  This  compound  is  soluble  in  water  and  shows  weak  x-ray 
patterns.  The  %  basic  sulfate,  [Al2(OH)5]2S04 ,  was  isolated  by  precipita- 
tion with  sodium  sulfate.  Denk  and  Bauer  also  found  that  the  basic  chloride 
reacts  slowly  with  more  aluminum  to  give  a  colloidal  product. 

Factors  Promoting  Olation 

Several  methods  have  been  suggested  for  the  measurement  of  the  degree 
of  olation,  but  none  is  entirely  accurate.  Stiasny  and  Kdnigfeld*  assumed 
that  olated  hydroxo  groups  do  not  readily  react  with  excess  acid  in  the 
cold,  but  do  react  when  boiled  for  an  hour  with  exec--  arid.  Back  titration 
of  the  excess  in  the  two  cases  measures  the  degree  of  olation.  Theis  and 
Serfass*  found  that  conduct ometric  titrations  give  more  accurate  and  more 

22.  Denk  and  Bauer,  Z.  mnorg.  allg*  »>.  Ckt  ».,  267,  SO  1 1951). 
-      Stiasny  and K5nigf eld  wi, 781,807    Ifl 

24.  Theis  and  Serfass,  J  iher  Chi  n 


156 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


reproducible  results  than  potentiometric  methods  or  those  using  indicators. 
Mitchell26  determined  the  Dumber  of  olated  groups  from  the  difference  be- 
tween the  number  of  equivalents  of  alkali  added  to  the  solution  when  first 
prepared,  and  the  number  of  equivalents  of  acid  needed  to  bring  the  pH 
to  3.3. 

In  recent  studies,28b  Mitchell  found  that  the  degree  of  olation  decreased 
from  100  to  50  per  cent  with  the  addition  of  increasing  amounts  of  sodium 
hydroxide  to  freshly  prepared  solutions  of  chromium  alum,  but  it  decreased 
only  to  75  per  cent  of  its  original  value  when  aged  solutions  of  the  alum 
were  used.  She  also  found  that  solutions  of  chromium  alum  boiled  with 
sodium  hydroxide  exhibited  100  per  cent  olation.  Complexes  of  the  compo- 
sition [Cr4(OH)3(S04)2(H20)i2]+++  were  formed  by  boiling,  cooling,  and 
aging  solutions  of  chromium  alum  for  fifteen  minutes.  In  these  solutions, 
there  was  a  stoichiometric  relationship  between  the  formation  of  olated  OH 
groups  and  the  entry  of  sulfate  groups  into  the  complexes.  If  the  hexaquo- 
chromium(III)  ion  was  heated  with  alkali  of  the  correct  concentration,  one 
ol  bridge  formed  and  one  sulfate  entered  the  complex.  However,  if  the  con- 
centration of  alkali  was  great  enough,  two  ol  bridges  formed  and  the  sulfate 
groups  were  eliminated  from  the  complex. 

The  process  of  olation  is  favored  by  an  increase  in  concentration,  an 
increase  in  temperature,  and  especially,  by  an  increase  in  basicity.  The  proc- 
ess reverses  very  slowly  when  solutions  of  olated  complexes  are  diluted,  or 
when  such  solutions  are  cooled;  i.e.,  olation  decreases  the  reactivity  of  co- 
ordinated OH  groups26. 


The  Oxolation  Process 

Stiasn}'  and  co-workers16, 27  observed  that  solutions  of  basic  chromium 
salts  become  more  acidic  and  the  salts  less  soluble  when  the  solutions  are 
heated.  When  the  solutions  are  cooled,  the  acidity  drops  to  the  original 
value,  but  only  after  a  long  time.  To  account  for  these  facts,  Stiasny  sug- 
gested the  process  of  oxolation;  i.e.,  conversion  of  ol  groups  to  oxo  groups  by 
the  loss  of  protons: 


on 
/  \ 

(H20)4Cr  Cr(H20) 

\    / 

on 


o 

/  \ 

(Ho())4Cr  Oidl.O), 

\    / 
0 


+  2H 


This  appears  to  be  a  resonable  explanation,  especially  in  view  of  the  acid 

25.  Mitchell, ./.  Soc.  Leather  Trades'  Chem.,  35,  154,  397  (1951). 

26.  Werner,  Ber.,  40,  1436  0907). 

27.  Stiasny  and  Grimm,  Collegium,  691,  505  (1927);  694,  49  (1928). 


OLATION  AND  KBLATBD  CHEMICAL  PROCESSES  467 

•cactinii  of  the  "erythro"  chromium  salts10;  the  equilibrium 


<)— C, All      J       ^  [(XII,) 


,  Ml     Ci — 0— Cr(NH  -O— Cr(NH 

may  be  involved. 
While  olation  and  oxolation  are  both  reversible,  the  long  time  required 

for  the  acidity  of  solutions,  which  have  been  hc;itcd  and  then  cooled,  to 
return  to  the  original  value  leads  to  the  conclusion  thai  de-oxolation  is 
extremely  slow.  In  general,  ol  compound-  are  more  readily  depolymerized 
than  oxo  compounds,  since  protons  react  more  rapidly  with  ol  groups  than 
with  oxo  groups. 

Jander  and  Jahr18  found  that  the  addit  ion  of  base  to  Bolul  ions  of  iron(III) 
perchlorate  cause-  the  formation  of  hydroxo  and  finally  oxolated  bi- 
molecular  hydrolysis  products,  which  they  formulated  as  follows: 

2[Fe(OH)(C104)(H20)]+^  [(C104)Fe— O— Fe(C104)(H20)]++  +  211  u 

The  addition  of  more  base  leads  to  such  products  as: 

[(C104)Fe— O— Fe— O— Fe— O— Fe— O— Fe(C104)2(H20)"|- 
III 
C104        C104        CIO4  J 

Jander  and  Jahr28  also  found  that  the  addition  of  one  mole  of  ammonia 
to  one  mole  of  aluminum  nitrate  causes  the  formation  of 

[Al(OH)(X03)2(H20)]m  . 

A  second  mole  of  base  causes  the  formation  of  the  oxolated  aggregate, 
[Al-0'(NOj)]«  .  These  reactions  were  represented  in  a  manner  analogous 
to  that  used  by  Thomas  and  Whitehead-'.  Jander  and  Jahr  report  that  the 
addition  of  more  base  does  not  cause  the  precipitation  of  aluminum  hy- 
droxide, but  increases  the  degree  of  aggregation: 

I  NO     a:  IL<)i-0-A1(X03)(OH)(H20)1  +  l(OH)Al  NO,  ,  IL<»d  — 

[H20(X03)2A1— O— A1'N< »     B,0)— O— Al(NO,),(B  I  I 

__    _    ion  would  lead  to  the  formation  of  such  condensation 
prod 


Al— O— Al  ..-  O— Al  ...  O— A1(\U;)2(H20)"| 
NO  NO  NO  J 


Similar  reactions  take  place  in  solutions  of  zirconium  perchloral 

:ider  and  Jahr,  KoUoid  BeihefU,  43,  :;.':;.  306    1936). 
ad  Whitehead,  •/.  /  16,  27      131  . 

30.  Refen-i       28        315. 


458  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Hall  and  Eyring31,  in  a  study  of  the  constitution  of  chromium  salts  in 
aqueous  solutions,  found  that  ammonium  paramolybdate,  (NH4)6Mo7024- 
4H20,  is  effective  in  precipitating  chromium  complexes.  They  report  that 
the  HMo04~  anion  penetrates  into  the  complex  and  displaces  the  OH  groups 
and  aquo  groups,  but  it  does  not  affect  the  ol  groups.  They  also  found  that 
the  process  of  oxolation  is  facilitated  by  the  addition  of  90  per  cent  ethanol. 
They  suggest  that  the  competition  between  the  alcohol  and  the  chromium 
for  the  aquo  and  hydroxo  groups  leads  to  the  loss  of  protons  from  the  ol 
bridges  with  the  formation  of  oxo  bridges.  Their  work  also  seems  to  show 
that  there  is  a  greater  amount  of  oxolation  than  Stiasny  postulated. 

Kuntzel32  is  also  in  partial  disagreement  with  Stiasny.  He  found  that  a 
J^  basic  chromium  chloride  solution  contains  only  single  ol  bridges  which 
give  rise  to  long  chain  colloidal  aggregates.  Upon  aging,  these  aggregates 
break  up  into  smaller  groups  which  contain  two  or  three  ol  bridges  joining 
each  pair  of  chromium  atoms.  Stiasny  proposed,  on  the  other  hand,  that 
the  aging  process  causes  oxolation  of  the  long,  large  aggregates. 

Anion  Penetration 

A  number  of  investigators  have  shown  that  the  addition  of  neutral 
salts  to  solutions  of  basic  chromium,  iron,  or  aluminum  sulfate  changes  the 
hydrogen  ion  concentration33.  Different  anions  were  found  to  differ  in  their 
effectiveness  in  this  respect.  Early  explanations  were  based  on  hydration 
and  the  formation  of  addition  compounds34.  Stiasny,  however,  explained 
the  phenomenon  by  postulating  "anion  penetration,"  i.e.,  replacement  of  a 
coordinated  group,  such  as  aquo,  hydroxo,  or  an  anion,  by  an  anion.  Re- 
actions of  this  type  are  common  among  complexes  of  low  ionic  weight. 
When  a  solution  of  the  violet  form  of  chromium  (III)  chloride  6-hydrate 
is  heated,  the  bright  green  form  (tetraquo)  is  produced35: 

[Cr(H20)6]Cl3  v  *"**  s  [Cr(H20)4Cl2]Cl  +  2H20 

•       1      i  C°°1 

violet  green 

In  pure  water  the  reaction  reverses  slowly  when  the  solution  is  cooled,  but 

31.  Hall  and  Eyring,  /.  Am.  Chem.  Soc,  72,  782  (1950). 

32.  Kuntzel,  Colloquimsber .  Insts.  Gerbereichem.  tech.  Hochschule  Darmstadt,  No.  2, 

31  (1948);  cf.,  Chem.  Abs.,  43,  1591a  (1949). 

33.  Stiasny  and  Szego,  Collegium,  670,  41  (1926);  Wilson  and  Kern,  ./.  Am.  Leather 

Chem.  Assoc.,  12,  450  (1917);  Wilson  and  Kuan,  ibid.,  25,  15  (1930);  Thomas, 
Paper  Trade  J.,  100,  36  (1935). 

34.  Wilson  and  Gallun,  J.  Am.  Leather  Chem.  Assoc,  15,  273  (1920);  Thomas  and 

Foster,  Ind.  Eng.  Chem.,  14,  132  (1922). 

35.  Ephraim,  "Inorganic  Chemistry,"  p.  291,  New  York,  Nordeman  Publishing  Co., 

Inc.,  1939;  Mellor,  "Modern  Inorganic  Chemistry",  p.  776,  New  York,  Long- 
mans Green  and  Co.,  1939. 


OLATIOh    AND  RELATED  CHEMICAL  PROCESSES 


I.V.I 


the  extent  of  reversal  is  decreased  by  sodium  chloride.  Lamb"  states  that 
all  of  the  chloride  can  be  precipitated  from  chromiumdll)  chloride  solu- 
tions by  silver  acetate,  but  not  by  silver  nitrate.  The  acetate  ion  can  dis- 
place chloro  groups  from  the  complex  chromium(III)  ion,  but  the  nitrate 

ion  cannot.  This  is  in  accordance  with  the  well-known  difference  in  the 
coordinating  power  of  these  groups. 

Stiasny  postulated  thai  an  equilibrium  exists  between  the  complex  cat- 
ion of  a  basic  chromium  salt  and  the  anion.  This  equilibrium  is  .shifted  by 
changing  the  relative  concentrations  of  anion  and  chromium  complex. 
The  following  examples  indicate  why  the  pll  is  changed  by  such  reactions: 

-|-         2Cf 


H,0 

H,0           " 

Cr 

H,0 

HzO          J 

H,0. 


H,0 
H*0 


■OH, 
OH 


H.O  _J 


-|-20H" 


(H20)4   Cr 


.OH. 


Cr(H?0). 


+       2C.-^^ 


-,   A  + 


(H20)4Cr^  ^Cr(H?0). 


-\-  20H" 


The  extent  to  which  anion  penetration  occurs  with  ol  complexes  is  de- 
termined by  the  relative  concentrations  of  the  reactants,  the  relative  co- 
ordinating tendencies  of  the  entering  anion  and  the  group  which  it  dis- 
places, and  the  length  of  time  which  the  solutions  are  allowed  to  stand37. 
Anions  that  can  enter  the  coordination  sphere  easily  and  displace  OH  groups 
effectively  prevent  olation.  Penetration  by  anions  decreases  in  the  order: 
oxalate  >  citrate  >  tartrate  >  glycolate  >  acetate  >  formate  >  sulfate. 
In  stock  solutions  of  basic  chromium(III)  sulfate,  however,  Serfass,  et  al.zl 
found  ionic  species  having  weights  of  68,000. 

Shuttleworth38,  in  studying  the  bond  forces  involved  in  chrome  tanning, 
examined  a  series  of  complex  chromium  ions  by  means  of  ion  exchange 
resins,  potentiometric  titrations,  and  spectrophotometries  curves.  He  ob- 
tained most  of  the  compounds  that  he  used  from  [CT,(H,())6(OH)2(S04)]++, 
which  will  be  called  (a)  in  the  following  discussion. 

By  boiling  a  solution  of  (a)  (96  g.  of  chromium  ion  per  liter)  with  stoichio- 
metric proportions  of  sodium  oxalate  and  then  aging  for  one  week  he  ob- 
tained 


[(Y,aiA)6(OH)2(C204)]+  +  ,  [Cr2(H20)4(()IlM(    «  I 


and 


|(VIU)),a)IlM(V>:>:;]  =  . 

36.  Lamb.  ./.  .1//-.  Ckem.  8oc.,  28,  1710  (1006);  Weinland  and  Koch,  Z.  anorg.  Chem., 

39,  2:><;    1904  . 

37.  Serfass,  Tin-is.  Thorstensen,  and  Agarwall,  •/.  Am.  Leather  Cfu  m.  Assoc.,  43,  132 

1948). 

38.  Shuttleworth,  J.  Am.  Leather  Chi  m.  Assoc.,  47,  :;s7  (1952). 


4(50 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


When  (a)  (a1  t  he  same  concentration)  was  warmed  at  37°  for  24  hours  with 
proportional   amounts  of  sodium  sulfite,   and  aged  for  one  week, 

|(  !r2(H20)6(OH)2(S03)]-"         and        [Cr2(H20)4(OH)2(S03)2]0 
were  obtained.  When  sodium  formate  instead  of  the  sulfite  was  used, 


[Cr2(H20)B(OH)2(HC02)J 


and         [Cr2(H20)4(OH)2(HC02)4]0. 


were  produced. 

Making  use  of  conductimetric,  potentiometric,  and  diffusion  measure- 
ments, Jander  and  Jahr  have  found  that  the  most  abundant  ionic  species 
present  in  solutions  of  beryllium  nitrate  is  [Be(H20)N03]+  39.  As  the  solu- 
t  ions  age,  the  pH  decreases,  apparently  due  to  the  replacement  of  the  nitrate 
in  the  complex  by  hydroxo  groups.  The  resulting  hydroxo  complex  was 
thought  to  be  capable  of  dimerizing: 

2[Be(H20)  (OH)]+^±  [(H20)Be— O— Be(H20)]++  +  H20 

The  concentration  of  this  condensation  product  increases  with  decreasing 
pH  until  almost  all  of  the  beryllium  is  present  in  the  form  of  dimeric  cations. 
Thorstensen  and  Theis40  have  studied  the  effect  of  adding  sodium  citrate 
to  solutions  of  basic  iron(III)  salts  used  in  iron  tannage.  They  found  com- 
pounds having  the  following  empirical  formulas: 

[Fe2  (S04)  (OH)  2]  •  Na-citrate 

[Fe2(OH)4]-Na-citrate 

[Fe2(S04)2(OH)2(OCH2C02Na)2]= 

[Fe2(S04)(OH)4(OCH2C02Na)2]= 

[Fe2(OH)6(OCH2C02Na)2]= 

Chelation  as  a  Factor  in  Anion  Penetration 

Since  displacement  of  OH  groups  from  the  complex  ion  involves  coordi- 
nation of  the  displacing  group  with  the  central  metal  ion,  the  reactivity  of 
various  anions  should  be  determined,  in  part,  by  the  number  of  donor 
groups  in  the  anion  and  their  relative  positions.  Thomas  and  Kremer41 
compared  the  effects  of  potassium  salts  of  aliphatic  monocarboxylic  acids, 
from  formate  to  valerate  inclusive,  and  of  aliphatic  dicarboxylic  acids, 
from  oxalate  to  pimelate  inclusive.  The  differences  in  effectiveness  oi  the 
homologous  monocarboxylic  anions  is  very  slight.  This  might  be  expected, 
since  coordination  of  these  anions  with  the  metal  of  the  complex  cation  is 
presumably  controlled  by  the  single  carboxyl  group. 

39.  Reference  28,  p.  301. 

m    Thorstensen  and  Theis,  •/.  .1///.  Leather  ('hem.  Assoc,  44,  841  (1949). 

II.  Thomas  and  Kremer,  J.  Am.  Chem.  Soc.,57,  1821,2538  (1935). 


OLATIOh   AND  RELATED  CHEMICAL  PROCESSES  Mil 

With  the  dicarboxylic  anions  the  order  of  reactivity  was  pimelate  < 
adipate  <  glutarate  <  succinate  <  malonate  <  oxalate.  Evidently  the 
carboxy]  groups  in  glutarate  and  higher  homologues  are  so  far  apart  thai 
these  anions  behave  like  the  monocarboxylates.  Conversely,  the  closer  to- 
gether the  carboxyl  groups  are.  the  more  reactive  the  anion  is,  as  would  be 
expected  from  the  fad  thai  chelate  rings  of  five  or  six  members  are  more 
stable  than  larger  ones.  As  might  be  expected,  no  measurable  difference  was 
found  in  the  effects  of  structural  isomers,  such  as  butyrate  and  isobutyrate, 
valerate  and  isovalerate.  With  cis-trans  isomers,  however,  the  effects  are 
quite  different .  Malate  is  more  effective  than  fumarate,  presumably  because 
of  chelation. 

Spectrophotometry  studies  <>n  penetration  of  anions  into  basic  chromium 
complexes  by  Serfass  and  his  co-workers42  indicated  that  the  order  of  de- 
creasing penetrating  ability  is  oxalate  >  glycinate  >  tartrate  >  citrate  > 
glycolate  >  acetate  >  monochloracet ate  >  formate.  This  order  is  the 
same  as  the  coordinating  ability  observed  for  the  anions  mentioned. 
Kubelka4*  found  that  pyrogallol  can  expel  sulfate,  hut  thai  resorcinol 
and  hydroquinone  cannot. 

An  investigation  of  the  effect  of  dicarboxylic  acids,  especially  phthalic 
acid,  on  chromium  complexes  has  been  carried  out  by  Plant44.  He  believes 
that  only  one  of  the  carboxyl  groups  can  readily  displace  another  anion  and 
coordinate  with  the  metal  ion.  He  found  a  drop  in  the  pH  of  the  solution 
after  the  addition  of  the  a<-id,  and  he  concluded  that  with  only  one  carboxyl 
group  coordinated  the  other  acid  group  becomes  stronger  and  approaches 
the  strength  of  benzoic  acid.  However,  Shuttleworth45  disputes  these  find- 
ings, lb-  asserts  that  the  dicarboxylic  acids  can  chelate  without  displacing 
anions.  Such  coordination  would  cause  the  formation  of  anionic  complexes. 

Shuttleworth4*  has  conducted  conductimetric  studies  on  chromium  com- 
plex compounds  which  are  used  in  tanning.  He  found  that  high  dilution  of 
chromium  sulfate  causes  the  formation  of  ol  bridges  and  the  expulsion  of 
sulfate  groups.  He  pointed  out  that  the  formation  of  sulfato  and  olated 
complexes  involves  the  formation  of  4-membered  rings,  while  oxalato  com- 
plexes involve  the  more  .-table  5-membered  rings.  lie  also  suggested  the 
presence  of  hydrogen  bonds  between  the  oxygens  of  a  hydroxo  group  and 
an  adjacent  aquo  group.  The  following  structures  were  suggested  for  such 

_    -        38,  Wilson,  and  Theis,  J.  Am.  Leather  Chem.  Assoc.,  44,  647    1949 

Kubelka,  Technicka Hlidka K  24,  97    1949  ;/.  Am.  Leather  Chem.  Assoc. , 
44,824    ! 

14.  Plant. ./.  g                               ■  Chem.,  32,  88    1948 

to.  Shuttleworth,  ./.  &  ides'  Chem.,  33,  112    L94S 

Shuttleworth,  ./.  Soc.  Leatiu  T  odes1  Chen      33,  92     1949  ;  34,  :;.   186     l- 

./.  .1       /.  44.  589    1949  :45,  II     1950);  46.  56    1951 


L62 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


compounds: 


H20 


H20 


V/°\i  /OH\l  /°\^° 


JS^       4(HC204Na) 


o         o 


\)H^|  V 


HoO 


HoO 


,,  2(HC204Na) 


H20  H20 

XX  Cr 

0=C-0      I  ^OH     I      T>-C=0 


H20 


H,0 


!f   3(NaOH) 


,HH  H  H-. 

HOxO^OHNY/OH 

HO     O     OH'  O     O 

•.  /  \  ii 

H  H  C-C 


ii      ij 
O    0 


o  o 

II     II 

c-c 
66 

I  I 

c-c 

II  II 

o  o 


,OH. 


OH' 


O  O 

ii    H 

<rj 

O  O 

o  o 
c-d 

II    II 

o  o 


HO' 
,0H^    | 


N  '  << 


to 

O-Cr 
Hc/rOH^ 

*.  OH 
H 


2(NaOH) 


OH 


^ 


0  c 


excess  NaOH  Cr(OH)3 


Mixed  Bridge  Formation 

Various  anions  can  function  as  bridging  groups  in  polynucleate  ions,  and 
dinucleate  compounds  containing  chloride47,  acetate,  sulfate,  and  selenate48 
as  bridging  groups  have  been  identified,  e.g.: 


CH3 

A 

(NHa^Co^— OH  -  -Co(lMH3)« 
^OH 


Moreover,  the  formation  of  the  jti-acetato-/A-diol  compound  by  the  action 
of  acetic  acid  on  the  triol  might  be  regarded  as  an  example  of  "anion  pene- 
tration," since,  whatever  the  mechanism,  an  acetate  radical  has  replaced 
an  ol  group. 

47.  Reference  5,  p,  1MI . 

18.  Reference  '.».  pp.  341,  343,  362,  366,  368. 


OLATION  AND  RELATED  CHEMICAL  PROCESSES 


463 


Kuntzel49  round  that  bidentate  anions  can  bridge  between  two  chromium 
atoms,  and  thai  carbonates,  sulfates,  sulfites  and  organic  anions  displace 
ol  groups  readily.  Kuntzel  has  proposed  the  following  structure  for  the 
fatty  acid-chromium  complexes: 


V° 


H,0 


R 


P       °\l 
— O OH— Cr— O-C 


H£> 


The  basic  acetato  complexes  may  be  formed  as  follows:5 


— ,Cr 


.OH. 

■OH 

'OH 


Cr 


OAc 


^OAc. 
Cr OH 


Cr 


OAc" 


Compounds  containing  three  bridging  acetate  anions  were  formed  by  heat- 
ing the  reactants  in  sealed  tubes.501* 


Hydrous  Metal  Oxides 

On  the  basis  of  the  results  of  extensive  investigations,  Thomas  and  co- 
workers have  concluded  that  the  formation  and  composition  of  colloidally 
dispersed  metal  oxides,  and  of  precipitated  hydrous  oxides,  may  be  ex- 
plained in  terms  of  olation,  oxolation  and  anion  penetration.  Whitehead51 
has  compared  this  ''complex  compound  theory  of  hydrous  oxides"  with 
other  theories  of  colloidal  behavior. 

Any  adequate  theory  of  the  stability  of  colloidal  oxides  must  account  for 
the  fact  that  the  presence  of  some  ion,  other  than  the  metal  ion,  hydrogen 
ion,  and  hydroxide  ion,  seems  to  be  necessary  for  the  stability  of  a  metal 
oxide  hydrosol.  For  example,  Graham  peptized  iron(III)  oxide  with  iron- 
(III)  chloride,  and  concluded  thai  pure  iron(III)  oxide  sols  cannot  be  pre- 
pared since  they  flocculate,  when  dialyzed,  before  all  the  chloride  is  re- 
moved. Apparently  all  investigators  except  Sorum52,  who  has  reported  the 
preparation  of  pure  iron(III)  oxide  sols,  are  in  agreement  on  this  point. 

19.  Kuntzel,  CoUoquinuiber.  Insts.  Gerbereichem.  tech.  Hochschuh  Darmstadt,  No.   1. 
19    1949);  cf.,  Chem.  Abs.,  43,  6861  i     1940). 

50.  Kuntzel,  Erdmann  and  Spahrkas,  Das  Leder  4,  73    1953  ;  3,  30    L952);cf.,CJ 

.  47,  12087  t    1953);  46,  5479  g    1052). 

51.  Whitehead,  Chen    ft        21,  I 

"_    Sorum,  •/.  .  50,  1263    L928). 


464  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

According  to  the  adsorption  theory,  which  has  been  developed  in  great 
detail  and  has  found  wide  acceptance,  the  "foreign"  ions  are  adsorbed  on 
the  surfaces  of  the  dispersed  particles.  Thus  the  dispersed  particles  are 
electrically  charged,  and  mutual  repulsion  of  the  similarly  charged  particles 
accounts  for  the  stability  of  the  sol.  Flocculation  is  caused  by  neutraliza- 
tion of  the  charges. 

In  the  opinion  of  Thomas  and  his  co-workers,  however,  the  colloidal 
particles  in  metal  oxide  sols  are  aggregates  of  definite  chemical  structure 
which  behave  according  to  the  same  principles  as  do  the  so-called  crystal- 
loids53. The  micelles  in  such  sols  are  considered  to  be  polymeric  ol  or  oxo 
compounds  in  which  a  variable  fraction  of  the  coordination  positions  may 
be  occupied  by  anions  rather  than  ol,  oxo,  or  hydroxo  groups.  Each  micelle 
is  thus  regarded  as  a  very  large  ion,  whose  charge  is  inherent  in  its  structure. 
What  has  been  regarded  as  an  "adsorbed"  ion  is  actually  a  part  of  the 
chemical  composition  of  the  micelle. 

On  the  basis  of  the  complex  compound  theory  of  colloidal  oxides,  the 
compounds  present  in  metal  oxide  hydrosols  may  be  regarded  as  oxy  salts, 
and  it  is  convenient  to  name  them  as  such.  For  example,  Thomas  desig- 
nates the  compounds  in  aluminum  oxide  sols  which  contain  chloride  ion  as 
"aluminum  oxychlorides."  This  terminology  will  be  used  in  the  following 
outline  of  Thomas'  work. 

Thomas  and  Whitehead29  prepared  aluminum  oxychloride  sols  by  peptiz- 
ing (with  HC1)  aluminum  hydroxide,  which  had  been  precipitated  from 
aluminum  chloride  solution  with  NH4OH  or  NaOH.  According  to  the  co- 
ordination theory,  this  caused  formation  of  larger  and  larger  olated  ions 
until  aggregates  of  zero  charge  precipitated.  Peptization  reversed  these 
processes  to  an  extent  sufficient  to  cause  dispersion  of  the  precipitates. 
These  sols  exhibited  the  usual  properties  of  colloids,  i.e.,  Tyndall  effect, 
migration  of  the  particles  in  an  electric  field  (in  this  case  to  the  cathode), 
and  failure  of  the  particles  to  diffuse  through  membranes.  Sedimentation 
was  not  effected  by  centrifuging.  Tests  for  aluminum  ions  were  negative, 
indicating  that  all  the  aluminum  was  bound  in  the  complex  micelle.  Nearly 
all  of  the  chloride  was  present  as  chloride  ion. 

The  changes  in  hydrogen  ion  concentration  in  aluminum  oxychloride 
sols  due  to  various  treatments  were  investigated  by  Thomas  and  White- 
held  '.  The  fact  that  sols  prepared  and  aged  at  room  temperature  became 
more  acidic  was  attributed  to  hydrolysis  of  the  highly  polymeric  ions.  Sols 
which  were  prepared  at  room  temperature  became  more  acidic  when  heated. 
The  reaction  reversed  very  slowly  niter  the  sols  were  cooled,  and  the  origi- 
nal pll  was  attained  after  several  weeks.  Heating  the  sols  evidently  caused 
increased  hydrolysis  followed  by  olatioD  and  oxolation,  while  the  reversal 

53.  Thomas, ./.  Chem.  Ed.}  2,  323  (1925). 


OLATION  AND  RELATED  CHEMICAL  PROCESSES 


465 


was  due  to  slow  conversion  of  oxo  groups  to  ol  groups,  according  to  tli<i 
scheme 


H20 


hUO 


K 


OH 


H20 


n-i 


+     H" 


OH 


— in  -  • 


+ 


H2O 


HO 


n  -1 


.  /0H\ 


2n-2 


+  2H20 


.OH. 


~Al 


■OH 


"Ale 


2n-2 


SLOW 


rAI 


:AI? 


2n-4 


-f-    2H 


Sols  which  were  prepared  at  elevated  temperature  slowly  became  more 
basic  when  aged  at  room  temperature.  Evidently,  the  complexes  in  such 
sols  initially  contained  oxo  groups  which  slowly  reacted  with  hydrogen  ions. 
The  pH  of  zirconium  oxide  sols54  was  found  to  decrease  less  rapidly  upon 
aging  than  did  the  pH  of  chromium  oxide  sols55.  The  pH  decreased  irre- 
versibly when  the  sols  were  boiled,  perhaps  because  of  the  strong  tendency 
of  zirconium  oxy  salt  complexes  to  oxolate. 

Anion  Penetration  in  Hydrosols 

The  addition  of  solutions  of  neutral  salts  to  aluminum  oxide  sols  in- 
creased the  pH  of  the  sols  in  all  cases29.  This  was  evidently  not  due  to  dilu- 
tion, since  there  was  practically  no  effect  on  the  pH  when  water  was  added 
in  quantities  equal  to  the  volume  of  the  salt  solutions  used.  The  magnitude 
of  the  effect  depended  on  the  salt  added.  This  phenomenon  may  be  ex- 
plained by  anion  penetration,  since  displacement  of  a  hydroxo  or  an  ol 
group  by  an  anion  would  increase  the  pH  of  the  hydrosol. 

The  increase  in  pH  accompanying  the  addition  of  a  given  amount  of  a 
particular  salt  was  much  less  if  the  sol  was  heated  before  the  salt  was  added. 
Heating  may  have  converted  many  of  the  ol  groups  to  oxo  groups  which 
are  much  less  reactive  and  more  difficult  to  replace.  Since  ol  group-  are 
less  reactive  than  hydroxo  groups,  the  effect  may  be  partially  due  to  in- 
creased olation  caused  by  heating  the  sols. 

The  order  of  decreasing  tendency  of  anions  to  penetrate  into  the  complex 

54.  Thomas  and  Owens,  J.  Am.  Chem.  Soc,  57,  1825,  2131  (1935). 

55.  Thomas  and  von  Wicklen,  /.  .1/".  Chem.  Soc,  56,  704  (1934). 


466  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

was  found  to  be  approximately  the  same  for  aluminum  oxide,  chromium 
oxide  and  thorium  oxide  sols29, 55, 41,  the  order  indicating  the  order  of  ability 
of  the  anions  to  coordinate. 

The  decrease  in  hydrogen  ion  concentration  on  addition  of  neutral  salts 
to  aluminum  oxychloride,  oxybromide,  oxyiodide  and  oxyacetate  sols  is  so 
great  in  some  cases  that  the  sols  become  quite  alkaline56.  The  order  of 
effectiveness  of  anions  is  nitrate  <  chloride  <  acetate  <  sulfate  <  oxalate. 
The  magnitude  of  the  effect  of  a  particular  salt  was  different  for  the  differ- 
ent sols,  the  order  being  oxyiodide  >  oxybromide  >  oxychloride  >  oxy- 
acetate. This  result  is  consistent  with  the  order  of  penetrating  ability  of  the 
ions.  Heating  such  sols  makes  them  less  sensitive  to  the  action  of  added 
salts. 

Whitehead  and  Clay57  applied  the  idea  of  anion  penetration  in  a  com- 
parison of  the  properties  of  true  solutions  and  colloidal  dispersions.  The 
addition  of  various  anions  decreases  the  hydrogen  ion  concentration  with 
both  types  of  substances  but  the  effect  is  greater  with  sols  than  with  true 
solutions.  This  is  to  be  expected  since  the  number  of  OH  groups  replace- 
able by  anions  depends  on  the  total  number  present,  which  will  increase 
with  the  degree  of  olation,  i.e.,  with  the  size  of  the  ion.  The  order  of  the 
effect  as  determined  by  these  investigators  is  A1C13  <  Al(OH)Cl2  < 
Al(OH)2Cl  <  sol,  wrhich  indicates  a  gradual  transition  from  crystalloidal 
to  colloidal  dispersion. 

Thomas  and  Miller58  investigated  the  effect  of  anions  on  the  conductivity 
of  beryllium  oxychloride  sols  by  titrating  the  sols  with  solutions  of  silver 
nitrate,  silver  acetate,  and  silver  tartrate  in  concentrations  so  small  that 
the  anions  could  not  displace  hydroxo  groups  to  any  great  extent,  but  could 
displace  chloro  and  aquo  groups.  In  each  case  there  was  an  initial  decrease 
in  the  conductivity  of  the  sol  (greatest  with  tartrate  and  least  with  nitrate) 
followed  by  an  abrupt  increase.  The  initial  decrease  was  due  to  the  dis- 
placement of  aquo  groups  from  the  complex  cationic  micelles  with  a  re- 
sultant decrease  in  net  charge  on  the  complex  cations.  The  magnitude  of 
this  charge  would  be  greatest  with  the  most  strongly  penetrating  anion 
(tartrate)  and  least  with  the  most  weakly  penetrating  anion  (nitrate). 

Extremely  interesting  results  were  obtained  by  Thomas  and  Kremer41 
with  anions  of  h}Tdroxy  acids.  The  addition  of  potassium  salts  of  such  acids 
to  thorium  oxychloride  sols  reverses  the  charge  on  the  particles.  Moreover, 
peptization  of  hydrous  thorium  oxide  by  salts  of  hydroxy  acids  produces 
hydrosols  in  which  the  micelles  are  anionic.  It  was  also  observed  that  con- 
centrated nitric  acid  reverses  the  charge  of  thorium  oxychloride  micelles, 
producing  short-lived  nitrato  thoreate  micelles. 

56.  Thomas  and  Tai,  ./.  Am.  Chem.  Soc,  54,  841  (1932). 

57.  Whitehead  and  Clay,  /.  Am.  Chem.  Soc,  56,  1844  (1934). 

58.  Thomas.and  Miller,  /.  Am.  Chem.  Soc.,  58,  2526  (1936). 


OLATION  AND  RELATED  CHEMICAL  PROCESSES 


if.; 


These  results  are  explained  by  the  change  in  charge  on  a  complex  ion 
when  an  anion  penetrates  the  complex: 


=  Th 


H20 
OH 


+  an 


.an 


OH 


n-i 


+  H20 


+ 


an 


sTh. 


/nH0\     / 


OH 


7^ 


n-i 


[f  enough  anions  enter,  the  complex  acquires  a  negative  charge.  This  re- 
versal of  charge  was  also  noted  with  zirconium  oxide  hydrosols64. 

Since  hydrolysis  (conversion  of  aquo  to  hydroxo  groups)  and  oxolation 
inversion  of  ol  to  o.vo  groups)  decrease  the  positive  charge  on  the  complex 
ions,  boiling  the  sols,  which  favors  both  processes,  should  decrease  the 
amount  of  added  anion  necessary  to  precipitate  the  micelles  or  reverse  their 
charge.  In  general,  this  was  found  to  be  the  case.  Zirconeate  sols  formed 
by  such  processes  are  very  stable. 

Acid  zirconeate  sols  were  also  prepared  by  the  action  of  acids  of  great 
coordinating  tendency  on  hydrated  zirconium  oxide.  Peptization  of  the 
oxide  by  tartaric  acid  produces  sols  containing  both  positive  and  negative 
micelles.  All  of  the  salts  effective  in  causing  the  reversal  of  charge  are  those 
containing  alpha  hydroxy  anions.  Two  types  of  combination  are  possible; 
(a    the  OH  group  coordinates  as  such,  (b)  it  acts  like  an  acidic  group: 


HO 


R 
■C— H 


"in 


%/c=0_ 

(a) 


9 

C  — H 


=  Zr  C-0 

(b) 


m 


( Chelation  of  type  i  b  is  twice  as  effective  in  reducing  the  ionic  charge  as 
that  of  type  (a).  Because  of  the  effectiveness  of  alpha  hydroxy  anions  in 
reversing  the  charge  of  zirconium  oxide  micelles,  Thomas  and  ( ►wens64  con- 
cluded thai  type  (b)  is  more  probable,  [f  this  is  true,  dissociation  of  the 
<>I1  groups  of  the  hydroxy  anion  will  produce  hydrogen  ions.  Evidence 
for  such  a  phenomenon  was  obtained  by  adding  sail  mixtures  to  the  zir- 
conium oxide  sols.  Mixtures  of  anions  which  do  not  reverse  the  charge 
produce  nearly  the  same  pll  values,  while  oxalate-lactate  and  oxalate- 
tartrate  mixture-  produce  lower  pll  values.  It  was  found  thai  oxalate  pre- 
cipitate.- basic  zirconium  oxide  sols  without    reversing  the  charge,   bu1 


L68 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


c 
k 

u  I 


subsequent  addition  of  a  salt  of  an  alpha  hydroxy  acid  peptizes  the  precipi- 
tate with  the  formation  of  a  complex  zirconeate  sol.  Moreover,  if  sufficient 
alpha  hydroxy  salt  is  first  added  to  a  zirconium  oxy chloride  sol,  the  addi- 
tioD  of  oxalate  does  not  cause  precipitation.  These  phenomena  are  entirely 
consistent  with  the  behavior  of  crystalloidal  zirconium  salts  which  usually 
form  stable  complexes  with  alpha  hydroxy  acids. 

Precipitation,    Peptization,    and    Dissolution    of    Hydrous    Metal 
Oxides 

It  is  well  known  that  metal  oxide  sols  can  be  flocculated  by  prolonged 
boiling  or  by  the  addition  of  alkali.  According  to  the  coordination  theory, 
flocculation  occurs  because  of  hydrolysis,  olation,  and  oxolation  of  the 
complex  cations.  Hydrolysis,  followed  by  olation,  leads  to  the  formation 
of  larger  aggregates.  The  loss  of  hydrogen  ions  by  aquo  groups  (hydroly- 
sis) and  by  ol  groups  (oxolation)  reduces  the  charge  on  the  cation,  the  sta- 
bility of  the  sol  decreasing  as  the  ratio  of  charge  of  the  micelle  to  its  mass 
decreases.  Beryllium  oxide  hydrosols  precipitate  immediately  when  boiled, 
and  in  about  two  hours  at  60°.  This  is  attributed  to  oxolation  and  the  conse- 
quent formation  of  complexes  of  zero  charge.  This  type  of  neutralization 
occurs  more  readily  with  beryllium  sols  than  with  ol  complexes  of  the  tri- 
valent  metals  whose  coordination  number  is  six,  simply  because  the  loss 
of  fewer  protons  is  required,  the  valence  and  coordination  number  of  beryl- 
lium being  only  two  and  four,  respectively. 

A  precipitated  hydrous  oxide  may  contain  such  complexes  as 


"       HO 
HgO 


OHv        ^OH 

Al^  Al 


OH- 


"Al 


OH 


H20 

L_        HO 


/\  \oh//\\oh^/\\oh^7\ 
ho  oh  ho  oh         ho  oh         ho  oh 

\  /^  OH^   \//OH\\//OH\\/ 

Al/  A  I  A  I  A  I       _ 


H20 


Al 


OH" 


.AK  AK 

\0H^      ^OH' 


•H20 


OH  _J 


which  are  not  appreciably  soluble  in  water.  In  the  presence  of  acid,  how- 
ever, a  number  of  reactions  occur,  i.e.,  conversion  of  hydroxo  to  aquo 
"roups,  penetration  of  anions  into  the  complex  nucleus,  and  deolation. 
The  final  result  depends  to  a  large  extent  on  the  penetrating  ability  of  the 
anion.  In  any  event,  the  complex  acquires  one  positive  charge  for  each 
hydroxo  group  converted  by  a  hydrogen  ion  to  an  aquo  group,  and  one  or 
more  negative  charges  (depending  on  the  anion)  for  each  anion  entering 
the  complex.  Deolation  also  occurs  to  some  extent.  Whether  the  oxide  is 
dissolved  or  peptized  depends  on  the  nature  of  the  anion,  since  this  deter- 
mines the  extent  of  anion  penetration  and  therefore,  of  deolation.  If  an 
acid  whose  anion  is  a  weak  penetrator  is  added,  anion  penetration  only 


OLATION  AND  RELATED  CHEMICAL  MtOCKSSKS 


469 


partly  neutralizes  any  positive  charge  which  the  complex  acquires  by  the 
conversion  of  hydroxo  groups  bo  aquo  groups  by  the  action  of  the  hydrogen 
ions.  When  the  ratio  of  charge  to  mass  becomes  large  enough,  peptization 
occurs,  provided  the  number  of  equivalents  of  acid  present  Is  much  less 
than  the  number  of  equivalents  of  aluminum. 

On  the  other  hand,  with  an  acid  whose  anion  is  a  powerful  penetrator, 
a  considerable  number  of  aquo  or  hydroxo  groups,  or  both,  arc  displaced 
by  anions.  This  offsets  the  increase  in  positive  charge  due  to  conversion  of 
hydroxo  to  aquo  groups.  With  a  small  ratio  of  acid  to  aluminum,  acid  dis- 
appears from  solution,  i.e.,  hydrogen  ions  and  anions  are  said  to  be  "ad- 
sorbed" by  the  alumina.  With  a  sufficiently  large  ratio  of  acid  to  alumina, 
complete  deolation  results  in  crystalloidal  dispersion  of  the  oxide,  provided 
it  were  not  oxolated. 

Experimental  results  are  in  accord  with  these  ideas59.  The  following  order 
of  effectiveness  of  acid  in  peptizing  hydrous  alumina  was  found:  trichloro- 
acetic >  dichloroacetic  >  nitric  >  hydrobromic  >  hydrochloric  >  mono- 
chloroacetic  >  formic  >  gly colic  >  acetic  >  oxalic  >  tartaric  >  sul- 
furic. With  the  exceptions  of  dichloroacetic,  sulfuric  and  tartaric 
(discrepancies  which  are  not  accounted  for),  the  peptizing  ability  of  the 
acids  approximates  the  reverse  of  the  order  of  the  effectiveness  of  their 
anions  in  raising  the  pH  of  hydrosols.  Both  orders  reflect  the  tendency  of 
the  anions  to  become  coordinately  bound  in  the  complex  cations.  The  acids 
having  strongly  penetrating  anions  were  removed  from  solution  as  indicated 
by  an  increase  in  pH.  To  the  extent  that  they  dispersed  hydrous  alumina, 
they  produced  a  large  proportion  of  crystalloidal  compounds  because  of 
their  deolating  effect. 

Thomas  and  Miller58  produced  stable  anionic  beryllium  oxide  hydrosols 
by  the  use  of  powerfully  coordinating  anions.  In  contrast  to  the  behavior 
of  cationic  hydrosols,  which  become  more  acid  on  aging  at  room  tempera- 
ture (due  to  oxolation  and  possibly  to  dissociation  of  aquo  groups),  these 
complex  beryllate  hydrosols  become  less  acid.  This  is  due  to  aquotization 
or  anation  and  may  be  exemplified  by  a  reaction  of  a  hypothetical  basic 
citrato  beryllate  (R  =  C6HAS): 


oil 

1 

OH              OH 

3 

RsEsBe— OH— Be— OH— Be — OH— Be—  II  <  1 

1                   1                   1 

+  H20  -» 

OH              OH              OH, 

oil               OH           II  <> 

1                   1                    1 

- 

R=He— oil      Be      OH— Be— OH— Be      II  <> 

1                    1                   1 

•    1  >ll 

OH              OH          II  0 

59.  Thomas  and  Yart  :mi:iii,  ./.  Am.  Chin,.  Sue..  57,  I  (1935 


470  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

The  conclusion  is  that,  in  general,  acids  with  anions  of  great  coordinating 
ability  are  poor  peptizers  of  hydrous  oxides  while  acids  of  weakly  coordinat- 
ing anions  are  good  peptizers. 

Other  Properties  of  Hydrous  Metal  Oxides 

According  to  the  coordination  theory,  precipitated  hydrous  oxides  are 
considered  polymeric  compounds  not  different  in  kind  from  those  existing 
in  crystalloidal  solutions  and  colloidal  dispersions33*1, 60.  They  are  regarded 
as  complexes  of  zero  charge  produced  by  a  continued  process  of  olation, 
accompanied  or  followed  by  oxolation  and  perhaps  by  anion  penetration. 
This  point  of  view  furnishes  an  explanation  of  two  well-known  character- 
istics of  hydrous  oxides,  such  as  those  of  aluminum  and  chromium — their 
decreased  chemical  reactivity  after  aging  or  heating  and  their  ability  to 
retain  certain  impurities  even  after  exhaustive  washing. 

As  to  the  first  of  these,  a  freshly  precipitated  hydroxide  may  consist  of 
complexes  of  relatively  low  aggregate  weight  containing  a  high  ratio  of  ol 
to  oxo  groups.  For  a  given  weight  of  hydroxide,  the  smaller  the  aggregates, 
the  more  "end  groups"  there  will  be,  i.e.,  hydroxo  and  aquo  groups.  The 
hydroxo  groups  are  easily  convertible  to  aquo  groups  by  the  action  of 
hydrogen  ions  and  may  easily  be  displaced  by  anions.  01  groups  are  not  so 
readily  attacked  by  hydrogen  ions  or  displaced  by  anions  but  do  react 
slowly.  Thus,  low  molecular  weight  aggregates,  which  are  not  too  highly 
oxolated,  may  be  dissolved  readily  in  acid60. 

However,  the  process  of  olation,  by  which  the  hydrous  oxide  was  pre- 
sumably formed,  may  continue  slowly  after  precipitation,  even  at  low  tem- 
perature. There  is  a  decrease  in  the  relative  number  of  hydroxo  groups,  and 
a  corresponding  increase  in  the  number  of  ol  and  oxo  groups60.  The  com- 
pletely oxolated  oxide  is  quite  inert. 

It  is  common  knowledge  that  precipitated  hydrous  oxides  almost  in- 
variably contain  the  anion  of  the  salt  from  which  the  oxide  was  formed, 
and  that  such  impurities  are  extremely  difficult  to  remove61.  The  explana- 
tion often  given  is  that  the  impurity  is  adsorbed,  or  occluded.  However, 
this  phenomenon  can  also  be  accounted  for  by  the  coordination  theory. 

If  anion  penetration  occurs  during  precipitation,  the  complexes  contain 
anions  as  an  integral  part  of  their  structure.  Washing  the  precipitate  may 
ultimately  cause  replacement  of  the  anions  by  aquo  groups.  On  this  basis, 
anions  of  greatest  coordinating  tendency  are  hardest  to  remove.  This  has 
been  found  to  be  the  case51.  The  facts  that  such  anions  are  displaced  by 
other  anions  of  greater  penetrating  ability611",  and  that  freshly  precipitated 

80.  Graham  and  Thomas, ./.  .1///.  Chem.  Soc,  69,  816  (1947). 

61.  Thomas  and  Frieden,  ./.  Am.  Chem.  Soc,  45,  2522  (1923);  Charriou,  Compt.  rend>, 
176,  679,  L890  (1923). 


t\ 


OLATION  AND  RELATED  CHEMICAL  PROCESSES  471 

hydrous  aluminum  oxide  liberates  hydroxide  ions  on  treatment  with  neu- 
tral Baits88  are  explainable  by  anion  penetration. 

In  addition  to  results  specifically  mentioned  in  the  foregoing  discussion, 
evidence  consistent   with  the  interpretation  given  has  been  obtained,  in 

investigations  of  titanium  oxide  sols68,  of  the  effect  of  anions  on  the  pi  I  of 
maximum  precipitation  of  aluminum  hydroxide84,  and  of  the  catalytic  ac- 
tivity of  aluminum  oxyiodide  sols  in  the  decomposition  of  hydrogen  per- 
oxide'''''. Summaries  of  the  coordination  theory  of  hydrous  oxides  have  been 
compiled  by  Whitehead''1,  Thomas88,  and  Perkins  and  Thomas87.  Other 
investigators,  notably  Pauli  and  co-workers88,  have  applied  the  coordinal  ion 
theory  to  colloidal  systems. 

Olation  and  oxolation  are  of  great  importance  in  leather  chemistry  as 
shown  by  Stiasny  and  other  investigators.69  In  tanning,  only  olated  com- 
pounds are  effective.  Briggs70  is  studying  the  separation  of  basic  chromium 
salts  by  means  of  aqueous  ethyl  alcohol.  His  work  shows  that  it  may  be 
possible  to  separate  such  compounds  fairly  simply  and  easily.  Basic  iron, 
aluminum  and  zirconium  compounds  are  also  of  interest  as  tanning  agents.71 

It  must  be  admitted  that  the  theory  is  controversial,  at  least  in  certain 
aspects.  Weiser  and  co-workers,  in  particular,  have  criticized  it  mainly  on 
the  basis  of  results  of  x-ray  studies  and  isobaric  and  isothermal  dehydra- 
tion studies72. 

62.  Sen,  /.  Phys.  Chem.,  31,  691  (1927). 

63.  Thomas  and  Stewart,  Koll.  Z.,  86,  279  (1939). 

64.  Marion  and  Thomas,  J.  Coll.  Sci.,  1,  221  (1946). 

65.  Thomas  and  Cohen,  ./.  Am.  Chem.  Soc.,  61,  401  (1939). 

66.  Thomas,  "Colloid  Chemistry,"  Chapt.  7,  New  York,  McGraw-Hill  Book  Com- 

pany, 1934. 

67.  Perkins  and  Thomas,  Stiasny  Festschr.,  307,  Darmstadt,  Ed.  Roether  Verlag,  1937. 
6S.  Pauli  and  Yalko,  "Elektrochemie  der  Kolloide,"  Vienna,  Julius  Springer,  1929. 
69.  Reference  17.  chapters  14-18;  McLaughlin  and  Theis,  "The  Chemistry  of  Leather 

Manufacture.''  chapters  14  16,  New    York,  Reinhold  Publishing  Corporation, 

L945,  Shuttleworth, /. Soc. Leather  Trades'  Chem. ,34,  410  (1950);  J.Am. Leatfo 

Chi  m.  Assoc  .  46,  582  (1951). 
7(i.   Briggs,  •/.  Soc.  Leatiu  r  Trades'  Chem.,  35,  235  (1951). 
71.  References  69b,  chapters  19,20,22. 
7_\  Weiser,  Milligan,  and  Coppoc,  ./.  Phys.  Chem.,  43,  1109  (1939);  Weiser  and  Milli- 

gan,  ibid.,  44,  KM  (1940);  Weiser,  Milligan,  and  Purcell,  Ind.  Eng.  Chem.,  33, 
I     1941);  Weiser,  Milligan  and  Simpson,  ./.  Phys.  Chem.,  46,  1051   (1942); 

Weiser  and  Milligan.  Chem.  Revs.,  25,  1  (1939  . 


14.   The  Poly-Acids 

Hans  B.  Jonassen 

Tulane  University,  New  Orleans,  Louisiana 

and 

Stanley  Kirschner 

Wayne  University,  Detroit,  Michigan 


The  poly-acids  are  characterized  by  the  fact  that  they  contain  more  than 
one  acid  anhydride  molecule  per  acid  anion1.  If  they  have  only  one  kind  of 
acid  anhydride,  they  are  called  isopoly-acids  (e.g.,  H2M04O13  or  H20- 
4Mo03);  if  they  contain  more  than  one  kind  of  acid  anhydride,  they  are 
called  heteropoly -acids  (e.g.,  H4SiWi204o  or  Si02-  (W03)i2-2H20). 

The  elements  whose  oxides  are  capable  of  undergoing  condensation  to 
form  isopoly-  and  heteropoly-acids  are  those  in  groups  V-B  (V,  Nb,  Ta) 
and  VI-B  (Cr,  Mo,  W,  and  U2)  of  the  periodic  table. 

k:  Early  Structural  Studies 

As  long  ago  as  1826  Berzelius3  described  ammonium  phosphomolybdate ; 
,and,  although  silicotungstates  were  known  as  early  as  18474, 5,  the  first  care- 
fid  determination  of  the  composition  of  a  silicotungstate  was  not  carried 
out  until  18626.  The  compositions  of  many  isopoly-  and  heteropoly-acids 
and  salts  were  subsequently  established,  but  very  few  structural  studies 
were  undertaken.  Klein7  attempted  to  explain  the  structure  of  the  para- 
tungstic  acid  prepared  by  Laurent8,  but  his  ideas  met  with  little  success 
after  the  discovery  of  many  other  more  complex  acids. 

1.  Rosenheim,  "Handbueh  der  Anorganischen  Chemie,"  Abegg  and  Auerbach,  Vol. 

4,  Part  1,  ii,  pp.  977-1065,  Leipzig,  Hirzel,  1921. 

2.  Wamser, ./.  .1///.  Chem.  Soc,  74,  1020  (1952). 

3.  Berzelius,  Pogg.  Ann.,  6,  369  (1826). 

1.  Laurent,  .1/'//.  chim.  phys.,  [3]  21,  54  (1847). 

5.  Riche,  Ann.  chim.  phys.,  [3]  50,  5  (1857). 

6.  Marignac,  Compt.  rend:,  55, 88  (1862). 

7.  Klein.  Bull.  80C.  chin,.,  [2]  36,  546  (1881). 
S    Laurent,  Compt.  red.,  31,  (i!)2  (1850). 

472 


THE  POL]     ICIDS  173 

Blomstrand9  l0  attempted  bo  explain  the  structure  of  fche  poly-acids  by 
assuming  a  chain  or  ring  configuration.  For  phosphomolybdic  acid,  for 
example,  he  proposed  a  straighl  chain  containing  twelve  MoOs  groups  with 
an  (  MI  group  at  one  end  and  an  IMM )::  group  at  t he  ol her: 

0  0  0  O  on 

/           /            /  /  / 

1 1  ( >— Mo— 0— Mo— 0— Mo— O— Mo— O— P=0 

\  \  \  \  \ 

0  O  O  O  OH 

However,  the  hypotheses  sel  forth  by  these  and  other  early  investigators11 
proved  to  be  unsatisfactory. 

Later  Structural  Studies 
The  W  ork  of  Copaux,  Werner,  Miolati,  and  Rosenheim 

In  1906,  Copaux'-1  attempted  a  classification  of  these  complex  acids  based 
upon  their  isomorphism,  and  he  concluded  that  the  isopoly-acids  were  quite 
similar  in  structure  to  the  heteropoly-acids.  For  the  isopoly-acids  he  as- 
sumed that  two  water  molecules  condensed  to  form  an  H4O2  unit  which 
then  behaved  as  an  anhydride  group;  thus  he  considered  these  acids  as 
heteropoly-acids,  in  which  the  H402  group  is  assumed  to  be  the  second  an- 
hydride. Although  it  is  now  regarded  as  incorrect,  Copaux's  hypothesis  is 
of  historical  importance,  since  it  started  later  workers  along  the  path  which 
ultimately  led  to  the  currently  accepted  structures  for  these  acids. 

Even  though  it  is  possible  to  form  condensed  aggregates  of  a  single  metal- 
loid anhydride  molecule  with  various  numbers  of  molecules  of  a  group  V-B 
or  VI-B  metal  anhydride,  two  types  of  aggregates  are  much  more  common 
than  any  of  the  others.  They  are  the  heteropoly-acids  (and  salts)  which  con- 
tain six  or  twelve  molecules  of  the  metal  anhydride  for  each  molecule  of  the 
metalloid  anhydride.  These  acids  are  called  limiting  acids  or  "Grenzsauren." 
Table  14.1  depict  s  1  hose  elements  which  have  been  reported  as  central  atoms 
of  the  '•metalloid"  anhydride. 

Tabic  1  L2  lists  a  few  examples  of  the  limiting  acids  and  their  salts. 

Werner"  applied  his  ideas  on  coord inat  ion  compounds  to  the  structure  of 
silicotungstic  acid  and  its  salts.  II<'  assumed  that  the  central  group  is  an 
Si();;  ion  surrounded  by  tour  (RW^Oe)*  groups  (II  =  a  unipositive  ion) 
which  are  linked  to  the  central  group  by  primary  valences.  In  addition,  he 

9.  Blomstrand, Z.  anorg.  Chem.,1,  10  (1892). 
in.  Rosenheim,  Z.  anorg.  Chem.,  75,  1  tl   1 1912). 

11.  Gibbs,  •/.  Am.  Chem.  80c. }  5,  391  (1883);  Friedheim  and  Castendyck,  Rev.,  33, 

1611    1900  . 

12.  Copaux,  Ann.ehim.phya.,  [8]7, 118  - 1906);  Bull. %oc.  chim.,  18, 820    1913 
L3.  Wen  10,  in    1907). 


474  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Table  14.1.  Elements  Reported  as  Central  Atoms  in  Heteropoly-Acids 

Group  Number  Elements 

I-A  II 

II-A  Be 

III-A  B,  Al,  Ce 

IV-B  Ti,  Zr,  Th 

V-B  V,  Nb,  Ta 

VI-B  Cr,  Mo,  W,  U 

VII-B  Mn 

VIII  Fe,  Co,  Ni,  Rh,  Os,  Ir,  Pt 

IB  Cu 

IV-A  C,  Si,  Ge,  Sn 

V-A  N,  P,  As,  Sb 

VI-A  S,  Se,  Te 

VII-A  I 

Table  14.2.  Examples  of  Limiting  Poly-Acids  and  Salts 
( jf 

Type  Formula  Name 

6-poly-acids  2Na2OP20512WOraq.  Sodium  phospho-6-tungstate 

2H20  •  Te03  •  6M0O3  •  aq.  Tellurium-6-molybdic  acid 

3H20-P205-12W03aq.  Phospho-6-tungstic  acid 

12-poly-acids  3K20-P205-24W03-aq.  Potassium  phospho-12-tungstate 

3H20-B203-24W03-aq.  Boro-12-tungstic  acid 

5(NH4)20-2P205-24V205  Ammonium  phospho-12-vanadate 

postulated  that  two  R2W207  groups  are  linked  by  secondary  valences  to 
this  same  central  group,  and  he  felt  that  this  would  result  in  an  octahedral 
configuration  for  the  poly-acids.  Although  this  structure  accounted  for  the 
behavior  of  some  of  the  limiting  poly-acids  containing  a  tetravalent  central 
ion,  difficulties  were  encountered  with  those  acids  having  a  central  ion 
with  a  valence  other  than  four,  and  with  those  containing  metal  anhydride 
aggregations  which  are  not  multiples  of  six. 

Miolati14  and  Rosenheim  and  co-workers1, 10,  15  extended  Werner's  ideas 
to  include  those  poly-acids  which  do  not  belong  to  a  limiting  acid  series 
and  attempted  to  explain  the  large  number  of  replaceable  hydrogens  in 
many  of  these  acids.  They  considered  the  poly-acids  as  being  formed  in  a 
manner  analogous  to  the  stepwise  displacement  of  hydroxyl  groups  by 
chloride  ions  from  platinic  acid,  H2[Pt(OH)6],  ultimately  yielding  hexa- 
chloroplatinic  acid,  H2[PtCl6].  Telluric  acid,  H6[Te06],  and  para-periodic 
acid,  H5[I06],  for  example,  were  regarded  as  parent  acids  which  show  six- 
coordination  and  which  possess  octahedral  structures.  They  were  thought 
to  form  heteropoly-acids  by  the  stepwise  displacement  of  the  oxygens  by 
WOr  groups  to  give  H6[Te(W04)6]  and  H5[I(W04)6],  respectively.  It  was 

14.  Miolati,  ./.  prakt.  Chem.,  77,  417  (1908). 

15.  Rosenheim,  Z.  anorg.  Chan.,  69,  247  (1910);  Rosenheim  and  Jaenicke,  ibid.,  100, 
304  (1912). 


b  » 


THE  POLY  ACIDS  475 

Table  14.3.  Rosenhbim-Miolati  Classification  oi    phe  6-Poli  Acids 

Valence  of 
Centra] 
item  Central  Atom    -   \  Parent  Acid  Typical  Heteropory  tall 

Mn,  Ni.Cu  BioPCOe]  NrH4)«H7[Mn(MoO«).]-3HtO 

3  Al,  Cr,  Co  H.pCOe]  K    Co  MoO4)«]xH«0 

6  Te  H.lTeOe]  C  ML        Te  Wt  >.,)f]-HU,<  > 

7  I  H»[IO«]  \;.,!I(\V()1)t,]-SH,() 

Table  lit    Rosenheim-Mioi  \n  Classify  ition  of  the  r_  Pols  Acids 

Valence  of 
tral 
Atom  Central  Atom  (-   \  Parent  Acid  T>  pual  Heteropoly-salt 

3  B  H9[B06]  llg9lB(W207)6]-12.-)]l  ,0 

1  Si,  Ge,  Sn,  Ti  H8[X06]  K4H4[Si(W207)6]-7H20 

5  P,  As,  Sb  H7[X06]  Ag7[Sb(Mo207)6J-15H20 

believed  that  a  W(  h  group  was  bonded  to  an  oxygen  at  each  corner  of  the 
octahedron  containing  the  central  atom. 

Rosenheim  and  Miolati  expanded  this  concept  by  postulating  an  entire 
series  of  hypothetical  parent  oxy-acids  showing  six-coordination  and  having 
oxygen  atoms  at  the  corners  of  the  octahedra  containing  the  central  metal 
atoms.  Each  oxygen  was  then  considered  to  be  coordinated  to  a  metal  an- 
hydride molecule.  Table  14.3  lists  some  of  the  parent  acids  postulated  by 
these  workers  for  the  6-poly-acid  series,  along  with  compounds  which  were 
thought  to  be  derived  from  them. 

In  a  similar  manner,  parent  acids  were  postulated  for  the  12-poly-acid 
scries,  and  Table  14.4  lists  some  of  these  along  with  examples  of  salts  of  the 
L2-heteropoly-acids. 

The  structures  of  the  unsaturated  heteropoly-acids  (i.e.,  those  which  do 
not  belong  to  the  six  or  twelve  limiting  acid  series)  can  be  explained,  ac- 
cording  to  Rosenheim10,  by  assuming  that  not  all  of  the  six  oxygens  are 
lisplaced  by  acid  anion  groups.  For  example,  if  only  five  of  the  oxygens  of 
the  parent  acid  H7[As06]  are  replaced  by  Mo20-=  groups,  then  the  arsenic- 
10-molybdic  acid  i  II7[A.-<  )^  M< »_( >7)5]  -acj.)  is  formed16.  Similarly,  the  un- 
saturated i  1-.  101  _;-.  and  9-poly-acids  of  the  phosphotungstic  series  can  be 
explained  by  Rosenheim's  postulates,  provided  that  polyoctahedral  aggre- 
gates  arc  assumed,  as  is  shown  in  Table  14.5. 

The  unsaturated  poly-acids  below  the  12-series  and  above  the  6-series  are 
formed  by  the  1"--  of  M  <  >7  groups  from  one  or  more  corners  of  the  octa- 
hedron with  the  resultant  formation  of  bridge  structures  of  different  types. 
For  the  unsaturated  acids  below  the  6-series,  Rosenheim  and  Pieck17  postu- 
lated thai  M<>;  groups  do  not  replace  all  of  the  oxygen  atoms  surrounding 

it',.  Rosenheim,  Z.  anorg.  Chem.,  91,  75    1916*  , 

17.  Rosenheim  and  Keck,  Z  ano  g   Chen   .96,  139  (1916). 


476 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


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'////•/  POLY  ACIDS  177 

Tabus  14.6.  Some  Polt  Amp-  lnd  Salts  i\  which  Tbtracoobdination 

i-   Lxiu  BI  i  'ED 
Class  of  Add  Acid 

3  acid  H    AbO  Mo04),]  iaO  MoO<     -711  0 

•J>_,  arid  II    P  <>    OH    M<>    i]    MO  Li,  IV  MO<    0    "II   ,]16Hg0 

the  (-(Mitral  atom.  They  proposed  that  in  the  case  of  manganol  I\'  l-5-tungsl  ic 
arid,  for  example,  the  parent  acid  has  the  formula  lK[MnIV0(\Y( ),),,]  and 
they  describe  the  salt  Xa6Ho[Mn()(W()4)5]. 

Difficulties  are  encountered  with  the  acids  and  salts  lower  than  the 
pies.  Only  by  assuming  tetracoordination  of  the  central  atom  of  these 
acids  was  Rosenheim16  able  to  include  them  in  his  system  of  classification. 
Some  typical  examples  of  such  compounds  are  given  in  Table  1  L6. 

Isopoly-Acids 

Since  the  isopoly  12-tungstates  are  isomorphous  with  the  22-hydrates  of 
the  boro-,  silico-,  and  phospho-12-tungstates,  and  since  12-tungstic  acid 
does  not  lose  all  of  its  water  on  ignition,  Rosenheim  and  Felix18  proposed 
that  these  isopoly-acids  be  considered  as  a  type  of  heteropoly-acid.  They 
postulated  the  hypothetical  parent  acid  (H20)6  or  Hio[H2Of],  with  the  H2++ 
group  acting  as  the  central  ion  of  the  heteropoly-acid.  The  six  oxygen 
atoms,  supposedly  octahedrally  located  about  the  H24+  central  group,  are 
then  replaced  by  W207=  groups  producing  the  hydrated  12-tungstic  acid, 
Hio[H2(\Y207)6].  The  6-acids  were  similarly  included  in  Rosenheim's  scheme17 
by  proposing  a  replacement  of  the  six  oxygens  by  W04=  groups  to  give  the 
hydrated  5-tungstic  acid,  Hi0[H2(WO4)6].  Rosenheim  treated  the  isopoly 
molybdic  acids  in  a  like  manner  by  postulating  the  replacement  of  the  six 
oxygens  of  the   'a<iiio  acid"  core,  H10[H2O6],  by  Mo207=  or  Mo04=  groups. 

The  vanadium  poly-acids  were  also  brought  into  this  classification  by 
Rosenheim  and  Pieck'7  who  proposed  the  existence  of  another  hypothetical 
aquo  acid,  H4[H>03].  By  replacing  each  oxygen  with  two  VO»~  groups,  the 
aquo-6-vanadic  acid,  H4[H2(  V03)6],  is  formed.  In  order  to  explain  the  penta- 
vanadates,  it  was  postulated  thai  aquo-6-vanadic  acid  undergoes  partial 
hydrolysis  with  the  replacement  of  a  Y<  I .-  group  by  an  OH~  group  to  give 
aquopentavanadic  acid.  II4[H2(V03)50H]. 

It  was  proposed  that  the  aggregation  processes  occurred  in  solution 
through  the  following  reaction  mechanisms: 

.molybdates: 
(MoO,)"  v     "J_  -  (Mo207)"  ,     "H.   -  (H2(Mo04)6)^  ^±  (H2(Mo207)6)^  ^±  (Mo03), 

18.  Rosenheim  and  Felix,  Z.  a  79,  202  (1913). 


478  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Polytung  states: 

(W04)=  'F===*  (WaOr)-  ^  (H2(W04),)>»-  ^  (H2(W207)6)io-  —  (WO3)* 


Polyvanadates 
"oil 
Critical  Discussion  of  Rosenheim's  Postulates 


<V04)-  ^==^  (Vs07)<-  ^±  (V30,)-  ^±  (H2(V03)c)4-  ^  (V205), 


Rosenheim's  work  was  based  upon  several  different  types  of  chemical  and 
physical  evidence,  but  he  did  not  have  access  to  methods,  such  as  x-ray 
analysis19,  which  were  developed  and  refined  several  years  after  his  ideas 
were  published.  As  a  result,  his  structural  theories  suffered  accordingly.  A 
very  important  part  of  his  work,  however,  involved  the  careful  preparation 
and  analysis  of  salts  with  the  accurate  determination  of  the  amount  of  con- 
stitutional water  which  cannot  be  removed  except  by  ignition  at  high  tem- 
peratures. His  work  in  this  field  was  extensive  and  carefully  carried  out. 

Determinations  of  the  maximum  basicities  of  the  different  salts  were 
also  made,  but  Rosenheim  was  able  to  isolate  only  a  few  salts  in  which  the 
maximum  basicity  of  his  hypothetical  acids  was  attained  (see  Tables  14.3 
to  14.6).  In  most  cases,  the  compounds  formed  could  be  explained  only  by 
postulating  a  partial  replacement  of  the  hydrogen  ions  by  basic  groups  to 
give  the  acid  salts.  Conductivity  measurements  and  conductometric  titra- 
tions were  also  utilized  by  Rosenheim,  but  the  results  obtained  by  these 
methods  can  easily  be  interpreted  to  fit  other  theories.  Furthermore,  later 


workers,  using  modern  experimental  techniques,  have  shown  that  several  of 
his  proposed  structures  (e.g.,  those  for  the  polyvanadates)  are  incorrect. 

One  of  the  most  important  objections  to  Rosenheim's  theory,  however,  is 
the  postulate  concerning  the  existence  of  M207  groups  in  solution.  Although 
such  "pyro"  radicals  have  definitely  been  shown  to  exist  in  acid  solution  in 
the  chromic  acid  series,  it  has  not  been  conclusively  demonstrated  that  such 
radicals  exist  in  other  acid  series.  (However,  Ripan  and  Poppei20  have  con- 
cluded that  the  W207=  group  may  exist  as  such  in  silico-12-tungstic  acid.) 

Another  objection  to  Rosenheim's  postulates  arises  when  one  considers 
that  almost  all  of  the  poly-acids  and  salts  reported  contain  a  great  deal  of 
water  of  hydration.  Rosenheim  proposed  that  the  12-acids  could  contain  up 
to  only  thirty  molecules  of  water  of  hydration  for  each  central  metalloid 
atom,  but  hydrates  containing  more  than  thirty  tightly  bound  waters  per 
central  atom  have  since  been  reported21.  It  becomes  impossible,  therefore, 
to  reconcile  the  large  numbers  of  water  molecules  with  the  structural  ideas 
proposed  by  Rosenheim  for  the  poly-acids. 

19.  Sturdivant,  J.  Am.  Chem.  Soc,  59,  530  (1937). 

20.  Ripan  and  Poppei,  Bui.  Soc.  Stunte  Cluj,  10,  85  (1948). 

21.  Kraus,  Z.  Krist,  91,  402  (1935);  93,  379  (1936). 


THE  POLY-ACIDS  17(> 


The  Work  of  Pfeiffer 


The  many  objections  to  Rosenheim's  postulates  brought  forth  by  differ- 
ent investigators  initiated  extensive  studies  in  this  field.  Various  experi- 
mental approaches  were  used,  among  them  x-ray  diffraction  techniques22. 

After  Lane,  Bragg,  Delize,  and  others  had  shown  that  crystals  follow  the 
crystallographic  coordination  laws,  Pfeiffer2* attempted  to  explain  the  struc- 
tures of  the  heteropoly  tungstates  by  utilizing  these  laws.   lie  accepted 

Rosenheim's  view  that  the  poly-acids  are  derived  from  hypothetical  parent 
acids  (i.e.,  IIu-,\"~06),  and  he  postulated  that  W03  groups,  for  example, 
coordinate  in  a  second  coordination  sphere  about  the  central  [X06]n~12 
group,  which  can  have  a  coordination  number  as  high  as  twelve.  Hence, 
phospho-12-tungstic  acid  should  really  be  formulated  as  H7[(P06)(W03)i2], 
according  to  Pfeiffer. 

He  proposed  an  imaginary  cube  containing  the  [XOe]12-71  group  at  the 
center  as  the  basis  for  the  structure  of  the  poly-acids,  since  this  would  allow 
coordination  numbers  of  various  magnitudes.  For  tetracoordination,  the 
four  W03  groups  would  occupy  alternate  corners  of  the  cube,  giving  a 
tetrahedral  type  of  structure  about  the  central  [X06]n_12  group.  For  a  co- 
ordination number  of  six,  the  W03  groups  would  be  at  the  face-centers  of 
the  cube,  giving  an  octahedral  structure,  and  for  twelve-coordination,  the 
W<  )3  groups  would  be  located  at  the  centers  of  the  edges  of  the  cube,  giving 
a  cubo-octahedral  arrangement. 

Although  the  structures  postulated  by  Pfeiffer  are  no  longer  believed 
correct,  his  ideas  foreshadowed  the  developments  made  by  Pauling,  Keggin, 
and  others  which  led  to  the  structures  accepted  today  for  many  of  the 
poly-acids. 

Later  Views  ox  the  Structure  of  the  Poly-Acids 

The  Work  of  Pauling 

In  1928,  Pauling24  and  later  Riesenfeld  and  Tobiank25  proposed  some 
ideas  concerning  the  structure  of  the  12-heteropoly-acids  which  are  quite 
different  from  those  of  Rosenheim,  but  which  bear  some  resemblance  to 
those  <>f  Pfeiffer.  Pauling  postulated  a  tetrahedral  [X()4]r,_s  central  ion, 
where  X  is  SnIV,  Pv,  etc.  (see  Table  14.1),  which  is  surrounded  by  twelve 
\V<  >,  octahedra,  each  octahedron  sharing  three  of  its  oxygens  with  three 
neighboring  octahedra — thus  forming  a  shell  of  these  octahedra  about  the 
central  tetrahedral  group.  Consequently,  a  total  of  eighteen  oxygen  atom- 
would  act  as  bridging  oxygens.  In  addition,  each  of  the  three  free  oxygens 

_'_>.  Groth,  "Chemische  Kristallography,"  Vol.  II,  Leipzig,  Englemann,  1908 

23.  Pfeiffer,  Z.  anorg.  aUgem.  Cfu  m.,  105,  20    1919). 

24.  Pauling,  •/.  A»<.  eh,,,,.  80c. ,  51,  2868  (1929  , 

25.  Riesenfeld  and  Tobiank,  Z    anorg.  allgem.  Chem.,  221,  287     I93fi 


480 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Table  11.7.  Pauling's  Formulation  of  Some  12-Polytungstic  Acids 

Formula  Name 

Silico-12-tungstic  acid 
Phospho-12-tungstic  acid 
12-Tungstic  acid 


H4[(Si04)W12018(OH)36] 
H3[(P04)W12018(OH)3e] 
H6[(H204)W12018(OH)36] 


1 1 


O  =  M06     OCTAHEDRA 
\^  =    X04    TETRAHEDRON 


Fig.  14.1.  Structure  of  the  12-heteropoly-acids  as  proposed  by  Pauling2^ 


on  every  octahedron  is  believed  to  take  up  a  proton  (making  a  total  of 
thirty-six  OH  groups)  which  results  in  compounds  such  as  shown  in  Table 
14.7.  The  isomorphous  isopoly-acids  were  postulated  as  having  a  similar 
structure  with  an  [H204]6_  ion  acting  as  the  central  group. 

It  was  felt  that  the  stability  of  these  ions  is  due  to  the  presence  of  the 
negative  central  group  surrounded  by  highly  charged  metal  cations  in  the 
octahedra,  and  to  the  completion  of  a  close-packed  structure,  which  is  due 
to  oxygen-oxygen  contacl  between  the  tetrahedrally  and  octahedrally  lo- 
cated  oxygons.  Figure  14.1  shows  the  location  of  the  octahedra  (each  octa- 
hedron having  an  oxygen  in  common  with  each  of  its  three  nearest  neigh- 
bors) abpui  the  central  tetrahedron,  as  proposed  by  Pauling24. 

Pauling's  structures  account  for  the  high  basicities  observed  in  the  alkali 
metal  salts  of  these  acids  quite  well.  In  addition,  those  salts  containing 
eighteen  or  more  molecules  of  water  per  acid  anion  can  readily  be  explained 


THE  POL}    ACIDS  181 

using  these  structures.  However,  Scroggie  and  Clark*6  and  Kahane  and 
Kahane-7  report  dehydrated  acids,  Buch  as  the  silico-12-tungstic  acid, 
HJSiWuO*],  and  an  8-hychate,  HJSiWuO*]  -SEW),  the  Btructures  of  which 
arc  quite  difficult  to  explain  od  the  basis  of  Pauling's  ideas,  since  t  hoc  acids 
contain  considerably  less  than  eighteen  molecules  of  water  per  acid  anion. 

Keggin's  Contributions 

Subsequently,  additional  Investigations  were  undertaken  by  Horan28  and 
Keggin28  using  x-ray  techniques.  Keggin29*  studied  the  phospho-12-tungstic 

acid  having  the  formula  1I;,|P\Y,,(  >;,  |-oII,(  >  and  found  that  the  [PWuOtf]" 
anion  has  the  following  structure  (see  Fig.  14.2):  a  central  PO*  tetrahedron 
is  surrounded  l>y  a  total  of  twelve  WOa  octahedra,  each  oxygen  of  the  PO4 
tetrahedron  being  common  to  three  of  the  W06  octahedra.  In  addition,  each 
\Y(  ),•  octahedron  has  four  of  its  remaining  five  oxygens  in  common  with  its 
four  nearest  neighbors,  while  one  oxygen  on  each  octahedron  remains  free 
(bonded  only  to  the  central  metal  atom  of  the  octahedron),  thus  making 
a  [PWrj04o]-  group. 

The  twelve  tungsten  atoms  lie  just  about  on  the  centers  of  the  edges  of 
a  large  cube  [a{)  =  12.14  A)  which  has  the  phosphorus  atom  at  the  center. 
It  can  be  seen  that  there  are  large  spaces  between  the  atoms  in  such  an 
arrangement,  which  accounts  for  the  existence  of  hydrates  containing  a 
large  number  of  water  molecules,  such  as  H^PWioO^]  -29H20 ?0.  Such  hy- 
drates should  be  readily  dehydrated  by  heat  without  undergoing  any  im- 
portant structural  changes  with  respect  to  the  framework  of  interconnected 
octahedra.  This  has  been  found  to  be  the  case  by  Signer  and  Gross81, 
Santos-,  and  .lander  and  his  co-workers33. 

A  large  number  of  heteropoly  12-tungstatea  have  been  prepared  by  Kraus 
and  his  co-workers22     ;  and  others85,  :,i  :,\  and  the  x-ray  data  for  these  salts 

_  -       ggie  and  Clark,  Proc.  Nat.  Acad.  Sri.,  Wash.,  15,  1  (1929). 

_'7.  Kahane  and  Kahane  Bull.  sac.  chim.,  [4]  49,  5.57  (1931). 

28.  Horan,  Z.  Krist.,  84,  217     1" 

29.  Keggin,  Nature,  131, 908    1933  :132,:'>51  (1933);  Proc.  Roy.  Soc.,A,  144, 75  (1934); 

DlingBworth  and  Keggin,  ./.  Chem.  Soc.,  1935,  575. 

30.  Bradley  and  OlingBworth,  P  ■      Roy.  8oc.,  .1.  157,  113  (1936). 

31.  Signer  and  Gross,  //-       Chim.  Acta,  17,  1076    L936  , 

-    P  oc.  Roy.  Soc.,  A.  150,  309    L935  . 
finder  and  Heukeshoven,  Z.  anorg.  all<i<m.  Chem.,  187,  tin  (1930  ;  Jander  and 
Banthieu,  ibid.,  225,  162    1935  ;  Jander  and  Exnei    Z  190,  195 

L942  . 
34.  Kraus,  Z.  Krist.,  A,  94,  256    L936  ;  96,  330    1937  ;  100,  394    (1939);   Kraus.   No 
27,  7io     1939  .  28,  304     1940  ;  Kraus  and  Musgnug,  ibid., 
28.  238    : 

rrari  and  Nanni,  Gaz.  chim.  Hal.,  69,  301    I  I 
Brintzinger,  Nairn  ■  if  ten,  18,   354     1930  ;   Brintzinger  and   Ratanarat, 

/.  anorg.  a  >  I  224,  97    1935  . 


t82 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Fig.  14.2A.  An  oxygen  of  the  central  tetrahedron  shown  in  common  with  three 
M06  octahedra290 


Fig.  14.2B.  The  structure  of  the  PWi2O40s  anion290 

indicate  that  they  possess  the  basic  [Xn+Wi204o]n_8  structure  proposed  for 
the  12-acids  by  Keggin,  so  it  can  now  be  considered  essentially  correct. 

Furthermore,  the  cage  structure  proposed  by  Keggin  is  complete  in  itself, 
even  if  the  four  innermost  oxygens  are  not  bonded  to  a  central  Xn+  atom. 
Therefore,  the  artificial  postulate  of  a  central  ion  formed  from  condensed 
water  molecules,  such  as  [H204]6_,  which  was  proposed  for  the  metatung- 
states,  mav  now  be  abandoned,  and  metaturigstic  acid  can  be  formulated  as 
H8[W12O40]. 

:<7.  Schulz  ;m<l  Jander,  '/..  anorg.  allgem.  Chem.,  162,  141  (1927);  Horan,  J.  .1///.  Chem. 
Soc,  61,  2022  (\<)W\;  J;.n<l(>r  and  Schulz,  Kolloid.  Z.,  36,  113  (1925). 


THE  POLY  MIPS 


483 


Fig.  14.3.  The  structure  of  the  [TeMo602<]6-  anion38 


Structural  Studies  on  the  6 -Poly -Acids 

The  6-heteropoly-acids,  such  as  Hi2_„[Xn+Mo6024],  and  the  para-isopoly- 
acids,  such  as  H6[Mo7024],  have  been  shown  to  possess  structures  which  are 
quite  different  from  those  of  the  12-poly-acids,  although  they  still  contain 
the  basic  octahedral  unit  in  their  structures. 

Anderson38  has  suggested  that  in  the  case  of  the  6-heteropoly  molyb- 
dates,  for  example,  six  Mo06  octahedra  are  located  at  the  corners  of  an 
imaginary  hexagon,  and  that  each  octahedron  shares  two  corners  (i.e.,  an 
edge)  with  each  of  its  two  nearest  neighbors,  giving  the  (Mo6024)  unit.  Such 
a  configuration  results  in  an  opening  at  the  center  of  the  hexagon  which  will 
just  accommodate  another  octahedron,  so  the  central  cation  Xn+  can  then 
be  centrally  placed  in  the  hexagon  where  it  will  share  the  six  nearest  oxygens 
of  the  (Mo60o4)  unit,  resulting  in  the  [Xn+Mo6024]n~12  anion  (see  Fig.  14.3). 

Evans39  was  able  to  verify  this  type  of  structure  for  (XH4)6[TeMo6024l  • 
7HjO,  and  it  is  interesting  to  note  that  only  those  elements  which  can  ex- 
hibit a  coordination  number  of  six  (with  valences  directed  octahedrally) 
have  been  reported  as  central  ions  in  the  6-poly-acidfl  (e.g.,  I,  Te,  Fe,  etc.  I 
lending  additional  support  to  the  above  structure. 

According  to  Lindqvist40,  the  para-molybdates,  R,-,(Mo7( )-..,],  have  a  struc- 
ture similar  to  that  of  the  heteropoly  molybdates.  In  this  case,  a  molyb- 
denum atom  is  centrally  located,  to  give  RelMoMogO*]. 

38.  Anders  e,  140,  850  (1937). 

.vans,  J.  .  Soc,  70,  1291  (1948). 

40.  Lindqvist,  Arkit .  F.  Kemi,  2,  32.5,  349  (1950). 


484  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

These  proposals  are  elaborated  upon  by  Wells41,  O'Daniel42,  and 
others43,  ll  who  include  other  acids  in  addition  to  the  limiting  6-  and  12-acid 
series. 

Additional  problems  remain  unsolved  in  this  field,  however,  especially 
with  regard  to  the  structures  of  the  unsaturated  acids  and  to  the  relation- 
ship between  the  structures  and  high  basicities  observed  for  these  com- 
pounds. 

Aggregation  Studies  of  the  Poly-Acids  in  Solution 
Methods  of  Investigation 

Although  the  structures  of  the  6-  and  12-poly-acids  in  the  solid  state  have 
been  fairly  well  established,  the  aggregation  and  degradation  phenomena  in 
solution  are  by  no  means  wTell  understood.  It  is  beyond  the  scope  of  this 
volume  to  describe  in  detail  the  investigations  carried  out  in  this  field, 
although  brief  mention  may  be  made  of  the  different  types  of  physical  and 
chemical  methods  employed  in  these  researches. 

In  attempting  to  determine  the  degree  of  aggregation  of  poly-anions  in 
solution,  Bjerrum45  and  others46- A1  utilized  pH  measurements,  but  met  with 
difficulties  due  to  the  simultaneous  occurrence  of  hydrolysis,  olation,  and 
other  poly-nuclear  aggregation  processes  (see  Chapter  13). 

Potentiometric,  conductometric,  and  thermometric  titration  methods 
have  also  been  employed333  - 33c- 48,  as  well  as  spectral  absorption  measure- 
ments, in  efforts  to  determine  the  extent  of  aggregation  of  these  acid  anions. 

Diffusion  measurements  were  used  in  an  attempt  to  obtain  the  molecular 

{•  41.  Wells,  Phil.  Mag.,  30,  103  (1940). 

42.  O'Daniel,  Z.  Krist.  A,  104,  225  (1942). 

43.  Jahr,  Naturwissenschaften,  29,  505  (1941). 

44.  Santos,  Rev .  faculdade  dene,  Univ.  Coimbra,  16,  5  (1947). 

45.  Bjerrum,  Z.  phys.  Chem.,  59,  350  (1907);  110,  657  (1924). 

46.  Souchay,  Ann.  chim.,  [11]  18,  61,  169  (1943);  Carpeni  and  Souchay,  /.  chim.  phys., 

42,  149  (1945);  Souchay  and  Carpeni,  Bull.  soc.  chim.,  13,  160  (1946);  Souchay 
and  Faucherre,  ibid.,  1951,  355;  Souchay,  ibid.,  1953,  395. 

47.  Britton,  J.  Chem.  Soc,  1930,  1249;  Vallance  and  Pritchett,  ibid.,  1935,  1586; 

Buchholz,  Z.  anorg.  allgem.  Chem.,  244,  168  (1940);  Bye,  Bull.  soc.  chim.,  9,  360 
(1942);  Britton  and  Wellford,  J.  Chem.  Soc,  1940,  764;  Ripan  and  Liteanu, 
Compt.  rend.,  224,  196  (1947). 

48.  Mayer  and  Fisch,  Z.  anal.  Chem.,  76,  418  (1929);  Bye,  Ann.  chim.,  [11]  20,  463 

(1945);  Britton,  Endeavor,  2,  148  (1943);  Ghosh  and  Biswas,  J.  Indian  Chem. 
Soc,  22,  287,  295  (1945);  Dullberg,  Z.  phys.  Chem.,  45,  119  (1903) ;  Murgulescu 
and  Alexa,  Z.  anal.  Chem.,  123,  341  (1942);  Carrier  and  Guiter,  Bull.  soc.  chim., 
12,  329  (1945) ;  Pierce  and  Yntema, J.  Phys.  Chem.,  34,  1822  (1930) ;  Britton  and 
Robinson,  /.  Chem.  Soc,  1932,  2265;  Bye,  Bull.  soc.  chim.,  1953,  390;  Hormann, 
Z.  anorg.  Chem.,  177,  145  (1928). 


THE  POLY  ACIDS  L85 

weights  of  tlui  poly-acids  in  solution.  Prytz49  and  Jander  and  Jahr  and  co 
workers  utilized  Riecke's  Law1'  thai  the  square  rool  of  the  molecu- 

lar weight  of  a  substance  is  inversely  proportional  to  its  diffusion  coefficient , 
ami  they  it'll  that  they  were  able  to  estimate  molecular  weights  with  an 

accuracy  of  about  5  per  cent. 

Brintzinger  and  his  co-workers88,  H  were  fairly  successful  in  utilizing  elec- 
trodialysis  methods  for  the  determination  of  molecular  weights,  and  tins 

method  was  later  used  by  Jander33c>  53  for  the  same  purpose. 

Gupta64  and  Theodoresco88  have  investigated  poly-acids  and  their  Baits 
in  solution  and  in  the  crystalline  state  by  means  of  Raman  spectra,  hut  it 
appears  difficult  to  draw  definite  conclusions  concerning  the  degree  of  ag- 
gregation of  these  materials  in  solution  from  thei^  spectra. 

Doucet86  and  other  workers4sb- 57  attempted  cryoscopic  determinations  of 
molecular  weights,  and  obtained  results  which  were  in  agreement  with  those 
obtained  polarographically  by  Souchay. 

Magneto-chemical  studies  were  carried  out  by  Das  and  Ray58,  who  noted 
changes  in  magnetic  susceptibility  with  changes  in  pH,  and  phase  studies 
were  performed  by  Kiehl  and  Maufredo59  and  Makarow  and  Repa60  which 
gave  evidence  for  the  existence  of  poly-anionic  aggregates. 

Preparations  of  the  Poly -Acids 

Many  other  investigations  have  been  conducted,  employing  variations  of 
oik1  or  more  of  the  above  methods.  Furthermore,  a  large  number  of  studies 

19.  Prytz,  Z.  anorg.  aUgem.  Chem.,  174,  360  (1928). 

.">i>.  Jander  and  Jahr,  Koll.  Beihefte,  41,  1  (1934);  Jander,  Mojert,  and  Aden,  Z.  anorg. 
aUgem.  Chem.,  162,  141  (1927);  Jahr  and  Witzmann,  ibid.,  208,  145  (1932);  Jander 
and  Jahr,  Koll.  Beihefte,  41,  297  (1935);  Jander  and  Drew,  Z.  phys.  Chem.,  190, 
217  1942  :  Jander  and  Jahr,  Z.  anorg.  allgem.  Chem.,  220,  201  (1934);  212,  1 
1933);  Jahr  and  Witzmann,  Z.  phys.  Chem.,  168,  283  (1934);  Jander  and  Aden, 
ibid.,  144,  197  (1929);  Jander  and  Schulz,  Z.  anorg.  allgem.  Chem.,  144,  225 
(1925). 

51.  Riecke,  Z.  phys.  Chem.,  6,  564  (1890). 

52.  Brintzinger,  Z.  anorg.  allgem.  Chem.,  196,  55  (1931);  Brintzinger  and  Wallok, 

ibid.,  224,  103  (1935). 
ader,  Z.  phys.  Chem.,  187,  149  (1940). 
54.  Gupta,  •/.  Indian  Chem.  8oc.,  12,  223  (1938). 

.v..  Theodoresco,  Compt.  rend.,  208,  1308  (1939);  210,  175  (1940);  210,  297  (1940);  211, 
28    L940  :214,  109    1 9  12  ;  215,  530  (1942);  216,  56  (1943). 

.  208,  :»77  - 1939);  •/.  phys.  radium,  [8]  4,  41  (1943). 
Ann.  chim.,  [11]  20,  74,  96    1945  ;  [12]  1,  232,  249  (1946);  [12]  2,  203,  229 
1947  . 
58.  Das  and  Ray,/.  Indian  Chem.  Soc.,21,  159  (1944  . 

Kiehl  and  Manfredo,  ./.  .1//'.  Chem.  8oc.,  59,  21  is    1933 
BO.  Makarow  and  Repa,  Bull.  ae.  sci.  U.R.S.S.,  1940,  349. 


486  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

concerned  with  the  preparation  and  properties  of  heteropoly-  and  isopoly- 
acids,  in  addition  to  those  already  mentioned,  have  been  carried  out  in 
recent  years61.  Among  these  are  reports610- 61p- 61y  of  some  interesting  com- 
pounds composed  of  heteropoly  anions  and  chelate-containing  cations,  such 
as  lCu(en)2]2[SiWi204o]-2H20. 

These  studies  have  greatly  increased  our  knowledge  of  the  poly-acids  and 
their  salts.  However,  much  remains  to  be  clarified,  especially  with  regard  to 
the  solution  chemistry  of  these  acids  and  salts,  and  it  is  hoped  that  research 
workers  will  continue  to  investigate  the  many  unsolved  problems  in  this 
field. 

61.  Bje,  Bull.  soc.  chim.,  10,  239  (1943) ;  Klason,  Ber.,  34,  153  (1901) ;  Junius,  Z.  anorg. 
allgem.  Chem.,  46,  428  (1905);  Wempe,  ibid.,  78,  298  (1912);  Sand  and  Eisen- 
lohr,  ibid.,  52,  68  (1907) ;  Jande*,  Jahr,  and  Heukeshoven,  ibid.,  194,  383  (1930) ; 
Ullik,  Ann.,  153,  373  (1870);  Travers  and  Malaprade,  Compt.  rend.,  183,  292, 
533  (1926);  Garelli  and  Tettamanzi,  Chem.  Abstr.,  29,  7864  (1935);  Ray  and 
Siddhanta,  J.  Indian  Chem.  Soc.,  18,  397  (1941);  Ray,  ibid.,  21,  139  (1944); 
Guiter,  Ann.  chim.,  [11]  15,  5  (1941);  Rosenheim,  Z.  anorg.  allgem.  Chem.,  96, 
139  (1916);  220,  73  (1934);  96,  139  (1916);  Lachartre,  Bull.  soc.  chim.,  35,  321 
(1924) ;  Parks  and  Prebluda,  J.  Am.  Chem.  Soc,  57,  1676  (1935) ;  Huffman,  ibid., 
60,  2227  (1938);  Guiter,  Compt  rend.,  209,  561  (1939);  Marignac,  Ann.  chim., 
[4]  8,  5  (1866) ;  Windmaisser,  Oster.  Chem.  Ztg.,  45,  201  (1942);  Balke  and  Smith, 
J.  Am.  Chem.  Soc,  30,  1651  (1908);  Russ,  Z.  anorg.  Chem.,  31,  60  (1902);  Sue, 
Ann.  chim.,  7,  493  (1937);  Sue,  Compt.  rend.,  208,  440  (1939);  Ferrari,  Cavelca, 
and  Nardelli,  Gazz.  chim.  ital.,  78,  551  (1948);  79,  61  (1949);  80,  352  (1950); 
Jean,  Ann.  chim.,  [12]  3,  470  (1948). 


(/ 


15 


Coordination  Compounds  of  Metal 
Ions  with  Olefins  and  Olefin-Like 
Substances 

Bodie  E.  Douglas 
The  University  of  Pittsburgh,  Pittsburgh,  Pennsylvania 

Coordination  compounds  of  olefins  with  compounds  of  the  heavy  metals 
were  discovered  before  the  advent  of  Werner's  theory,  but  the  problem  of 
explaining  how  they  are  formed  and  why  they  are  stable  is  still  perplexing. 
Ethylene  has  no  unshared  pair  of  electrons  which  it  can  share  with  the 
metal  as  do  ammonia  and  other  common  ligands.  Olefinic  complexes  take 
on  added  importance  since  some  workers  believe  that  they  supply  a  crucial 
test  of  the  generally  accepted  view  that  the  coordinate  covalent  bond  re- 
sults from  the  sharing  of  a  "lone  pair"  of  electrons  furnished  by  the  ligand. 
Excellent  reviews  on  these  compounds  have  been  written  by  Keller1  and 
by  Chatt*. 

The  complexes  of  platinum  with  unsaturated  molecules  are  generally 
more  stable  than  those  of  other  metals,  and  the  olefins  generally  form  more 
stable  complexes  than  do  unsaturated  alcohols,  aldehydes,  acids,  esters, 
halogenated  hydrocarbons,  and  aromatic  substances.  Because  of  their  sta- 
bility, the  platinum-olefin  complexes  have  been  studied  most  extensively. 

Compounds  That  Have  Been  Reported 

Platinum  -olefin  Compounds 

The  first  report  of  a  platinum-olefin  compound  was  published  by  Zeise3 
in  1827.  The  work  was  further  described  in  later  publications4.  In  1830 
Berzelius5  announced  that  by  refluxing  a  mixture  of  alcohol  and  sodium 
hexachloroplatinate(IV),  a  very  acid  solution  was  formed;  this  yielded 

1.  Keller,  Chen  ,  &     .  28,  229  (1941). 

1.  C'hatt,  J.  Chem.Soc,  1949,  33-40. 

3.  Zeise,  Pogg.  Ann.,  9,  632  (1827). 

4.  Zei.se,  Magaz.  f.  Phar,,,.,  35,  105  (1830);  Pogg.  Ann.,  21,  497  (1931);  Schweig*, 

Journal  der  Chemic  v.  Physik,  62,  303  (1831);  63,  121  (1831). 

5.  Berzelius,  Jahresber,  9,  162  (1830). 

487 


488  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

yellow  crystals  when  concentrated  and  treated  with  potassium  chloride. 
The  analysis  of  this  compound  conformed  to  the  composition  reported  by 
Zeise.  Zeise  had  prepared  the  compound  (reported  on  an  anhydrous  basis 
as  KCl-PtCl2-C2H4)  by  boiling  platinum (IV)  chloride  with  alcohol  and 
adding  potassium  chloride.  The  analyses  were  challenged  by  Liebig6,  but 
Zeise7  repeated  them  and  confirmed  the  presence  of  ethylene.  The  potas- 
sium and  ammonium  salts  usually  obtained  by  such  a  procedure  are  the 
1 -hydrates,  which  probably  accounts  for  Liebig's  insistence  that  the  radical 
C4H10O  was  present  and  that  the  correct  formula  was  2KC1  -2PtCl2-  C4Hi0O. 

Zeise  also  prepared  a  compound  reported  as  PtCl2  •  C2H4 ,  but  it  is  more 
likely  that  this  was  impure  H[PtC2H4Cl3],  now  known  as  "Zeise's  acid." 
The  nonionic  compound  [Pt(NH3)(C2H4)Cl2]  was  also  reported  by  Zeise. 

Zeise's  formula  was  confirmed  by  Griess  and  Martius8,  who  also  demon- 
strated that  ethylene  was  liberated  during  the  thermal  decomposition  of 
Zeise's  salt.  Some  doubt  concerning  the  presence  of  ethylene  in  the  original 
compound  still  existed,  however,  since  appreciable  amounts  of  platinum 
and  carbonaceous  substances  were  among  the  decomposition  products. 
Birnbaum9  proved  the  presence  of  ethylene  when  he  synthesized  Zeise's 
salt  by  treating  platinum (II)  chloride  in  hydrochloric  acid  solution  with 
ethylene,  followed  by  the  addition  of  potassium  chloride.  Birnbaum  also 
prepared  the  propylene  and  amylene  analogs  of  Zeise's  salt.  He  described 
Zeise's  preparation  by  the  equation 

PtCl4  +  2C2H5OH  -*  PtCl2.C2H4  +  CH3CHO  +  H20  +  2HC1 

Allyl  alcohol10  and  unsaturated  acids11  with  the  double  bond  in  the  /3-po- 
sition,  or  farther  from  the  carboxyl  group,  form  compounds  similar  to  those 
of  ethylene.  Additional  analogs  of  Zeise's  salt,  containing  unsaturated 
acids,  esters,  alcohols,  and  aldehydes,  have  been  prepared12. 

The  compound  containing  only  platinum (II)  chloride  and  ethylene, 
PtCl2-C2H4  (actually  shown  later  to  be  a  dimer),  was  prepared  by  An- 
derson13 by  reducing  sodium  hexachloroplatinate(IY)  with  alcohol.  The 
resulting  solution  was  evaporated  in  a  high  vacuum  and  the  ethylene- 
platinum(II)  chloride  was  extracted  with  chloroform  from  the  tarry, 
strongly  acid  mass.  Anderson14  was  also  able  to  isolate  PtCl2  •  C6H5CII=CII2 

6.  Liebig,  Ann.,  9,  1  (1834);  23,  12  (1837). 

7.  Zeise,  Ann.,  23,  1  (1837);  Pogg.  Ann.,  40,  234  (1837). 

8.  Griess  and  Martius,  Ann.,  120,  324  (1861);  Compt.  rend.,  53,  122  (1861). 

9.  Birnbaum,  Ann.,  145,  67  (1869). 

10.  Biilmann,  Ber.,  33,  2196  (1900). 

11.  Biilmann  and  Hoff,  Rec.  trav.  chim.,  36,  306  (1916). 

12.  Pfeiffer  and  Hoyer,  Z.  anorg.  allgem.  chem.,  211,  241  (1933). 

13.  Anderson,  J.  Chem.  Soc,  1934,  971. 

14.  Anderson,  J.  Chem.  Soc.,  1936,  1042. 


MP0UND8  OF  METAL  IONS  WITH  OLEFIN 8  189 

by  the  essentially  quantitative  displacement  of  ethylene  by  styrene 
from  Ptl'l.-C-jHi  .  By  the  Bame  method,  Anderson  prepared 
K(Pt(CeHiCH==CHi)Cls]  from  Zeise's  salt  and  styrene.  He  established  an 

order  of  stability  for  the  complexes  based  on  the  displacement  reaction- 
and  considerations  o\  the  relative  volatility  of  the  hydrocarbons.  The  sta- 
bility decreased  from  ethylene  in  the  order  CH^CHa  >  C6H5CH=CH2  > 
indene  >  cyclohexene  >   (Cf,II5)2C=CH2  ,  (C6H5)(CH3)C=CH2  . 

A  fairly  general  method  of  preparation  of  the  olefin  complexes  was  de- 
vised by  Kharasch  and  Ashford'5,  who  treated  anhydrous  plat  inum(IV) 
chloride  or  bromide  with  the  unsaturated  substance  in  an  anhydrous 
solvent.  Chloro-substitnted  olefins  react  satisfactorily,  but  unsaturated  acids 
and  esters  do  not  yield  complexes  by  this  method. 

A  variety  of  compounds  of  platinum  with  unsaturated  substances  has 
been  prepared  by  Russian  workers.  Chernyaev  and  Hel'man16  prepared 
Zeise's  salt  by  passing  ethylene,  for  15  days,  through  a  concentrated 
aqueous  solution  of  potassium  tetrachloroplatinate(II)  containing  3  to  5 
per  cent  of  hydrochloric  acid,  followed  by  precipitation  of  [Pt(XH3)4] 
[PtC2H4Cl3]2  •  Compounds  of  the  type  [Pt  R  C2H4  X2]  were  also  prepared. 
The  stability163  of  these  compounds  was  reported  to  decrease  in  the  order: 
R  =  quinoline  >  pyridine  >  ammonia  >  thiourea  and  X  =  CI-  >  Br~  > 
I~~  >  NOj"  >  NCS~  >  CX~.  From  the  study  of  a  series  of  complexes 
containing  several  unsaturated  substances,  Hel'man17  arrived  at  a  stability 
series  differing  from  Anderson's  in  that  styrene  was  placed  above  ethylene. 
The  order  given  by  her  is  XO  >  CO  >  styrene  >  butadiene,  ethylene  > 
propene  >  butene.  The  difference  is  probably  due  to  the  qualitative  nature 
of  the  work,  since  relative  volatilities  and  solubilities  were  not  considered. 
Butadiene  was  found  to  occupy  only  one  coordination  position  per  metal 
ion  instead  of  forming  a  chelate  ring,  although  a  compound  was  isolated 
in  which  one  butadiene  was  coordinated  to  two  platinum  atoms,  forming 
the  bridged  (XH4)2[(PtCl3)2C4H6]18.  The  bridged  butadiene  complex  was 
found  to  react  with  ethylenediamine  to  give  a  long-chain  polymer, 
[— CH2XH2PtCl2— CH2=CHCH=CH2— PtCl2XH2CH2— ]B19.  Similarly 
Zeise's  salt  was  found  to  react  with  ethylenediamine  to  give  a  bridged  com- 
pound, [C2H4Cl2Pt— XH2C2H4XH2— PtCl2C2H4],  rather  than  the  expected 
chelate  compound. 

15.  Kharasch  and  Ashford,  J.  Am.  Chem.  Soc,  58,  1733  (1936). 

16.  Chernyaev  and  Hel'man,  Ann.  secteur  platine,  Inst.  chim.  gen.  (U.S.S.R.) ,  Xo.  14, 

77  (1937);  Herman,  Sci.  Repts.  Leningrad  State  Univ.,  2,  No.  2,  5  (1936); 
Chernyaev  and  Bel'man,  Compt.  rend.  Acad.  Sci.,  U.R.S.S.(N.S.),  4,  181 
(1936). 

17.  Hel'man,  Compt.  rend.  aead.  *<■[.,  UJt.8.S.t  20,  307  (1938);  32,  347  (1941). 

18.  Hel'man,  Compt.  rend.  acad.  sci.,  U.R.S.S.,  23,  532  (1939). 

19.  Hel'man,  Doklady  Akad.  Nauk  S.S.S.R.,  38,  272  (1943). 


490 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Cationic  complexes20  have  been  prepared  by  the  following  reactions: 

cis-[PtNH3C2H4Cl2]    AgN°3)  [PtNH3C2H4ClN03l  -^* 

[PtH2ONH3C2H4Cl]  N03     pyridine  )  [Pt  py  NH3C2H4C1]N03  . 

Hel'man  reported  that  the  final  compound  was  a  white  crystalline  sub- 
stance which  was  very  soluble  in  water  and  which  decomposed  on  standing 
in  air.  It  reacted  with  chloride  ion  to  give  the  original  starting  material. 
All  three  possible  isomers  of  the  compound  [PtNH3C2H4ClBr]  were  isolated 
by  Hel'man  and  co-workers21.  The  compound  with  the  halides  in  trans 
positions  was  obtained  by  treating  Zeise's  salt  with  potassium  bromide  and 
then  with  ammonia.  The  other  isomers  were  prepared  as  follows: 


"C2H4             CI" 
\    / 

NH4[PtNH3Cl3]  -^>  trans  K[PtNH3Br2Cl]  -^^U 

Pt 

/    \ 
_  NH3              Br. 

cis-[PtNH3C2H4Br2]    A^°3> 

"C2H4             Br 

\    / 

Pt 

/    \ 
_  NH3             H20_ 

N03  -^> 

'C2H4              Br 

\    / 
Pt 

/    \ 
_  NH3             CI 

Cis  and  trans  isomers  of  the  compounds  [Pt  R  C2H4  Cl2],  where  R  is  am- 
monia or  pyridine,  have  also  been  obtained  by  Chernyaev  and  Hel'man22. 
In  the  preparation  of  the  isomers  of  the  platinum  ethylene  compounds,  the 
Russian  workers  have  taken  advantage  of  the  high  trans  effect  of  ethylene, 
resulting  in  easy  substitution  in  the  position  trans  to  ethylene.  The  trans 
compounds  result  from  the  addition  of  an  amine  to  Zeise's  salt,  while  the 
cis  isomers  are  formed  by  the  addition  of  ethylene  to  compounds  of  the 
type  K[PtNH3Cl3].  (Chapter  4) 

Hel'man  and  Essen23  studied  the  complexes  of  allylamine  with  platinum. 
Addition  of  allylamine  to  K2PtCl6  gave  [Pt(C3H7N)2Cl2]  in  which  the  allyl- 
amine was  said  to  coordinate  only  through  the  nitrogen.  A  similar  reaction 
carried  out  in  strongly  acidic  solution  produced  [Pt(C3H7N-HCl)Cl2]2  in 
which  the  coordination  presumably  involved  only  the  double  bond.  This 
product  was  converted  to  H[Pt(C3H7N-HCl)Cl3](I)  by  heating  with  10  per 
cent  hydrochloric  acid.  Careful  neutralization  of  (I)  with  5  per  cent  alkali 


20.  Hel'man  and  Meilakh,  Compt.  rend.  acad.  sci.  U.S.S.R.,  51,  207  (1946). 

21.  Herman,  Doklady  Akad.  Nauk  S.S.S.R.,  38,  327  (1943);  Hel'man  and  Gorush- 

kina,  Compt.  rend.  acad.  sci.  U.S.S.R.,  55,  33  (1947). 

22.  Chernyaev  and  Hel'man,  Ann.  secteur  platine,  Inst.  chim.  gen.  U.S.S.R.,  No.  15, 

5  (1938). 

23.  Hel'man  and  Kssen,  Doklady  Akad.  Nauk  S.S.S.R.,  77,  273  (1951). 


I  OMPOUNDS  OF  METAL  IONS  11777/  OLEFINS  l«.M 

produced  [Pt(C8H7N)Cli](II)  in  which  the  allylamine  was  presumed  to 
function  as  a  bidentate  group,  coordinating  through  the  nitrogen  and  the 
double  bond.  Hel'man  stated  that  this  was  proved  by  the  fact  that  allyl- 
amine hydrochloride  displaced  ethylene  from  XH4[PtC2H4Cl;{]  to  produce 
the  ammonium  salt  of  (I),  which  produced  (II)  on  neutralization.  Actually 
these  reactions  do  not  eliminate  a  dimeric  structure  for  (ID  similar  to  that 
of  [PtCjHUCIJa  ,  in  which  the  ethylene  is  monodentate.  The  platinum  com- 
plexes of  diallylamine  (abbreviated  dim)  have  been  studied  by  other  in- 
vestigators14 who  report  that  the  action  of  two  moles  of  diallylamine  on  one 
mole  of  ammonium  tetrachloroplatinate(II)  gave  a  dark  precipitate  and 
more  slowly  a  light-yellow  precipitate  of  the  same  empirical  composition, 
PtCli'dlm.  The  light-yellow  material  was  shown  to  be  a  dimer  by  the  fact 
that  it  could  be  prepared  by  the  addition  of  (NH4)2[PtCl4]  to  a  solution  of 
[Pt(dlm)2]Cl2  (prepared  from  [Pt  dim  Cl2]  and  an  excess  of  dim)  to  give 
[Pt(dlm)2][PtCl4].  The  dark  precipitate  could  be  converted  to 
[Pt(XH3)2  dlm]Cl2  by  treatment  w^ith  ammonium  hydroxide.  Thus,  in  each 
compound  the  diallylamine  apparently  occupies  two  coordination  positions, 
at  least  one  of  which  must  be  filled  by  an  olefinic  linkage.  It  is  unlikely  that 
both  double  bonds  function  as  donor  groups  since  large  chelate  rings  are 
not  frequently  encountered  and  the  ability  of  the  nitrogen  to  coordinate  is 
doubtless  greater  than  that  of  the  olefinic  linkage.  The  data  reported  for 
the  diallylamine  complexes  lend  support  to  the  structure  proposed  by 
Hel'man  for  the  allylamine  complexes. 

Chatt  and  Wilkins25  prepared  the  first  compound  containing  two  double 
bonds  linked  to  the  same  platinum  atom,  although  Anderson14  had  found 
some  evidence  for  the  existence  of  the  compound  PtCl2  •  2C6H5CH=CH2 , 
which  he  could  not  isolate.  Hel'man26  disputed  the  existence  of  such  a  com- 
pound on  theoretical  grounds.  The  compound  described  by  Chatt  and 
Wilkins,  [Pt(C2H4)2Cl2],  was  prepared  by  passing  ethylene  through  a  solu- 
tion of  [PtC2H4Cl2]2  in  acetone  at  —70°.  It  dissociates  at  —6°  in  an  ethylene 
atmosphere  and  probably  has  a  trans  configuration.  Chatt  and  Wilkins 
considered  the  low  stability  of  the  compound  to  be  due  to  the  high  trans 
effect  of  ethylene  and  the  relatively  weak  bond  between  platinum  and 
ethylene.  They  were  able  to  prepare  two  complexes  of  platinum  with  di- 
pentene,  both  of  which  had  the  same  empirical  composition,  Pt(CioHi6)Cl2 . 
One  of  these  was  monomeric  and  must  have  been  a  complex  in  which  the 
dipentene  functioned  as  a  chelate  group  unless  it  was  simply  an  addition 
compound.  Kharasch  and  Ashford15  had  prepared  a  dipentene  compound 
of  the  same  composition,  but  assumed  it  to  be  a  dimer. 

24.  Etabinshtein  and  Derbisher,  ibid.,  74,  283  (1950). 

25.  Chatt  and  Wilkins,  Nature,  165,  860  (I960);  J.  Chem.  Soc.t  1952,  2622. 

26.  Hel'man,  Compt.  rend.  aead.  set.  U.R.S.S.,  24,  540  (1939). 


192  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Hel'man  and  her  co-workers27  prepared  a  compound  analogous  to  Zeise's 
salt  containing  an  acetylene  derivative,  2,5-dimethyl-3-hexyne-2,5-diol. 
The  product  was  treated  with  pyridine  to  form  [Pt  C8Hi402  py  CI2].  The 
molecular  weight  of  the  pyridine  compound,  determined  cryoscopically  in 
benzene  solution,  indicated  it  to  be  a  monomer.  Its  properties  led  the 
authors  to  assume  that  it  had  a  trans  configuration. 

The  properties  of  the  olefinic  complexes  of  platinum  are  extremely  in- 
teresting. The  simplest  stable  compounds,  [PtUnCl2]2  (Un  represents  an 
unsaturated  group),  are  decomposed  by  water,  but  are  soluble  in  the  com- 
mon organic  solvents  except  glacial  acetic  acid,  and  only  moderately  soluble 
in  cold  benzene.  Most  of  the  compounds  are  thermally  unstable  and  de- 
compose before  melting.  Some  decompose  on  standing  for  several  days, 
but  the  dipentene  complex  remains  unchanged  after  standing  in  air  for  ten 
ir  months15. 

The  olefins  in  most  olefinic  complexes  can  be  substituted  readily  by  other 
olefins14  or  by  coordinating  agents  such  as  pyridine15  or  chloride  ion  (when 
treated  with  concentrated  hydrochloric  acid).  These  reactions  liberate  the 
coordinated  olefin  unchanged.  Bromine  decomposes  the  complexes  with  the 
formation  of  the  brominated  olefin.  The  ethylene  complex  is  rapidly  and 
quantitatively  reduced  by  hydrogen  at  room  temperature  to  platinum, 
hydrogen  chloride,  and  ethane13. 

Zeise's  salt  reacts  with  potassium  cyanide  to  liberate  ethylene  quanti- 
tatively, and  other  complexing  agents,  such  as  pyridine,  tend  to  react  simi- 
larly13. Hot  water  decomposes  the  salt  according  to  the  equation 

K[PtC2H4Cl3]  +  H20  ->  KC1  +  2HC1  +  Pt  +  CH3CHO. 


Anderson's  stability  series14,  as  well  as  the  results  of  Kharasch  and  Ash- 
ford15,  indicate  that  in  general  the  stability  of  the  platinum-olefin  com- 
pounds decreases  with  increasing  substitution  adjacent  to  the  double  bond. 
The  effect  seems  to  be  largely  steric.  However,  the  behavior  of  cis-trans 
isomers  does  not  appear  to  be  completely  consistent.  Kharasch  and  Ashford 
were  able  to  isolate  complexes  with  cyclohexene,  dipentene,  pinene,  ethyl- 
ene, isobutylene,  styrene,  and  frans-dichloroethylene.  The  first  three  com- 
pounds have  a  cis  configuration,  but  cis-dichloroethylene  and  czs-diphenyl- 
ethylene  have  not  yet  yielded  complexes,  although  those  of  the  trans 
compounds  are  known.  Anderson  isolated  the  indene  (a  cis  compound) 
complex  and  reported  that  a  crystalline  complex  formed  with  a  compound 
which  he  stated  to  be  presumably  ^mns-2-pentene.  Oppegard28  prepared  a 
crystalline  complex  with  m-2-pentene,  but  obtained  only  a  red  oil  with 
/rans-2-pentcne. 


27.  Herman,  Bukhovets  and  Meilakh,  ibid.,  ±6,  105  (1945). 

28.  I  tppegard,  thesis,  University  of  Illinois  (1946). 


I  OMPOUNDS  OF  METAL  IONS  WITH  OLEFINS  193 

Palladimii-olHiii  ( lompounds 

The  first  palladium-olefin  compound  reported  was  PdGls'CfHio  which 

was  said  to  be. formed  when  palladium(II)  chloride,  trimethylethylene  and 
a  trace  oi  some  basic  substance  were  allowed  to  react-'1.  I  [owever,  Kharasch, 
Seyler,  and  Mayo'1  were  not  able  to  repeat  this  work.  Although  they  were 
not  able  to  cause  palladium!  1 1 1  chloride  to  react  directly  with  unsaturated 
compounds,  they  found  that  bis-benzonitrile  palladium(II)  chloride  reacted 
readily  with  olefins.  Palladium(II)  complexes  of  the  type  [PdClj'Un]j  were 
prepared  with  cyclohexene,  ethylene,  styrene,  butylene,  pinene  and  cam- 
phorene.  The  stability  of  the  complexes  decreased  in  the  order  given  and 
when  a  less  stable  compound  was  treated  with  the  olefin  substituent  of  a 
more  stable  one,  the  latter  compound  was  formed  by  replacement.  The 
complexes  wen1  colored,  unstable,  and  rather  insoluble  in  the  common 
organic  solvents.  They  were  less  stable  than  the  corresponding  platinum 
compounds. 

Iron -olefin  Compounds 

The  compound  FeCVCoH^HoO  was  reported  by  Kachler31  to  be 
formed  by  the  reaction  of  iron(III)  chloride  with  ether  in  the  presence  of  a 
small  amount  of  phosphorus  in  a  sealed  tube.  The  equation  was  given  as 

2(  JI5OC0H5  +  2FeCl3  ->  2FeCl.-CoH4  +  2C2H5OH  +  Cl2  . 

Alcohol  did  not  give  the  same  product  under  similar  conditions.  Chojnacki '•'•- 
was  unable  to  prepare  Kachler's  compound  from  iron(II)  chloride  and 
ethylene,  but  did  prepare  the  bromide,  FeBr2-C2H4-2H20.  He  reported 
that,  when  treated  with  potassium  bromide,  a  solution  of  this  compound 
gave  almost  colorless  crystals  containing  iron,  bromine,  potassium,  and 
ethylene.  Manchot  and  Haas33  were  unable  to  duplicate  the  work  of  Kachler 
and  Chojnacki  and  felt  that  Kachler's  compound  was  a  partially  decom- 
posed ether  addition  compound. 

The  compound  Fe(CO)3-C4H6  has  been  reported34  to  be  formed  by  long 
heating  of  iron  pentacarbonyl  with  butadiene.  Less  well-defined  compounds 
were  obtained  with  other  olefins. 

The  most  interesting  olefinic  compound  of  iron  was  reported  only  re- 
cently. Kealy  and  Pauson*8  added  a  solution  of  iron(lll)  chloride  in  an- 
hydrous ether  to  a  benzene  solution  of  cyclopentadieny]  magnesium  bro- 

29.  Kondakov,  Bolaa,  and  Vit,  Chi  m.  List;/,  23,  579    L929);24,  1,  26  (1930). 

30.  Kharaach,  Seyler,  and  Mayo,  /.  Am.  Ch       -       80,882(1938). 

31.  Kachler,  Ber.,  2,  510  (1869);  ./.  prakt.  ch  m.t  107,  315  (1869). 

32.  Chojnacki.  Jahrcsber.,  23,  510  (1870);  Z.  Chem.,  2,  6,  419    1870 
Manchot  and  Haas,  Ber.,  45,  3052  (1912). 

34.  Reihlen,  Gruhl,  Heading,  and  Pfrengle,  .1/.//..  482,  nil     1 

35,  Kealy  and  Pan*  •  .  168,  1039    1951). 


494  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

mide.  The  solution  was  allowed  to  stand  overnight,  was  refluxed  for  an 
hour,  and  was  then  treated  with  an  ice-cold  solution  of  ammonium  chloride, 
after  which  evaporation  gave  an  orange  solid  which  melted  at  173-174°C 
with  sublimation.  The  composition  of  the  solid  was  FeCioHio  .  Miller, 
Tebboth,  and  Tremain36  found  that  reduced  iron,  in  the  presence  of  po- 
tassium oxide,  reacted  with  cyclopentadiene  in  nitrogen  at  300°C  to  give  a 
yellow  solid,  FeCioH]0 ,  which  melted  at  172.5-173°C  with  sublimation. 
Bis(cyclopentadienyl)iron(II)  is  soluble  in  alcohol,  ether,  and  benzene.  It  is 
insoluble  in,  and  unattacked  by  water,  10  per  cent  sodium  hydroxide,  or 
concentrated  hydrochloric  acid.  It  dissolves  in  dilute  nitric  acid  or  con- 
centrated sulfuric  acid  to  give  a  deep  red  solution  with  strong  blue  fluo- 
rescence. It  decolorizes  permanganate.  Wilkinson  and  co-workers37  found 
the  compound  to  be  diamagnetic.  It  is  easily  oxidized  to  a  blue  cation 
Fe(C5H5)2+  (polarographic  half -wave  potential,  —0.59  volt),  which  is  para- 
magnetic with  a  magnetic  moment  suggesting  the  presence  of  one  unpaired 
electron.  The  structure  of  the  compound  will  be  considered  later  (page  507). 

Iridium-olefin  Compounds 

Several  iridium-olefin  compounds  have  been  reported38.  Treatment  of 
iridium(III)  chloride  with  absolute  alcohol  produced  IrCU^EU  which, 
when  treated  with  ammonium  or  potassium  chloride,  gave  mixtures  of 
other  products.  Formulas,  for  the  products  isolated,  indicated  the  presence 
of  iridium  chloride,  ammonium  or  potassium  chloride,  ethylene,  and  some- 
times water.  No  compounds  of  iridium  could  be  obtained  from  ethylene 
and  iridium(III)  chloride  or  a  solution  of  iridium(III)  chloride. 


Copper -olefin  Compounds 

The  absorption  of  ethylene  and  propylene  by  a  hydrochloric  acid  solu- 
tion of  copper  (I)  chloride  was  observed  by  Berthelot39.  The  mole  ratio  of 
ethylene  to  copper  (I)  chloride  was  0.17  and  of  propylene  to  copper  (I) 
chloride,  0.25.  An  unstable  compound,  CuCl-C2H4 ,  was  reported  by  Man- 
chot  and  Brandt40,  although  they  could  not  isolate  it.  It  has,  however,  been 
isolated  from  the  reaction  of  ethylene  under  pressure  with  solid  copper(I) 
chloride41.  It  is  not  known  whether  this  substance  is  a  coordination  com- 
pound or  only  an  addition  compound.  The  absorption  of  propylene  and 

36.  Miller,  Tebboth,  and  Tremaine,  J.  Chem.  Soc,  1952,  632. 

37.  Wilkinson,  Rosenblum,  Whiting,  and  Woodward,  ./.  Am.  Chem.  Soc,  74,  2125 

(1952). 

38.  Sadtier,  Chem.  News,  24,  280  (1871);  Bull.  soc.  chim.,  17,  54  (1872). 

39.  Berthelot,  Ann.  chim.  phys.,  23,  32  (1901). 
10.  Man.hot  and  Brandt,  Ann.,  370,  286  (1909). 

41.  Tropsch  and  Mattox,  J.  Am.  Chem.  Soc,  57,  1102  (1935). 


I  OMPOUNDS  OF  METAL  TONS  11/77/  OLEFINS  495 

isobutylene43  and  butadiene4"  by  solid  copper(I)  chloride  has  also  been 
demonstrated.  Gilliland  and  co-workers44  prepared  a  complex  containing 
two  moles  of  copper(I)  chloride  and  one  mole  of  butadiene.  Prom  the 
studies  of  vapor  pressures  of  olefins  over  copper(I)  chloride,  they  found 
that  one  mole  of  copper(I)  chloride  absorbed  0.336  mole  of  isoprene,  0.62 

molt4  of  isobutylene,  and  formed  1:1  complexes  with  ethylene  and  pro- 
pylene. Neither  cyclopentadiene  nor  amylene  reacted.  Ward  and  Makin41 
characterized  complexes  containing  one  mole  of  1 ,3-pentadiene  or  isoprene 

to  two  moles  of  copper(I)  chloride. 

Osterlof46  identified  two  compounds,  3CuClC2H2  and  2CuClC2H2, 
formed  from  copper(I)  chloride  in  acid  solution  with  acetylene  at  pressures 
up  to  2  atmospheres.  However,  from  the  x-ray  powder  photograms,  he  con- 
cluded that  they  were  interstitial  compounds. 

On  the  basis  of  studies  involving  the  distribution  of  copper(I)  chloride 
between  water  and  an  organic  solvent  in  the  presence  of  an  unsaturated 
substance,  Andrews  and  co-workers  have  obtained  formation  constants  for 
a  variety  of  copper(I)  complexes.  Only  1:1  complexes  were  indicated  with 
all  the  unsaturated  alcohols47  and  acids48  investigated.  The  compounds 
formed  by  the  unsaturated  alcohols  were  generally  more  stable  than  those 
with  the  acids,  asone  might  expect,  since  the  carboxyl  group  should  decrease 
the  electron  density  in  the  vicinity  of  the  double  bond.  Substitution  of  H 
by  — CH3  or  — C02H  decreased  stability,  probably  due  also  to  steric  effects. 
Of  the  two  complexes  generally  formed,  Cu  -IJn+  and  CuCl  -Un,  the  cationic 
complexes  were  the  more  stable. 

Silver -olefin  Compounds 

Most  of  the  silver-olefin  complexes  are  too  unstable  to  be  isolated  and 
much  of  the  available  information  has  been  obtained  from  distribution 
studies.  Lucas  and  co-workers  used  this  method  for  the  study  of  silver  com- 
plexes containing  isobutylene49,  a  series  of  mono-  and  diolefins50  and  a  few 

42.  Gilliland,  Seebold,  Fitzhugh,  and  Morgan,  ibid.,  61,  1960  (1939). 

43.  Lur'e,  Marushkin,  Afanas'ev,  and  Pimenov,  Sintet.  Kauchuk,  3,  Xo.  6,  19  (1934). 

44.  Gilliland,  Bliss,  and  Kip,  ./.  Am.  Chem.  Soc,  63,  2088  (1941). 

45.  Ward  and  Makin,  ibid.,  69,  657  (1947). 

46.  Osterlof,  Acta  Chem.  Scand.,  4,  374  (1950). 

47.  Kepner  and  Andrews,  J.  Org.  Chem.,  13,  208  (1948);  ./.  Am.  Chem.  Soc,  71,  1723 

(1949);  Keefer,  Andrews,  and  Kepner,  ibid.,  71,  3906  (1949). 

48.  Andrews  and  Keefer,  ibid.,  70,  3261  (1948) ;  71,  2379  (1949) ;  Keefer,  Andrews,  and 

Kepner,  ibid.,  71,  2381  (1949). 

49.  Eberz,  Wilge,  Yost,  and  Lucas,  ibid.,  59,  45  (1937). 

50.  Winstein  and  Lucas,  ibid.,  60,  836  (1938);  Lucas,  Moore,  and  Pressman,  ibid.,  65, 

227  (1943);  Hepner,  Trueblood,  and  Lucas,  ibid.,  74,  1333  (1952);  Trueblood 
and  Lucas,  ibid.,  74,  1338  (1952). 


496  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

unsaturated  oxygenated  compounds5011.  Compounds  with  a  1:1  mole  ratio 
were  observed  in  all  cases  and  several  unsaturated  molecules  gave  ratios  of 
I  wo  unsaturated  groups  to  one  silver  ion.  Most  of  the  systems  showed  evi- 
dence for  compounds  containing  two  silver  ions  and  one  unsaturated  group 
at  high  silver  ion  concentrations. 

cis-2-Pentene  gave  a  more  stable  complex  than  the  trans  isomer  and  the 
stability  of  the  compounds  of  the  isomeric  butenes  indicated  that  steric 
effects  were  very  important  and  that  substitution  around  the  double  bond 
decreased  the  stability  of  the  complexes.  Similarly,  Nichols51  found  that 
the  silver  complex  of  the  methyl  ester  of  oleic  acid  (cis  form)  was  more 
stable  than  that  of  the  methyl  ester  of  elaidic  acid  (trans  form).  Lucas  et  al. 
observed  no  isomerization  or  polymerization  when  any  of  the  organic 
molecules  combined  with  silver  ion. 

Keefer,  Andrews,  and  Kepner47c  studied  the  silver  complexes  formed  by 
a  series  of  unsaturated  alcohols  and  found  them  to  be  much  less  stable  than 
the  corresponding  copper(I)  complexes.  The  stability  trends  within  the 
series  were  similar. 

Andrews  and  Keefer52  obtained  formation  constants  for  a  series  of  silver 
complexes  with  aromatic  substances  by  the  distribution  method.  They  ob- 
served that  most  simple  aromatic  systems  formed  complexes  containing 
one  silver  ion  and  one  aromatic  molecule  as  well  as  a  less  stable  complex 
containing  two  silver  ions  and  one  aromatic  molecule.  The  relative  sta- 
bilities of  the  complexes  were  associated  primarily  with  the  inductive  effects 
of  ring  substituents  and  steric  factors.  Thus,  the  substitution  of  a  methyl 
group  on  benzene  increases  its  basicity  and  also  the  stability  of  the  silver 
complex.  However,  further  substitution  of  methyl  groups  on  toluene  in- 
creases the  basicity,  but  the  stability  of  the  silver  complexes  decreases  or 
increases  only  slightly  while  the  increase  in  basicity  is  great.  Allowing  for 
the  very  important  steric  effects,  the  stability  of  the  aromatic  complexes 
generally  increases  with  the  basicity  of  the  aromatic  nucleus53. 

Andrews  and  Keefer54  found  that  aromatic  and  olefinic  iodides  gave  far 
more  stable  silver  complexes  than  related  substances,  presumably  because 
the  coordination  occurs  through  the  iodine  atom. 

Mercury -olefin  Compounds 

The  mercury-olefin  compounds  have  been  studied  extensively  and  excel- 
lent reviews  are  available1  • 55.  Lucas,  Hepner,  and  Winstein56  used  the 

51.  Nichols,  ibid. ,74,  1091  (1952). 

52.  Andrews  and  Keefer,  ibid.,  71,  3644  (1949);  72,  3113  (1950);  74,  640  (1952). 

53.  Brown  and  Brady,  ibid.,  71,  3573  (1949);  McCaulay  and  Lien,  ibid.,  73,  2013 

(1951). 

54.  Andrews  and  Keefer,  ibid.,  73,  5733  (1951). 

55.  Chatt,  Chevi.  Rev.,  48,  7  (1951). 

56.  Lucas,  Hepner,  and  Winstein,  J.  Am.  Cheni.  Soc,  61,  3102  (1939). 


COMPOUNDS  OP  METAL  TONS  WITH  OLEFINS  P.»7 

distribution  method  to  study  the  complexes  of  mercury  (I  I)  ion  with  cyclo- 

hexene.  They  obtained  equilibrium  constants  for  two  reactions: 

(     II  II.  •  CM     Ik 

CeHifl  +  Hg++  +  HoO  -*  C6H10HgOH+  +  H+ 

The  equilibrium  constant  for  the  second  reaction  is  slightly  greater  than 
that  for  the  first,  and  other  slower  reactions  were  said  to  proceed  concur- 
rently with  these  two.  The  first  reaction  is  probably  analogous  to  the  com- 
plex formation  by  silver(I)  ion,  but  the  second  reaction  seems  to  be  more 
characteristic  of  mercury (II). 

Some  of  the  mercury-olefin  compounds  probably  exist  as  coordination 
compounds,  at  least  as  intermediates.  However,  the  structure  in  which 
there  is  addition  across  the  double  bond 

\  / 

C— C 

/I       |\ 
HO    HgX 

is  generally  accepted  for  these  compounds57.  The  existence  of  optically- 
active  mercury  compounds  with  olefins  of  the  type  RR'C=CRR'  58  rather 
conclusively  supports  this  structure. 

Miscellaneous  Compounds 

Some  evidence29  • 59  is  available  for  the  existence  of  addition  compounds 
of  zinc  chloride  and  amylene,  but  the  exact  nature  of  the  compounds  is  not 
clear. 

Unstable  aluminum  compounds  with  ethylene,  other  unsaturated  hydro- 
carbons, acids,  aldehydes,  and  alcohols  have  been  isolated60,  but  the  com- 
position of  such  materials  is  difficult  to  determine  because  of  their  insta- 
bility and  hygroscopic  character.  Aluminum  compounds  with  acetylene6015, 
benzene61,  and  substituted  benzenes62  have  also  been  prepared. 

Winstein  and  Lucas50a  found  that  olefins  failed  to  form  complexes  in 
aqueous  solution  with  Cd++,  Co++,  Cr+++  Cu++,  Fe+++,  Ni++  Pb++,  T1+ 
and  Zn"1"*.  However,  Jura  and  his  co-workers63  found  that  the  reaction  of 

Adams,  Roman,  and  Sperry,  ibid.,  44,  1781  (1922). 

58.  Sandborn  and  Marvel,  ibid.,  48,  1409  (1926). 

59.  KondakofT,  J.  Russ.  Phys.-Chem.  Soc,  24,  309  (1892);  25,  345,  456  (1893);  Bull. 

soc.  chim  [3]  7,  576  (1892). 

60.  GanglofT  and  Henderson,  J.  Am.  Chem.  Soc,  39,  1420  (1917);  Henderson  and 

gioff,  ibid.,  38,  1382  (1916). 

61.  Weinland,  "Einfuhrung  in  die  Chemie  der  Komplex-Verbindungen,"  p.  340, 

Stuttgart,  Verlag  von  Ferdinand  Enke,  1924. 

62.  Xorris  and  Ingraham,  ./.  Am.  Chem.  Soc,  62,  1298  (1940). 

63.  Jura,  Grotz,  and  Hildebrand,  Abstracts  of  Papers  presented  at  the  118th  Mtg 

of  A.C.S.,  Chicago,  Sept.  1950. 


498  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

metal  ions  with  aromatic  hydrocarbons  is  quite  general.  On  a  silica  gel 
surface,  mesitylene  was  found  to  react  with  the  ions  of  most  heavy  metals. 
Naphthalene  reacted  to  about  the  same  extent  as  mesitylene,  cyclohexa- 
none  to  a  lesser  extent,  xylene  and  toluene  only  very  weakly,  and  benzene 
showed  no  effect.  This  order  is  essentially  the  same  as  that  found  by 
Andrews  and  Keefer52  for  silver  and  by  Brown  and  Brady53a  for  the  basicity 
of  aromatic  hydrocarbons. 

The  compound  Ni(CN)2-NH3-C6H664  which  has  been  considered  as  a 
coordination  compound,  has  been  shown  to  be  a  clathrate  compound65  in 
which  the  nickel  is  coordinated  only  to  ammonia  and  cyanide  ion  with  the 
benzene  trapped  in  the  lattice  (page  378). 

The  interesting  and  unusual  character  of  bis(cyclopentadienyl)iron(II) 
led  to  the  investigation  of  other  metal  derivatives  of  the  cyclopentadienyl 
radical.  Wilkinson66  prepared  the  analogous  bis(cyclopentadienyl)ruthe- 
nium(II)  which  could  be  oxidized  to  the  cationic  ruthenium(III)  compound 
and  isolated  as  a  salt.  Wilkinson67  was  also  able  to  prepare  the  monovalent 
bis(cyclopentadienyl)cobalt(III)  ion  which  could  be  reduced  to  the  easily 
oxidizable,  neutral  cobalt(II)  compound68,  which  could  also  be  prepared 
from  Co2(CO)8  and  cyclopentadiene  in  the  vapor  phase  at  300°C.  The  cor- 
responding rhodium(III)  and  iridium(III)  compounds  were  also  prepared69. 
The  rhodium  (III)  compound  could  be  reduced  polarographically  although 
at  a  higher  potential  than  that  required  for  the  reduction  of  the  cobalt  (III) 
compound.  The  iridium  compound  showed  no  clear  cut  polarographic  wave. 

The  neutral  bis(cyclopentadienyl)nickel(II)  compound  was  prepared,  but 
it  slowly  decomposed70.  It  could  be  oxidized  to  the  cationic  nickel  (III)  com- 
pound, but  the  latter  decomposed  in  water.  The  neutral  palladium (II)  com- 
pound68 was  obtained  in  solution,  but  it  was  less  stable  than  the  nickel (II) 
compound.  No  copper (II)  derivative  was  obtained. 

Moving  in  the  other  direction  in  the  periodic  table,  Wilkinson  and  co- 
workers68 obtained  evidence  for  a  neutral  cyclopentadienyl  derivative  of 
manganese,  but  the  material  was  oxidized  rapidly  in  air.  Bis  (cyclopentadi- 
enyl) chromium  (II)  was  prepared  from  chromium  hexacarbonyl  and  cyclo- 
pentadiene in  a  hot  tube68b.  The  corresponding  molybdenum  compound  was 
prepared  in  small  yield.  The  compounds  CioHi0TiBr2 ,  CioHi0ZrBr2 , 
CioHioVCl2,  and  Ci0Hi0NbBr3  were  also  obtained68-70.  The  titanium (IV) 

64.  Hoffmann  and  Kiispert,  Z.  anorg.  Chem.,  15,  203  (1897). 

65.  Powell  and  Rayner,  Nature,  163,  567  (1949). 

66.  Wilkinson,  J.  Am.  Chem.  Soc,  74,  6146  (1952). 

67.  Wilkinson,  ibid.,  6148. 

68.  Wilkinson,  Private  communication,  July,  1953;  /.  Am.  Chem.  Soc.,  76,  209  (1954) ; 

Pauson  and  Wilkinson,  J.  Am.  Chem.  Soc.,  76,  2024  (1954). 

69.  Cotton,  Whipple,  and  Wilkinson,  J.  Am.  Chem.  Soc,  75,  3586  (1953). 

70.  Wilkinson,  Pauson,  Birmingham,  and  Cotton,  ibid.,  1011. 


COMPOUNDS  OF  METAL  TONS  WITH  OLEFINS  499 

compound  could  be  reduced  in  solution  to  the  CioHioTi"1  lod  and  there  was 
some  polarographic  evidence  for  the  neutral  compound. 

Wilkinson  and  co-workers  have  shown  thai  the  formation  of  compounds 
with  the  cyclopentadieny]  radical  is  quite  general  for  the  transition  metals, 

but  not  for  the  metals  with  filled  d  orbitals.  The  maximum  stability  ifi 
achieved  for  those  metals  such  as  iron (II)  which  can  complete  the  d  ortibals 
through  bonding  to  two  cyclopentadieny]  radicals.  It  is  possible  to  prepare 
compounds  with  only  one  cyclopentadienyl  ring  attached  to  a  metal  ion  if 
the  metal  can  be  satisfied  with  groups  on  the  side  opposite  to  the  ring. 
Wilkinson6811  prepared  the  compounds  C5H5Mo(CO)bMoC5H5  and 
C*HiW(CO)eWCiHi  in  which  the  metals  are  bridged  by  the  carbonyl 
groups.  Pauson  and  Wilkinson6Sc  prepared  bis(indenyl)iron(II)  and  salts  of 
bis(indenyl)cobalt(III)  from  indenyllithium  and  indenylmagnesium  bro- 
mide, respectively. 

The  well-known  metal  complexes  of  the  azo  and  azomethine  dyes  cer- 
tainly involve  bond  formation  between  some  part  of  the  — N=N —  or 
— CH=X —  system,  but  it  is  not  known  whether  coordination  is  through 
the  double  bond  or  through  the  nitrogen  (Chapter  22). 

Practical  Importance  of  Metal-Olefin  Compounds 

The  exact  role  of  many  metal  salts  in  reactions  involving  olefins  is  not 
known,  but  it  is  significant  that  the  most  important  metal  salts  used  to 
polymerize  or  otherwise  change  olefins  are  those  known  to  form  metal-olefin 
compounds. 

In  the  presence  of  aluminum  chloride,  olefins  are  reported  to  potymerize, 
isomerize,  cyclize,  and  form  paraffins  and  more  highly  unsaturated  com- 
pounds71. Aluminum  chloride  has  been  used  for  converting  gaseous  and 
high-boiling  olefins  into  low-boiling  liquids72,  viscous  oils73,  synthetic  lubri- 
cating oils74,  and  synthetic  resins75.  The  preparation  of  a  compound  of 
aluminum  chloride  with  ethylene,  used  for  condensing  hydrocarbons,  has 
been  patented.  It  is  likely  that  the  Friedel-Crafts  reactions  involve  alumi- 

71.  Egloff,  Wilson,  Hulla,  and  Van  Arsdell,  Chem.  Rev.,  20,  345  (1937);  National 

Research  Council,  "Twelfth  Report  of  the  Committee  on  Catalysis,"  pp. 
182-3,  New  York,  John  Wiley  &  Sons,  Inc.,  1940. 

72.  Ricard  (to  Soc.  Ricard,  Allenet  et  Cie),  U.  S.  Patent  1,745,028  (Jan.  28,  1930); 

cf.  Chem.  Abst.,  24,  1390  (1930). 

73.  N.  V.  de  Bataafsche  Petroleum  Maatschappij,  British  Patent  479,632  (Feb.  9, 

1938);cf.  Chi  m.  Afo.,82,  5197  l938);Sixt  (to  Consortium  fur  elektrochemische 
Industrie  G.  m.  b.  H.),  I'.  S.  Patent  2,183,154  (Dec.  12,  1939);  cf.  Chem.  Abs., 
34,  2302  (1940). 

74.  Perquin  (to  Shell  Development  Co.),  Canadian  Patent  380,056  (Mar.  14,  1939); 

cf.  Chem.  Abs.,  33,  1016    1039). 
Dayton  Synthetic  Chemicals,  Enc.,  German  Patent  061,668  (Oct.  18,  1937);  cf. 
As.,  32,  680  (1938). 


500  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

num  chloride  complexes;  indeed,  some  of  the  supposed  intermediate  alumi- 
num halide  complexes  have  been  isolated76. 

Heavy  metal  carbonyls  have  served  to  convert  high-boiling  hydrocarbons 
into  lower  boiling  forms  by  high-pressure  hydrogenation77. 

The  polymerization  of  butadiene  is  effected  by  boron  fluoride78,  alumi- 
num chloride79,  heavy  metal  carbonyls80,  and  the  iron  phthalocyanine  sul- 
fonic acid  complex81.  Vinylacetylene  is  prepared  by  the  dimerization  of 
acetylene  by  copper(I)  chloride  solutions82. 

Many  complex-forming  metal  salts  have  been  found  to  be  effective  in 
the  hydration  of  olefins  in  acid  solutions83. 

Gaseous  olefins  may  be  extracted  from  mixtures  with  saturated  hydro- 
carbons by  aqueous  solutions  of  copper(I),  silver,  mercury(II),  and  plati- 
num (II)  salts84.  The  olefins  can  be  subsequently  recovered  by  heating  the 
solutions  or  by  reducing  the  pressure.  Diolefins  can  be  separated  from 
monoolefins  as  a  result  of  the  formation  of  insoluble  complexes  by  the 
diolefins  and  certain  heavy-metal  salts85. 
5 

76.  Norris  and  Wood,  J.  Am.  Chem.  Soc,  62,  1428  (1940). 

77.  I.   G.   Farbenindustrie   A.-G.    (Zorn   and  Vogel,   inventors),   German   Patent 

579,565  (June  29,  1933);  cf.  Chem.  Abs.,  28,  1045  (1934). 

78.  Harmon  (to  E.  I.  du  Pont  de  Nemours  and  Co.),  U.  S.  Patent  2,151,382  (Mar.  21, 

1939);  cf.  Chem.  Abs.,  33,  5096  (1939). 

79.  Zelinshil,  Densienko,  Eventova,  and  Khromov,  Sintet  Kauchuk,  1933,  No.  4,  11. 

80.  Ambros,  Reindel,  Eisele,  and  Stoehrel  (to  I.  G.  Farbenindustrie  A.-G.),  U.  S. 

Patent  1,891,203  (Dec.  13,  1932);  cf.  Chem.  Abs.,  27,  1893  (1933);  I.  G.  Farben- 
industrie A.-G.,  British  Patent  340,004  (Aug.  12,  1929);  cf.  Chem.  Abs.,  25, 
Si  2878  (1931). 

81.  I.   G.   Farbenindustrie  A.-G.    (Gumlich   and  Dennstedt,   inventors),   German 

Patent  679,587  (Aug.  9,  1939);  cf.  Chem.  Abs.,  33,  9328  (1939). 

82.  Burk,  Thompson,  Weith,  and  Williams,  "Polymerization  and  its  Applications 

in  the  Fields  of  Rubber,  Synthetic  Resins  and  Petroleum,"  p.  76,  New  York, 
Reinhold  Publishing  Corp.,  1937;  Klebanskii,  Tzyurikh,  and  Dolgopol'shil, 
Bull.  acad.  sci.  U.R.S.S.,  1935,  No.  2,  189;  J.  Research  Assoc.  Brit.  Rubber 
Mfrs.,  4,  505  (1935). 

83.  Dreyfus,  British  Patent  397,187  (Aug.  21,  1933);  cf.  Chem.  Abs.,  28,  777  (1934); 

Standard  Alcohol  Co.,  British  Patent  493,884  (Oct.  17,  1938);  cf.  Chem.  Abs., 
33,  2533  (1939). 

84.  Ellis,  "The  Chemistry  of  Petroleum  Derivatives,"  p.  142,  New  York,  The  Chemi- 

cal Catalog  Co.,  Inc.,  (Reinhold  Publishing  Corp.),  1934;  N.  V.  de  Bataafsche 
Petroleum  Maatschappi j ,  German  Patent  622,965  (Dec.  10,  1935);  cf.  Chem. 
Abs.,  30,  3442  (1936);  Gilliland  (to  Standard  Oil  Development  Co.),  U.  S. 
Patent  2,209,452  (July  30,  1940)  and  2,289,773  (July  14,  1942);  cf.  Chem.  Abs., 
35,  134  (1941)  and  37,  386  (1943)  resp. ;  Gilliland  and  Seebold,  Ind.  Eng.  Chem., 
33,  1143  (1941);  Imperial  Chemical  Industries,  Ltd.,  French  Patent  662,099 
(Mar.  12,  1928);  cf.  Chem.  Abs.,  24,  376  (1930);  Stern,  Reichsant  Wirtschafts- 
aubau,  Pruf-Nr.,  43,  (PB52003)  15-56  (1940);  cf.  Chem.  Abs.,  41,  6490  (1947). 

85.  Hebbard  and  Lloyd  (to  Dow  Chemical  Co.),  U.  S.  Patents  2,188,899  and  2,189,173 

Feb.  6,  1940);  cf.  Chem.  Abs.,  34,  3760  (1940). 


(/ 


I  OMPOUNDS  OF  METAL  TONS  n  I  ill  OLEFINS  50] 

The  Structure  of  Metal-Olefin  Compounds 
Although  many  structures  have  been  proposed  for  the  metal-olefin  com- 
pounds, satisfactory  structures  have  been  proposed  only  recently.  Various 
suggested  structures  have  been  reviewed  by  Keller1  and  more  recently  by 
Chatt*.  Although  mo>t  of  the  proposed  structures  and  some  structural  data 
can  be  elminated  on  the  basis  of  the  evidence,  much  remains  to  be  learned 
about  the  structure  of  metal-olefin  compounds. 

The  compound  [PtClo-OiHs]*  is  known  to  be  dimeric  on  the  basis  of  an 
accurate  molecular  weight  determination  in  benzene15.  An  approximate 
molecular  weight  determination  for  ethylene-platinum(II)  chloride  indi- 
cated it  to  be  a  dimer1*.  Styrene-palladium(II)  chloride  is  probably  dimeric, 
although  an  exact  molecular  weight  could  not  be  obtained  by  the  freezing- 
point  method18. 

Pfieffers6  proposed  formula  (I)  for  the  ethylene-platinum(II)  chloride 
complex,  although  he  did  not  indicate  the  nature  of  the  Pt-Un  bond. 
Kharasch  and  Ashford15  objected  to  (I)  because  of  the  formation  of  two 
coordinate  bonds  by  the  same  chloride  ion.  They  proposed  structure  (II), 

H2   H2 
Un  CI  CI  CI  C— C  CI 

\   /    \    /  \   /  \    / 

PI  Pt  Pt  Pt 

/    \    /    \  /\/\ 

CI  CI  Un  CI  C— C  CI 

XX2      H2 

(I)  (ID 

in  which  the  double  bond  is  broken  to  permit  the  olefin  to  act  as  the  bridge. 
Halide  ions  act  as  bridges  in  many  stable  polymeric  complexes87  so  the 
objection  of  Kharasch  and  Ashford  is  without  foundation.  The  represen- 
tation of  the  platinum-olefin  complexes  as  metal-alkyls  seems  objectionable 
on  the  basis  of  the  ready  displacement  of  one  olefin  by  another14  or  by  other 
coordinating  groups  such  as  pyridine  and  cyanide  ion13. 

Although  most  complexes  of  the  type  [Pt  a  C2H4C12]  (a  =  ammonia  or 
pyridine)  arc  too  insoluble  for  molecular  weight  determinations,  ( Jhatt2  was 
able  to  establish  that  the  corresponding  p-toluidine  complex  is  monomeric 
Oppegard28  found  the  complexes  [PtCJI^  quinoline  C1J  and  [Pt  styrene 
quinoline  Cl«]  to  be  monomeric  in  benzene.  Thus,  an  olefin  bridge  cannot 
be  used  to  explain  the  structure  of  these  complexes  and  there  is  do  reason 
to  suppose  thai  such  a  bridge  exists  in  other  platinum-olefin  compounds. 

86.  Pfeiffer,  "Organische  Molekulverbindungen,"  i>.   161,  Stuttgart,  Verlag  von 

linand  Enke,  ; 

87.  Gibson  and  Simonsen,  /.  Chem.  Soc. ,  1930,  2531;  Mann  and  Purdie,  Tbid.t  1936, 

^7:;:  Palmer  and  Elliott,/.  An  Soc  .  60,  1852  (1938  ;  Wells,  Z,  KrUt., 

100,  180  (1938). 


502 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


From  an  x-ray  structure  analysis,  Bokii  and  co-workers88  reported  the 
compound  cfs-[PtC2H4NH3Cl2]  to  be  dimeric  with  a  platinum-platinum 
bond  length  of  1.4  A.;  however,  the  results  mentioned  above  indicate  that 
a  dimeric  structure  is  unlikely  and  there  seems  to  be  no  other  evidence  for 
a  platinum-platinum  bond.  Apparently  the  interpretation  of  the  x-ray  data 
was  erroneous. 

Bennett  and  Willis89  proposed  structure  (III),  in  which  one  pair  of  elec- 
trons from  the  double  bond  migrates  to  one  carbon  to  be  shared  with  the 
platinum  atom.  This  leaves  the  other  carbon  as  a  carbonium  ion,  which 
should  be  very  reactive.  Similarly,  Stiegman90  proposed  structure  (IV)  in 


"     H  H 

H:C:C:PtCl3 
+  •• 

H 
(III) 


"      H  H 
H:C:C:PtCl3 
.      "  H 

(IV) 


which  the  double  bond  is  broken,  but  the  carbonium  ion  shares  a  pair  of 
electrons  furnished  by  the  platinum.  Here  the  remaining  carbon  would  be  a 
carbanion  which  should  also  be  very  reactive.  In  addition,  if  the  platinum, 
and  not  the  ethylene,  is  the  donor,  one  would  not  expect  the  ethylene  to 
behave  as  a  typical  ligand  and  be  readily  replaced  by  ligands  such  as 
chloride  ion  and  ammonia.  These  structures  seem  unlikely. 

Drew,  Pinkard,  Wardlaw,  and  Cox91  proposed  structure  (V)  (written  as 
(VI)  by  Chatt)  for  the  ion  [PtC2H4Cl3]-.  It  is  objectionable  on  the  same 


'C1CH2CH< 


CI 
(V) 


Pt— CI 


[H,C— MCliT 
I        I 
H2C— CI      J 


(VI) 


grounds  as  a  platinum-alkyl  structure.  Chatt2  mentioned  that  an  attempt 
to  prepare  2-benzoylethyl  chloride  by  heating  ethylene-platinum(II)  chlo- 
ride with  an  excess  of  benzoyl  chloride  was  unsuccessful.  He  believed  that 
this  reaction  should  proceed  if  the  olefin  complexes  had  structure  (V). 

Chatt92  emphasized  the  similarity  between  the  platinum  complexes  with 
olefins  and  those  with  carbon  monoxide.  Both  groups,  unlike  most  neutral 

88.  Bokii,  Usikov,  and  Trusevich,  Bull.  acad.  sci.,  U.R.S.S.,  Classe  sci.  Chan.,  1942, 

413;  Bokii  and  Baishteil,  Doklady  Akad.  Nauk  S.S.S.R.,  38,  323  (1943);  Bokii 
and  Vainshtein,  Compt.  rend.  acad.  sci.  U.R.S.S.,  38,  307  (1943). 

89.  Bennett  and  Willis,  J.  Chem.  Soc,  1929,  259. 

90.  Stiegman,  thesis,  University  of  Illinois,  1937. 

91.  Drew,  Pinkard,  Wardlaw,  and  Cox,  /.  Chem.  Soc,  1932,  897. 

92.  Chatt,  Nature,  165,  637  (1950). 


COMPOUNDS  OF  METAL  IONS  117  77/  OLEFINS  503 

ligands,  show  a  very  marked  trans  effect,  which  Chatl  staled  Is  probably 
associated  with  double  bond  character  between  the  metal  and  donor  group 
as  suggested  by  Pauling91  for  the  metal  carbonyls. 

Hel'man*  found  that  Zeise's  salt  resists  oxidation  by  permanganate, 
giving  an  initial  potential  in  an  electrometric  titration  of  650  to  700  m.v., 
comparable  to  that  observed  for  typical  platinum(IV)  complexes.  Plati- 
num(II')  salts  are  readily  oxidized  by  permanganate  at  a  lower  potential. 
She  considered  this  to  be  evidence  that  the  platinum  is  present  as  plati- 
num(IV)  as  a  result  of  the  sharing  of  a  pair  of  d  electrons  from  the  platinum 
with  the  ethylene  which  in  turn  shares  a  pair  of  its  electrons  with  the 
platinum  to  form  a  four  electron  bond26.  Hel'man  did  not  specify  the  nature 
of  the  four  electron  bond,  show  how  the  ethylene  accommodates  the  two 
electrons  from  the  platinum,  or  what  happens  to  the  carbon-carbon  double 
bond.  She  believed  that  only  one  ethylene  molecule  could  be  coordinated 
to  a  platinum  atom,  since  the  platinum  would  be  required  to  furnish  a 
pair  of  electrons  for  each  ethylene  coordinated.  Chatt25  discredited 
Hel'man's  structure  by  preparing  the  compound  [Pt^H^Cy.  However, 
this  would  require  only  a  slight  modification  by  Hel'man,  since  the  con- 
sideration of  the  oxidation  state  of  the  platinum  is  purely  formal. 

The  bulk  of  the  evidence  is  in  favor  of  the  view^  that  the  platinum-olefin 
compounds  are  derivatives  of  platinum (II).  This  is  indicated  by  the  fact 
that  the  olefins  readily  replace  other  ligands  in  platinum(II)  compounds  or 
are  readily  replaced  by  other  ligands  to  give  platinum(II)  compounds. 
However,  such  an  argument  tells  only  what  is  put  into  and  wThat  is  obtained 
from  olefin  complexes  and  ignores  the  fact  that  the  assignment  of  the  oxi- 
dation state  of  the  platinum  is  purely  formal  if  the  bond  order  differs  in  any 
case. 

Chatt95  proposed  the  structure 

'CH3CH  CI' 

\    / 
Pt 

/    \ 

CI  CI. 

representing  the  ethylene  compound  as  a  substituted  ethylidene  complex 
formed  as  a  result  of  migration  of  a  hydrogen  atom  on  coordination.  How- 
ever,  Chatt91  no  longer  believes  this  structure  to  be  correct.  Objection.-  toil 

wrere  cited  by  Douglas97  and  by  Chatt96. 

93.  Pauling,  "Nature  of  the  Chemical  Bond,"  2nd    ed.,  pp.  251  el  Beq.,  [thaca, 
Cornell  1  Iniversii  y  Press,  1940. 

'it.  Hel'man  and  Ryabchikov,  Compt.  rend,  acad  set.  &.R.S.S.,  33,  162  (1941). 
''.V  Chatt,  Research,  4,  L80  (1961). 

96.  Chatt,  •/.  Chem.  8oe.}  1953,  2939. 

97.  Douglas,  J.  Am.  Chem    Soc.,  75,  4836    L953 


504  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Oppegard88  found  that  as-2-pentene  gave  a  crystalline  complex  with 
platinum,  while  /rans-2-pentene  gave  a  red  oil.  The  infrared  spectra  for  the 
two  compounds  were  also  found  to  differ.  This  is  in  agreement  with  the 
observations  of  Winstein  and  Lucas50  that  the  silver-olefin  complexes  give 
no  rearrangements  and  that  cis  and  trans  isomers  possess  different  coordi- 
nating properties  with  respect  to  silver.  On  the  basis  of  the  ethylidene 
>i  ructure,  one  wrould  predict  the  isomerization  of  cis-trans  isomers  during 
coordination  to  and  subsequent  liberation  from  platinum(II)  salts. 

Oppegard  also  found  that  the  ultraviolet  spectra  of  £rcms-stilbene  and 
the  complex,  [Pt  stilbene  Cl2]2 ,  were  almost  identical,  indicating  that  the 

\     / 

resonance  of  stilbene,  involving       C=C     ,  w^as  not  greatly  disturbed. 

/     \ 

The  results  wTere  not  conclusive  because  the  spectra  for  the  styrene  and 
2-pentene  complexes  could  not  be  interpreted  so  simply.  The  infrared  data 
indicated  that  the  carbon-carbon  distance  in  the  olefinic  complexes  was 
lengthened  considerably,  although  the  different  spectra  obtained  with  the 
isomeric  2-pentenes  indicated  that  free  rotation  was  not  permitted. 

Chatt96  has  found  from  infrared  data  that  the  olefin  retains  its  double 
bond  in  platinum  complexes  and  that  the  double  bond  is  symmetrically 
coordinated  to  the  platinum.  The  greater  lowering  of  the  double  bond 
stretching  band  for  the  platinum  complexes  as  compared  with  those  of 
silver  was  attributed  to  the  stronger  bonding  in  the  platinum  complexes. 
Wunderlich  and  Mellor98  obtained  x-ray  structural  data  for  Zeise's  salt  and 
determined  that  the  C-C  axis  is  approximately  perpendicular  to  the  plane 
of  the  PtCl3  group  and  probably  symmetrically  arranged  with  respect  to 
the  platinum  atom.  The  distance  between  platinum  and  the  chloride  trans 
to  the  ethylene  molecule  is  abnormally  great. 

Dempsey  and  Baenziger98a  determined  the  crystal  structure  of 
(PdCl2C2H4)2  by  x-ray  diffraction  methods.  The  dimer  has  the  trans 
bridged  structure  similar  to  structure  I  (p.  501)  for  the  corresponding 
platinum  compound.  The  axis  of  the  ethylene  molecule  is  perpendicular 
to  the  plane  of  the  dimer  and  the  center  of  the  ethylene  bond  lies  in  the 
plane  of  the  dimer.  Holden  and  Baenziger98a  obtained  the  structure  of  the 
corresponding  styrene  complex  since  the  carbons  of  the  ethylene  molecule 
could  not  be  resolved.  The  general  features  of  the  structure  are  the  same 
as  those  of  the  ethylene  complex  except  that  the  palladium  is  slightly  off 
center  with  respect  to  the  carbon-carbon  double  bond  in  the  styrene  com- 
plex. The  Pd-Cl  bonds  opposite  the  Pd-olefin  bonds  are  somewhat  longer 
than  the  other  Pd-Cl  bonds. 

98.  Wunderlich  and  Mellor,  Acta  Cnjst.,  7,  130  (1954);  8,  57  (1955). 

D<  mpsey  and  Baenziger,  J.  Am.  Chem.  Soc,  77, 4984  (1955) ;  Holden  and  Baen- 
aiger,  ibid.,  77,  1987  (1965). 


COMPOUNDS  OF  METAL  IONS  WITH  OLEFINS  505 

Winstein  and  Lucas"  proposed  a  structure  for  the  Bttver-olefhi  complexes 
based  on  resonance  involving  three  forms. 


/ 

c— c 

V\ 

\          / 

c=c 

/     \ 

Ag+ 

\ 

c— c 

/v 

Ag 

(VII) 

(VIII) 

(IX) 

The  resonance  hybrid  would  not  have  the  properties  of  a  molecule  contain- 
ing a  carbonium  ion,  nor  would  the  double  bond  need  to  be  activated  suffi- 
ciently to  lead  to  polymerization  or  rearrangement  of  cis-lrans  isomers. 
They  stated  that  the  C — C — Ag  bond  angle  would  be  greater  than  the  60° 
angle  for  cyclopropane  and  that  the  resonance  energy  could  compensate 
for  the  strain. 

Pitzer"  indicated  that  the  protonated  double  bond  type  of  structure 
which  he  proposed  for  the  boron  hydrides  can  be  applied  to  the  silver-olefin 
complexes.  He  pointed  out  that  silver  has  an  s  orbital  which  it  can  use  for 
bond  formation  with  the  olefin. 

Dewar100  and  Walsh101  stated  that  bonding  electrons  can,  under  certain 
conditions,  be  utilized  in  the  formation  of  a  coordinate  covalent  bond. 
Walsh  pointed  out  that  the  x  electrons  of  ethylene  lie  in  an  orbital  of  ion- 
ization potential  10.45  volts,  almost  equal  to  that  (10.8  volts)  of  the  am- 
monia lone  pair.  Werner102  and  Bateman103  related  these  views  to  the  olefin 
complexes  and  Bateman  mentioned  that  they  were  essentially  those  ex- 
pressed by  Winstein  and  Lucas  and  restated  more  precisely  by  Pitzer. 

Dewar104  described  the  structure  of  the  silver-olefin  complexes  in  terms 
of  molecular  orbitals.  The  structure  involved  the  combination  of  the  vacant 
s  orbital  of  silver  with  the  7r-orbital  of  the  olefin  and  the  combination  of  a 
filled  4<7  orbital  of  silver  with  the  p  orbital  of  the  olefin. 

Chatt*  discarded  the  Pitzer  structure  for  the  platinum  complexes  since 
platinum  does  not  have  a  vacant  s  orbital  (see  footnote  p.  506).  However, 
in  view  of  more  recent  data,  Chatt96  considers  a  similar  structure  to  be 
correct. 

Chatt2  found  no  evidence  for  association  between  ethylene  and  tri- 
methylborine  and  interpreted  this  to  mean  that  "the  donation  of  electrons 
in  any  manner  from  the  ethylene  molecule  to  the  metal  cannot,  of  itself, 
be  responsible  for  the  coordination  of  ethylene."  He  felt  that  the  distin- 

99.  Pitzer,  ./.  Am.  Chem.  Soc,  67,  1127  (1045). 

100.  Dewar,  ./  toe.,  1946,  408. 

101.  Walsh,  ibid.,  1947,  89. 

102.  Werner,  Nai  i  ■  .  160,  644    1947). 
i  kit  email,  ibid.,  56. 

104.  Dewar.  Bull.  soc.  chiui.,  18,  C70  (1 


506 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


guishing  feature  of  platinum  as  compared  to  boron  is  the  ability  to  donate 
d  elect  ions  to  form  a  double  bond.  However,  he  did  allow  that  Pitzer's 
structure  mighl  apply  to  the  silver-olefin  complexes.  He  considered  the 
structure  of  the  silver  complexes  to  differ  from  that  of  the  platinum-olefin 
complexes,  since  it  is  known  that  olefins  existing  as  cis-trans  isomers  do  not 
rearrange  in  the  silver  complexes  and  because  of  the  presence  of  a  vacant 
s  orbital  in  the  case  of  silver.  Since  new  evidence  indicates  that  cis-trans 
isomers  should  not  rearrange  in  the  platinum  complexes,  this  distinction 
between  the  silver  and  the  platinum  complexes  cannot  be  made. 

Professor  Pitzer105  has  been  kind  enough  to  make  a  statement*  which  re- 
moves the  misconception  that  he  has  excluded  the  possibility  that  a  metal 
ion  without  a  vacant  s  orbital  could  form  a  complex  with  the  protonated 
double  bond  type  of  structure. 

Douglas97  has  proposed  a  modification  of  the  Winstein-Lucas  structure, 
(VII),  (VIII),  and  (IX),  by  adding  two  resonance  forms,  (X)  and  (XI), 
involving  the  sharing  of  a  pair  of  d  electrons  from  the  platinum. 


XC 


V 

\ 

£l3 


AND 


>s 


PtCI3 


21  21 

This  is  similar  to  the  molecular  orbital  structure  proposed  by  Dewar  for  the 
silver-olefin  complexes.  Chatt96  has  made  the  similarity  even  greater  by 
extending  Dewar's  structure  to  include  the  platinum-olefin  compounds.  He 
considers  the  sharing  of  electrons  from  the  olefin  to  occur  through  the 
overlap  of  a  5c?6s6p2  hybrid  orbital  of  the  platinum  atom  with  the  7r-orbital 
of  the  olefin  and  the  sharing  of  electrons  from  the  platinum  to  occur  by  the 
overlap  of  a  hybridized  5d6p  orbital  with  the  antibonding  orbitals  of  the 
olefin.  This  is  essentially  the  same  as  the  resonance  structure  proposed, 
but  is  more  detailed  in  terms  of  the  orbitals  involved.  The  structures  of 
the  palladium  and  platinum  complexes  determined  by  x-ray  methods98-  98a 
seem  to  be  consistent  with  the  orbital  assignment  given  by  Chatt. 

105.  Pitzer,  private  communication,  Sept.  17,  1952. 

*  "Because  of  their  non-directional  property,  s  orbitals  can  be  combined  into  the 
protonated  double  bond  type  of  orbitals  better  than  p  or  d  orbitals.  This  is  not  to 
imply  that  it  is  impossible  to  use  p  or  d  or  hybrid  orbitals  for  this  purpose — indeed  I 
now  feel  that  there  is  adequate  evidence  in  favor  of  bridge  bonds  of  this  type. 

''1  believe  we  should  use  some  caution  in  assuming  larger  and  more  complex  groups 
to  be  bounded  to  a  pair  of  electrons  in  a  double  bond.  However,  I  do  not  pretend  to 
prescribe  any  particular  limit  and  I  feel  it  probable  that  a  limitation  to  single  atoms 
with  8  orbitals  available  would  be  incorrect." 


COMPOUNDS  OF  METAL  TONS  WITH  OLEFINS 


507 


Andrews  and  Reefer"  suggested  that  a  likely  structure  for  the  silver- 
benzene  complexes  is  one  with  the  silver  ion  above  the  ring  on  the  six-fold 
axis  of  Bymmetry;  in  the  disilver  complexes,  there  would  be  one  silver  ion 
on  eaeh  side  of  the  ring.  X-ray  analysis  of  the  solid  silver  perchlorate- 
benzene  complex  shows  thai  each  silver  is  bonded  equally  to  two  carbon 

atoms  of  each  of  two  rings  lying  above  and  below  the  rings,  suggesting  t 
bonding106.  However,  the  structure  in  solution  might  differ  from  this.  No 
conclusions  could  be  reached  concerning  the  bonding  between  silver  and 
toluene107. 

Interesting  developments  in  the  structure  determination  of  bis(cyclo- 
pentadienvl)iron(II)  have  been  presented.  The  compound  almost  certainly 
contains  iron (II)  since  it  is  diamagnetic  and  is  readily  oxidized  to  a  blue 
cation  Fe(C5H5)2+  which  has  a  magnetic  moment  corresponding  to  one  un- 
paired electron37.  The  structure  was  originally  assumed  to  be  one  repre- 
sented by  two  resonance  forms  (XII)36,  but  the  diamagnetic  character 
suggests  structure  (XIII), 


=>—-<= 


AND 
(+1    M 


=\       +)    M    v/= 
)(-)   Fe   (-)( 


zn 


"YTTT 


as  does  the  fact  that  the  infrared  absorption  spectrum  contains,  in  the  3  to 
4  ii  region,  a  single  sharp  band  which  indicates  the  presence  of  only  one 
type  of  C — H  bond37.  This  does  not  exclude  the  prismatic  structure  with 
the  rings  lined  up  above  one  another.  The  dipole  moment  is  effectively  zero. 

A  structure  in  which  the  iron  atom  is  symmetrically  placed  between 
two  cyclopentadienyl  rings  (XIII)  was  confirmed  by  x-ray  analysis108.  The 
x-ray  data  support  the  antiprismatic  structure  (XIII)  in  the  solid  state. 
However,  the  isomers  of  derivatives  of  ferrocene  are  those  to  be  expected 
if  free  rotation  of  the  rings  occurs  in  solution109. 

The  structure  of  bis(cyclopentadienyl)  compounds  has  been  presented  in 

106.  Rundle  and  Goring,  /.  Am.  Chem.  Soc,  72,  5337  (1950). 

107.  Murrav  and  Cleveland,  ibid.,  65,  2110  (1943). 

108.  Kiland  and  Pepinsky,  J.  Am.  Chem.  Soc,  74,  4971  (1952);  Fisher  and  Pfab,  Z. 

Xaturforschung,  7B,  377  (1052);  Dunitz  and  Orgel,  Nature,  171,  121  (1953). 

109.  Woodward  and  Rosenblum,  private  communication,  August,  1953. 


508  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

terms  of  molecular  orbitals  by  Moffitt110.  The  bonding  is  described  as  a 
delocalized  two  electron  covalent  bond  between  the  metal  ion  and  each 
cyclopentadienyl  ring.  Such  bonding  is  consistent  with  free  rotation  of  the 
rings  and  with  the  magnetic  data.  It  also  explains  the  absence  of  a  copper 
compound  and  the  fact  that  Ti(C5H5)2+  can  exist  although  there  are  only 
two  metal  electrons  which  can  bond  with  the  unpaired  tt  electrons  of  each 
ring.  Since  only  one  ir  electron  of  each  cyclopentadienyl  ring  is  used  in 
bonding,  the  rings  have  aromatic  character. 

110.  Moffitt,  J.  Am.  Chem.  Soc,  76,  3386  (1954). 


10.   Metal  Carbonyls  and  Nitrosyls 

J.  A.  Mattern 

University  of  Buffalo,  Buffalo,  New  York 

and 

Stanley  J.  Gill 

University  of  Illinois,  Urbana,  Illinois 

Early  History 

Upon  observing  that  nickel  valves  were  corroded  by  hot  gases  containing 
carbon  monoxide,  Mond  and  his  co-workers1  studied  the  action  of  carbon 
monoxide  upon  nickel  under  various  conditions.  They  found  that  a  stream 
of  carbon  monoxide,  after  passing  over  finely  divided  nickel,  burned  with 
a  luminous  flame  which  deposited  metallic  spots  upon  a  cold  surface.  From 
such  a  stream  of  gas  they  isolated  a  colorless  liquid  with  a  musty  odor  and 
remarkably  high  refractive  index  and  coefficient  of  expansion.  This  com- 
pound has  the  formula  Ni(CO)4 .  In  1834  von  Liebig2  prepared  a  compound 
having  the  empirical  formula  KCO  by  passing  carbon  monoxide  over 
molten  potassium;  this  however,  is  the  potassium  salt  of  hexahydroxyben- 
zene3  and  is  quite  different  from  the  covalent  carbonyls  discussed  in  this 
chapter. 

A  volatile  iron  carbonyl  was  discovered  in  18914  and  was  shown  to  have 
the  formula  Fe(CO)55.  Dewar  and  Jones6  showed  the  photodecomposition 
product  of  the  pentacarbonyl  to  be  the  enneacarbonyl,  Fe2(CO)9 ,  and 
demonstrated  the  existence  of  a  third  carbonyl,  Fe3(CO)i2 . 

The  known  mononuclear  and  polynuclear  metal  carbonyls  and  their 
hydrides  are  listed  in  Table  16.1. 

1.  Mond,  Langer,  and  Quincke,  /.  Chem.  Soc,  57,  749  (1890);  Mond,  /.  Soc.  Chem. 

Ind.,  14,945  (1895). 

2.  Liebig,  Fogg.  Ann.,  30,  90  (1834). 

3.  Xietski  and  Benckiser,  Ber.,  18,  499,  1833  (1885). 

4.  Berthelot,  Compt.  rend.,  112,  1343   (1891);  Mond  and  Quincke,  Ber.,  24,  2248 

(1891); ./.  Chem.  Soc.,  59,  604  (1891);  Chem.  News,  63,  301  (1891). 
:>.  Mond  and  Langer,  J.  Chem.  Soc.,  59,  1090  (1891). 
6.  Dewar  and  Jones,  Proc.  Roy.  Soc,  (London),  A76,  558  (1905);  A79,  66  (1906). 

509 


Table  16.1.  Metal  Carbonyls  and  Carbonyl  Hydrides7 


Met- 
als 


Cr 


Mn 
Fe 


Co 


Ni 


Mo 

Tc 
Ru 


Rh 


Pd 
W 


Re 


Os 


II 


Pt 


Monomeric  Carbonyls  with  Rare  Gas  Coring. 
Volatile,  Soluble  in  Organic  Liquid 


Carbonyls 


Cr(CO)6  color- 
less, rhomb., 
sublimes 

Fe(CO)5       yel., 
volatile,  M.P. 
-20°C.  B.P. 
103  °C 


Ni(CO)4     color- 
less, volatile 
M.P.  -25°C. 
B.P.  43°C 

Mo  (CO)  6    color- 
less,   sublimes 

Ru(CO)5    color- 
less,       M.P. 
-22°C 


W(CO)6colorless 
rhomb.,  sub- 
limes 


Os(CO)5     color- 
less, volatile 
M.P.  ca.  -18°C 


Carbonyl  Hydrides 


Fe(CO)4H2  col- 
orless, volatile 
M.P.  -70°C 

Co(CO)4H   light 
yel.,  volatile 
M.P.  -26°C 


rhomb., 


volatile 


Polynuclear  Carbonyls,  Less  Volatile  or 
Non-volatile,  Less  or  Not  Soluble 


Rh(CO)4H  dark 
yel.,  volatile 
M.P.  -12°C 


Re(CO)5H* 


Os(CO)4H2(?) 


Ir(CO)4H+ 


Dinuclear  Carbonyls 


Mn2(CO)10 

Fe2 (CO)  9  gold-yel- 
low, pseudo-hex- 
agonal, dec. 
100°C 

Co2(CO)8    orange, 
cryst.M.P.51°C 


Ru2(CO)9    orange 
monoclinic  pris- 
matic, sublimes 
Rh2(CO)8  yel. -red, 
dec.  76°C 


Re2(CO)i0  color- 
less, monocl. 
prismatic,  sub- 
limes M.P.  177°C 

Os2(CO)9  light  yel- 
low, pseudo-hex- 
agonal       M.P. 
224°C 

Ir2(CO)8  green- 
yel.  cryst.,  sub- 
limes 


Higher  Carbonyl 
Polymers 


Fe3(CO)i2  green 
monocl.  pris- 
matic, dec. 
140°C 

Co4(CO)12    black, 
cryst.  dec.  60°C 


Ru3(CO)i2   green, 
insoluble 

Rhft(CO)3»t    dark 
red  crystl.,  sub- 
limes 150° 
Rh4(CO)n 
black,  dec. 
200°C 


Ir„(CO)3»t    ca- 
nary yel.   tri- 
gonal, dec. 
210°C 


*  Formula  qualitatively  established. 

t  Degree  of  polymerization  greater  than  4  not  definitely  established. 

7.  Hieber,  FIAT  Rev.  German  Sci.,  1939-46,  Inorg.  Chem.,  Pt.  II,  108  (1948). 

510 


METAL  CARBONYLS  AND  NITROSYLS  511 

Methods  of  Preparation 
Direct  Combination;  xM  +  yCO  — >  Mx(CO)y 

Passage  of  carbon  monoxide  over  the  finely  divided  metal  at  suitable 
temperatures  and  pressures  has  been  used  for  the  preparation  of  Xi(CO)4 , 
Fe(CO)5,  [Co(CO)J*8,  Mo(CO)e8'u,  W(CO)6u,  Hu(CO)58'9  and 
[Rh(CO)4]210.  Pressure  greater  than  atmospheric  is  required  in  the  prepa- 
ration of  all  except  nickel  carbonyl,  and  the  yields  are  small  except  for  the 
carbonyls  of  iron  and  nickel.  In  general,  the  metal  must  be  in  a  finely- 
divided,  active  state.  In  the  case  of  nickel,  the  metal  has  been  prepared  by 
reduction  of  the  oxide  by  hydrogen  at  400°C  or  of  the  oxalate  at  300°C. 
The  lower  the  temperature  of  reduction,  the  more  active  is  the  resulting 
metal.  The  presence  of  copper  or  iron  in  the  nickel  increases  the  rate  of 
formation  of  nickel  carbonyl.  A  very  active  metal  has  been  prepared  by 
electrolysis  of  a  solution  of  nickel  sulfate  with  a  mercury  cathode  and  sub- 
sequent low  temperature  distillation  of  the  mercury12. 

Nickel  carbonyl  may  be  formed  at  atmospheric  pressure  and  a  tempera- 
ture of  30  to  100°  13.  Processes  have  been  developed  for  the  preparation  by 
passing  carbon  monoxide  through  suspensions  of  nickel  in  inert  liquids, 
such  as  paraffin  oils. 

The  preparation  of  iron  pentacarbonyl  employs  a  pressure  of  20  to  200 
atmospheres  and  a  temperature  of  200°C.  The  presence  of  oxygen  or  an 
oxide  coating  on  the  iron  hinders  the  reaction,  but  the  presence  of  finely 
divided  alumina,  bismuth,  nickel,  or  copper  accelerates  it,  as  do  ammonia, 
hydrogen,  and  small  quantities  of  sulfur  compounds. 

Preparation  from  Grignard  Reagents 

The  hexacarbonyls  of  chromium,  molybdenum,  and  tungsten,  as  well  as 
the  carbonyl  of  nickel,  have  been  prepared  by  the  reaction  between  carbon 
monoxide  and  Grignard  reagents  in  the  presence  of  the  anhydrous  chloride 
of  the  metal14.  Hieber  and  Romberg14b,  studying  the  mechanism  of  the 

-    Mond,  Hirtz,  and  Cowap,  /.  Chem.  Soc,  97,  798  (1910). 
9.  Manchot  and  Manchot,  Z.  anorg.  Chem.,  226,  385  (1936). 

10.  Hieber  and  Lagally,  Z.  anorg.  Chem.,  251,  96  (1943). 

11.  I.  (1.  Farbenindustrie,  AC.,  German  Patents  531402  (Jan.  21,  1930)-  cf.  Chan. 

Ah,  .  25,  5523  (1931)-  and  531479   (Feb.  13,  1930)-  cf.  Chem.  Abs.,  25,  5521 
L931   ;  French  Patents  708209    Dec.  23,  L930     cf.  Chem.  Abe.,  26,  1399  (1932)- 
and  708379    Dec   26,  1930)-  cf.  Chem.  Abe.,  26,  1401  (1932). 

12.  Bennetl  (to  Catalyst  Research  Corporation),  U.  S.  Patenl  1975076    October  2, 

1934).- cf.  Caen      '        28,7439    I 

13.  Gilliland  and  Blanchard,  In*  ]   ■■•■  ■     2,234    1941 

14.  Job,  etal.,Compt.  rend.,  177, 1439  (1923  ;  183,  58,  392    1926) ;  137, 564  (1928) ; Bull. 

Soc.  cMm.,  41,  1041  d(.t27;;  Hieber  and  Romberg,  Z    anorg.  Chem.,  221,  321 
(1935). 


512  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

process,  showed  that  no  chromium  carbonyl  is  formed  before  the  hydrolysis 
of  the  Grignard  reagent.  Presumably  an  organic  carbonyl  derivative,  such 
as  Cr(CO)2l\4 ,  is  an  intermediate  product. 

The  hexacarbonyls  are  colorless,  crystalline  solids,  much  more  stable 
than  the  carbonyls  of  iron  or  nickel.  They  are  not  oxidized  in  air,  and  they 
may  be  sublimed  without  decomposition.  (Chromium  hexacarbonyl  de- 
posits some  chromium  above  140°C.) 

High -pressure  Synthesis 

Almost  all  of  the  known  carbonyls  have  been  prepared  by  reactions  be- 
tween metallic  halides,  sulfides,  or  oxides  and  carbon  monoxide  under 
pressure.  Such  reactions  are  especially  useful  in  cases  in  which  the  metallic 
compounds  are  largely  covalent.  For  example,  CoS  (NiAs  structure)  is 
quantitatively  converted  into  [Co(CO)4]2  at  200°  and  200  atmospheres 
pressure,  but  cobalt  oxide  does  not  react15.  Generally,  some  free  metal  must 
be  present  to  act  as  an  acceptor  for  the  nonmetal.  If  no  such  acceptor  is 
present,  the  lining  metal  of  the  autoclave  (for  example,  copper)  may  enter 
into  the  reaction: 

2CoS  +  8CO  +  4Cu  ->  [Co(CO)4]2  +  2Cu2S 


For  the  reaction 

2CoX2  +  4Cu  +  8CO  ->  [Co(CO)4]2  +  4CuX, 

at  250°  and  200  atmospheres  in  a  copper  lined  autoclave,  the  percentages  of 
conversion  into  the  carbonyl  are16: 

X  =  F       CI         Br       I 

%  conversion        0         3.5        9  100 

A  volatile  carbonyl  halide,  such  as  Co(CO)I2 ,  is  assumed  to  be  an  inter- 
mediate: 

CoI2  +  CO  -*  Co(CO)I2 

2Co(CO)I2  +  4Cu  +  6CO  ->  4CuI  +  [Co(CO)4]2 

The  increase  in  reactivity  with  increasing  covalency  of  the  cobalt  halide  is 
explained  by  an  increase  in  the  ease  of  formation  of  the  carbonyl  halide  in 
the  order  chloride-bromide-iodide. 

In  .some  cases  (e.g.,  iridium  halides  at  110°  and  atmospheric  pressure) 
the  order  of  reactivity  is  reversed17;  this  suggests  a  different  mechanism, 
such  as 

L5.  Hieber,  Schulten,  and  Marin,/,  anorg.  Chem.,  240,  261  (1939). 

L6.  Hieber  and  Schulten,  Z.  anorg.  Chem.,  243,  145  (1939). 

17.   Hieber,  et  al.t  Z.  anorg.  Chen,.,  245,  321  (1940) ;  246,  138  (1940). 


METAL  CARB0NYL8  AND  NITROS]  L8  513 

2IrX,       5C0     •  2Ir(C0)»Xi   ;    COX 
2Ir(CO)»X1  +  3C0     ►  2Ir(CO),X    I   COX 
2Ir(CO)aX  +  CO  —  2[Ir(CO)«]B  +  COX 

It  is  assumed  thai  the  formation  of  a  stable  compound  COX2  is  necessary 
for  the  completion  of  these  reactions.  Carbony]  iodide  is  not  known  and 
the  reaction  iMrl,  +  ICO  — »  2Ir(CO)sIj  +  Is  takes  place,  bul  there  is  no 

further  reaction.  The  chloride  is  the  only  halide  of  iridium  that  gives  appre- 
ciable yields  of  the  carbonyl  by  this  method;  even  here  the  chief  product 
is  Ir(CO)jCl.  However,  with  iridium  halides  at  high  pressure  in  the  presence 

of  a  halogen  acceptor,  the  order  of  reactivity  is  as  originally  given. 

The  use  of  such  nonmetals  as  iodine7,  18  and  sulfur19  (or  their  compounds) 
as  catalysts  in  the  synthesis  of  carbonyls  can  be  understood  in  terms  of 
these  reactions.  Sulfur,  for  example,  may  form  metal  carbonyl  sulfides 
which  upon  further  reaction  with  carbon  monoxide  produce  the  metal 
carbonyl: 

3Fe  +  2S  +  8CO  -+  Fe3S2(CO)8 

Fe3S2(CO)8  +  7CO  -»  3Fe(CO)5  +  2S 

This  mechanism  is  given  support  by  Hieber's  isolation7  of  both  Fe3S2(CO)8 
and  Fe3Se2(CO)8 . 

It  is  not  often  that  oxides  can  be  used  for  the  preparation  of  carbonyls. 
However,  the  best  synthesis  of  osmium  carbonyl  is  the  reaction  of  carbon 
monoxide  and  the  covalent  oxide  Os04  : 

Os04  +  9CO  -»  Os(CO)5  +  4C022°. 

In  some  cases  the  extreme  stability  of  the  intermediates  makes  the  prep- 
aration of  the  simple  carbonyls  difficult.  For  example,  rhenium  carbonyl 
halides  are  the  only  products  of  the  reaction  of  rhenium  halides  or  complex 
halides  with  carbon  monoxide.  Their  stability  is  demonstrated  by  such  re- 
actions as 

2Re  +  NiX,  +  14CO  ->  2Rc(CO)5X  +  Xi(C())4 

in  which  rhenium  acts  as  the  halogen  acceptor  for  the  formation  of  nickel 
carbonyl.  and21 

KR.o.  +  ((  1.   -    s(o  ->  KC1  +  Re(CO)iCl  +  COClj  +  3COj  . 
In  order  to  obtain  a  simple  rhenium  carbonyl  by  this  method  it  is  necessary 

18.  Geisenberger,  unpublished  experiments. 

19.  Mittasch,  Z.  angew.  Chem.,  41,  587,  827    L928). 

20.  Bieber,  et  al.,  Z   EUktrochem.,  49,  288  (1943  ;  Ber.,  75,  1172  C1942 

21.  Hi. -he.  et  al.,  /   anorg.  Chem.,  243,  164  (1939);  248,  243  (1941);  348,  256  (1941   . 


c 


514  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

to  use  Re2S7 ,  Re207  or  KRe04  as  the  starting  material,  the  reaction  being 
carried  out  in  the  absence  of  halogens. 

The  nature  of  the  metal  used  as  the  acceptor  influences  the  extent  to 
which  these  reactions  go.  If  cobalt  bromide  is  heated  with  silver,  copper, 
cadmium  or  zinc  in  an  inert  atmosphere,  the  extent  to  which  free  cobalt  is 
liberated  increases  in  the  order  Ag,  Cu,  Cd,  Zn.  When  the  inert  atmosphere 
is  replaced  by  carbon  monoxide,  the  extent  to  which  carbonyls  are  formed 
increases  in  the  same  order.  The  product  in  the  case  of  zinc  or  cadmium  is 
not  [Co(CO)4]2  but  a  mixed  carbonyl,  [Co(CO)4]2M.  This  tendency  of  the 
more  active  metals  to  form  mixed  compounds  must  be  considered  in  select- 
ing the  acceptor. 

In  the  experiment  just  described,  the  extent  of  carbonyl  formation  is 
much  greater  than  the  extent  of  the  corresponding  displacement  reaction  in 
the  absence  of  carbon  monoxide,  and  the  high  pressure  synthesis  may  not 
actually  involve  reduction  to  the  free  metal  followed  by  combination  to 
form  the  carbonyl.  This  is  supported  by  the  fact  that  iridium  and  osmium, 
which  are  inert  toward  carbon  monoxide,  form  carbonyls  by  the  high  pres- 
sure synthesis. 

Formation  by  Disproportionation  Reactions 

When  nickel(I)  cyanide  is  treated  with  carbon  monoxide,  nickel  car- 
bonyl and  nickel(II)  cyanide  are  formed22: 

2NiCN  +  4CO  ->  Ni(CN)2  +  Ni(CO)4 

A  similar  reaction  takes  place  when  a  complex  of  univalent  nickel  is  em- 
ployed, an  intermediate  probably  being  formed: 

K2Ni(CN)3  +  CO  ->  K2[Ni(CN)3CO] 

2K2[Ni(CN)3CO]  +  2CO  ->  Ni(CO)4  +  K2Ni(CN)4  +  2KCN 

Nickel  carbonyl  is  also  produced  when  carbon  monoxide  is  passed  into 
an  alkaline  mixture  of  a  nickel (II)  salt  and  etlryl  mercaptan  or  potassium 
hydrogen  sulfide  in  water;  the  formation  of  a  univalent  carbonyl  com- 
pound, followed  by  disproportionation,  is  postulated22 

2Ni(SH)2  +  2nCO  ->  2NiSH(CO)«  +  H8S2  (absorbed  by  alkali) 

2NiSH(CO)„  +  (4  -  2n)CO  ->  Ni(CO)4  +  Ni(SH)2 

Disproportionations  are  also  responsible  for  the  preparation  of  certain 
carbonyls  from  carbonyl  derivatives23: 

22.  Manchot  and  Gall,  Ber.,  59,  1060  (1926);  Ber.,  62,  678  (1929) ;  Beducci,  Z.  anorg. 
Chem.,  86,  88  (1914);  Blanchard,  Rafter,  and  Adams,  J".  Am.  Chem.  Soc,  56, 
16  (1934). 

23.  Hieber  et  al.,  Ber.,  63,  1405  (1930);  Z.  anorg.  Chem.,  221,  337  (1935). 


METAL  cMiBONYLS  AND  NITROSYLS  515 

3Fe(C0),CH,0H  +  4H+  ->  Fe(CO)5  +  2Fe++  +  3CII3OH  +  211,  +  4C0 

2[Fe(CO)4]i  +  3py  -*  3Fe(C0),py  +  3Fc(CO)6 

Cr(CO)«pyi  +  py  -•  Cr(CO)»pyi  +  CO 

!0)tpyi  +  15HC1  +  2H,0  -»  Cr(CO). 

+  2[CrCl»H»0]  (pyH)j  +  5pyHCl  +  3CO  +  311, 

Similar  reactions  arc  shown  by  some  of  the  carbonyls  themselves.  For 
example,  iron  enneacarbony]  is  formed  from  the  pentacarbony]  by  the  ac- 
tion  of  light  of  wave  Length  shorter  than  4100A. 

2Fe(CO)6  ->  Fe2(CO)9  +  CO 

The  product  undergoes  disproportionation  when  heated  in  benzene  or  ether 
solution. 

3Fe2(CO)9  -»  Fe3(CO)i2  +  3Fe(CO)5 

The  Formation  of  Carbonyl  Hydrides 

The  High-pressure  Synthesis 

Carbonyl  hydrides  sometimes  form  as  byproducts  of  the  high  pressure 
synthesis  of  carbonyls.  If  moist  cobalt  sulfide  or  iodide  is  treated  with 
carbon  monoxide  under  high  pressure  and  in  the  presence  of  an  acceptor, 
cobalt  carbonyl  hydride  forms15.  The  reaction  is  probably  2CoS  +  H20  + 
9CO  +  4Cu  ->  2Co(CO)4H  +  C02  +  2Cu2S.  This  method  has  also  been 
used  to  prepare  Rh(CO)4H,  Ir(CO)4H,  and  Os(CO)4H2 .  Cobalt  carbonyl 
hydride  also  results  when  cobalt  carbonyl  is  heated  with  hydrogen  and 
carbon  monoxide  (to  prevent  decomposition)  by  the  reversible  reaction 
[Co(CO)4]2  +  H2  =  2Co(CO)4H.  Some  cobalt  carbonyl  hydride  forms  when 
cobalt  or  cobalt  sulfide  is  heated  with  hydrogen  and  carbon  monoxide. 

2Co  +  8CO  +  H2  -►  2Co(CO)4H 

2CoS  +  8CO  +  H2  +  4Cu  ->  2Co(CO)4H  +  2Cu2S 

The  Bame  methods  have  been  used  for  the  preparation  of  rhodium  carbonyl 
hydride,  but  attempts  to  produce  iron  carbonyl  hydride  always  result  in 
the  formation  of  the  pentacarbonyl. 

Hydrolysis  of  Carbonyls 

Hieber  and  his  co-workers24  reported  the  formation  of  an  unstable  iron 
carbonyl  hydride  by  the  action  of  bases  upon  iron  pentacarbonyl: 

Fe(CO)5  +  Ba(OH)2  ->  Fe(CO)4H2  +  BaC03 

24.  Hieber  and  Leutert,  Z.  anorg.  Chem.,  204,  145  (1932);  Hieber  and  Z.  Vetter, 

anorg.  Chem.,  212,  145  (1933);  Hieber,  Mllhlbauer,  and  Khmaim,  Ber.,  65,  1090 
(1932). 


516  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Treatment  of  certain  derivatives  of  iron  carbonyl  with  acid  also  produces 
the  carbonyl  hydride 

Fe2(CO)4en3  +  8H+  ->  Fe(CO)4H2  +  Fe++  +  3(enH2)++ 
Disproportionation  Reactions 

Reactions  similar  to  those  used  to  prepare  carbonyls  may  be  used  to 
.     prepare  carbonyl  hydrides.  An  alkaline  solution  of  a  cysteine  cobalt(II) 
complex  absorbs  carbon  monoxide25,  presumably  forming  a  carbonyl  inter- 
mediate which  disproportionates  to  form  cobalt  carbonyl  hydride  and  a 
cobalt(III)  complex: 

9[Cocy2]=  +  8CO  +  2H20  -»  6[Cocy»]a  +  Co(OH)2  +  2Co(CO)4H 

Further  treatment  with  carbon  monoxide  produces  more  carbonyl  hydride 
and  regenerates  the  cysteine 

U  [Cocy8]-  +  6CO  +  70H-  -*  2C03=  +  3Cy=  +  3H20  +  Co(CO)4H 

The  carbonyl  hydrides  behave  as  very  weak  acids.  Hieber  and  co- 
workers26 give  the  following  data: 

2[Co(CO)4]-^  [Co(CO)4]2  +  2e-  E2°93°  =  -0.40 

3[Fe(CO)4]=;=±  [Fe(CO)4]3  +  6e~  E2°93  =  -0.74 

3[Fe(CO)4H]-^±  [Fe(CO)4]3  +  3H+  +  6e~  E2°93  =  -0.35 

Fe(CO)4H2  -  dibasic  acid  at  0° 

Kx  =  3.6  X  10-5 

K2  =  1.10  X  10"14 


True  salts  of  the  carbonyl  hydrides  are  formed  only  with  alkali  and  alkaline 
earth  metals  and  large  ammine  cations.  Compounds  with  other  metals  do 
not  have  the  properties  of  salts  and  are  discussed  under  mixed  carbonyls. 
Behrens27  prepared  carbonyl  salts  directly  in  liquid  ammonia: 

[M(CO)n]x  +  xyNa  <=*  x  Nay[M(CO)n] 

Attempts  to  prepare  a  chromium  carbonyl  hydride  by  means  of  this  reac- 
tion have  been  unsuccessful2613. 

Metal  Cakbonyl  Halides  and  Related  Compounds 

Some  metal  carbonyl  halides  have  been  isolated  as  intermediates  in  the 
preparation  of  metal  carbonyls  by  high  pressure  synthesis;  in  other  cases 

25.  Schubert,  /.  Am.  Chrtn.  Soc,  55,  4563  (1933). 

26.  Hieber  and  Ilubcl,  Z.  Naturforschung,  7b,  322  (1952);  Hieber  and  Abeck,  Z. 

Naturforschung,  7b,  320  (1952). 

27.  Behrens,  Z.  Naturforschung,  7b,  321-22  (1952). 


METAL  CARBONYLS  AND  NITR08YL8  517 

their  existence  is  only  postulated.  The  list  of  elements  which  form  car- 
bony]  halides  is  qoI  the  same  as  the  list  of  those  which  form  simple  car- 
bonyls.  For  palladium,  platinum,  copper,  and  gold,  which  form  no  simple 
carbonyls,  the  stability  of  the  carbonyl  halides  appears  to  be  iodide  <  bro- 
mide <  chloride".  The  stability,  ease  of  formation,  and  volatility  of  the 
compounds  of  the  carbonyl-forming  metals,  however,  all  show  trends  in  the 

opposite  direction. 

Carbonyl  halides  are  obtained  by  the  action  of  halogen  upon  carbonyl 
hydrides,  mixed  carbonyls,  simple  carbonyls,  or  other  carbonyl  halides:  For 
example29, 

Yel,  +  4CO  -»  Fe(CO)Js 

Fe(CO)5  +  I2  ->  Fe(CO)4I2  +  CO 
Fe(CO)4ll.  +  2IS  -»  Fe(CO)4Ij  +  2HI 
Fe(CO)4Hg  +  21,  -»  Fe(CO)Js  +  Hgla 

Mixed  Carbonyls 

Mixed  carbonyls,  such  as  [Co(CO)4]2Zn,  are  covalent  compounds  and  are 
soluble  in  organic  solvents;  they  are  therefore  not  to  be  classed  with  the 
salts  of  the  carbonyl  hydrides.  Typical  reactions  which  produce  these  com- 
pounds are  illustrated  by  the  equations: 

2CoBr2  +  3Zn  +  8C0  ->  2ZnBr2  +  [Co(CO)4]2Zn 

2Co  +  Zd  +  8CO  ->  [Co(CO)4]2Zn 

[Co(CO)4]2  +  Zn  ->  [Co(CO)4]2Zn 

Fe(CO)4H2  +  HgCl2  ->  2HC1  +  [Fe(CO)4]Hg 

Mercury  forms  mixed  carbonyls  most  readily;  among  the  other  metals 
which  form  them  are  zinc,  cadmium,  indium,  thallium,  and  tin. 

Structure  of  the  Carbonyls  and  Their  Derivatives 

Bond  Type 

The  carbonyl  group,  at  least  in  the  mononuclear  carbonyl,  may  be  con- 
sidered to  be  a  carbon  monoxide  molecule  (not  greatly  modified)  coordi- 
nated to  a  central  metal  atom  in  much  the  same  way  that  other  neutral 
molecules  or  ions  are  coordinated  to  central  cations.  This  postulate  is  the 
most  consistent  w  it h  the  energetics  involved  and  with  the  properties  of  the 
compounds,  thus  excluding  the  possibility  of  important  contributions  from 

Wagner,  Z.  arwrg.  Chun.,  196,  364  (1931). 
29.  Ilieber  et  al..  Ber.,  61,  1717  (1928);  Z.  anorg.  Chem.,  245,  296  (1940);  245,  305 
(1940). 


518  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

van  der  Waals  bonding30.  The  evidence  supporting  this  view  may  be  sum- 
marized as  follows: 

(1)  spectroscopic  data,  showing  thai  the  pairing  of  d  electrons  requires 
energy  of  the  order  of  50  kcal; 

(2)  the  nonpolar  character  of  simple1  carbonyls  as  shown  by  their  vola- 
tility; 

(3)  the  liberation  of  carbon  monoxide,  either  by  decomposition  or  by 
stepwise  replacement  with  neutral  molecules; 

(4)  the  diamagnetic  character  of  the  simple  carbonyls; 

(5)  the  C — O  bond  distance  (from  electron  diffract  ion  data)  of  between 
1.13  and  1.15  A.,  which  is  very  close  to  that  in  carbon  monoxide  itself 
(1.13  A); 

(6)  the  strongest  Raman  frequency  of  nickel  carbonyl  (2039  cm-1)  com- 
pares favorably  with  that  in  carbon  monoxide  itself  (2155  cm-1); 

(7)  the  analogy  between  the  simplest  carbonyl  compound-borine  car- 
bonyl BH3CO-  and  BF3NH3,  and  that  between  [PtCl2-PR3]  and 
[PtCl2-CO];and 

(8)  the  relation  between  the  position  of  a  metal  in  the  periodic  table  and 
the  composition  of  the  carbonyls  it  forms. 

Such  evidence  leads  to  the  conclusion  that  the  bonding  between  the 
metallic  element  and  the  carbonyl  group  in  the  mononuclear  compounds  is 
essentially  an  electron  pair  bond.  The  supposition  of  a  higher  electron 

i  density  than  that  supplied  by  a  two-electron  bond  finds  support  from  both 

resonance  considerations  and  a  shortening  of  bond  distance  observed  in 
diffraction  studies.  Spectroscopic  analyses  of  all  of  the  mononuclear  com- 
pounds show  that  the  bond  between  the  carbon  and  oxygen  in  the  carbonyl 
group  retains  the  characteristics  of  carbon  monoxide.  However,  with  the 
polynuclear  carbonyls  there  is  evidence  suggesting  a  similarity  in  structure 
between  the  carbonyls  and  aldehydes  or  ketones.  This  evidence  has  been 
studied  in  particular  with  the  iron  carbonyls. 

It  should  be  noted  that  elements  of  odd  atomic  number  form  no  mono- 
nuclear carbonyls,  whereas  elements  of  even  atomic  number,  in  forming 
mononuclear  carbonyls,  acquire  enough  carbonyl  groups  to  give  the  effec- 
tive atomic  number  of  the  next  inert  gas. 

Structure  of  the  Mononuclear  Carbonyls 

There  was  an  early  tendency  to  regard  the  carbonyls  as  ring  compounds. 
Werner  first  suggested  that  all  the  carbonyl  groups  are  attached  directly  to 
the  metal  atoms,  leading  to  the  supposition  by  Langmuir  that  in  these 
compounds  the  central  atom  attains  the  number  of  electrons  of  the  next 

30.  Syrian  and  Dyatkina,  "Structure  of  Molecules  and  the  Chemical  Bond,"  p. 
358,  New  York,  [nterscience  Publishers,  Inc.,  1950. 


METAL  CARBONYLS  AND  NITR08YL8  519 

Table  L6.2.  Compounds  with  ran  Cr(CO)(  Configuration 

Cr  U*2s'2ptts13pa3dl  forma  0  covalenl  bonds  (3d<4t*4p( 
||(— CN  M      CO  m      NO 

Mn(CN).»-  Mn(CN)iNO 

I  e  <'X)64-  Fe(CN)»CO  F.   CN)»NO- 

inert  gas11.  Sidgwick  termed  this  total  Qumber  of  electrons  the  "Effective 
Atomic  Number"  (E.A.V  -'.  Langmuir's  suggestion  has  been  found  to  hold 
without  exception  for  the  simple  carbonyls8*.  It  is  assumed  thai  each  carbon 
monoxide  molecule  donates  two  electrons  to  the  centra]  metal  atom;  thus, 
chromium,  iron,  and  nickel,  having  12,  10,  and  8  fewer  electrons  than 
krypton,  add  6,  5,  and  4  molecules  of  carbon  monoxide,  respectively.  It  is 
interesting  that  similar  electronic  configurations  result  with  several  differ- 
ent complexing  groups  to  give  the  same  E.A.N.,  as  shown  in  Table  16. 234. 

Numerous  methods  have  been  employed  in  the  determination  of  the 
structures  of  these  compounds.  Perhaps  the  most  conclusive  are  x-ray  and 
electron  diffraction  methods,  which  are  in  turn  supported  by  applications 
of  Raman  spectra,  infrared  spectra,  dipole  moments,  and  magnetochemical 
techniques.  The  metal  atom  is  surrounded  by  the  carbonyl  groups;  bonding 
to  the  metal  occurs  through  the  carbon  atom,  and  the  metal,  carbon,  and 
oxygen  atoms  are  collinear. 

The  structural  determination  of  nickel  tetracarbonyl  illustrates  the  con- 
clusions and  adds  insight  into  the  possible  electronic  configuration.  Early 
evidence  from  Raman  spectra  was  interpreted  to  indicate  a  planar  con- 
figuration35, but  electron  diffraction  studies  by  Brockway  and  Cross36  led 
to  the  conclusion  that  the  molecule  is  tetrahedral.  Further  study  by  means 
of  infrared  absorption37,  Raman  spectra38  and  the  observation  that  the 
compounds  show  no  dipole  moment39  add  support  to  the  tetrahedral  con- 
figuration. According  to  Pauling's  theory  of  directed  valence,  Xi++  has  the 
configuration  (3s23p63d8) .  The  eight  added  electrons  go  into  the  states 
&P4**4p*,  giving  rise  to  dsp2  hybrid  bonds,  which  are  planar.  The  atom  Ni° 
has  the  configuration  (3s23p63d84s2) .  Degeneration  of  the  4s  electrons  to  the 
M  level  permits  the  formation  of  sp3  hybrid  bonds,  which  are  tetrahedral 

31.  Langmuir,  Science,  54,  65  (1921). 

32.  Sidgwick,  "Electronic  Theory  of  Valency,"  p.  163,  Oxford  Press,  1927. 

33.  Bl&nchard,  Chem.  Reos.t  26,  409  (1940). 

34.  Hieber,  Z.  angeu  .  Ckem.t  55,  7  '1942). 

35.  Duncan  and  Murray,  ./.  Chem.  Phy8.s  2,  636  (1934). 

36.  Brockway  and  Croat     /    Cht  m.  JJi<!J*-,  3,  828  (1935). 

37.  Crawford  and  ( IroBS,  ./ .  (  ft  n  .  Ph  -    .  6.  525    L938). 

38.  Crawford  and  Horiwits,  /.  Chem.  Phys.,  16,  1  17    l'.MS). 

39.  Sutton,  New,  and  Bentley,  ./.  Chem.  Soc.,  1933,  652. 


520  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Table  16.3.  Interatomic  Distances  from  Electron  Diffraction 

Data  (A) 


Bond 

c-o 

M— C 

N-0 

M-N 

M— C(calc) 

Shortening 

Ni(CO)4 

1.15 

1.82 

1.98 

0.16 

Fe(CO)5 

1.15 

1.84 

2.00 

0.16 

Cr(CO)6 

1.15 

1.92 

2.02 

0.10 

Mo  (CO)  6 

1.15 

2.08 

W(CO)6 

1.13 

2.06 

Co(CO)3NO 

1.14 

1.83 

1.10 

1.76 

1.99 

0.16 

Fe(CO)2(NO)s 

1.15 

1.84 

1.12 

1.77 

2.00 

0.16 

Co(CO)4H 

1.16 

1.83* 

1.99 

0.16 

Fe(CO)4H2 

1.15 

1.82* 

2.00 

0.18 

*  Average  bond 

lengths 

and  commensurate  with  the  known  configuration  for  the  nickel  carbonyl40. 
This  situation  is  comparable  to  that  found  in  [Cu(CN)4]~  and  [Zn(CN)4]=, 
which  are  tetrahedral. 

For  the  hexacarbonyls  of  chromium,  molybdenum,  and  tungsten  the 
octahedral  configuration  of  the  carbonyl  groups  has  been  established  by 
x-ray41,  electron  diffraction42,  and  infrared  spectra  studies43. 

Because  it  shows  the  unusual  coordination  number  of  five,  iron  penta- 
carbonyl  has  inspired  a  great  deal  of  experimental  work  and  many  theo- 
retical speculations.  Most  of  the  evidence  supports  the  trigonal  bipyramid 
structure  (as  in  PF5)  proposed  by  Ewens  and  Lister  (based  on  their  elec- 
tron diffraction  study)44.  The  small  dipole  moment  has  been  interpreted  to 
indicate  a  nonequivalence  of  bonds45  but  experimental  conditions  or  a 
polarization  in  the  molecule  may  account  for  the  observed  dipole  moment46. 
Infrared  spectra  add  evidence  for  the  trigonal  bipyramid  (dsp*)  structure46. 

Table  16.3  summarizes  electron  diffraction  determinations  of  interatomic 
distances  in  some  carbonyls,  carbonyl  hydrides,  and  nitrosyls47. 

The  M — C  bond  distance  is,  in  each  case,  shorter  by  approximately 
0.16A  than  the  sum  of  the  corresponding  covalent  radii.  Brockway  and  his 
co-workers  attributed  this  bond  shortening  to  the  contribution  of  a  double 

40.  Pauling,  ''Nature  of  the  Chemical  Bond,"  2nd  ed.,  p.  251,  Ithaca,  N.  Y.,  Cornell 
University  Press,  1940. 

41.  Rudorff  and  Hofmann,  Z.  phys.  Chem.,  B28,  351  (1935). 

42.  Brockway,  Ewens,  and  Lister,  Trans.  Faraday  Soc,  34,  1350  (1938). 

43.  Sheli  ne,  J.Am.  Chem.  Soc,  72,  5761  (1950). 

44.  Ewens  and  Lister,  Trans.  Faraday  Soc,  35,  681  (1939);  Ann.  Reports,  36,  166 
(1939). 

45.  Bergmann  and  Engel,  Z.  phys.  Chem.,  B13,  232  (1931) ;  GrafTunder  and  Heymann, 
Z.  phys.  Chem.,  B15,  377  (1932). 

46.  Sheline  and  Pitzer,  /.  Am.  Chem.  Soc,  72,  1107  (1950). 

47.  Anderson,  Quart.  Revs.,  1,  331  (1947). 


METAL  CARBONYLS  AND  NITROSYLS  521 

bond  structure,  assuming  the  resonance  forms36- 42 

M  «-  C=eO         and         M=C=0 

Similar  and  extended  considerations  arc  to  be  found  in  other  sources40,48. 
Hieber  has  suggested  that  the  decrease  in  bond  distance  may  be  due  to 
secondary  interactions  between  the  it  electrons  of  the  C=()  bond  and  the 
3d  electrons  of  the  metal  atom49.  Whichever  explanation  is  invoked,  the 
C — 0  distance  corresponding  to  the  carbon  monoxide  triple  bond  character 
must  be  preserved  to  conform  with  experimental  evidence.  Thus  the  two- 
electron  bond  structure  is  dominant.  This  conclusion  is  borne  out  by  the 
calculated  force  constants  of  the  Fe — C  bond  in  iron  pentacarbonyl,  given 
by  Sheline46  as  nearly  the  same  as  those  found  in  the  metal  alkyls50. 

Structure  of  the  Polynuclear  Carbonyls 

Since  elements  of  odd  atomic  numbers  cannot  attain  the  rare  gas  con- 
figuration by  simple  coordination  of  electron  pairs,  polymerization  occurs 
in  carbonyl  formation.  This  polynuclearity  is  also  evidenced  in  the  lower 
carbonyls  of  the  even  numbered  elements.  The  postulation  of  the  structures 
of  the  polynuclear  compounds  presents  greater  problems  than  in  the  case  of 
the  mononuclear  carbonyls,  and  these  problems  have  not  as  yet  been  com- 
pletely solved. 

Sidgwick  and  Bailey51  proposed  to  account  for  the  formulas  of  poly- 
nuclear carbonyls  on  the  assumptions  that  (1)  the  metal  atoms  acquire  the 
configuration  of  the  next  inert  gas,  and  (2)  the  carbon  monoxide  molecule 
is  able  to  join  two  metal  atoms  by  linking  through  carbon  to  one  and 
through  oxygen  to  the  other.  Iron  enneacarbonyl  was  represented  as 
(CO)4Fe  <—  C=0  — >  Fe(CO)4  ,  each  iron  achieving  the  krypton  configu- 
ration by  accepting  five  pairs  of  electrons.  Cobalt  carbonyl  was  pictured  as 
(CO)4Co — CO — Co(CO)3 ,  in  which  one  cobalt  has  an  effective  atomic  num- 
ber of  37  and  the  other  35;  the  excess  electron  on  the  former  is  passed  to 
the  latter  to  give  a  krypton  structure.  A  similar  formulation  was  suggested 
for  [Co(CO)3]4  in  which  the  cobalt  atoms  are  assumed  to  be  linked  to  each 
other  by  carbonyl  groups  in  the  form  of  a  tetrahedron;  an  electron  transfer 
between  two  cobalt  atoms  effects  an  inert  gas  structure.  Such  an  unsym- 
metrical  structure  appears  somewhat  tenuous.  Brill52  inferred  a  trigonal 
symmetry  of  iron  enneacarbonyl  from  x-ray  studies.  Powell  and  Ewens33 

48.  Syrkin  and  Dyatkina,  Acta  Physicochim.  U.R.S.S.,  20,  137  (1945);  Long  and 

Walsh,  Trans.  Faraday  Soc,  43,  342  (1947). 

49.  Hieber,  Die  Chemie,  55,  25  (1942). 

50.  Gutowsky,  ./.  Chan.  Phys.,  17,  128  (1949). 

51.  Sidgwick  and  Bailey,  Proc.  Roy.  Soc,  A,  144,  521  (1934). 

52.  Brill,  Z.  Krist,  65,  85  (1927). 

53.  Powell  and  Ewens,  J.  Chem.  Soc,  1939,  286. 


522  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

confirmed  this  by  means  of  Patterson  and  Fourier  analysis,  ascribing  struc- 
ture (I). 


//\ 

o=c   c=o  c=o 

V 

in 

o 
CO 

Thus  the  Sidgwick-Bailey  rule  does  not  apply  here.  In  order  to  account  for 
the  observed  diamagnetism,  Klemm54  suggested  spin  coupling  between  the 
unpaired  electrons  of  each  iron  atom.  Powell  and  Ewens  support  this  view, 
noting  that  the  iron-iron  distance  is  only  2.46  A.  Three  of  the  CO  groups 
are  predicted  to  be  ketonic  in  character,  while  the  terminal  CO  groups  are 
linear  and  are  true  carbon  monoxide  types.  These  assignments  are  sup- 
ported by  the  spectral  data  of  Sheline46.  An  alternative  viewpoint  is  that  of 
Jensen55,  who  thinks  of  Fe2(CO)9  as  a  hybrid  of  the  resonance  forms  (II) 
and  (III). 

i 

OHC-*Fe — -*C -Fe^GHO 

ii 
o 

00  Cm) 

Ewens  has  criticized  these  resonance  structures,  stating  that  they  contain 
the  equivalent  of  a  covalent  iron-iron  bond  but  the  compounds  do  not 
have  the  color  expected  from  an  iron-iron  bond56.  The  assumption  of  a 
metal-metal  bond  appears  logical  in  view  of  these  findings  and  other  recent 
studies  on  intermetallic  bonding.  The  postulation  of  structures  for  the 

54.  Klemm,  Jacobi,  and  Tilk,  Z.  anorg.  Chem.,  201,  1  (1931). 

55.  Jensen  and  Asmussen,  Z.  anorg.  Chem.,  252,  234  (1944). 

56.  Ewens,  Nature,  161,  530  (1948). 


9 

o,cx   /  \^   ^ 
c 

=0 
=0 
=0 

II 

o 

METAL  CARBONYLS  AND  NITR0S1  LS 

polynuclear  compounds  of  elements  such  as  cobalt  presents  the  same  diffi- 
culties, it*  the  two  cobalt  atoms  have  the  same  effective  atomic  Dumber.  Ii 
they  do,  however,  this  number  is  35,  and  other  hypotheses  are  necessary 
to  account  for  the  absence  of  paramagnetism.  Spectral  studies  have  not 

yet   continued  the  presence  of  a  ketonic  group  ill  Co    CO)g  .  There  is  the 

possibility  of  direct  metal-metal  bonding  without  the  ketonic  bridge  struc- 
ture I  0  (  :(  CO)4 ,  but  in  view  of  the  presence  «>!'  ketonic  bonds 
in  iron  enneacarbonyl,  the  bridge-like  structure  appears  more  plausible, 

perhaps  coupled  with  the  intermetallic  bond. 

Similar  bridge-like  structures  have  been  suggested  for  [Cu(CO)s]i*7 
[Re(CO)JjM  and  other  dinuclear  compounds06. 

I  >-mium  enneacarbonyl,  which  is  soluble  in  benzene  and  which  sublimes, 
differs  markedly  from  the  corresponding  iron  compound,  which  is  insoluble 
in  benzene  and  does  not  sublime.  Such  properties  might  indicate  a  differ- 
ence in  structure,  though  the  enhanced  covalent  character  of  the  osmium 
compound  may  arise  simply  from  the  larger  size  of  the  metal  atoms,  per- 
mitting a  more  strictly  covalent  intermetallic  bond. 

Few  of  the  more  complex  polynuclear  carbonyls  have  been  examined, 
only  the  structure  of  the  iron  tetracarbonyl  having  been  studied  in  detail. 
In  1930,  Hieber  and  Becker59  proposed  the  following  structures,  which  are 
based  on  the  properties  and  reactions  of  the  material: 


CO 

OC           CO 

ocx|  JZO 

0=C. C=0/C= 

\/ 

=  0 

OC Fe CO 

/  \ 

o=c           c=o 

=  0 

\  / 

oc^No 

Fe 

/\ 

o=c           c=o 

CO 

\  / 

m 

OC Fe CO 

OC  CO 

Brill60  performed  the  only  x-ray  diffraction  studies  yet  made  on  this  com- 

57.  Robinson  and  Btainthorpe,  Suture,  153,  24  (1944). 

58.  Hieber  and  Fuchs,  Z.  anorg.  Chem.,  248,  2.56  (1941). 
50.  Hieber  and  Becker,  Ber.,  63B,  1406  (1930). 

60.  Brill,  Z.  Krist.,  77,  36  (1931). 


524  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

l)ound  and  found  that  although  all  of  Hieber's  structures  find  correspond- 
ence with  the  crystal  structure  determination,  structure  (VI)  is  the  most 
logical.  He  depicted  it  in  (VII). 


(sn)  Csnr) 

Sidgwick  and  Bailey51  represented  the  structure  as  shown  in  (VIII).  Such 
a  structure  does  not  appear  likely  from  the  preceding  discussions  of  the 
dinuclear  compound.  The  central  iron  atom  of  Brill's  structure  would  be 
expected  to  exhibit  paramagnetism  unless  a  form  of  metallic  bond  exists 
between  the  iron  atoms;  such  a  bond  appears  quite  reasonable.  The  spectra 
of  this  compound61  show  both  infrared  and  ultra  violet  bands  corresponding 
to  the  known  frequencies  of  carbon  monoxide  and  of  theketonic  or  aldehydic 
group. 

The  high  solubility  of  the  tetracarbonyl  in  organic  solvents  has  been  in- 
terpreted to  mean  that  the  three  empty  4p  orbitals  of  the  central  atom 
furnish  convenient  sets  of  empty  orbitals  through  which  the  Fe3(CO)i2 
molecule  can  solvate61.  The  solubility  might  also  be  explained  by  the  in- 
crease of  metallic  covalent  bonding.  In  any  case,  the  assignments  of  elec- 
trons to  specific  locations  is  tenuous.  Electron  densities  may  be  depicted  by 
t  he  possible  resonance  structures.  Syrkin62  has  suggested  that  perhaps  one 
of  the  main  resonance  forms  for  the  tetracarbonyl  is  (IX). 

61.  Sheline,  /.  Am.  Chem.  Soc,  73,  1615  (1951). 

62.  Syrkin  and  Dyatkina,  "Structure  of  Molecules  and  the  Chemical  Bond,"  p.  364, 

New  York,  [nterscience  Publishers  Inc.,  1950. 


METAL  CARBOXYLS  AM)  MTh'OSYLS 


IX 
Such  a  representation,  though  differing  from  the  above  structures,  satisfies 
the  genera]  properties  and  observations  previously  made 

By  analogy  with  the  iron  carbonyls,  similar  rules  and  theories  should 
apply  to  other  polymeric  carbonyls.  Higher  degrees  of  polymerization  lead 
to  structures  which  give  the  molecules  low  solubility  and  nonvolatility.  An 
example  is  Rh4(C())ii10.  Ormont63  has  studied  the  conditions  of  formation 
and  the  stability  of  the  carbonyls.  His  conclusions  are  summarized  in  several 
rules  which  relate  stability  to  effective  atomic  number  and  steric  configu- 
ration. From  heat  of  formation  data,  Ormont  advances  the  idea  that  metals 
of  the  zinc  group  should  form  tricarbonyls. 

Pospekhov64  has  outlined  a  principle  of  formation  for  the  polynuclear 
carbonyls  which  stems  largely  from  Ormont's  considerations  and  is  mark- 
edly similar  to  the  Sidgwick  analysis.  It  is  general  enough  that  it  does  not 
necessitate  the  hypothesis  of  bonding  through  both  oxygen  and  carbon. 
An  intermetallic  bond  accounts  for  the  observed  diamagnetism.  Assuming 
that  each  CO  molecule  supplies  two  electrons  to  the  metal  atom,  a  quantity 
A  is  defined  as  the  effective  atomic  number  of  the  central  atom,  minus  the 
atomic  number  of  the  next  inert  gas.  A  metal  carbonyl  will  be  polymeric 
if  A  <  0.  The  degree  of  polymerization  is  equal  to  1  —  A.  The  resulting 
polymers  are  assumed  to  be  bonded  through  the  metal  atoms.  This  rule, 
though  not  in  strict  accord  with  the  ketonic  bridge  structures,  accounts  for 
all  the  known  formulas  for  metal  carbonyl  polymers.  The  rule  predicts 
formulas  for  materials  the  molecular  weights  of  which  have  not  yet  been 
determined,  such  as  [Ru(CO)4]3  ,  [Ir(CO)3]4  ,  [Ag(CO)3]2 ,  and  [Cu(CO),]a  . 
Mechanisms  for  formation  have  been  suggested65.  The  recently  prepared 
manganese  carbonyl61  has  the  predicted  composition,  [Mn(CO)»]a . 

63.  Ormont,  Acta  Physicockim.,  U.R.S.S.,  11,  585  (1939);  ./.  Phye,  Chem.    I  88R  . 

12,  256    L938  ;  Acta  Physicochim.,  U.R.S.S.,  19,  :>71    1944   ;  Acta  Physicochim, 
I    /,'.>  .v.  21,  H3    1946). 

64.  Pospekhov,  •/.  Phys.  Chem.  (U.S.S.R.  .21,  II     1947  .  Zhur.  Obshekei  Khim,  18, 

2045    L948 
Pospekhov,  Zkur.  Obshekei  Khan.,  18,  610    1948 
66.  Brimm,  private  communication;  see  also  Elund,  Sentell,  and  Norton,  ./.  Am. 
c  .  71,  1806    I'M1- 


526 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


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METAL  CARBON YLS  AND  NITROSYLS  .">'_>  7 

The  Metal  Carbonyl  Halides;  Their  Derivatives  and  Properties  Re- 
lated to  Structupe 

The  most  common  metallic  carbonyl  halides  are  listed  in  Table  L6.4. 
Iron  pentacarbonyl  adds  free  halogens  at  low  temperatures  to  form 
Fe(C<  I  \  .  which  in  turn  decomposes  In-low  0°  to  give  Fe((  '<  )>1X/'7.  This 
suggests  thai  there  Is  a  tendency  for  the  iron  to  acquire  a  coordination 

number  of  six,  though  the  tendency  LS  lessened  by  the  size  of  the  carbonyl 
groups.  Mixed  halides  such  as  Fe(CO)4ICl  form,  but  decompose  to  mixtures 
of  the  symmetrical  compounds,  e.g.,  Fe(CO)4l2  and  Fe(CO)4Cl2 .  The  dia- 
magnetic  compounds  Fe(CO)4SbCl5  and  Fe(CO)4SnCl4  have  been  shown, 
both  by  molecular  weight  determination  in  benzene  and  nitrophenol,  and 
by  conductivity  measurements,  to  be  nonelectrolytes,  represented  by  the 
structures68 

CI  CI 

/  \  /  \ 

(OC)4Fe      SbCl3  and  (OC)4Fe      SnCl2 . 

\  /  \  / 

CI  CI 

The  lower  carbonyl  halides  are  probably  dimeric,  containing  halogen 
bridges,  as  in  [Fe(CO)2I]27- 68. 

I 

/\ 
(OC)oFe  Fe(CO)2. 

\/ 

I 

This  compound  reduces  silver  nitrate  in  nitric  acid  and  reacts  with  water 
to  give  iron  (II)  hydroxide  and  hydrogen.  The  only  other  carbonyl  halides 
of  the  first  transition  series  are  the  unstable  cobalt  iodide  monocarbonyl 
and  the  tetrameric  copper  carbonyls;  the  latter  are  thought  to  be  structural 
analogs  of  [(C2H5)3As-CuI]469. 

The  osmium  halides  show  an  increasing  tendency  towards  the  formation 
of  the  dimeric  [(^(CO)^]^  (Table  16.5).  Again  a  halogen  bridge  appears 

67.  Hieber  and  Lagally,  Z.  anorg.  Chem.,  245,  305  (1940);  Hieber  and  Bader,  Z. 

anorg.  Chem.,  190,  193,  215  (1930);  Z.  anorg.  Chan.,  201,  329  (1931). 

68.  Hieber  and  Lagally,  Z.  anorg.  Chem.,  245,  295  (1940). 

69.  Mann,  Purdie,  and  Wells,  ./.  Chem.  Soc,  1936,  1503;  Emeleus  and  Anderson, 

"Modern  Aspects  of  Inorganic  Chemistry,"  p.  117,  New  York,  D.  Van  Nos- 
trand  Company,  Inc.,  1938. 


528 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Table  16.5.  Stability  and  Composition  of  Osmium  Carbonyl  Halides 


Type 

Os(CO)4X2 

Os(CO)3X2 
i 

CI    ' 

colorless 

colorless 

Br 

colorless        and 
light -yellow 

3rellow 

I 

yellow          and 
dark-yellow 

dark  yellow 

maximL 

m  stability 

Os(CO)2X2 


[Os(CO)iX]2 


light  yellow  canary  yellow 

light  yellow  orange  yellow 


in  the  most  logical  structures,  as  in 


(OC)4Os 


Os(CO)4     and     (OC)2Rh 


CI 

/     \ 


I  CI 

The  rhenium  compound  Re(CO)5X  illustrates  the  significance  of  the 
inert  gas  type  of  structure  in  determining  the  formulas  for  carbonyl  hal- 
ides71. An  increase  in  volatility  and  color  from  the  chloride  through  the 
iodide  implies  that  the  iodide  is  essentially  nonpolar;  the  chloride,  however, 
has  been  shown  to  have  partial  ionic  character  in  dioxane72. 

The  carbonyl  halides  show  typical  carbonyl  character  in  volatility,  solu- 
bility in  organic  solvents,  and  other  properties.  The  ease  of  formation  in- 
creases down  the  groups  of  the  periodic  table,  maximum  CO  contents 
being  found  in  Re(CO)5X,  Os(CO)4X,  Ir(CO)3X,  and  Pt(CO)2X2 .  The 
last  two  have  an  incomplete  rare  gas  configuration,  involving  sixteen  elec- 
trons, which  is  also  found  in  [Ni(CN)4]=. 

In  relation  to  the  structure  of  the  carbonyls,  it  is  interesting  that  the  CO 
groups  may  be  replaced  by  molecules  of  ammonia,  pyridine,  or  alcohol,  and 
two  CO  groups  may  be  replaced  by  one  bidentate  chelating  group  like  eth- 
ylenediamine  or  o-phenanthroline,  yielding  Fe(CO)3(NH3)2 ,  Cr(CO)3py3 , 
Fe2(CO)5en2 ,  or  Ni(CO)2(o-phen)73b.  The  compound 

CO 

I 

(O  C)3Fe— S— Fe— S— Fe(  CO)3 


CO 

is  similar  in  some  respects  to  iron  tetracarbonyl74.  Analogs  are  known  in 
which  sulfur  atoms  are  replaced  by  selenium,  and  the  CO  groups  by  pyridine 

70.  ilieber  and  Stallman,  A.  Electrochem.  angew.,  49,  288  (1942). 

71.  Hieber  and  Schulten,  Z.  anorg.  Chem.,  243,  164  (1939). 

72.  Schuh,  Z.  anorg.  Chem.,  248,  276  (1941). 

73.  Hieber,  Z.  Elektrochem.,  43,  390  (1937);  Hein,  Z.  angew.  Chem.,  62A,  205  (1950). 

74.  Ilieber  and  Geisenberger,  Z.  anorg.  Chem.,  262,  15  (1950). 


HBTAL  CARBONYLS  AND  NITR08YL8  529 

molecules.  Examples  of  mercapto  forms71  are  the  monomelic  (OC)gFe — 

S— C6II5  and  the  dimeric  [(OC)3Fe— S— C2H5]276.  These  compounds  indi- 
cate the  influence  of  steric  hindrance  in  the  formation  of  carbony]  deriva- 
tives. A  number  of  other  carbony]  derivatives  with  organic  bases,  phos- 
phonium  and  arsonium  compounds77  and  organometallic  bases78  have  been 
prepared.  The  structure  of  such  nonsalt-like  heavy  metal  derivatives  as 
[Co(CO)4]i_3M,  where  M  =  T1+,  Zn++  Cd++,  Hg++,  Ga+++,  In+++  or 
Tl+++,  is  best  represented*4  by  a  bridge-like  form: 

CO 
OC— Co— CO 
CO  \ 

Hg 

\  CO 

OC— Co— CO 
CO 

Theoretical  considerations  have  been  applied  by  Ormont79  to  the  forma- 
tion of  the  metal  carbonyl  halides  and  their  derivatives.  With  the  halide 
forms  such  as  Fe(CO)sX2  the  conclusion  was  reached  that  an  energetically 
unstable  compound  forms,  independent  of  the  value  of  A.  (see  p.  525). 
This  accounts  for  the  fact  that  the  compounds  Fe (CO)4X2  are  thermally 
unstable  at  298°K,  whereas  CuCl2-2CO  and  CuBr2-2CO  are  stable  at  this 
temperature.  The  argument  is  further  advanced  that  elements  with  valence 
electrons  in  different  quantum  levels  must  form  halides  with  a  small  num- 
ber of  carbonyl  groups  although  A  is  often  quite  different  from  zero.  This 
tendency  has  been  noted  above  with  platinum  and  iridium  compounds. 
This  explanation  is  useful  in  interpreting  the  properties  of  these  compounds 
when  the  rules  of  effective  atomic  number  are  inapplicable. 

Pospekhov80  has  concluded  that  the  volatility  and  the  color  of  the  car- 
bonyls  and  nitrosylcarbonyls  are  determined  by  A,  calculated  on  the  basis 
that  the  carbonyl  group  supplies  two  electrons  and  the  nitrosyl  group, 
three  electrons.  For  A  =  0  the  properties  of  high  volatility  and  the  absence 
of  color  are  observed.  A  more  negative  value  of  A  is  accompanied  by  deeper 
color  unless  the  formation  of  polymers  counteracts  the  effect.  When  the 
carbonyl  molecules  are  replaced  by  amines  and  other  groups,  the  intensi- 
fication of  color  is  attributable  to  dissymmetry  in  the  electron  cloud.  In 

75.  Hieber  and  Spacu,  Z   anorg.  Chem.,  233,  353  (1937). 

76.  Reihlen,  et  al,  Ann.,  465,  95  (1928). 

77.  Reppe,  et  al.,  Ann.,  560,  104,  108  (1948). 

78.  Hein  and  Heuser,  Z.  anorg.  Chem.,  249,  293  (1942). 

79.  Ormont,  Acta  Physicochim.  U.R.S.S.,  21,  741  (1946) ;  Acta  Physicochim.  U.R.S.S.. 

12,  757  (1940). 

80.  Pospekhov,  Zhur.  Obshekei  Khim.,  20,  1737  (1950);  J.  Gen.  Chem.  U.S.S.R.,  20, 

1797  (1950). 


530  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

nonvolatile  molecules  of  Fe2(CO)9  and  Ru2(CO)9  it  is  postulated  that  dis- 
symmetry leads  to  crowding  the  carbonyl  and  the  formation  of  closed 
cycles,  wherein  the  number  of  electrons  supplied  to  the  metal  atom  per 
carbonyl  group  is  less  than  two  (A  <  0). 

The  Structures  of  Carbonyl  Hydrides 

The  comparison  of  the  formulas  of  the  carbonyl  hydrides  with  the  formu- 
las of  mononuclear  carbonyls  for  the  series  iron,  cobalt,  nickel 

Fe(CO)5  —  Ni(CO)4 

Fe(CO)4H2  Co(CO)4H 

shows  that  the  effective  atomic  number  of  36  is  achieved  in  each  case  if 
each  carbonyl  group  contributes  two  electrons  and  each  hydrogen  atom, 
one.  However,  the  hydrogen  atom  does  not  appear  to  contribute  to  the 
spatial  arrangement,  since  both  of  the  above  hydrides,  like  nickel  carbonyl, 
are  tetrahedral.  Two  proposals  have  been  made  to  account  for  the  struc- 
ture. Hieber's  postulation24a » 49  of  a  structure  into  which  the  hydrogen  atoms 
are  incorporated  as  protons  is  similar  to  the  diborane  structure  proposed 
by  Pitzer81.  Ewens  and  Lister  proposed82  that  an  electron  from  hydrogen  is 
transferred  to  the  metal  atom  and  that  the  resulting  proton  is  coordinated 
to  the  oxygen  atom  of  a  carbonyl  group.  The  resulting  group  (:C: :  :0:H+) 
would  be  isoelectronic  with  the  nitrosyl  group  ( :  N : : :  0 : +),  and  the  formula 
for  cobalt  carbonyl  hydride,  for  example,  should  be  written  Co(CO)3(COH). 
Similarities  between  carbonyl  hydrides  and  nitrosylcarbonyls  will  be 
pointed  out  later.  Hieber49  has  pointed  out  that  this  proposal  is  equivalent 
to  the  proposal  of  a  quaternary  oxonium  ion  (with  a  formal  charge  of  2+ 
on  the  oxygen  atom),  which  is  unlikely.  Although  such  a  group  ought  to 
be  stablized  by  alkylation,  no  alkyl  derivatives  have  yet  been  formed. 
Moreover,  evidence  for  the  existence  of  two  different  M — C  and  C — 0 
bond  distances  within  the  molecule  is  lacking. 

However,  by  reviewing  some  of  the  properties  of  the  carbonyl  halides, 
a  logical  structure  can  be  proposed.  The  existence  of  mixed  carbonyls  such 
as  [Co(CO)4]2Zn  suggests  the  possibility  of  such  anions  as  [Co(CO)4]~  and 
[Fe(CO)4]=.  The  existence  of  these  anionic  forms  has  been  shown  in  the 
determination  of  acid  equilibrium  constants  and  electrode  potential  values. 
The  conductivity  of  M(CO)4Hn  in  pyridine  is  similar  to  that  of  a  strong- 
electrolyte.  The  hydrides  are  soluble  in  liquid  ammonia,  forming  low- 
melting  ammonia  derivatives  like  [Fe(CO)4]  (NH4)2  and  [Co(CO)4]NH4 . 
These  compounds  behave  as  acids  in  liquid  ammonia7  • 83.  Typical  acid  re- 

81.  Pitzer,  J.  Am.  Chem.  Soc,  67,  1126  (1945). 

82.  Ewens  and  Lister,  Trans.  Faraday  Soc,  35,  681  (1939). 

83.  Hieber  and  Schulten,  Z.  anorg.  Chem.,  232,  17  (1937) ;  Hieber  and  Fack,  Z.  anorg. 

Chem.,  236,  83  (1938). 


METAL  CARBONYLS  AND  NITR08YL8 


531 


actions  are  to  be  found  in  titrations,  salt  formation,  and  liberation  of  hy- 
drogen by  alkali  metals.  Ionic  properties  arc  found  in  all  of  the  dcrivat  i\  efi 
containing  alkali  and  alkaline  earth  metals. 
Thus  the  most  likely  resonance  forms7  may  be  depicted  as 


[m-<4:Hm 


C  ^  0: 


II 


:ili(l 


M 


++       "I 

O— H 


As  noted  above,  there  is  disagreement  as  to  the  last  of  these. 

Ilieber  has  offered  a  reaction  mechanism  to  explain  the  formation  of 
these  hydrides: 


o< 

:    o 

\c 

Fe«^-  C= 

=0-^-> 

/c 

OC     0 

H 

\ 

OC     0 

0 

\c 

i 

Fe<- 

C— 0 

/c 

T 

OC     0 

0 

/ 

H 

_ 

OH- 


"OC     O         OH 

\c      I 

Fe  <-  C=0 

/C 

OC     o 


H 
O 
C 
OC— Fe— CO  +  CO- 
C 
O 
H 


The  structures  of  the  low-melting  ammonia  derivatives  are  postulated 
to  contain  hydrogen  bonds. 

H3N-H        CO 

OC— Fe— CO 

CO       H-.-NH3 


Coordination  Compounds  Containing  the  Nitrosyl  Group 

Nitric  oxide  is  able  to  form  coordination  compounds  in  much  the  same 
way  that  carbon  monoxide  does.  However,  nitric  oxide  differs  from  carbon 
monoxide  in  one  important  respect — it  is  an  odd  molecule.  It  may  therefore 
be  expected  to  form  coordination  compounds  in  three  different  ways:  (1) 
loss  of  the  odd  electron  followed  by  coordination  of  the  resultant  \'<  > 
group,  (2)  the  gain  of  an  electron  followed  by  the  coordination  of  the  result- 
ant NO-  group,  (3)  coordination  of  the  neutral  N( )  group84.  To  these  must 
be  added  the  possibility  that  the  nitrosyl  group  forms  a  double  bond  with 
the  metal  atom;  this  will  be  considered  later. 

The  fact  that  reduction  of  [Fe(CN)»NO]~  yields  an  ammine 

st.  Moeller,  J.  Ch  m .  E<L,  23,  441  (1946) ;  23,  .542  (1046) ;  24, 1 1'.t    !«.» 17 1 ;  Bee! .  /.  anorg. 
Chem.,2i9,  321  (1942). 


532  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

[Fe(CN)6NH3]= 

and  that  treatment  with  alkali  yields  a  nitro  compound  [Fe(CN)5N02]4~, 
indicates  that  nitrogen  is  the  donor  atom51- 85. 

There  is  little  experimental  evidence  that  nitric  oxide  coordinates  as  a 
neutral  group.  As  an  odd  molecule  it  should  contribute  paramagnetism  to 
such  complexes  as  Fe(NO)2(CO)2  and  Co(NO)(CO)3 ,  but  these  are  dia- 
magnetic*8. Hiickel87  states  that  the  black  paramagnetic  form  of 

[Co(NH3)5NO]++ 

exemplifies  the  coordination  of  nitric  oxide  as  a  neutral  group.  Loose 
addition  compounds  such  as  Fe(NO)S04  may  be  of  the  same  type  but 
magnetochemical  evidence  is  lacking.  The  formation  of  the  unstable, 
paramagnetic  pentacyanonitrosyl  compounds,  M3[Fe(CN)6(NO)],  by  the 
reaction  [Fe(CN)6NH8]-  +  NO  -»  [Fe(CN)5(NO)]s  +  NH3  may  be  an 
example  of  coordination  of  the  nitrosyl  group  as  a  neutral  molecule,  al- 
though Sidgwick88  thinks  these  substances  are  true  nitroso  compounds. 

In  very  few  cases  is  there  any  indication  that  nitric  oxide  may  coordinate 
as  the  ion  NO-89.  The  only  simple  compound  containing  the  NO- group  is 
NaNO90.  Its  reactions  are  entirely  distinct  from  those  of  sodium  hypo- 
nitrite,  which  has  the  same  empirical  formula.  It  is  diamagnetic91,  as  would 
be  expected  if  it  contains  the  NO-  ion. 

The  pink  diamagnetic  form  of  [Co(NH3)5NO]++  is  believed  to  be  an 
example  of  a  case  in  which  NO-  is  present  and  plays  the  same  role  as 
CI"  in  [Co(NH3)5Cl]++92.  The  neutral  molecule  [Co(CO)3NO]  allows  a 
thorough  analysis  of  the  NO  coordination.  This  compound  is  monomeric, 
diamagnetic,  and  pyridine  does  not  replace  the  NO93.  Since  the  compound 
is  diamagnetic,  the  NO  group  does  not  function  as  a  neutral  molecule.  If 
X( )  were  functioning  as  a  negative  group,  corresponding  halides,  Co(CO)3X, 
would  be  expected,  but  these  are  not  known.  Finally,  these  compounds 
cannot  be  derivatives  of  hyponitrous  acid,  because  the  mononitrosyls  are 

85.  Emel^us  and  Anderson,   "Modern  Aspects  of  Inorganic  Chemistry,"  p.  414, 

New  York,  D.  Van  Nostrand  Company,  Inc.,  1938. 

86.  Reiff,  Z.  anorg.  allgem.  Chem.,  202,  375  (1931). 

87.  Hiickel,  "Structural  Chemistry  of  Inorganic  Compounds,"  translated  by  L.  H. 

Long,  Vol.  II,  p.  516.,  Amsterdam,  Elsevier  Publishing  Company,  1951;  Ray 
and  Hliar,  /.  Indian  Chem.  Soc.,  5,  499  (1928). 

88.  Sidgwick,  "Chemical  Elements  and  Their  Compounds,"  Vol.  II,  p.  1360,  London, 

Oxford  University  Press,  1950. 

89.  Cambi,  Z.  anorg.  Chem.,  247,  22  (1941);  Hieber  and  Nast,  Z.  anorg.  Chem.,  247, 

31  (1941). 
Zinil  and  Harder,  Ber.,  66B,  760  (1933). 

91.  Frazer  and  Long,  ./.  Chem.  Phye.,  6,  462  (1938). 

92.  Mellor  and  Craig,  J.  Proc.  Roy.  Soc,  N.  S.  Wales,  78,  25  (1944). 

93.  Hieber  and  Anderson,  Z.  anorg.  Chem.,  208,  238  (1932);  211,  132  (1932). 


M  E  T.  \L  CA  KBONYLS  AND  NITROSYLS  533 

not  dimers,  and  the  dinitrosylfl  do  not  correspond  to  the  balides.  The  sug- 
gestion has  also  been  made  that  nit  lie  oxide  functions  as  N( )"  in  a  complex 
cation  but  functions  as  NO+  in  a  complex  anion. 

It  is  well  established  that  nitric  oxide  can  coordinate  as  the  NO4"  ion. 
This  ion  is  isosteric  with  carbon  monoxide  and  with  cyanide  ion: 

:N=0:+  :C=0:  :C=N:" 

Isonitrile  complexes,  in  which  C=N — R  groups  replace  CO  groups  in 
carbonyl  structures,  have  been  prepared  (p.  92);  [Ni(CNCH3)3CO]  and 
[Co2(CXC6H5)5(CO)3]  are  examples94.  In  such  series  as  K4[Fe(CN)6], 
K3[Fe(CX)5CO],  K2[Fe(CX)5X0]  the  differences  in  the  charge  of  the  anion 
are  as  expected  if  a  cyanide  group  in  the  first  is  replaced  by  a  neutral 
carbonyl  group  in  the  second  or  by  a  positive  nitrosyl  group  in  the  third. 
That  the  last  compound,  potassium  nitroprusside,  represents  an  oxidation 
state  of  2+  for  iron  is  shown  by  its  diamagnetism  and  its  conversion  by 
alkali  to  K4[Fe(CN)5N02]. 

In  calculating  the  effective  atomic  number  of  the  central  atom  in  these 
nitrosyl  or  nitrosyl-carbonyl  complexes  one  must  assume  that  the  nitrosyl 
group  contributes  three  electrons  to  the  central  atom.  With  this  stipulation, 
the  effective  atomic  number  for  most  of  the  nitrosyls  is  that  of  an  inert  gas. 

However,  the  case  of  a  positive  group  (instead  of  a  neutral  or  negative 
group)  donating  an  electron  pair  to  a  metal  atom  or  ion  presents  a  difficulty 
in  that  a  certain  amount  of  negative  charge  is  imparted  to  the  metal 
:M~:X+: :  :0: .  Pauling  points  out  that  the  accumulation  of  such  negative 
charge  is  unlikely.  An  alternative  suggestion  is  that  the  metal  also  contrib- 
utes   two    electrons    for    the    combination,    producing    a    double    bond 

M:  :N+:  :0.  Hel'man95  has  suggested  that  nitric  oxide,  as  well  as  carbon 

monoxide  and  ethylene,  forms  bonds  of  this  type  with  platinum.  An  anal- 
ogy is  noted  between  [PtXOCl3]-  and  [PtC2H4Cl3]- 

Evidence  for  considerable  double  bond  character  also  comes  from  esti- 
mation of  bond  distances  by  electron  diffraction  methods96.  In  Co(NO)- 
(CO)3  and  Fe(X"0)2(CO)2  the  metal-nitrogen  bond  is  shorter  than  that 
calculated  for  a  single  bond,  and  the  nitrogen-oxygen  bond  distance  is  in- 
termediate between  those  for  N=0  and  X'=0.  (Table  16.6).  Both  of  the 
above  compounds,  like  nickel  carbonyl,  are  tetrahedral.  Xeither  the  contri- 
bution of  three  electrons  by  the  nitric  oxide  nor  the  possibility  of  double- 
bond   character  changes   the  structure.   Furthermore,   the  possibility  of 

94.  Hieber  and  Bockly,  Z.  anorg.  Che,,,.,  262,  344  (1950);  Hieber,  Z.  Natarforsch., 

56,  129  (1950). 

95.  BeTman,  Com pt.  rend.  ucad.  sci.  U.irS.S..  24,  549  l  L939). 

96.  Brockway  and  Anderson,  Trans.  Faraday  Soc,  33,  1233  (1937). 


534  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Table  16.6.  Bond  Distances  in  Nitrosyl-Carbonyl  Compounds 


M- 

-N  (A.) 

N— O  (A.) 

Obs. 

Calc. 

Obs. 

Co(CO)3NO 

1.76 

1.95 

1.10 

Calculated  for  N=0,  1.15  A 

Fe(CO)2(NO)2 

1.77 

1.93 

1.12 

Calculated  for  N^O,  1.05  A 

double-bond  character  does  not  disturb  the  effective  atomic  number 
relationship. 

Assuming  that  the  nitrosyl  group  contributes  three  electrons  to  the  metal 
atom,  and  the  hydrogen  atom  in  carbonyl  hydrides  contributes  one  elec- 
tron, one  notes  the  existence  of  isoelectronic  series: 

Fe(CO)2(NO)2  Co(CO)3NO  Ni(CO)4 

Fe(CO)4H2  Co(CO)4H  Ni(CO)4 

The  replacement  of  the  nucleus  2gNi  by  27C0  corresponds  electronically  to 
the  formation  of  the  ion  [Co(CO)4]~.  Replacement  of  one  CO  group  in  this 
with  one  NO+  group  forms  a  neutral  molecule.  The  process  may  also  be 
represented  by  NO  — »  NO+  +  e~,  the  metal  atom  gaining  the  additional 
electron.  The  acquisition  of  a  negative  charge  by  the  central  atom  makes 
it  understandable  that  only  a  limited  number  of  NO  molecules  can  be 
bound,  and  the  stability  of  such  compounds  decreases  in  the  order 
Ni(NO)Cl,  Co(NO)2Cl,  Fe(NO)3CF.  Ewens56  believes  the  structure  of 
[Fe(NO)2X]2  and  other  dimeric  forms  to  be: 

X 

•    \ 
(ON)2Fe Fe(NO)2 

\    / 
X 

Similar  postulations  have  been  made  by  Seel84d  concerning  the  Roussin 
salts  [(NO)2FeS]K,  [(NO)7Fe4S3]K-H20,  and  [(NO)2Fe-S-C2H5]2 .  It  is 
noteworthy  that  the  sum  of  the  atomic  number  and  the  maximum  number 
of  bonded  NO  molecules  has  the  constant  value  29  with  the  metals  of  the 
first  transition  series. 

Preparation  and  Properties  of  Nitrosyls 

Preparation  by  the  Action  of  Nitric  Oxide  Upon  Metallic  Salts 

The  familiar  brown  ring  test  for  nitrites  and  nitrates  is  based  on  the 
absorption  of  nitric  oxide  by  solutions  of  iron(II)  salts97.  The  reaction  is 

97.  Kohlschutter  and  Kutscheroff,  Ber.,  40,  873  (1907);  Kohlschutter  and  SazanofT, 
Ber.,  44,  1423  (1911);  Manchot,  Ber.,  47,  1601  (1914);  Manchot  and  Zechent- 
mayer,  Ann.,  350,  368  (1906). 


UBTAL  CARBONYLS  AND  NITROSYLS  535 

readily  reversed  by  heating,  nitric  oxide  being  evolved  and  the  iron (II) 
salt  recovered97**98.  It  is  difficult  to  isolate  solid  compounds,  especially 
since  most  solid  salts  do  not  absorb  nitric  oxide  extensively.  However,  such 
compounds  as  Fe(NO)HP04Wo  and  Fe(NO)Se< >,■  UT20"  have  been  isolated 

from  solution.  Such  solutions  may  be  red,  green  or  brown100.  More  than 
one  species  is  present,  as  shown  by  absorption  spectra  data101  and  trans- 
ference studies97*,  (which  indicate  that  the  complexes  may  be  cationic, 
anionic,  or  neutral).  An  iron(III)  derivative,  Fe2(NO)2(S04)3  ,  has  also 
been  reported978. 

Copper (II)  salt  solutions  in  the  presence  of  free  acid  absorb  a  molar 
quantity  of  nitric  oxide  to  form  deep  blue-violet  solutions9715 • 97c>  102.  Com- 
parable reactions  result  in  the  formation  of  palladium(II)  nitrosyl  deriva- 
tives, Pd(NO)jCli  and  Pd(XO)2S04103.  A  chromium(II)  nitrosyl  dithio- 
carbamate  can  be  prepared  by  treating  chromium(II)  acetate  with  alcoholic 
RjNCSjNa  (R  =  ethyl  or  propyl)  and  nitric  oxide  at  0°104. 

In  these  nitrosyl  compounds  (except  the  iron(III)  salt)  the  oxidation 
state  of  the  metal  is  presumably  2+,  but  there  is  no  confirmatory  experi- 
mental evidence.  Many  examples  are  known  of  the  formation  of  nitrosyl 
derivatives  of  normally  divalent  metals  in  the  univalent  state.  Iron(II) 
chloride  forms  the  derivative  Fe(NO)3Cl  when  treated  with  nitric  oxide 
in  the  presence  of  zinc105.  In  a  similar  manner  anhydrous  cobalt  halides 
form  Co(XO)2X106,  and  nickel  halides  form  Ni(XO)X105,  the  ease  of  forma- 
tion decreasing  in  the  orders  Fe  >  Co  >  XTi  and  I  >  Br  >  CI.  These 
compounds  are  characterized  by  thermal  instability,  coordinate  unsatura- 
tion,  and  extreme  reducing  ability.  Most  of  them  react  readily  with  such 
donors  as  pyridine  and  o-phenanthroline. 

The  number  of  combined  nitric  oxide  molecules  decreases  in  the  order 
Fe  >  Co  >  Ni.  Seel107  has  suggested  a  nitrosyl  displacement  series  com- 
parable with  Grimm's  hydride  displacement  series,  in  which  the  addition 
of  n  molecules  of  nitric  oxide  is  supposed  to  convert  a  metal  atom  into  a 
pseudo  atom  n  groups  to  the  right  in  the  Periodic  Table.  This  series  would 

98.  Manchot  and  Haunschild,  Z.  anorg.  allgem.  Chem.,  140,  22  (1924). 
Manchol  and  Linckh,  Z.  anorg.  allgem.  Chem.,  140,  37  (1924). 

100.  Manchol  and  Huttner,  Ann.,  372,  153  (1910). 

101.  Manchol  and  Linckh,  Bar.,  59B,  406  (1926);  Schlesinger  and  Salathe,  ./.  Am. 

§  ...  45,  L863  (1923). 
L02.  Manchot,  Ann.t  375,  308  (1910). 

103.  Manchot  and  Waldmuller,  Ber.,  59B,  2363  (1926). 

104.  Malatesta,  Gazz.  chim.  iud.,  70,  729,  734  (1940  , 

105.  Biebei  and  Nast,  Z.  anorg.  allgem.  Chem.,  244,  23  (1940). 

106.  Biebei  and  Marin,  Z.  anorg.  allgem.  Chem.,  240, 241    I 

107.  Seel,  Z  anorg.  allg\  249,  308    L942). 


536  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

contain  such  pseudo  atoms  as 

Fe  Co  Ni  Cu 

Fe(NO)  Co  (NO)  Ni(NO) 

Fe(NO)2  Co(NO)2 

Fe(NO)3 

These  monovalent  halides  correspond  to  copper(I)  halides.  In  order  to 
achieve  a  coordination  number  of  four,  the  compound  Fe(NO)3Cl  is  rep- 
resented as  monomeric  (analogous  to  Cu(NH3)3I),  the  compounds  Co(NO)2X 
(analogous  to  [Cu(PEt3)2I]2)  as  dimeric,  and  the  compounds  Ni(NO)X 
(analogous  to[CuAsEt3I]4)  astetrameric.  Other  known  compounds  fitting  in- 
to the  series  are  Fe(NO)I,  Fe(NO)2I  (see  p.  538)  and  Co(NO)I,  which  is 
known  only  in  addition  compounds  such  as  Co(NO)I-6py106. 

A  number  of  nitrosyl  thio  compounds  are  known,  but  further  work  is 
necessary  to  establish  their  structures.  The  best  known  of  these  compounds 
are  the  so-called  red  and  black  salts  of  Roussin,  who  first  prepared  them 
in  1857.  Upon  treatment  with  Fe(NO)S04 ,  the  red  salts,  M^FefNO^S], 
are  converted  to  the  more  stable  black  salts,  MI[Fe4(NO)7S3],  which  may 
be  reconverted  to  the  red  salts  by  the  action  of  alkali108 

3Na[Fe(NO)2Sl  +  Fe(NO)S04  -»  Na[Fe4(NO)7S3]  +  Na2S04 

2Na[Fe4(NO)7S3]  +  4NaOH  -*  6Na[Fe(NO)2S]  +  Fe203  +  N20  +  2H20 

According  to  Seel's  scheme  the  red  compounds  must  be  dimeric  and  Ewens56 
reported  that  they  have  the  same  structure  as  Fe2(CO)9  with  a  direct  link 
between  iron  atoms. 

X 

•    \ 
(ON)2Fe Fe(NO)2 

\    / 
X 

Whereas  iron  forms  the  series  Fe(NO)2SA,  cobalt  and  nickel  form  the 
series  Co(NO)2(SA)2  and  Ni(NO)(SA)2 .  Thiosulfate  derivatives, 

K3[Co(NO)2(S203)2] 

and  K3[Ni(NO)(S203)2],  have  been  prepared  by  the  action  of  nitric  oxide 
and  potassium  thiosulfate  upon  solutions  of  cobalt (II)  acetate  and  nickel 
(II)  acetate109.  Ethyl  mercaptan  derivatives  have  also  been  prepared"  •  n0 

108.  Marchlewski  and  Sachs,  Z.  anorg.  Chem.,  2,  175  (1892);  Hofman  and  Wiede,  Z. 

anorg.  Chem.,  9,  295  (1895). 
L09.  Manchot,  Ber.,  59B,  2445  (1926). 
110.  Manchot  and  Kaess,  Ber.,  60B,  2175  (1927). 


METAL  CARBONYLS  AND  NITROSYLS  537 

Co(SR)2  +  3N0  ->  Co(NO),SR  +  NOSR 
Ni(SR)2  +  2NO  -»  Ni(NO)SR  +  NOSR 

The  exact  structure  of  the  [Fe4(NO)7S3]~  ion  has  not  been  determined.  It 
is  believed  that  each  iron  is  tet incoordinate,  with  sulfur  atoms  acting  as 
bridging  groups;  nitrosy]  groups  occupy  the  remaining  positions84d. 

Solutions  of  cobalt(ll)  salts  containing  ammonia  absorb  nitric  oxide  to 
form  the  complex  ion  [Co(NH3)6(NO)]++  m.  Such  compounds  exist  in 
isomeric  forms.  The  black  compounds  (of  which  only  the  chloride  and  iodate 
have  been  reported)  are  unstable  and  paramagnetic.  The  pink  compounds 
are  diamagnetic  and  do  not  evolve  nitric  oxide  upon  treatment  with  acids. 
The  pink  compounds  probably  contain  the  NO-  group  whereas  the  black 
compounds  may  contain  cobalt  in  the  divalent  state  with  nitric  oxide 
coordinating  as  a  neutral  group. 

Treatment  of  saturated  ammonium  or  potassium  tetrachloroplatinate- 
(II)  solution  with  nitric  oxide  yields  a  green  solution  from  which  [Pt(NH3)4] 
[Pt(XO)Cl3]  is  precipitated  by  a  solution  of  tetrammineplatinum(II) 
chloride112.  The  addition  of  pyridine  to  the  green  solution  precipitates 
fran*-[Pt(NO)pyCl2].  The  nitric  oxide  group  therefore  appears  to  be  trans 
directing.  Such  compounds  show  a  marked  resemblance  to  the  correspond- 
ing ethylene  and  carbonyl  compounds. 

Pentacyanoiron(II)  complexes  such  as  Na3[Fe(CN)5NH3]  react  with 
nitric  oxide  to  form  Na3[Fe(CN)5(XO)]113.  This  is  one  of  the  few  cases  in 
which  nitric  oxide  replaces  a  neutral  group  without  change  of  charge.  Such 
compounds  are  entirely  distinct  from  the  nitroprussides.  They  are  dark 
yellow  in  neutral  solution  but  violet  in  acid  solution.  Baudisch114  reports 
that  such  complexes  also  result  from  the  action  of  light  upon  a  nitroprus- 
side,  the  nitrosyl  group  being  activated.  Thus,  sodium  nitroprusside,  in 
the  presence  of  light  and  hydrogen  peroxide,  is  able  to  convert  benzene 
into  o-nitrosophenol.  Light  also  catalyzes  the  reaction  of  sodium  nitro- 
prusside with  cupferron,  with  thiourea,  and  with  a  mixture  of  hydrogen 
peroxide  and  sodium  azide. 

Preparation  by  the  Action  of  Nitric  Oxide  upon  Carbonyls  or  Re- 
lated Compounds 

The  nitrosyl  carbonyls  of  cobalt  and  iron  are  generally  prepared  by  the 
action  of  nitric  oxide  upon  the  carbonyls.  The  cobalt  compound,  Co(XO)- 

111.  Sand  and  Genssler,  Ber.,  36,  2083  (1903) ;  Werner  and  Karrer,  Helv.  Chim.  Acta, 

1,  54  (1918). 

112.  Hel'man  and  Maximova,  Compt.  rend.  acad.  sci.  U.R.S.S.,  24,  549  (1939). 

113.  Manchot,  Merry,  and  Woringer,  Bcr.,  45,  2869  (1912) 
111.  Baudisch,  Science,  108,  443  (1948). 


538  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

(C0)3 ,  was  first  obtained  by  Mond  and  Wallis115  by  reaction  of  dry  nitric 
oxide  with  cobalt  tetracarbonyl.  It  has  also  been  prepared  by  treating 
alkaline  suspensions  of  cobalt(II)  cyanide  with  carbon  monoxide,  followed 
by  saturation  with  nitric  oxide22c- 96-  116.  This  probably  involves  the  inter- 
mediate formation  of  cobalt  carbonyl  hydride  or  its  salt.  The  nitrosyl- 
carbonyl  is  a  yellow  gas  which  condenses  to  a  red  liquid.  The  iron  compound, 
Fe(NO)2(CO)2 ,  has  been  obtained  in  similar  manner  by  the  action  of  nitric 
oxide  upon  iron  tetracarbonyl93a. 

Reactions  of  the  nitrosyl-carbonyls  indicate  that  the  nitrosyl  group  is 
more  tightly  bound  than  the  carbonyl  group.  Treatment  of  iron  nitrosyl  - 
carbonyl  with  pyridine  (py)  and  with  o-phenanthroline  (o-phen)  produces 
[Fe2(NO)4(py)3]  and  [Fe(NO)2(o-phen)],  respectively,  and  treatment  of 
cobalt  nitrosyl-carbonyl  in  the  same  way  yields  [Co2(NO)2(CO)(py)2]  and 
[Co(NO) (CO) (o-phen)].  Further  evidence  is  given  by  the  formation  of 
Fe(NO)2I  from  iron  nitrosyl-carbonyl  and  iodine93b. 

Other  nitrosyls  have  been  prepared  from  carbonyls.  Nitric  oxide  reacts 
with  iron  pentacarbonyl  under  pressure  to  form  the  interesting  compound, 
iron  tetranitrosyl,  Fe(NO)4117.  This  black  crystalline  substance  is  converted 
into  Fe(NO)3NH3  by  liquid  ammonia,  into  Fe(NO)S04  by  dilute  sulfuric 
acid,  into  K[Fe(NO)2S203]  by  potassium  thiosulfate,  and  into 

K[Fe4(NO)7S3] 

by  potassium  bisulfide.  Manchot  and  Manchot9  have  reported  that  a  similar 
reaction  with  ruthenium  enneacarbonyl  produces  a  pentanitrosyl, 
Ru(NO)o ,  as  well  as  a  tetranitrosyl,  Ru(NO)4  .  These  results  have  been 
questioned  by  Emeleus  and  Anderson69. 

Nickel  carbonyl  reacts  with  nitric  oxide  in  the  presence  of  a  trace  of 
moisture  to  form  a  water-soluble  nitrosyl-hydroxide,  Ni(NO)OH.  This 
blue  basic  compound  shows  the  reducing  power  expected  for  univalent 
nickel118. 

Carbonyl  derivatives  also  react  with  nitric  oxide  in  some  cases.  An 
unusual  nitrosyl  iodide,  Fe2(NO)4I3 ,  results  from  the  treatment  of  the 
tetracarbonyl  iodide,  Fe(CO)4I2 ,  with  nitric  oxide.  This  compound  is 
presumed  to  contain  both  univalent  and  divalent  iron67b. 

Miscellaneous  Methods  of  Preparation 

Nitrosyls  may  be  prepared  by  reactions  involving  the  oxidation  or  re- 
duction of  some  nitrogen  compound  other  than  nitric  oxide.  The  nitro 

115.  Mond  and  Wallis,  /.  Chem.  Soc,  121,  32  (1922). 

116.  Blanchard  and  Gilmdnt,  J.  Am.  Chen,.  Soc,  62,  1192  (1940). 

117.  Manchol  and  Enk,  Ann.,  470,  275  (1929), 

118.  Anderson,  Z.  anorg.  allgem.  Chem.,  229,  357  (1936). 


METAL  CARBONYLS  AND  NITR0S1  Lfl  539 

prussides,  Mi[Fe  CN  sNO],  were  firsl  prepared"9  by  the  action  of  30% 

nitric  acid  upon  a  fcrrocyanidc  or  iVrricyanidc,  a  complicated  and  violent 

reaction  which  is  still  used  for  their  preparation.  Another  method  involves 
the  action  of  nitrite  ion  upon  ferrocyanide  ion1*-0. 

,1  .   (\  •    NO,    .  ■  [Fe(CN)»N0    ;     ;   CN 

[Fe(CN)»NO,]*-  +  Ho()  ^  [Fe(CN)5NO]-  +  2<  >ll 

These  reactions  are  reversible,  hut  may  be  brought  to  completion  by  adding 
acid  to  combine  with  t he  cyanide  ion  or  hydroxide  ion.  The  corresponding 
ruthenium  compound,  Iv2[Rii(CN)5(XO)]-2H20,  has  been  prepared  by  the 
action  of  nitric  acid  upon  the  ruthenocyanide,  K4[Rii(CN)6]121,  and  the 
manganese  compound,  K3[Mn(CX)5(XO)],  by  the  action  of  nitric  oxide 
upon  manganous  salts  in  the  presence  of  cyanide  ion122. 

The  nitroprussides  develop  intense  violet  colorations  when  treated  with 
alkali  sulfides  (Gmelin  reaction)  but  not  with  hydrogen  sulfide123.  Intense 
red  colorations  with  alkali  sulfites  (Bodecker  reaction)  are  due  perhaps  to 
the  formation  of  [Fe(CX)5(XOS03)]4-  124.  The  insolubility  of  mercury(II) 
nitroprusside  has  been  suggested  as  a  basis  for  the  quantitative  determina- 
tion of  the  radical125.  Recent  work126  has  confirmed  the  dipositive  state  of 
iron  in  the  nitroprussides  and  has  indicated  that  one  cyanide  group  is 
attached  to  iron  through  nitrogen  and  the  other  four  through  carbon. 

Osmium  nitrosyl  compounds  K2[OsCl5(NO)]  and  K2[OsBr5(XO)],  result 
when  the  hexanitro  compound,  K2[Os(N02)6L  is  heated  with  hydrochloric 
or  hydrobromic  acid127.  The  ruthenium  compound,  K2[RuCloXO],  is  ob- 
tained when  metallic  ruthenium  is  dissolved  in  a  molten  mixture  of  potas- 
sium hydroxide  and  potassium  nitrate  or  nitrite  and  the  resulting  mass 
treated  with  hydrochloric  acid128. 

Hydroxylamine  can  be  used  to  introduce  a  nitrosyl  group  into  a  com- 
plex129. The  nickel  compound  K2[Xi(CX)3(NO)]  has  also  been  prepared  by 

119.  Playfair,  Phil.  Mag.,  [3]  36,  197  (1850);  Ann.,  (Liebig's),  74,  317  (1850). 

120.  Shwarzkopf,  Abhandl.  deut.  Xatunv.  Med.  Ver.  Bohmen,  3,  1  (1911). 

121.  Manchot  and  Dusing,  Ber.,  63B,  1226  (1930). 

122.  Blanchard  and  Magnusson,  ./.  Am.  Chem.  Soc,  63,  2236  (1941);  Manchot  and 

Schmid,  Ber.,  59B,  2360  (1926). 

123.  Sas,  A  miles  soc.  espan.  fis.  quint.,  34,  419  (1936);  Scagliarini,  Atti  congr.  naz. 

ckim.  pura  applicaia  4th  Cong.,  1933,  597  (1932). 

124.  Scagliarini,  Atti  accad.  Lined,  22,  155  (1935);  Morgan  and  Burst  all,  "Inorganic 

survey  of  Modern  Developments,"  p.  364,  Cambridge,  W.  1 
and  Sons,  Ltd.,  1936. 

125.  Tomicek  and  Kubik,  Collection  Czechoslov.  CI  rnun.,  9,  377  (19 

126.  Sas,  Analesfis.  quim.  (Spain),  39,  55  (1943). 

127.  Wintrebert,  Ann.  chim.  phys.,  [7]  28,  15  (1903). 

128.  Joly,  Compt.  rend.,  107,  994  (1888). 

129.  EBeber,  Nasi  and  Gehring,  Z  anorg.  allgem.  Chem.t  256,  150,  169  (1948);  Bieber 

and  N  .-•    /    Xnturforsch.,  2b,  321  (1947). 


540  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

the  action  of  nitric  oxide  upon  the  complex  cyanide  K2Ni(CN)3  in  liquid 
ammonia  or  absolute  alcohol130. 

Industrial  Significance  of  Metal  Carbonyls 
The  Metallurgy  of  Nickel 

The  discovery  that  nickel  readily  forms  a  volatile  carbonyl  was  utilized 
by  Ludwig  Mond131  for  the  separation  of  nickel  from  ores  containing  cobalt 
and  other  metals.  He  built  an  experimental  plant  for  separating  nickel 
from  Canadian  matte.  The  plant  was  torn  down  and  rebuilt  several  times, 
but  within  five  years  from  the  discovery  of  nickel  carbonyl  it  was  success- 
fully producing  1.5  tons  of  nickel  per  week. 

For  the  Mond  process,  the  ore  is  heated  with  coke  and  limestone  with 
the  result  that  some  of  the  iron  sulfide  is  converted  to  oxide.  The  matte 
is  further  concentrated  in  a  Bessemer  converter  until  it  contains  about 
80  per  cent  nickel  and  copper.  The  finely  ground  matte  is  subject  to  calci- 
nation at  700  to  800°C  and  extracted  with  dilute  sulfuric  acid,  which  dis- 
solves most  of  the  copper  oxide  but  attacks  the  nickel  oxide  only  slightly. 
The  nickel  oxide  is  then  led  through  a  series  of  reducers  and  volatilizers. 
The  reducing  agent  is  water  gas  at  330  to  350°C ;  97  per  cent  of  the  reduc- 
tion results  from  the  hydrogen,  while  the  carbon  monoxide  acts  upon  the 
metallic  nickel  in  the  volatilizer  at  a  temperature  of  50°C  to  form  the  car- 
bonyl. 

The  gases  from  the  volatilizers  are  passed  into  decomposers,  where  they 
come  into  contact  with  nickel  pellets  at  180°C,  whereupon  the  carbonyl 
is  decomposed  and  the  nickel  deposits  on  the  pellets.  From  time  to  time 
the  pellets  are  sorted,  the  smaller  ones  being  returned  to  the  decomposers. 

Carbonyls  as  Antiknock  Agents 

Antidetonants,  or  antiknocks,  are  now  added  to  most  gasolines.  The 
most  widely  used  antiknock  agent  is  lead  tetraethyl ;  however,  the  carbonyls 
of  iron,  cobalt,  and  nickel  have  been  found  to  be  almost  as  effective.  The 
substitution  of  a  carbonyl  for  lead  tetraethyl  may  result  in  a  considerable 
increase  in  maximum  power  output.  In  one  process  the  carbonyl  is  heated 
with  an  unsaturated  hydrocarbon,  such  as  butadiene,  and  the  resulting 
complex  is  added  to  the  gasoline132. 

Iron  pentacarbonyl  has  been  most  often  suggested  as  a  replacement  for 

130.  Hieber,  Nast  and  Proeschel,  Z.  anorg.  allgem.  Chem.,  256,  145  (1948). 

131.  Trout,  J.  Chem.  Ed.,  15,  113  (1938);  Mond,  J.  Soc.  Chem.  Ind.,  T49,  271,  283,  287 

(1930). 

132.  Johnson  (to  Texaco  Development  Corp.),  U.  S.  Patent  2406544  (Aug.  27,  1946) 

cf.  Chem.  Abs.,  41,  266  (1947);  Veltman  (to  Texaco  Development  Corp.),  U.  S. 
Patent  2409167  (Oct.  8,  1946)  cf.  Chem.  Abs.,  41,  595  (1947). 


METAL  CARB0NYL8  AND  NITR08YLS  541 

lead  tetraethyl.  Although  iron  carbony]  is  poisonous,  it  probably  docs  not 
have  the  cumulative  effect  that  is  associated  with  lead  compounds  and  the 

products  of  its  combustion  arc  less  toxic.  It  is  soluble  in  all  proportions  in 
gasoline  and  vaporizes  readily  in  the  carburetor.  There  are,  however,  two 
serious  disadvantages  in  the  use  of  iron  pentacarbonyl.  Iron(III)  oxide 
produced  by  combustion  tends  to  foul  the  combustion  chamber  and  its 
decomposition  to  Fe2(CO)9  is  light  catalyzed.  Lead  tetraethyl  alone  also 
fouls  the  combustion  chamber,  but  the  addition  of  small  amounts  of  ethyl- 
ene dibromide  prevents  lead  oxide  from  accumulating.  The  decomposition 
of  iron  pentacarbonyl  is  not  a  serious  problem,  since  a  number  of  stabilizers 
are  known133.  In  alcohol  fuels,  iron  pentacarbonyl  is  a  good  antiknock  agent, 
while  lead  tetraethyl  is  said  to  have  a  negative  effect  and  actually  depresses 
the  octane  rating134. 

King135  describes  experiments  to  show  that  the  oxidation  of  hydrocarbons 
in  the  presence  of  iron  carbonyl  is  a  heterogeneous  reaction  on  the  surface 
of  iron  which  results  from  decomposition  of  the  carbonyl.  The  fuel  is  there- 
fore partly  oxidized  to  carbon  dioxide  and  steam  prior  to  ignition.  The 
consequent  dilution  of  the  fuel  causes  a  reduction  of  inflammability  which 
is  sufficient  to  prevent  the  completion  of  combustion  by  detonation. 

The  Preparation  of  "Carbonyl  Metals" 

Nickel  produced  by  the  decomposition  of  the  carbonyl  is  remarkably 
pure,  and  Mond131d  suggested  that  nickel  carbonyl  may  be  used  for  the 
deposition  of  metallic  mirrors  (as  in  the  preparation  of  Dewar  flasks)  or  to 
build  up  nickel  articles  by  decomposing  the  carbonyl  in  contact  with  a 
suitably  shaped  mold.  Carbonyl  nickel  has  been  used  as  a  hydrogenation 
catalyst136. 

In  similar  maimer,  iron  pentacarbonyl  has  been  used  to  prepare  metallic 
iron.  By  varying  the  conditions,  it  is  possible  to  prepare  iron  as  scales, 
grains,  sponge,  or  powder.  "Carbonyl  iron"  is  remarkably  free  of  impurities 
except  for  small  amounts  of  carbon  and  oxygen.  Its  grains  are  nearly 
spherical  and  quite  uniform  in  size.  When  the  powder  is  subjected  to 
mechanical  pressure  in  hydrogen  or  in  vacuum  at  a  temperature  below  its 
melting  point,  it  may  be  compressed  into  a  solid  without  pores.  Most  of  the 
carbon  and  oxygen  are  driven  off  as  carbon  monoxide  and  carbon  dioxide, 
leaving  a  pure,  fresh  iron  surface  which  sinters  readily.  The  iron  thus  pre- 
pared is  soft,  ductile,  and  resistant  to  corrosion.  The  chief  use  of  carbonyl 
iron  is  in  the  making  of  magnetic  cores  for  electronic  equipment.  It  is  ex- 

133.  Leahy,  Refiner  Natural  Gasoline  Mfr.,  14,  82  (1935). 

134.  Pitesky  and  Wiebe,  Ind.  Eng.  Ckem.,  37,  577  (1045). 

135.  King,  Canadian  J.  Research,  26F,  125  (1946). 

136.  Shukoff,  German  Patent  241823  (Jan.  18,  1910),  cf.  Chem.  Abs.,  6,  2146  (1912 


542  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

(•client  for  that  purpose  because  of  its  uniform  particle  size  and  shape  as 
well  as  its  purity. 

Nickel-iron  alloys  and  cobalt-molybdenum  alloys  have  been  prepared 
by  the  sintering  of  powders  obtained  from  the  decomposition  of  the  respec- 
tive  carbonyls.  These  alloys  have  electromagnetic  properties  which  com- 
pare favorably  with  alloys  prepared  by  other  methods. 

Carbonyl  iron  has  been  entered  in  The  National  Formulary  as  a  substi- 
tute for  iron  reduced  by  hydrogen137. 

Preparation  of  Oxides 

Very  finely  divided  iron  oxide  may  be  obtained  by  heating  iron  carbonyl 
below  100°C  under  carefully  controlled  conditions.  Catalysts  may  be  used 
to  accelerate  the  formation  of  the  oxide.  This  oxide  is  suitable  for  use  as  a 
coloring  agent,  polishing  powder,  or  decarbonizing  agent  for  cast  iron  or 
steel138. 

Carbonyls  in  Synthesis 

Much  work  has  been  done  on  the  use  of  carbonyls  of  iron,  cobalt,  and 
nickel  as  catalysts,  particularly  when  carbon  monoxide  is  a  reactant.  In 
some  of  these  reactions  the  carbonyl  functions  as  a  homogeneous  catalyst. 
In  others  the  carbonyls  are  added  in  stoichiometric  amounts  and  may  or 
may  not  be  regenerated  in  the  course  of  subsequent  reactions. 

Reppe139  has  carried  out  carboxylation  reactions  with  acetylene  or  ethyl- 
ene at  high  pressure  for  the  preparation  of  various  types  of  organic  com- 
pounds. Some  typical  reactions  are 

1.  (a)  Preparation  of  acrylic  acid  from  acetylene: 

Ni(CO)4  +  4C2H2  +  2HC1  +  4H20  -+  4CH2=CHCOOH  +  NiCl2  +  H2 
(b)  Regeneration  of  the  carbonyl: 

NiCla  +  2XH3  +  H20  +  5CO  -+  Ni(CO)4  +  2NH4C1  +  C02 
(Cobalt  carbonyl  can  also  be  used  in  this  reaction,  but  iron  carbonyl  cannot.) 

2.  (a)  Preparation  of  n-propyl  alcohol  from  iron  carbonyl  hydride: 

Fe(CO)4H2  +  2C2H4  +  4H20  ->  2CH3CH2CH2OH  +  Fe(HC03)2 
(b)  Preparation  of  the  carbonyl  hydride: 

Fe(CO)5  +  H20  ->  Fe(CO)4H,  +  C02 

3.  Preparation  of  hydroquinone  from  acetylene  (in  the  presence  of  iron  carbonyl 

hydride  or  cobalt  carbonyl  hydride) : 

2C2H2  +  3CO  +  H20  ->  C6H4(OH)2  +  C02 

Reppe77  has  also  used  carbonyls  for  the  polymerization  of  acetylene  to 

137.  Bull.  Nat.  Formulary  Comm.,  18,  87  (1950). 

138.  Ehrmann,  Rev.  chim.  ind.,  44,  10  (1935). 

L39.  Reppe,  Modern  Plastics,  23,  162  (1945);  U.  S.  Dept.  of  Commerce  OTS  PB1112 
(Jan.  25, 1946) ;  Bigelow,  Chem.  Eng.  News,  25, 1038  (1947) ;  Hanford  and  Fuller, 
Ind.  Eng.  Chew.,  40,  1171  (1948). 


METAL  I  ARBONYLS  AND  NITROSYLS  543 

benzene  and  the  polymerization  of  vinyl  compounds  to  the  corresponding 
trimers.  Possible  catalysts  are  of  the  types  (1)  (R;iP),MX2,  (2)  (RgP)Ni- 

(C0)3 ,  and  (3)  (R3P)2Ni(CO)2  (R  is  an  alky]  or  aryl  radical;  iron  or  cobalt 
may  be  substituted  tor  nickel).  Types  (2)  and  (3)  are  made  by  the  action 
of  the  carbonyl  upon  one  or  two  moles  of  R3P,  or  the  action  of  the  carbonyl 
upon  compounds  of  type  (1).  The  catalysts  arc  first  treated  with  acetylene 
under  pressure  at  100-120°C,  and  the  polymerization  of  acetylene  is 
carried  out  at  a  temperature  of  6O-70°C.  The  polymerization  of  acetylene 
to  cycloctatetraene  (which  was  accomplished  by  Reppe,  using  a  catalyst 
of  nickel  cyanide)  has  been  carried  out  by  Cech140  using  nickel  carbonyl  in 
tetrahydrofuran  at  60-70°C. 

According  to  Lopez-Rubio  and  Pacheco141,  the  activity  of  iron,  cobalt, 
and  nickel  in  the  Fischer-Tropsch  hydrocarbon  synthesis  is  due  to  the  for- 
mation of  carbonyls  as  intermediates.  They  postulate  such  reactions  as 

20  CO  +  4Fe->  [Fe(CO)5]4 
[Fe(CO)5]4  +  33H2  ->  2C8H18  +  15H20  +  C02  +  3CO  +  4Fe 

The  so-called  Oxo  Process142  involves  the  addition  of  carbon  monoxide 
and  hydrogen  to  olefins  in  the  presence  of  solid  catalysts  (e.g.,  metallic 
cobalt)  to  produce  aldehydes.  Adkins  and  Krsek143  came  to  the  conclusion 
that  the  real  catalyst  is  cobalt  carbonyl.  They  found,  in  fact,  that  the  re- 
action proceeded  more  rapidly  and  at  a  lower  temperature  with  dicobalt 
octacarbonyl  as  a  catalyst  than  with  a  solid  catalyst.  The  reactions  they 
propose  (with  ethylene)  are 

2Co  +  8CO  -^  [Co(CO)4]2 

[Co(CO)4]2  +  H2  -*  2Co(CO)4H 

4Co(CO)4H  +  4C2H4  +  2H2  ->  4CH3CH2CHO  +  [Co(CO)3]4 

[Co(CO)3]4  +  4CO  -♦  2[Co(CO)4]2 

The  reaction  has  been  extended  to  produce  compounds  other  than  alde- 
hydes by  the  use  of  water  or  alcohols  instead  of  hydrogen.  Du  Pont,  Pig- 
anion,  and  Vialle144  consider  that  the  carbonyl  first  reacts  with  an  active 
compound  API  (H2 ,  H20,  ROH,  etc.)  to  form  a  complex,  which  reacts  with 
the  olefin  in  the  presence  of  carbon  monoxide  to  regenerate  the  metal 
carbonyl  and  give  the  corresponding  organic  carbonyl  derivative.   For 

MD.  Cech, Chi  "   i  .    Prague),**  '    L948 

141.  Lopes-Rubio  and  Pacheco,  Ion,  8,  86    L948 

142.  Roelen,  I  .  B.  Patenl  2327066  (Aug.  17,  1943  ;  cf.  Chem.  Aba.,  38,  550  (1944  . 

143.  Adkins  and  Krsek,  ./.  Am.  Chem.  Soc,  70,  383  (1948);  71,  :io:>l    l'.i49). 

144.  Du  Pont,  Piganion,  and  Vialle,  Bull.  soc.  ehim.,  France,  1948,  5' 


544  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

example,  with  nickel  carbonyl: 

Ni(CO)4  +  AH  ->  (CO)3Ni— C=0 

I        I 
H      A 

(CO)3Ni—  C=0  +  RCH=CHR  +  CO  ->  Ni(CO)4  +  RCHCH2R 

I         I  I 

H      A  AC=0 

Sternberg  and  his  co-workers145  have  used  cobalt  carbonyl  as  a  catalyst 
for  the  conversion  of  dimethylamine  to  dimethylf ormamide : 

(1)  3[Co(CO)4]2  +  20(CH3)2NH  ->  2[Co{(CH3)2NH}6]++ 

+  4[Co(CO)4]-  +  8HCON(CH3)2 

(2)  2[Co{(CH3)2NH}6]++  +  4[Co(CO)4]-  +  8CO  ->  3[Co(CO)4l2   +    12(CH3)2NH 

The  Presence  of  Carbonyls  in  Industrial  Gases 

Since  carbonyls,  particularly  those  of  nickel  and  iron,  may  be  formed 
when  gases  containing  carbon  monoxide  are  brought  into  contact  with  the 
metal,  they  may  be  present  as  adulterants  in  industrial  gases.  The  forma- 
tion of  iron  carbonyl  in  this  way  is  of  some  significance  in  dealing  with 
artificial  gases.  The  carbonyl  is  not  formed  during  the  manufacture  of  the 
gases  but  only  at  temperatures  below  250°C  in  purifying  boxes,  distributing 
pipes  and  gas  meters.  Mittasch146  found  almost  500  ml.  of  liquid  iron  pen- 
tacarbonyl  in  an  iron  tank  containing  illuminating  gas.  The  carbonyl  has 
also  been  found  in  tanks  of  hydrogen  which  contains  carbon  monoxide  as 
an  impurity147. 

Blueprints 

The  instability  of  iron  pentacarbonyl  toward  light  has  been  used  for 
the  preparation  of  blueprints148.  Paper  is  soaked  in  iron  pentacarbonyl  in 
the  dark.  After  exposure  to  light  and  washing  with  water,  the  exposed 
part  has  a  brown  deposit  of  Fe2(CO)9  .  This  is  converted  to  Prussian  blue 
by  an  acid  solution  of  potassium  ferrocyanide. 

The  Physiological  Action  of  Metal  Carbonyls 

The  increasing  use  of  metallic  carbonyls  makes  it  imperative  that  investi- 
gators realize  their  poisonous  nature149.  The  highly  volatile  nickel  carbonyl 
is  particularly  hazardous,   but  any  volatile  carbonyl   is  dangerous.  The 

145.  Sternberg,  Wender,  Friedel,  and  Urchin,  J.  Am.  Chem.  Soc,  76,  3148  (1953). 

1  W>.  Mittasch,  Z.  angew.  Chem.,  41,  831  (1928). 

1  17.    King  and  Sutchliffe,  ./.  Soc.  Chem.  hid.,  T47,  356  (1928). 

us.  Frankenburger,  German  Patenl  416996  (1924). 

1 19.  Trout,  ./.  Chem.  Educ,  15,  77  (1938). 


METAL  CARBONYLS  AND  NITROSYLS  545 

danger  with  nickel  carbony]  may  be  emphasized  by  the  example  of  the 

chemist,  who,  in  the  process  of  pouring  nickel  carbonyl  from  one  container 
to  another,  inhaled  enough  to  cause  his  death150. 

Although  a  study  of  the  toxicology  of  nickel  carbonyl  was  made  as  early 
as  1890  by  McKendrick  and  Snodgrass151  and  precautions  were  taken  by 
the  Mond  Nickel  Company  to  avoid  poisoning  of  its  employees,  an  accident 
took  place  in  which  ten  men  were  poisoned,  two  of  them  fatally.  Immedi- 
ately, Armit152  was  employed  to  study  the  problem  anew,  and  his  sugges- 
tions have  enabled  the  company  to  reduce  the  danger. 

The  assumption  that  metallic  carbonyls  are  poisonous  because  of  the 
carbon  monoxide  they  produce  upon  decomposition  is  not  valid.  Nickel 
carbonyl  is  at  least  five  times  as  deadly  as  carbon  monoxide.  Armit  found 
that  a  rabbit  is  killed  by  an  exposure  of  one  hour  to  air  containing  0.018 
per  cent  by  volume  of  the  carbonyl.  On  the  other  hand,  he  has  shown  that 
a  rabbit  would  not  absorb  harmful  amounts  of  cobalt  carbonyl  in  the  course 
of  two  hours'  exposure  even  if  the  atmosphere  were  saturated  with  this 
carbonyl153.  This  is  not  to  say,  however,  that  continued  exposure  to  cobalt 
carbonyl  would  not  be  injurious. 

Immediately  after  being  exposed  to  the  fumes  of  nickel  carbonyl,  a 
person  has  a  sensation  of  giddiness,  a  throbbing  headache,  and  nausea, 
sometimes  with  vomiting154.  If  the  carbonyl  is  mixed  with  carbon  monoxide, 
unconsciousness  may  result.  If  the  amount  of  carbonyl  in  the  air  is  very 
small,  exposure  of  the  person  for  some  time  may  result  only  in  a  throbbing 
headache.  These  symptoms  may  disappear  rather  quickly.  This  period, 
however,  is  frequently  followed  by  such  symptoms  as  difficult  breathing, 
pain  in  the  chest,  and  cyanosis.  The  skin  may  be  pale,  the  forehead  cold 
and  clammy,  and  the  general  expression  one  of  anxiety.  A  trace  of  nickel 
may  be  found  in  the  urine,  and  the  blood  may  show  the  presence  of  car- 
boxy  hemoglobin.  Post  mortem  examinations  of  fatal  cases  show  that  tissues 
of  the  lungs  and  brain  are  severely  damaged. 

The  treatment  depends  upon  the  severity  and  the  presence  or  absence  of 
poisoning  by  carbon  monoxide.  The  patient  must  be  kept  warm  and  should, 
if  necessary,  be  given  stimulants  to  aid  respiration  and  heart  action.  Abso- 
lute rest  is  necessary  to  relieve  the  heart  and  lungs  of  undue  strain.  The 
effects  of  the  poisoning  are  not  chronic;  persons  who  have  received  non- 
fatal doses  have  shown  complete  recovery. 

Persons  working  with  carbonyls  must  use  the  same  precautions  which 

150.  Brandes,  ./.  Am.  Med.  Assoc,  102,  1204  (1934). 

151.  McKendrick  and  Snodgrass,  proc.  Phil.  Soc,  Glasgow,  22,  204  (1890-91). 

152.  Annit .  ./.  Hyg.,  7,  526    L907   ;  8.  665    1908). 

153.  Armit.  ./    Hyg.,  9,  249    1909). 

154.  Amor,  ./.  Ind.  Hyg.,  14,  216    L932). 


546  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

are  used  for  any  deadly  gas  or  vapor.  A  well-ventilated  hood  must  be 
used  for  all  experiments.  The  compounds  must  be  kept  in  strong  glass  or 
steel  containers,  preferably  under  carbon  dioxide  or  nitrogen.  Continual 
teste  for  leaks  should  be  made.  One  part  of  nickel  carbonyl  in  80,000  parts 
of  air  may  be  detected  by  the  luminosity  which  it  adds  to  a  flame. 

(ioneral    Bibliography   on   Carbonyls    and   Nitrosyls 

1.  Welch,  Ann.  Repts.  Progr.  Chem.,  38,  71  (1941). 

2.  Blanchard,  Chem.  Revs.,  26,  409  (1940). 

3.  Anderson,  Quart.  Revs.,  1,  331  (1947). 

4.  Hieber,  FIAT  Rev.  German  Sci.,  1939-1946,  Inorg.  Chem.,  Pt.  II,  p.  108  (1948). 

5.  Hieber,  Z.  Elektrochem.,  43,  390  (1937). 

6.  Hieber,  Z.  angew.  Chem.,  55,  11  (1942). 

7.  Smith,  Science  Progr.  35,  283  (1947). 

8.  Trout,  /.  Chem.  Education,  14,  573,  575  (1937);  15,  77,  113,  145  (1938). 

9.  Emel£us  and  Anderson,  "Modern  Aspects  of  Inorganic  Chemistry,"  Second  Edi- 

tion, Chapter  XIV,  New  York,  D.  Van  Nostrand  Co.,  Inc.,  1952. 

Nitrosyls 

1.  Hieber  and  Nast,  FIAT  Rev.  German  Sci.,  Pt  II,  p.  148  (1948). 

2.  Moeller,  /.  Chem.  Ed.,  23,  441,  542  (1946). 


I/.  Organic  Molecular  Compounds 

Leallyn  B.  Clapp 

Brown  University,  Providence,  Rhode  Island 

A  molecular  compound  is  a  substance  formed  from  two  different  com- 
ponents each  of  which  may  have  an  independent  crystal  structure  and 
which,  in  solution  (or  the  vapor  state),  decomposes  into  its  components 
according  to  the  law  of  mass  action.  The  force  which  holds  them  together 
in  the  molecular  compound  has  been  called  secondary  valence  or  residual 
affinity." 

This  translation  of  a  paragraph  from  Hertel1  is  a  working  definition  of 
the  term  ''molecular  compound."  Modifications  necessary  to  fit  more  recent 
concepts  will  pervade  the  text  to  follow. 

One  early  idea  associated  with  the  words  "molecular  compound"  indi- 
cated that  there  was  a  center  of  addition  in  each  component.  The  work  of 
Werner  and  Pfeiffer  led  them  to  suggest  that  the  center  of  addition  in  a 
molecule  could  be  precisely  located  on  a  particular  atom.  The  hypothesis 
of  a  directed  valence  in  molecular  compounds  has  been  attenuated  con- 
siderably by  modern  talk  of  "electron  smears"  and  by  the  ideas  associated 
with  the  word  "resonance." 

The  concept  of  a  center  of  addition  may  be  put  into  symbols1  in  the 
following  way:  if  A  is  an  addition  center  in  molecule  M  which  contains  a 
reactive  group  R,  then  a  true  molecular  compound  is  formed,  if  the  product 
in  Equation  (1)  results  from  the  reaction.  On  the  other  hand,  if 

Ai— Mx— Ri  +  A2— M2— R2  -♦  Ri— Mi— Ai  .  .  .  A2— M2— R2  (1) 

the  reaction  takes  place  according  to  Equation  (2),  the  primary  valences 
are  involved.  The  products  in  the  two  reactions  are,  of  course,  isomers.  As 
an  example,  the  reaction  of  2,4,6-trinitroanisole  and  dimethylaniline2 

\      Mr- R,  +  A  a— M2— R2  ->  Ai— Mi— Ri— R2— M2— A2  (2) 

gives  two  isomeric  products,  one  a  molecular  compound  (Equation  3)  and 
the  other  a  salt,  a  substituted  ammonium  picrate   (Equation  4)3.  The 

1.  Hertel  and  Romer,  Ber.,  63B,  2446  (1930). 

2.  Hertel  and  van  Cleef,  Ber.,  61,  1545  (1928);  Hertel,  Ber.,  57,  1559  (1924). 

3.  Hertel  and  Schneider,  Z.  phys.  Chem.,  151A,  413  (1930);  13B,  387  (1931). 

547 


548 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


product  from  the  reaction  shown  in  Equation  (3)  occurs  as  unstable  red 
needles  and  is  made  by  cooling  a  solution  of  trinitroanisole  in  dimethyl- 
aniline.  If  the  solution  is  heated,  it  turns  yellow  and  deposits  yellow  needles 
on  cooling.  This  is  the  substituted  ammonium  salt  (Equation  4) ;  it  is  solu- 
ble in  water  and  exhibits  other  salt-like  properties. 


NO: 


02N    \T     /"OCH3   •    'CH3— N 


) 


NO, 


CH, 


CH, 


-rO 


chK 


+   r 


NO- 


0-0- 

NO? 


NO- 


In  the  product  from  (3)  the  centers  of  addition  cannot  be  precisely  located 
on  particular  atoms  but  rather  exist  throughout  the  aromatic  ring  in  each 
moiety.  The  linkage  (designated  by  a  dotted  line  in  Equation  1)  may, 
perhaps,  best  be  described  as  a  weak  coordinate  covalent  bond  arising  from 
resonance  conditions  in  the  two  rings. 

This  discussion  of  organic  molecular  compounds  will  be  limited  to  the 
first  three  of  the  following  classes: 

1.  Products  formed  from  benzoquinone,  substituted  quinone,  or  closely 
related  compounds  with  aromatic  hydrocarbons,  amines,  phenols,  and 
aromatic  ethers.  Quinhydrone  is  an  example  known  to  all  chemists. 

2.  Products  of  nitro  compounds  (generally  polynitro)  with  aromatic 
hydrocarbons,  halides,  amines,  and  phenols.  Picrates  of  aromatic  hydro- 
carbons are  well  known  in  this  group. 

3.  Compounds  of  the  bile  acids  (desoxycholic  and  apocholic,  for  example) 
with  fatty  acids,  esters,  paraffins,  and  a  few  other  compounds,  of  importance 
in  biochemistry.  The  clathrates  and  other  occlusion  compounds  are  in- 
cluded in  this  group. 

4.  Compounds  containing  a  hydrogen  bond. 


General  Properties  of  Organic  Molecular  Compounds 

Many  properties  of  organic  molecular  compounds  are  held  in  common 
by  the  first  two  of  these  classes.  Students  of  organic  chemistry  are  familiar 
with  these  compounds  since  they  are  useful  in  identifying  a  number  of  sub- 
stances, particularly  aromatic  hydrocarbons,  ethers,  and  tertiary  amines. 


ORGANIC  MOLECULAR  COMPOUNDS  549 

The  pit-rates1,  especially,  and  sonic  other  molecular  compounds1  have  found 

wide  usage  for  this  purpose.  Many  of  them  are  readily  prepared  merely  by 

mixing  alcohol  solutions  of  the  two  components.  The  stability  of  organic 

molecular  Compounds  varies  but  most  of  them  decompose  rather  than  melt. 

Many  of  them  cannot   be  recrystallized  from  any  solvent   after  they  have 
been  precipitated  because  they  dissociate  into  their  components  in  solution. 

The  influence  of  the  solvent6  is  quite  important.  If  either  component  is 
insoluble  in  a  given  solvent,  the  compound  will  always  decompose.  This 
indicates  that  the  bonding  in  such  compounds  is  (mite  weak.  In  general, 
the  strength  of  the  bond  is  somewhat  less  than  that  of  a  hydrogen  bond; 
it  is,  perhaps,  o  kcal  per  mole  and  certainly  never  more  than  10  kcal  per 
mole7. 

In  a  series  of  fifty  papers,  the  last  of  which  appeared  in  1925,  Kremann8 
and  his  coworkers  reported  studies  of  the  formation  of  a  large  group  of 
organic  compounds  from  binary  mixtures.  They  concluded  that  the  ease  of 
formation  (some  measure  of  stability)  depends  on  an  interaction  of  a  number 
of  factors.  By  far  the  most  important  of  these  is  what  might  now  be  called 
the  difference  in  electronegativity  (electron  affinity)  of  the  two  components. 
If  the  threshold  value  of  this  primary  affinity  is  exceeded,  then  the  ease  of 
formation  of  the  molecular  compound  depends  on  the  positions  of  the 
groups  in  the  aromatic  ring  (asymmetry  of  the  molecule)  and  steric  hin- 
drance. In  this  way  Kremann  accounted  for  the  fact  that  frequently  not 
all  members  of  a  given  homologous  series  nor  all  ortho,  meta,  and  para 
isomers  of  the  same  compound  will  form  a  given  molecular  compound. 

Quinhydrones  and  Related  Compounds 

If  an  alcohol  solution  of  hydroquinone  is  mixed  with  an  alcohol  solution 
of  quinone,  the  solution  turns  brown-red,  and  dark  green  crystals  with  a 
metallic  luster  form.  The  original  hydroquinone  solution  is  colorless  and  the 
quinone  solution  is  yellow.  This  profound  change  is  due  to  the  formation 

4.  Dermer  and  Dermer,  J .  Org.  Chem.,  3,  289  (1938);  Baril  and  Megrdichian,  J.  Am. 

(hem.  Soc,  58,  1415  (1936);  Wang,  J.  Chinese  Chem.  Soc,  1,  59  (1933);  Brown 
and  Campbell,  J.  Chem.  Soc,  1937,  1699;  Mason  and  Manning,  J.  Am.  Chem. 
Soc,  62,  1639  (1940). 

5.  Stephens,  Hargis,  and  Entrikin,  Proc  Louisiana  Acad.  Sci.,  10,  210  (1947);  cf., 

Chem.  Abs.,  42,  1921  (1948),  Reichstein,  Helv.  chim.  Acta,  9,  799  (1926);  Sut- 
ter, II<  h.  chim.  Acta  21,  1266  (1938);  Buehler,  Wood,  Hull,  and  Irwin,  ./.  .1///. 
Chem.  Soc,  54,  2398  (1932). 
ti.  Dimroth,  Ann.,  438,  58  (1924);  Dimroth  and  Bamberger,  Ann.,  438,  67  (1924). 

7.  Wheland,  "The  Theory  of  Resonance,"  p.  4(5,  New  York,  John  Wiley  &  Sons, 

Inc.,  1944. 

8.  PfeifTer,  "Organiache  Verbindungen,"  2nd  Eld.,  p.  272,  Stuttgart,  Ferdinand 

Ilnke,  1927.  (See  author  index  in  PfeifTer  for  original  references  to  Kremann 's 
work.) 


550  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

of  quinhydrone,  a  molecular  compound,  from  equivalent  amounts  of  hydro- 
quinone  and  quinone. 

In  solution,  quinhydrone  dissociates  into  its  two  components  to  an  equi- 
librium point.  The  oxidation  and  reduction  of  quinhydrone  to  quinone  and 
hydroquinone,  respectively,  is  quantitative,  reversible,  and  rapid  enough 
to  be  used  as  an  organic  half-cell  with  a  reproducible  electrode  potential  of 
0.699  volts9  (for  system  from  quinone).  It  is  a  useful  half-cell  for  determin- 
ing pH  values  below  8. 

Both  components  of  the  quinhydrone  molecule  may  be  considerably 
modified  and  still  yield  a  molecular  compound.  Pfeiffer  found  that  aromatic 
ethers  and  even  aromatic  hydrocarbons,  such  as  durene  or  hexamethyl- 
benzene,  could  be  used  with  certain  quinones,  (chloranil,  etc.)  to  give  deeply 
colored  molecular  compounds.  Although  the  phenolic  group  is  unnecessary, 
the  presence  of  the  unsaturated  carbons  of  the  benzenoid  nucleus  is  essen- 
tial. Hexahydrodurene,  for  example,  does  not  give  a  colored  product  with 
any  quinone. 

Only  one  olefinic  double  bond  is  necessary  for  the  quinone  moiety.  As  a 
general  formula,  RCOCH=CHCOR  may  be  substituted  for  the  quinone 
and,  even  here,  the  R  groups  may  be  substituted  by  a  bridging  oxygen 
atom,  for  example,  in  3 , 4 , 5 , 6-tetrachlorophthalic  anhydride.  Quinhydrone 
itself,  then,  is  a  special  case  of  a  more  general  type  of  molecular  compound. 

In  the  solid  state  the  ratio  of  phenolic  component  to  quinone  component 
may  be  1:1,  1:2,  or  2:1,  but  in  solution  the  ratio  is  always  1:1,  regardless 
of  substitutions  on  hydroxyl  groups  in  the  phenolic  part.  Michaelis  and 
Granick10  have  pointed  out  that  crystalline  quinhydrones  have  been  iso- 
lated only  when  there  was  at  least  one  free  hydrogen  on  a  hydroxyl  group 
in  the  phenolic  component.  Yet,  in  solution,  the  affinity  of  the  phenolic 
component  for  the  quinone  is  not  changed  by  alkylation  of  the  phenol  to 
an  ether,  so  a  hydrogen  bond  cannot  play  an  essential  role  in  forming  the 
compounds.  However,  even  as  recently  as  1944,  Pfeiffer11  clung  to  the 
opinion  that  there  is  probably  a  hydrogen  bond  (carbonyl  oxygen  to  hy- 
drogen) in  quinhydrone  itself,  although  workers  in  the  field  of  x-ray  analy- 
sis have  since  rejected  the  notion  that  it  plays  any  important  rcle  in  holding 
the  compound  together. 

Gradation  in  the  color  of  organic  molecular  compounds  has  been  found 
to  be  a  qualitative  measure  of  the  stabilities  of  these  compounds.  The  more 
deeply  colored  compounds  are  usually  more  stable.  In  the  benzenoid  part 
of  the  quinhydrone,  the  groups  — CH3 ,  — OH,  — OCH3 ,  — NH2 ,  and 
— N(CH3)2  deepen  the  colors  of  the  molecular  compounds  while  halogens 

9.  Lammert  and  Morgan,  J.  Am.  Chem.  Soc.t  54,  910  (1932). 

10.  Michaelis  and  Granick,  J.  Am.  Chem.  Soc,  66,  1023  (1944). 

11.  Pfeiffer,  Ber.,  77 A,  59  (1944). 


ORGANIC  MOLECULAR  <  <>MPOUNDS  551 

Table  17. l.  Colob  Gradation  i\  Compounds  Related  k>  Quinhtdboni 

Quinone  Component 
Hvilroquinoiu-  Component  quinone  chloranil  duroquinonc 

Benzene  green  yellow  green-yellow 

Bexamethylbenzene  orange-yellow  red-violet  pure  yellow 

Phenol  orange  blood  red  deep  yellow 

Aniline  blood  red  violet  bright  orange 

Dimethylaniline  violet  red  deep  blue  orange  red 

Anisole  yellow  orange  red 

have  a  hypsochromic  effect.  Substitution  of  halogens  in  the  quinoid  part, 

on  the  other  hand,  deepens  the  color  of  the  molecular  compound  and  sub- 
stitution of  — CH3 ,  — OH,  and  — OCH3  attenuates  the  colors.  These  effects 
are  shown  qualitatively  in  Table  17.1. 

Some  properties  of  a  number  of  compounds  related  to  quinhydrone  are 
shown  in  Table  17.2. 

Picrates  and  Related  Compounds 

Picric  acid  is  an  organic  acid  of  strength  comparable  to  that  of  the  short 
chain  carboxylic  acids.  With  strong  organic  bases  it  forms  picrates  having 
some  of  the  properties  of  substituted  ammonium  salts.  In  many  cases 
these  salts  may  be  recrystallized  from  water  without  decomposition  and 
differ  only  slightly  in  color  from  the  bright  yellow  of  picric  acid,  itself. 
But  with  very  weak  bases,  picric  acid  forms  molecular  compounds  which 
show  pronounced  color  deepening  and  none  of  the  properties  commonly 
•iated  with  salts  (Table  17.3).  One  of  the  satisfying  evidences  that  these 
two  kinds  of  picrates  are  of  different  character  is  that,  in  a  few  cases,  a 
single  amine  can  be  made  to  form  two  picrates,  one  having  salt-like  charac- 
ter and  the  other  exhibiting  molecular  character.  (Table  17.3)  It  was  once 
suggested12  that  the  existence  and  colors  of  these  isomeric  amine  picrates 
could  be  accounted  for  on  a  purely  ionic  basis,  the  formulas  of  the  two 
picrate  ions  being : 

N02  NO: 


1   \1J/    -:  AND  °=  \        )=N> 


N02  N02 


PICRATE    ION  PICRATE    ION    FOR 

FOR  SALT  MOLECULAR    COMPOUND 

While  this  may  be  a  reasonable  picture,  and  might  account  for  the  colors 
of  the  two  kinds  of  picrates,  it  cannot  account  for  the  picrates  of  aromatic 
hydrocarbons,  ethers,  phenols,  and  amine  oxides,  or  the  closely  related 
derivatives  of  polynitro  compounds. 

12.  Bennett  and  Willis,  /.  Chem.  Soc,  1929,  256. 


Table  17.2. 

Properties  of  Some  Compounds  Related  to  Quinhydrone 

Components 

Properties 

Ref. 

Quinone 

thiophenol 

ratio  1:2;  dark  bronze  plates  sol.  benzene, 
ligroin. 

a 

Quinone 

phenol 

ratio  1:2;  red  needles. 

b 

Chloranil 

p-phenyl- 
enediam- 
ine 

blue-black  needles. 

c 

Chloranil 

V  C6H4- 

(NMe2)2 

red  needles,  m.p.  80°,  sol.  hot  alcohol. 

c 

Fluorenone 

benzidine 

yellow  prisms,  m.p.  126  to  127°,  sol.  hot  alco- 
hol, 
dark  red  cryst.,  m.p.  89  to  90°. 

c 

Quinone 

2-nitrohy- 

d 

droqui- 

none 

Quinone 

p-phenyl- 
enediam- 
ine 

dark  brown  ppt.  from  acetic  acid,  insol.  H20. 

e 

Naphthoqui- 

hydroqui- 

dark  green  cryst.  refl.  light,  ruby  red  trans- 

f 

none 

none 

mitted  light,  m.p.  123°. 

Fluorenone 

a-naphthol 

short  red  cryst.  from  benzene,  m.p.  89° 

g 

Chloranil 

acenaph- 

violet  mass  by  melting  components  together, 

h 

thene 

sol.  benzene. 

Bromanil 

durene 

red  needles  from  acetic  acid,  decomp.  in  air 
on  standing,  decomp.  rapidly  80  to  90°. 

i 

Chloranil 

diethoxydi- 
naphtho- 
stilbene 

ratio  1:2,  heavy  black  cryst.  from  benzene. 

i 

Dibenzalace- 

resorcinol 

yellow  needles  from  benzene,  m.p.  95°. 

J 

tone 

2,5-Dichloro- 

hexamethyl- 

bright  red  needles  from  acetic  acid,  m.p.  132 

k 

quinone 

benzene 

to  136°,  stable  a  few  days  in  a  desiccator. 

Chloranil 

hexamethyl- 

fine,  long,  brown-violet  needles  from  acetic 

k 

benzene 

acid,  stable  for  a  long  time. 

Tetrachloro- 

benzene 

ratio  1:3,  benzene  sol.  slowly  evaporated  in 

k 

quinone 

a  vacuum  desiccator  gives  dark  red  cryst., 
m.p.  37  to  42°,  decomp.  in  air  in  a  few  min- 
utes. 

Tetrachloro- 

p-xylene 

dark  red  prisms  from  xylene  sol.  in  vacuum, 

k 

quinone 

m.p.  near  83°,  stable  in  air  few  minutes. 

a  Troeger  and  Eggert,  J.  prakt.  Chem.,  [2]  53,  478  (1896). 

b  Nietzki,  Ann.,  215,  125  (1882). 

c  Schlenk  and  Knorr,  Ann.,  368,  277  (1909). 

d  Richter,  Ber.,  46,  3434  (1913). 

e  Erdmann,  Z.  angew.  Chem.,  8,  424  (1895). 

1  Urban,  Monatsh.,  28,  299  (1907). 

«  Meyer,  Ber.,  43,  157  (1910). 

h  Haakh,  Ber.,  42,  4594  (1909). 

1  Pfeiffer,  Ann.,  404,  1  (1914). 

1  PfeifTer,  Goebel,  and  Angern,  Ann.,  440,  241  (1925). 

k  Pfeiffer,  Jowleff,  Fischer,  Monti,  and  Mully,  Ann.,  412,  253  (1917). 


552 


ORGANIC  MOLECULAR  COMPOUNDS  553 

Table  17.3.  The  Types  of  Picrates  Formed  with  Various  Amines 


Compound  with  Picric  Acid 

Amine 

Ref. 

Salt-like 

Molecular 

a-Naphthyl  amine 

green  yellow  161° 

a,  p.  343 

Methylamine 

yellow  207° 

l» 

0-Naph.tby]  amine 

yellow  198  to  199° 

c 

Carbasole 

red 

d 

Indene 

red 

a,  p.  344 

p,p'-dimethylaminodi- 

straw  yellow  185° 

a,  p.  343 

phenylmethane 

C6H5— CH=N— NHC6H6 

dark  violet  117° 

a,  p.  344 

m-02NC6H4— CH=N— 

dark  red  118° 

a,  p.  344 

NHC6H6 

o-Bromoaniline 

yellow  trans,  pt.  85° 

orange-red  128° 

a,  p.  347 

o-Iodoaniline 

yellow  trans,  pt.  90° 

deep  orange  112° 

a,  p.  347 

l-Chloro-2-aminonaph- 

yellow  trans,  pt.  130° 

dark  red  174° 

a,  p.  347 

thalene 

1  -Bromo-2-aminonaph- 

yellow  trans,  pt.  114° 

violet-red  178° 

a,  p.  347 

thalene 

a  Pfeiffer,  "Organische  Verbindungen,"  2nd  Ed.,  Stuttgart,  Ferdinand  Enke,  1927- 

b  Jerusalem,  J.  Chem.  Soc,  95,  1275  (1909). 

c  Liebermann  and  Scheiding,  Ann.,  183,  258  (1876). 

d  Graebe  and  Glaser,  Ann.,  163,  343  (1872). 

The  introduction  of  radicals  into  the  polynitro  unit  of  the  molecule  or 
into  the  hydrocarbon  part  has  color  effects  comparable  to  those  shown  by 
the  quinhydrone  compounds.  In  the  nitro  part  of  the  molecular  compound, 
an  alkyl  group  in  the  ring  has  a  hypsochromic  effect,  as  it  does  in  the 
quinoid  kernel  of  quinhydrones.  Halogens,  methoxyl,  and  amino  groups  in 

Table  17.4.  Influence  on  Color  of  Substituents  in  the  Nitro  Components 

of  Molecular  Compounds 


Benzenoid  Component 

With  Nitro  Component 

With  Substituted  Nitro  Component 

Hydroquinone 

p-dinitrobenzene,  red- 
orange 

dinitrodurene,    bright   yellow 

Dimethylaniline 

p-dinitrobenzene,  deep 
orange-red 

dinitrodurene,    greenish   yellow 

Durene 

p-dinitrobenzene,  greenish 
yellow 

dinitrodurene,  almost  colorless 

.Viphthalene 

s/ym-trinitrobenzene, 
yellow 

picryl    chloride,   canary   yellow 

a-Xaphthyl  amine 

.s/yw-trinitrobenzene,  red 

picryl  chloride,  brown 

a-Xaphthyl  amine 

picramide,  red 

a-Xaphthyl  amine 

2,4,6-trinitroanisole,  red 

554  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

the  nitro  derivative  (already  strongly  electronegative  due  to  the  presence 
of  the  nitro  group)  have  very  little  influence  on  the  color.  This  will  be 
evident  from  the  data  in  Table  17.4. 

In  the  benzenoid  component  of  the  picrates  and  related  compounds, 
alkyl  groups,  fused  rings,  double  and  triple  bonds,  hydroxyl,  methoxyl,  and 
amino  groups  all  act  as  bathochromes.  Alkyl-  and  aryl-amino  groups  have 
an  even  more  marked  effect  in  deepening  the  colors  while  an  acyl  group 
lessens  the  effect  slightly.  Halogens  in  the  benzenoid  component  have  a 
hypsochromic  effect. 

Structures  of  Molecular  Compounds 

Three  theories  have  been  advanced  to  account  for  the  structures  of  or- 
ganic molecular  compounds.  None  of  the  three  has  attained  complete  ac- 
ceptance, and  none  of  the  three  has  been  completely  discarded. 

1.  Formation  of  a  coordinate  covalent  bond  between  the  two  components. 

2.  Formation  of  polarization  aggregates  which  mutually  saturate  the 
residual  valences  in  the  two  parts. 

3.  Formation  of  an  essentially  ionic  bond  by  transfer  of  an  electron  from 
one  component  to  the  other. 

Coordination  Theory 

The  first  proponents  of  the  theory  of  formation  of  a  coordinate  covalent 
bond  between  the  two  components  of  an  organic  molecular  compound 
were  Bennett  and  Willis12- 13,  closely  followed  by  Moore,  Shepherd,  and 


Fig.  17.1.  Molecular  addition  compound  of  quinoline  with  sym-trinitrobenzene. 

N02 


<y-&K 


N02 

Fig.  17.2.  Molecular  addition  compound  of  sym-trinitrobenzene  with  an  aromatic 
hydrocarbon. 

Goodall14.  In  the  molecular  compound  formed  from  quinoline  and  sym- 
trinitrobenzene,  the  bonding  was  represented  as  shown  in  Fig.  17.1.  If  the 

13.  Bennett  and  Wain,  /.  Chem.  Soc,  1936,  1108. 

14.  Moore,  Shepherd,  and  Goodall,  /.  Chem.  Soc,  1931,  1447. 


ORGANIC  MOLECULAR  COMPOUNDS  555 

amine  is  replaced  by  an  aromatic  hydrocarbon,  it  becomes  more  difficult 
to  locate  the  donor  (Fig,  L7.2)  and  acceptor  atoms.  Further  difficulty  must 
be  faced  in  some  of  the  quinhydrone  type  molecular  compounds  in  having 
to  draw  unfavorable  electronic  distributions  in  some  canonical  forms. 
However,  if  one  pair  of  the  w  electrons  of  a  double  bond  in  an  aromatic 
hydrocarbon  may  be  considered  as  the  donor  pair,  then  the  theory  is  still 
tenable  and  such  pictures  as  Fig.  17.12  will  account  for  the  color  of  such 
molecular  compounds.  The  bathochromic  and  hypsochromic  effects,  de- 
scribed previously  for  the  quinhydrone  type  (see  page  550)  and  the 
picrates  and  related  compounds  (see  page  553),  when  functional  groups 
are  substituted  in  the  aromatic  nucleus,  are  all  plausible  in  terms  of  modern 
electronic  concepts  of  electron  withdrawal  from  (and  electron  supply  to) 
the  ring. 

Polarization  Theory 

The  second  theory,  the  saturation  of  residual  valences,  was  proposed  by 
PfeitYer15  as  a  means  of  accounting  for  the  colors  and  other  properties  of 
organic  molecular  compounds.  Briegleb16  expressed  the  view  that  the  re- 
sidual valences  arise  from  an  inductive  effect.  In  a  compound  of  sym- 
trinitrobenzene  and  an  aromatic  hydrocarbon,  for  example,  the  polar 
groups  (nitro)  induce  an  electric  dipole  in  the  polarizable  aromatic  hydro- 
carbon. The  resulting  electrostatic  attraction  between  the  two  aromatic 
nuclei  maintains  the  compound. 

In  compounds  containing  completely  conjugated  rings,  there  are  two 
types  of  polarization — that  induced  in  the  localized  a  bonds  of  the  hydro- 
carbon and  that  due  to  distortion  of  charge  distribution  of  the  tt  electrons 
(double  bonds).  Briegleb  determined  these  polarizations  spectroscopically. 
The  heats  of  formation  of  a  number  of  molecular  compounds  calculated 
from  the  polarization  values  agreed  with  those  found  experimentally.  Since 
the  heats  of  formation  of  organic  molecular  compounds  are  of  the  order  of 
1  to  5  kcals  per  mole  and  the  force  between  components  of  the  system  (if 
electrostatic)  varies  as  the  inverse  sixth  power,  Briegleb  infers  that  the 
components  cannot  approach  each  other  closely  enough  (1  to  2  A)  to  form 
a  chemical  bond. 

The  chief  objection  to  the  concept  of  polarization  aggregates  due  to 
electrostatic  interactions  is  that  it  does  not  account  for  the  simple  ratios 
of  the  components  which  form  molecular  compounds.  Even  though  one 
would  be  inclined  to  consider  residual  valences  as  integral  since  they  arise 
from   electrons,   the  field  about  the  components  could   not  be    uniform. 

15.  Ref.  8,  Chapt.  I 

16.  Briegloh,  "Zwisrhenmolekiilare  Krafte  und  Molckulst ruktur,"  Stuttgart,  Ferdi- 

nand Enke,  L937    Ahrens  Vortrage,  Vol.  37,  1937);  Briegleb,  Z.  Elektrochem., 

50,  35  (1944). 


\l 


556  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

N02 


^NR2  02N<f  \— OH 


electron  drift  tvt/-v 
>                   JNU: 


Fig.  17.3.  Molecular  addition  compound  of  picric  acid  with  a  tertiary  amine,  ac- 
cording to  the  polarization  theory. 

Hence  one  would  not  expect  molecules  of  greatly  different  sizes  (such  as 
benzene  and  anthracene)  to  form  molecular  compounds  with  a  second 
component  in  which  the  ratios  were  the  same;  but  the  contrary  is  the  case. 
Rheinboldt17  has  compiled  statistics  which  show  that  of  598  organic  molec- 
ular compounds  recorded  in  the  literature,  85.3  per  cent  have  the  1:1 
ratio  of  components  and  98.2  per  cent  bear  the  ratios  1:1,  1:2,  or  2:1. 
Compounds  in  which  the  ratios  do  not  appear  to  be  whole  numbers18  are 
not  numerous  enough  to  remove  the  objection  to  the  theory  of  Pfeiffer 
and  Briegleb. 

As  an  example  of  a  colored  molecular  compound  we  may  take  a  tertiary 
amine  picrate.  From  the  standpoint  of  the  theory  of  residual  valences,  the 
color  in  the  picrate  of  a  weak  base  may  be  thought  of  as  due  to  the  reces- 
sion of  electrons  into  the  picric  acid  end  of  the  pair,  that  is,  in  the  direc- 
tion indicated  by  the  arrow  in  Fig.  17.3. 

Ionization  Theory 

The  polarization  mechanism  for  the  production  of  color19  is  the  primary 
step  in  the  incipient  oxidation-reduction  mechanism  (the  basis  for  the 
third  theory)  proposed  by  Gibson  and  Loeffler20.  They  suggested  that 

[primary  inductomeric  or  electromeric  polarized  associations  (and  not 
simply  Briegleb 's  dipole  aggregates)  do  occur  and  that  they  account  for 
the  color  change.  They  suggested  that  there  is  an  electron  drift  in  the  di- 
rection indicated  in  Fig.  17.3  and  that  the  components  are  brought  into 
close  contact  in  solution  by  thermal  agitation.  The  fact  that  poly  nit  ro 
compounds  give  more  deeply  colored  molecular  compounds  than  mononitro 
compounds  is  accounted  for,  since  the  former  would  promote  a  greater 
electron  drift. 

The  point  of  distinction  between  the  second  theory  and  the  third  theory 
is  just  the  difference  (an  important  one)  between  an  electrostatic  bond 
and  a  chemical  bond. 

Weiss21  has  modified  this  theory  of  the  bonding  in  molecular  compounds 

17.  Rheinboldt,  Z.angew.  Chem.,  39,  765  (1926). 

18.  Emmert,  Schneider,  and  Koberne,  Ber.,  64,  950  (1931). 

19.  Hammick  and  Sixsmith,  /.  Chem.  Soc,  1935,  580. 

20.  GibBOD  and  Loeffler,  /.  Am.  Chem.  Soc.,  62,  1324  (1940). 

21.  Weiss,  J.  Chem.  Soc,  1942,  245. 


ORGANIC  MOLECULAI!  CUMl'OCXDS 


55 1 


NO 


CH- 


-O4-  «+<Z 


NO- 


CH 


■o 


+     r 


FlQ.  17.4.  Transition  Bt&te  in  the 
formation  of  a  molecular  addition  com- 
pound. 


Fig.  17.5.  Ionic  bonding  in  :i  molecu 

l:ir  addition  compound. 


to  the  point  where  it  amounts  to  assuming  an  ionic  bond,  though  this,  of 
course,  la  the  limiting  case.  His  suggestion  is  that  the  bonding  elect  ion 
pair  is  transferred  to  some  extent.  This  really  amounts  to  a  difference  in 
degree  rather  than  kind  since  Weiss'  theory  does  not  suppose  100  per  cent 
ionic  character  for  the  bond.  Molecular  compound  formation  is  represented 
in  Equation  (5), 

A  +  B  ^=±  (AB)t  ->  Ai+B>-  (5) 

where  (AB)t  is  a  transition  complex  probably  resulting  from  dipole  and 
dispersion  interactions.  The  formation  of  the  transition  complex  is  followed 
by  the  actual  electron  transfer,  A  being  the  donor  and  B  the  acceptor. 

The  quantum  mechanical  picture  derived  from  the  potential  energy 
curves  for  the  ionic  state  of  these  organic  molecular  compounds  is  con- 
sistent with  the  observation  that  the  formation  of  such  compounds  is  rapid 
and  often  reversible  and  that  only  a  low  heat  of  activation  is  necessary. 

The  transition  state  in  equation  5  might  be  represented  in  an  early 
stage  by  Fig.  17.4,  the  partial  negative  charge  representing  a  position  of 
high  electron  density  and  the  partial  positive  charge,  a  position  of  electron 
deficiency  as  a  result  of  the  positions  of  the  methyl  and  nitro  groups.  After 
the  electron  transfer  is  consummated,  it  is  probably  best  to  consider  the 
extra  electron  in  the  negative  ion  (Fig.  17.5)  as  "smeared  out"  over  the 
whole  radical.  The  electron  deficiency  in  the  cation  likewise  cannot  be 
precisely  located.  The  conductivities  of  solutions  of  polynitro  compounds 
in  liquid  ammonia22  and  of  aromatic  hydrocarbons  in  sulfur  dioxide23  in- 
dicate that  the  polynitro  compounds  may  act  as  electron  acceptors  and  the 
aromatic  hydrocarbons  may  act  as  electron  donors. 

It  was  found  that  ???-dinitrobenzene  is  a  much  better  conductor  than  the 
ortho  and  para  derivatives,  which  fits  in  with  the  present  electronic  con- 
cepts. In  addition  to  the  conductometric  evidence  for  the  existence  of 
ionic  entities  in  solution,  the  dielectric  properties  of  some  solid  molecular 

22.  Franklin  and  Krans.  Am.  Chem.  ./.,  23,  277  (1000);  ./.  Am.  Chem.  8oc.,  27,  197 

(1905);  Franklin,  Z.  pays.  Chem.,99, 272  (1909);Kraua  and  Bray, ./.  Am.  Chem. 
Soc.,35,  1315  (1913);  Field, Garner,  and  Smith,/.  Chem.  8oe.,W,  1227  (1925); 
Garner  and  Gfflbe,  J.  Chem.  Soc,  1928,  2889. 

23.  Walden,  Z.  pays.  Chem.,  43,  385  (1903). 


558  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

compounds  have  been  measured24.  Weiss  suggests  that  deviations  from 
strict  additivity  of  the  polarizations  of  components  is  additional  evidence 
of  ionic  character. 

To  Weiss,  a  molecule  having  completely  conjugated  double  bonds  repre- 
sents an  electronic  system  similar  to  a  metal  and  so  interaction  between 
two  such  molecules  could  correspond  to  "alloy  formation."  If  the  two  com- 
pounds are  similar  in  electronic  character,  one  would  expect  only  solid 
solutions  of  the  two  "metals,"  whereas  if  there  are  loosely  bound  electrons 
in  one  and  relatively  large  electron  affinity  in  the  other,  molecular  com- 
pound formation  will  result.  This  corresponds  to  intermetallic  compound 
formation.  One  group  of  organic  molecular  compounds  which  show  some 
analogies  to  the  intermetallic  compounds  consists  of  the  colored  compounds 
of  sym-trinitrobenzene  with  unsaturated  ketones25. 
.r  However,  there  is  evidence  against  the  assumption  of  ionic  structures 

for  these  compounds.  Work  in  x-ray  analysis26  of  organic  molecular  com- 
pounds points  to  the  nonexistence  of  ions  in  the  lattice.  Powell  and  co- 
workers point  out  that  ionic  bonds  should  mean  stronger  crystal  lattice 
structures,  which  would  result  in  increased  hardness  and  higher  melting 
points  for  the  complex.  They  list  a  number  of  molecular  compounds  in 
which  the  melting  points  are  lower  than  that  of  one  or  both  components. 
This  has  been  noted  previously27.  The  occurrence  of  diffuse  x-ray  reflections 
in  some  compounds,  e.g.,  that  of  picryl  chloride  with  hexamethylbenzene28, 
shows  that  the  bonds  in  the  crystal  are  not  stronger  than  the  bonds  be- 
tween molecules  of  picryl  chloride  itself,  where  electron  transfer  is  not 
postulated. 

Cook29  voiced  the  opinion  that  further  experimental  verification  is  needed 
before  the  ionic  theory  of  binding  in  organic  molecular  compounds  can  be 
accepted.  Anderson30  has  stated  that  the  constitutions  of  organic  molecu- 
lar compounds  is  the  major  unsolved  problem  confronting  the  theory  of 
valency. 

Occlusion  Compounds 

The  third  class  of  organic  molecular  compounds  is  a  group  in  which  the 
chemical  properties  of  the  components  play  a  secondary  role  to  the  sizes 
and  geometries  of  the  molecules. 

24.  Kronberger  and  Weiss,  J.  Chem.  Soc,  1944,  464. 

25.  Weiss,  J.  Chem.  Soc,  1943,  462;  Reddelien,  J.  prakt.  Chem.,  91,  213  (1915). 

26.  Powell,  Huse,  and  Cooke,  J.  Chem.  Soc.,  1943,  153;  Powell  and  Huse,  J.  Chan. 
Soc.,  1943,  435;  Ann.  Repts.  Chem.  Soc.,  40,  93  (1943). 

27.  Buehler,  Hisey,  and  Wood,  J.  Am.  Chem.  Soc,  52,  1939  (1930). 

28.  Powell  and  Huse,  Nature,  144,  77  (1939);  Ann.  Repts.  Chem.  Soc,  36,  184  (1939). 

29.  Cook,  Ann.  Repts.  Chem.  Soc,  39,  167  (1942). 

30.  Anderson,  Aust.  Chem.  Inst.  J.,  Proc,  6,  232  (1939). 


ORGAX/c  MOLECULAU  COMPOUNDS  559 

Choleic  Acids 

The  choleic  acids  are  a  group  of  water  soluble  molecular  compounds  of 
the  bile  acids  (the  mosl  prominent  being  desoxycholic  acid)  with  a  variety 
or  organic  compounds  such  as  fatty  acids81,  esters82,  ketones  which  enoli: 
camphor14,  long  chain  paraffins88,  polycyclic  aromatic  compounds88,  and 

unsaturated  acids'1.  They  may  also  coordinate  solvent  molecules'57,  such 

as  ether,  ethanol,  benzene,  or  dioxane,  to  form  less  stable  lattices  contain- 
ing solvent  of  crystallization. 

It  is  remarkable  that  the  numbers  of  molecules  of  desoxycholic  acid 
which  coordinate  with  one  molecule  of  a  fatty  acid  are  also  the  coordination 
numbers  commonly  found  in  inorganic  complexes,  namely,  4,  6,  and  8; 
in  a  few  cases,  other  numbers  are  found.  The  coordination  number  ex- 
hibited toward  desoxycholic  acid  (and  apocholic  acid)  by  formic  acid  is 
zero;  by  acetic  acid,  one;  by  propionic  acid,  three;  by  acids  containing 
carbon  chains  C4  to  C8 ,  four;  C9  to  Cm  ,  six;  and  C15  to  C29 ,  eight.  In 
branch-chain  acids,  such  as  isobutyric,  trimethylacetic,  and  isovaleric, 
the  coordination  number  drops  to  two,  while  in  the  unsaturated  long  chain 
acids  (both  cis  and  trans)  such  as  oleic,  erucic,  brassidic,  and  elaidic,  the 
coordination  number  is  eight.  In  dicarboxylic  acids,  the  coordination  num- 
bers are  as  follows:  C4  ,  two;  C6 ,  three;  C7  to  Cn  ,  four;  and  C12  to  C20 ,  six. 
In  esters  of  the  fatty  acids,  the  length  of  the  acid  part  of  the  ester  still 
determines  the  coordination  number  unless  the  alcohol  part  is  long  in  com- 
parison with  the  alkyl  group  of  the  acid. 

Sobotka38  was  led  to  suggest  that,  since  desoxycholic  and  apocholic  acid 
both  have  hydroxyl  groups  at  C3  and  C12 ,  in  contrast  to  the  bile  acids, 
which  do  not  form  choleic  acids,  their  coordinating  abilities  must  be  due 
to  these  two  groups  and  the  shapes  of  these  molecules.  Soon  after,  Kratky, 
Go,  and  Giacomello39,  from  a  series  of  x-ray  studies,  concluded  that  the 

31.  Rheinboldt,  Pieper,  and  Zervas,  Ann.,  451,  256  (1927). 

32.  Rheinboldt,  Konig,  and  Otten,  Ann.,  473,  249  (1929). 

33.  Sobotka  and  Kahn,  Biochem.  J.,  26,  898  (1932);  Ber.,  65B,  227  (1932). 

34.  Rheinboldt,  Konig,  and  Flume,  Z.  physiol.  Chem.,  184,  219  (1929). 

35.  Rheinboldt,  Braun,  Flume,  Konig,  and  Lauber,  J.  prakt.  Chem.,  [2]  153,  313  (1939). 

36.  Marx  and  Sobotka,  J.  Org.  Chem.,  1,  275  (1936) ;  Fieser  and  Newman,  J.  Am.  Chem. 

Soc.,  57,  1602  (1935). 

37.  Wieland  and  Sorge,  Z.  physiol.  Chem.,  97,  1   (1916);  Boedecker,  Ber.,  53,  1852 

(1920). 

38.  Sobotka,  Chem.  Rets.,  15,  311  (1934). 

39.  Herzog,  Kratky,  and  Kurijama,  Xnturwissenschaften,  19,  524   (1931);  Go  and 

Kratky,  Z.  phtjs.  Chem.,  26B,  439  (1934  I;  Go,  IX  Congr.  inU  rn.  quint,  pura  apli- 
cada,  4,  193  L934  ;  cf,  Chem.  Ah,.,  30,  5091  (1936);  Kratky  and  Giacomello, 
Afonatea.,  09, 427  (1936) ;  Go  and  Kratky,Z.  Krtit.,  92A,  310  (1936) ;  Giacomello 
and  Kratky,  /.  K  1st.,  95A,  459  (1935);  Caglioti  and  Giacomello,  Gazz.  chim. 
Hal.,  69,  245  (1939);  Giacomello,  Gazz.  chim.  Hal.,  69,  790  (1939);  Giacomello 
and  Romeo,  Gazz.  chim.  ital.,  73,  285  (1943). 


560  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

crystal  structure  of  desoxycholic  acid  acts  as  an  enveloping  shell  leaving 
a  "channel"  parallel  to  the  longitudinal  carbon  axis  in  which  the  coordin- 
ating molecules  can  lie.  The  unit  cell,  then,  is  cylindrical. 

The  fact  that  space  relations  play  an  important  part  in  the  formation 
of  this  kind  of  organic  molecular  compound  suggests  the  possibility  of  re- 
solving optical  antipodes  by  the  use  of  molecular  compounds.  Although 
only  partial  resolutions  have  been  accomplished  by  this  method,  as  yet, 
it  is  important  because  it  allows  the  resolution  of  compounds  containing 
no  functional  groups.  Windaus,  et  al.40,  were  able  to  resolve  dZ-a-terpineol 
with  digitonin.  Weiss  and  Abeles41  resolved  dZ-sec-butylpicramide  by 
forming  a  molecular  compound  with  d-j8-naphthylcamphylamine,  and 
c?Z-resorcylmethyl  carbinol  has  been  resolved  with  brucine42.  Partial  resolu- 
tions of  methylethylacetic  acid43,  a-phenylbutanol,  dipentene,  and  cam- 
phor44 have  been  accomplished  by  the  use  of  desoxycholic  acid. 

Other  Molecular  Compounds  Involving  a  Channel  Type  Lattice 

Closely  related  to  the  choleic  acids  from  the  standpoint  of  structure  are 
the  colored  molecular  compounds  of  4 , 4/-dinitrobiphenyl  with  various 
adducts,  such  as  benzidene,  4-bromobiphenyl,  4-hydroxybiphenyl,  and 
4-aminobiphenyl.  The  ratios  of  the  components  in  these  compounds  are 
respectively,  4:1,  7:2,  3:1,  and  3:1,  depending  in  large  measure  on  the 
length  of  the  rod-like  molecules  which  fill  in  the  cylindrical  channels  in  the 
4 , 4'-dinitrobiphenyl   lattice45. 

Other  known  molecular  compounds  which  may  be  described  as  having 
a  channel  type  lattice  are  the  urea46  adducts  with  paraffins  and  other 
compounds,  and  the  thiourea47  adducts  with  the  same  wide  variety  of  com- 
ponents. Both  urea  and  thiourea  furnish  a  loose  hexagonal  lattice  for  the 
second  component.  The  ratios48  of  adduct  to  urea  vary  from  1:4.0  with 
butyric  acid  to  1:21.4  with  octaeicosane,  the  ratios  not  necessarily  being 
integral.  The  calculated  length  of  the  holes  in  the  lattice  approximate 
very  closely  the  calculated  lengths  of  fully  extended   adduct  molecules. 

40.  Windaus,  Klanhardt,  and  Weinhold,  Z.  physiol.  Chem.,  126,  308  (1923). 

41.  Weiss  and  Abeles,  Monatsh.,  59,  238  (1932). 

42.  Eisenlohr  and  Meier,  Ber.,  71B,  1005  (1938). 

43.  Sobotka,  Naturwissenschaften,  19,  595  (1931). 

44.  Sobotka  and  Goldberg,  Biochem.  J.,  26,  905  (1932). 

45.  Rapson,  Saunder,  Stewart,  J.  Chem.  Soc,  1946,  1110;  Saunder,  Proc  Roy.  Soc, 

188A,  31  (1946);  190A,  508  (1947);  James  and  Saunder,  Proc.  Roy.  Soc,  190A, 
518  (1947). 

46.  Schlenk,    Ann.,   565,    204    (1949);   Zimmerschied,   Dinerstein,   Weitkamp,    and 

Marschner,  Ind.  Eng.  Chem.,  42,  1300  (1950) ; J.  Am.  Chem.  Soc,  71,  2947  (1949). 

47.  Schlenk,  Experientia,  6,  292  (1950);  Ann.,  573,  142  (1951);  Angla,  Ann.  chim., 

[12]  4,  639  (1949);  Bengen  and  Schlenk,  Experientia,  5,  200  (1949). 

48.  Smith,  Science  Progress,  36,  656  (1948) ;  38,  698  (1950). 


ORGA.MC   UOLECULMS  COM  POUNDS 


561 


O        Ni  ©     NH3 

OO  CN  O     CH 

Fig.  17.6.  "Cage"  lattice  structure  of  a  clathrate  of  benzene,  ammonia,  and 
nickel  cyanide  of  formula  [Ni(C6H6)(NH3)(CN)2]. 


For  example,  in  urea-M-nonane,  the  molecular  ratio  7.7:1  allows  a  hole  of 
14. 1A  in  the  lattice  and  in  urea-n-tetraeicosane,  the  ratio  18.0:1  allows 
a  hole  of  33A.  The  fully  extended  n-nonane  and  n-tetraeicosane  molecules 
should  measure  11.7A  and  30. 6A  respectively. 

Schlenk49  has  reviewed  the  chemistry  of  the  organic  occlusion  compounds, 
including  in  the  channel  type  the  zeolite  adsorption  compounds  which 
have  remarkable  powers  of  adsorbing  straight  chain  hydrocarbons  and 
rejecting  branch  chains  of  the  same  number  of  carbons.  Chabasite,  for 
example,  can  be  used  to  separate  n-butane  from  isobutane  rather  effec- 
tively. 

Clathrates 

Another  group  of  molecular  compounds  in  which  the  geometry  of  the 
crystal  lattice  is  of  prime  importance  is  the  clathrates50.  These  are  com- 
pounds in  which  one  component  is  trapped  in  a  "cage"  lattice  structure 
of  the  second  component.  It  is  evident  that  the  ratio  of  the  two  components 
might  be  integral  only  in  the  limiting  case,  that  is,  in  the  event  of  a  perfect 
lattice  where  every  cage  is  filled  with  the  requisite  number  of  molecules 
of  the  other  component. 

Schlenk,  Fortschr.  Chem.  Forsch.,  2,  92  (1951). 
50.  Powell,  Endeavor,  9,  154  (1950). 


562  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

In  these  compounds,  the  nature  of  the  trapped  component  depends  not 
at  all  on  chemical  properties  but  only  on  molecular  size.  This  is  illustrated 
very  sharply  by  the  clathrates  which  hydroquinone51  forms  with  such 
chemically  unrelated  substances  as  H2S,  S02 ,  CH3OH,  CH3CN,  HCOOH, 
C02 ,  HC1,  HBr,  HC  =  CH,  A,  Kr,  and  Xe.  The  three  inert  gases  emphasize 
the  point  that  chemical  bonds  cannot  be  involved  in  the  formation  of  these 
compounds.  The  x-ray  work  of  Powell  has  been  instrumental  in  elucidating 
the  structures  of  clathrate  compounds.  The  framework  of  the  clathrate 
formed  by  benzene  and  ammonia  with  nickel  cyanide,  [Ni(C6H6)(NH3)- 
(CN)2]52  is  shown  in  Fig.  17.6. 

Water  and  the  aliphatic  hydrocarbons  found  in  natural  gas  form  crystal- 
line clathrates  which  sometimes  cause  considerable  trouble  in  pipeline 
transportation   systems. 

Occlusion  Compounds  Involving  a  Layer  Type  of  Lattice 

A  third  group  of  occlusion  compounds49  is  formed  from  substances  which 
are  trapped  in  the  lattice  of  a  second  component  by  being  caught  between 
layers  of  molecules  forming  the  lattice.  As  examples,  the  following  may  be 
cited:  mineral  clay  adsorbates,  such  as  montmorillonite  with  alcohols, 
glycols,  and  aromatic  hydrocarbons;  basic  zinc  salts  of  organic  acids,  such 
as  naphthol  yellow,  with  water,  alcohols,  and  nitriles;  and  the  liquid  of 
crystallization  adsorbed  in  certain  protein  molecules,  such  as  haemoglobin 
and  horse  methaemoglobin. 

51.  Palin  and  Powell,  /.  Chem.  Soc,  1947,  208;  1948,  571,  817;  Powell,  /.  Chem.  Soc, 

1948,  61;  1950,  298,  300,  468;  Proc.  Intern.  Congr.  Pure  and  Applied  Chem.,  11, 
585  (1947) ;  Powell  and  Guter,  Nature,  164,  240  (1949). 

52.  Powell  and  Rayner,  Nature,  163,  566  (1949);  Rayner  and  Powell,  J.  Chem.  Soc., 

1952,  319. 


lO.  Physical  Methods  in  Coordination 
Chemistry 

Robert  C.  Brasted 

University  of  Minnesota,  Minneapolis,  Minnesota 

and 

William  E.  Cooley* 

University  of  Illinois,  Urbana,  Illinois 

The  study  of  coordination  compounds  has  benefited  greatly  from  data 
accumulated  through  the  use  of  physical  methods.  These  methods  are 
quite  numerous,  and  they  vary  widely  in  degree  of  usefulness  and  breadth 
of  application.  This  chapter  describes  briefly  the  nature  of  the  more  im- 
portant methods,  and  cites  examples  of  their  application. 

Spectrophotometry  Methods 

The  spectra  of  metal  complexes  may  be  broadly  classified  as  absorptions 
due  to  election  vibrations,  absorptions  due  to  molecular  vibrations,  and 
•tra  characterized  by  emitted  frequencies  different  from  a  given  single 
irradiating  frequency.  The  first  type  of  absorption  is  found  in  the  ultra- 
violet and  visible  ranges;  the  second,  in  the  infrared.  The  third  is  due  to 
the  Raman  effect  and  is  a  shifting  of  frequencies.  The  Raman  effect  is 
also  produced  by  molecular  vibrations. 

Correct  interpretation  of  the  absorption  and  Raman  spectra  of  com- 
plexes may  lead  to  conclusions  regarding  their  formulas,  relative  stabili- 
ties, mechanisms  and  rates  of  their  formations,  their  configurations,  and 
in  certain  cases,  their  coordination  numbers.  Raman  spectra  serve  also  as 
a  tool  for  the  measurement  of  the  homopolarity  of  the  coordination  link 
and  of  valence  bond  angles,  and  as  a  basis  for  certain  deduction-  concern- 
ing spatial  arrangements. 

1.  ,<t  and  complete  interpretation  of  visible  and  ultraviolel  spectra  is 
u>nally  not  attempted.  Instead,  comparisons  are  made  between  spectra  to 
be   analyzed   and   standard    spectra    of   known    compounds    Variations   in 

*  Now  at  Procter  and  Gamble  Co.,  Cincinnati,  Ohio. 

563 


564  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

positions  of  absorption  maxima  may  often  be  given  semi-quantitative 
interpretations  with  respect  to  stability  or  displacement  of  one  ligand  by 
another. 

Color  and  Absorption  Spectra 

Before  the  announcement  of  Werner's  theory,  attempts  were  made  to 
relate  the  color  of  complex  compounds  to  the  presence  of  certain  groups. 
Color  was  seen  to  be  related  to  composition,  but  the  presence  of  a  given 
group  in  a  complex  was  found  not  to  be  uniquely  correspondent  to  a  specific 
color.  Kastle1  and  Houston2  were  among  the  first  to  note  a  relationship 
between  color  and  the  positions  of  constituent  elements  in  the  periodic 
table,  as  well  as  the  effect  of  temperature  on  colored  compounds.  In  gen- 
eral, heating  a  compound  having  a  color  in  the  list  below  was  found  to 
produce  successively  the  colors  to  the  right,  while  cooling  was  found  to 
reverse  the  process. 

White  ^=±  Violet  ^±  Blue  ^±  Green  ^  Yellow  ^  Orange  ^  Red  ^  Brown  ^  Black. 
* 

Violet,  blue,  and  green  may  often  be  omitted  because  of  greater  absorption 
of  the  more  refrangible  visible  rays,  and  the  presence  of  white  in  the  list 
refers  only  to  cooling  of  normally  colored  salts. 

According  to  Connelly3,  if  the  mass  of  a  molecule  is  small,  its  period  of 
vibration  in  the  presence  of  light  energy  will  be  small,  leading  to  absorption 
in  the  ultraviolet.  An  increase  in  mass  causes  a  slower  vibrational  period 
and  shifts  absorption  to  the  visible.  Connelly's  interpretation  of  the  effect 
of  temperature  was  based  on  the  concept  of  vibration  of  molecules  about 
a  mean  position.  He  suggested  that  a  rising  temperature  increases  the 
amplitude  of  vibration  and  thus  results  in  a  weakened  restoring  force, 
hence  a  longer  period  of  vibration  and  a  lower  frequency. 

The  first  systematic  study  of  the  color  of  complex  compounds  was  made 
by  Werner4,  who  concluded  that  color  is  more  a  function  of  arrangement  of 
groups  about  the  central  metal  atom  than  of  composition. 

Shibata5,  while  studying  the  spectra  of  complexes  of  cobalt,  nickel,  and 
chromium,  concluded  that  color  is  a  function  of  bonding  and  structural 
arrangement.  He  noted  that  a  complex  may  show  color  even  though  its 
constituents  are  transparent  to  visible  and  ultraviolet  light.  He  related 
the  positions  of  metals  in  the  periodic  table  to  their  color-forming  ability 
in  complexes.  The  metals  of  Groups  I,  II,  and  III  tend  to  form  simple 

1.  Kastle,  Am.  Chem.  J.,  23,  500  (1900). 

2.  Houston,  /.  Franklin  Inst.,  62,  115  (1871). 

3.  Connelly,  Phil.  Mag.,  (5)  18,  130  (1884);  Nichols  and  Snow,  Phil.  Mag.,  (5)  32, 

401  (1891). 

4.  Werner,  Z.  anorg.  Chem.,  22,  91  (1900). 

5.  Shibata,  ./.  Tokyo  Chem.  Soc.,  40,  463  (1919). 


PHYSICAL  METHODS  IN  COORDINATION  CHEMISTRY  565 

ions,  but   in  the  higher  groups,  in  which  COmplexing  tendencies  are  more 

pronounced,  most  Baits  are  colored.  There  are  such  apparent  exceptions 
as  titanium  tetrachloride  and  tetrammine  platinum(II)  chloride;  however, 

the  former  shows  color  upon  aquation,  and  the  latter  absorbs  strongly  in 
the  aear  ultraviolet.  Shibata  attributed  all  color  in  inorganic  compounds 

to  completing,  the  color  resulting  from  molecular  vibrations  or  vibrations 
oi  small  localizations  of  electrons. 

Theories  of  Absorption 

The  origin  of  the  absorption  bands  characteristic  of  coordinated  struc- 
tures is  thought  to  be  in  the  electronic  vibrations  occurring  within  the 
metal  ion,  within  the  coordinated  groups,  and  between  the  metal  and 
ligands.  There  is  no  general  agreement  as  to  the  number  of  absorption 
bands  which  should  be  considered  significant  in  structural  studies.  Since 
a  large  number  of  authors  have  interpreted  structures  in  terms  of  three 
bands  in  the  visible  and  ultraviolet,  these  bands  will  be  considered  stand- 
ard in  this  discussion.  The  first  band  is  usually  found  in  the  range  450  to 
550  mp,  the  second  in  the  range  320  to  400  ma,  and  the  third  in  the  range 
195  to  250  m/i. 

In  1913  Luther  and  Xikolopulos6  postulated  that  the  first  band  arises 
from  the  metal-ligand  bond.  Pauling7  and  Mead8  have  modified  this  by 
attributing  the  band  to  a  combination  of  the  translational  energy  of  the 
bonding  electrons  and  the  vibrational  energies  of  the  central  ion  and  co- 
ordinated groups.  It  is  now  frequently  assumed  that  the  greatest  single 
factor  leading  to  absorption  in  the  first  range  is  vibration  of  the  nonbonding 
electrons  of  the  metal  ion. 

The  coordinate-bond  electrons  are  generally  thought  to  be  responsible 
for  the  second  absorption  band.  Although  there  is  evidence  that  both  the 
first  and  second  bands  result  from  energy  differences  in  excited  states  of 
the  bonding  electrons8, 9,  there  are  dissimilarities  in  the  behaviors  of  the 
two  bands10.  In  the  nitroammine  cobalt(III)  series,  the  substitution  of  a 
nit ro  group  for  an  ammine  group  has  a  hypsochromic  effect  (shift  toward 
the  violet)  on  the  first  band  and  a  bathochromic  effect  (shift  Inward  the 
red)  on  the  second.  For  this  reason  Tsuchida  supports  the  idea  that  these 
bands  have  different  sources. 

The  work  of  Kiss  and  Czegledy11  with  cobalt(III)  complexes  leads  them 
to  conclude  that  any  assignment  of  absorption  bands  to  particular  elec- 

•  i.  Luther  and  NTikolopulos,  Z.  phyaik.  Chem.,  82,  361  (1913 

7.  Pauling,  ./.  Am.  Chem.  Soc,  63,  1367  (1931). 

8.  Mead,  T  an*.  Faraday  Soc.,  30,  1052  (1934 
Mathieu,  Bull. soc. chim.,  (5)3,463  (1936). 

10.  Tsuchida,  Bull  v.  Japan,   ■',   13,  388    i 

11.  Kiss  and  Czeglch  .  Z.  anorg.  all<i*m.  Chem.,  235,  107  (1938). 


566  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

tronic  influences  is  only  approximate.  Accordingly,  they  attribute  the 
first  band  to  the  general  nature  of  the  complex,  rather  than  any  specific 
group  of  electrons.  Their  data  show  that  complexes  of  similar  type,  such 
as  [Co(NH3)6]+++  and  [Co  en3]+++,  have  absorption  curves  of  similar  shapes. 
Successive  replacement  of  ammine  groups  by  nitro  groups  in  the  hexam- 
mine  changes  the  magnitude  of  the  extinction  at  the  maxima.  This  effect 
is  additive  with  respect  to  the  number  of  nitro  groups  present,  and  is  typi- 
cal of  changes  in  the  spectra  of  complexes  having  varying  numbers  of  like 
groups. 

Some  coordinating  groups  have  characteristic  absorption  bands: 

Group  Xmax  of  free  ligand,  m/x 

NOr  366 

NO  3-  302 

S203-  216 

SO  3-  300 

SON"  215 

CN~  220 

C5H5N  250 

These  bands  may  or  may  not  be  shifted  upon  coordination.  The  absorption 
of  the  nitrite  group,  for  example,  is  shifted  on  coordination  to  give  values 
ranging  from  330  to  350  nnx,  which  fall  within  the  limits  of  the  second 
band. 

Two  absorption  maxima  corresponding  to  the  second  and  third  bands 
are  shown  by  K2[HgI4]12.  No  "first  band"  maximum  is  present.  Since  co- 
ordination electrons  are  certainly  involved  in  the  structure  of  this  complex, 
Tsuchida  concludes  that  the  first  band  does  not  necessarily  appear  because 
of  the  formation  of  coordinate  bonds.  Similar  observations  made  with  other 
complexes13  suggest  that  the  first  band  cannot  result  from  vibrations  of 
bonding  electrons.  Tsuchida  suggests  that  it  arises  from  vibrations  in  an 
incomplete  electron  subshell.  The  second  band,  however,  seems  to  be  a 
function  of  bonding,  and  this  band  is  considered  by  Tsuchida  to  be  the 
most  general  absorption  characteristic  of  complexes.  This  conclusion  is 
supported  by  the  fact  that  incident  light  of  the  same  frequency  as  the 
second  band  maximum  may  weaken  or  break  coordinate  bonds. 

Among  the  cobalt  ammine  complexes  containing  ligands  in  addition  to 
ammonia,  Tsuchida  has  assigned  the  following  order  of  stability,  based  on 
hypsochromic  effects  in  the  second  band:  Most  hypsochromic,  most 
stable— N02-  ONO-,  H20,  SCN"  OH"  N03",  Cl~  CO-r,  Br"— least 
hypsochromic,  least  stable. 

A  number  of  studies  of  complexes  have  shown  more  than  three  absorp- 

12.  Tsuchida,  Bull.  Chem.  Soc.  Japan,  (5)  13,  392  (1938). 

13.  Kashimoto  and  Tsuchida,  J.  Chem.  Soc.  Japan,  60,  347  (1939). 


PHYSICAL  METHODS  IN  COORDINATION  CHEMISTRY  567 

tion  hands.  Thus  Csokan  and  Nyiri14,  working  with  inner  complexes  con- 
taining  the  SchifTs  base  of  salicylaldehyde  and  ethylenediamine,  observed 
more  than  three  bands  and  concluded  thai  hydrogen  bonding,  aromatic 
character,  and  polarization  of  molecules,  as  well  as  electronic  shifts,  are 
source's  of  absorption.  Czegledy"  noted  four  distinct  hands  between  200 
and  700  m/j  in  studying  a  number  of  cobalt  complexes. 

Babaeva1'1  1T  lias  noted  the  effects  on  hand  maxima  of  successive  sub- 
stitution of  ammine  groups  in  platinum  complexes.  Nearly  all  the  com- 
plexes studied  show  a  maximum  in  the  range  280  to  290  m/z.  This  range  ifl 
also  common  to  cyanide  complexes  of  cobalt,  chromium,  ruthenium, 
rhodium,  and  palladium.  Platinum  complexes  containing  anionic  ligands 
with  nitrogen  donors  show  another  maximum  in  the  range  256  to  2G7  iriju. 
Substitution  of  nitro  and  chloro  groups  for  ammonia  produces  a  maximum 
in  the  range  330  to  340  niju.  Replacement  of  two  or  more  ammonia  groups 
results  in  complete  absorption  above  about  450  mju.  Extensive  substitution 
by  several  different  groups,  such  as  chloro,  nitro,  and  amido,  increases 
the  number  of  bands  to  six  or  more. 

In  studying  chloro  complexes  of  the  platinum  group,  Babaeva18  concluded 
that  when  two  complexes  are  identical  except  for  the  metal,  the  complex 
of  the  metal  of  lower  atomic  number  shows  absorption  bands  at  greater 
wave  lengths.  This  relation  applies  only  to  metals  of  the  same  periodic 
group.  Babaeva  attributes  the  effect  to  differences  in  excitation  ener- 
gies of  d  electrons. 

It  has  been  generally  assumed  that  groups  outside  the  coordination 
sphere  do  not  contribute  to  the  spectrum  of  the  complex,  but  this  assump- 
tion seems  unjustified.  Linhard19  observed  cobalt(III)  and  chromium(III) 
ammines  and  ethylenediamine  complexes  in  the  presence  of  halide,  per- 
chlorate,  and  nitrate  ions,  and  found  that  weak  associations  yielding  ions 
of  the  type  [Co(XH3)6]I++  produce  absorption  bands. 

The  Third  Band 

The  complex  absorption  maximum  of  shortest  wave  length  wras  first 
given  systematic  consideration  by  Shibata  and  Tsuchida  and  their  co- 
workers20- M«  -.  Data  accumulated  by  these  authors  for  the  cobalt   nitro- 

14.  Csokan  and  Nyiri,  Magyar  Cfu  m.  /■'<>! yoiral,  47,  149  (1941). 

15.  Czegledy,  Acta  Lit.  Set.  Regiai  Univ.  Hung.  Frencsico-Josephinae,  Sect.  Chem., 

Minimi.  Pkys.,  6,  121  (1937). 

16.  Babaeva,  <  end.  acad.  aci.t  U.R.S.S., 20, 366  (1938). 

17.  Babaeva,  Compt.  rend.  acad.  aci.,  U.R.S.S.,  40,  61  (1943  . 

18.  Babaeva,  />'///.  acad.  sci.  U.R.S.S.,  <'/>iss<  .•«■;.  chim.}  171  1,1943  . 

19.  Linhard,  /.  Elektrochem.,  50,  224    1944  . 

20.  Shi!, at;,.  ./.  Coll.  Sri.  Jmp.  Univ.  Tokyo,  37,  Am.  2,  1  28  (1915  ;  37.  An.  8,  1-12 

L916). 
Shibata  and  Qrbain,  Compt.  rend.,  157,  503  5  (1914). 


568  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

ammine  complexes  showed  that  a  third  band  was  consistently  found  when 
two  nitro  groups  occupied  trans  positions  in  the  complex.  Tsuchida23  also 
found  a  third  band  for  ^rans-[Co(NH3)4Cl2]Cl.  Extension  of  these  studies 
showed  that  the  presence  of  a  third  band  could  be  quite  generally  related 
to  a  frans-diacido  structure.  Tsuchida  noted  that  the  presence  of  a  third 
band  seemed  independent  of  the  configuration  of  the  complex,  the  identity 
of  the  ligands,  and  the  ionic  charge  of  the  complex,  so  long  as  two  negative 
groups  occupied  trans  positions.  Tsuchida's  explanation  of  the  presence  of 
the  third  band  describes  it  as  a  polarization  phenomenon  possible  only 
when  two  negative  groups  occupy  antipodal  positions  in  the  coordination 
sphere. 

More  recent  spectral  studies  by  Basolo24  have  shown  that  cis-diacido 
complexes  also  show  absorption  in  the  third  band  region.  Older  investiga- 
tions generally  extended  only  to  a  lower  limit  of  250  mu.  Basolo  has  found 
that  the  cis  complexes  absorb  at  wave  lengths  which  are  usually  less  than 
250  m/x,  and  these  absorptions  were  undetected  by  Shibata,  Tsuchida,  and 
others.  The  hypotheses  attributing  the  third  band  to  phenomena  peculiar 
to  trans  structures  are  therefore  disproved.  Nevertheless,  Basolo's  data 
point  out  that  the  cis  and  trans  forms  of  a  given  complex  do  exhibit  con- 
sistent differences  in  the  positions  of  absorption  maxima  in  the  second  and 
third  bands,  as  shown  below: 

Complex 

czs-[Co(NH3)4(N02)*]+  238 

*mns-[Co(NH3)<(N02)2]+ 

cis-[Coen2  (N02)2]+  240 

itrans-[Co  en2  (N02)2]+ 
The  positions  of  the  second  and  third  maxima  are  therefore  useful  in  de- 
termining geometric  configurations  when  the  maxima  are  known  for 
analogous   complexes. 

Special  Bands 

Complexes  containing  certain  ligands,  among  them  chromate,  isothio- 
cyanate,  and  dimethylglyoxime,  sometimes  show  absorption  maxima 
which  are  not  attributable  to  the  causes  previously  discussed.  Tsuchida25, 26 
has  classified  these  special  bands  into  two  types:  those  which  are  charac- 
teristic of  the  ligands,  whether  coordinated  or  free,  and  those  appearing 
only  on  coordination.  The  ion  [Co(NH3)5Cr04]+  shows  special  band  ab- 

21.  Shibata  and  Matsuno,  J.  Tokyo  Chem.  Soc,  39,  661  (1918). 

22.  Tsuchida  and  Kashimoto,  Bull.  Chem.  Soc.  Japan,  11,  785  (1936). 

23.  Tsuchida,  Bull.  Chem.  Soc.  Japan ,11,  721  (1936). 

24.  Basolo,  ./.  .1///.  Chem.  Soc,  72,  1393  (1950). 

25.  Tsuchida  and  Kibayashi,  Bull.  Chem.  Soc.  Japan,  (7)  13,474  (1938). 

26.  Tsuchida,  Bull.  Chem.  Soc.  Japan,  (6)  13,  437  (1938). 


Xmax,  ni/i 

327 

255 

356 

325 

250 

345 

PHYSICAL   METHODS  Ih   COORDINATION  CHBMISTR)  569 

sorption  of  the  firsl  type.  A  complex  having  special  band  absorption  of 
the  second  type  is  [Cr  Ml  .,WS]?f.  This  absorption  is  present  also  when 
more  than  one  isothiocyanato  group  is  present,  and  the  extinction  is  ad- 
ditive with  respect  to  the  number  of  these  groups. 

Determinations  of  the  Nature  and  Stability  of  Complexes 

Complexes  In  Solution.  The  spectrophotometric  method  is  especially 

well  suited  to  the  study  of  complexes  not  sufficiently  stable  to  permit 

their  isolation  from  solution.  Work  of  this  type  has  been  done  by  Job*7, 
who  developed  the  M<(hod  of  Continuous  Variations.  This  method  makes 
use  of  any  measurable  additive  property  of  two  species  in  solution,  so 
long  as  the  property  has  different  values  for  the  two  species.  Any  complex 
formed  by  the  two  species  must  give  a  value  for  the  same  property  which  is 
different  from  the  weighted  mean  of  the  values  for  the  separate  species. 
The  simplest  application  of  the  method  involves  an  equilibrium  of  the 
type  A  +  ?iB  ^±  ABn  ,  where  A  represents  a  metal,  B  a  coordinating  group, 
and  AB,  a  complex.  Solutions  are  prepared  in  which  the  mole  fractions  of 
the  components  are  varied  and  the  total  number  of  moles  of  both  together 
is  kept  constant.  Volume  changes  are  usually  ignored,  unless  they  are  so 
great  that  the  volume  may  be  used  as  the  additive  property.  The  extinc- 
tion coefficients  of  the  solutions  are  measured,  using  a  monochromatic 
li<dit  source.  If  there  is  no  complexing,  the  plot  of  extinction  coefficient 
against  mole  fraction  of  one  component  is  a  straight  line.  But  if  a  complex 
is  formed,  the  plot  deviates  from  linearity,  the  deviation  being  a  maximum 
at  the  mole  fraction  corresponding  to  the  composition  of  the  complex. 
When  the  deviation  is  plotted  against  mole  fractions,  the  maximum  point 
gives  the  desired  composition.  The  conclusion  may  be  verified  by  repeating 
the  process  at  other  wave  lengths,  since  the  position  of  the  maximum  is 
independent  of  wave  length. 

A  good  example  of  the  use  of  the  method  is  given  by  a  study  of  complexes 
of  iron(III)  with  various  anions28.  The  data  showing  formation  of  a  citrate 
complex  are  given  in  Figure  18.1.  The  dotted  lines  represent  solutions  ten 
times  as  concentrated  as  those  plotted  with  solid  lines.  The  single  maxima 
support  the  conclusion  that  only  one  complex  is  formed. 

The  Job  method  has  been  extended  by  Vbsburgh  and  his  associates 
particularly  to  deal  with  the  formation  of  more  than  one  complex.  In  work- 
ing with  o-phenanthroline  complexes  of  nickel (II),  Vosburgh  and  Cooper29 

27.  Job,  Ann.  chim.,  9,  113  (1928). 

28.  Lanford  and  Quinan.  J.  Am.  Chem.  Soc,  70,  2!KX)  (1948). 

29.  Vosburgh  and  Cooper,  ./.  Am.  Chem.  Soc,  63,  437  (10  tl 

30.  Gould  and  Vosburgh,./.  Am.  Chun.  8(H   .  64,  L630    L942  . 


570 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


0.2       03        04       0.5        0.6 
MOLE       FRACTION,   Fe+++ 
Fig.  18.1.  Deviations  of  extinction  coefficients  from  additivity,  iron  (III) -citrate 

solutions. 


first  determined  the  optical  densities  of  solutions  of  the  components  hav- 
ing mole  fractions  of  nickel  ion  equal  to  0.50,  0.33,  and  0.25.  A  range  of 
wave  lengths  between  500  and  650  m^u  was  used.  Mathematical  analysis 
shows  that  if  complexes  are  formed  with  molar  ratios  of  1:1,  1:2,  and  1:3, 
determination  of  the  first  complex  is  most  conveniently  made  at  a  wave 
Length  corresponding  to  nearly  equal  extinction  coefficients  of  the  first 
two  complexes.  Similarly,  the  second  is  determined  by  use  of  a  wave  length 


PHYSICAL  METHODS  IN  COORDINATION  CHEMISTRY  571 

giving  oearly  equal  extinction  coefficients  for  the  second  and  third.  For 
determination  of  the  third  complex,  its  extinction  coefficienl  should  be 
much  greater  than  that  of  the  second,  provided  no  fourth  complex  is 
formed.  The  appropriate  wave  Lengths  in  each  case  were  found  from  the 

optical  density  curves  for  the  1:1,  1:2,  and  1  \'.\  solutions.  It  is  assumed 
that  formation  of  the  first  complex,  having  a  1:1  ratio,  consumes  all  the 
free  metal  ion.  Then  the  linear  plot  (assuming  no  reaction)  of  extinction 
coefficient  against  mole  fraction  is  made  between  pure  1:1  complex  and 
pure  complexing  agent.  Accordingly,  Vosburgh  and  Cooper  used  light  at 
620  mjj  to  establish  the  existence  of  [Xi(o-phen)]++.  This  complex  was 
then  assumed  to  he  mixed  with  o-phenanthroline  in  the  solutions  of  greater 
concentration  of  the  latter;  no  uncomplexed  nickel  was  considered  to  be 
present.  A  new  linear  plot,  of  a  different  slope,  was  next  required,  and  the 
existence  of  [Xi(o-phen)2]++  was  demonstrated  with  light  at  580  m/x.  Finally, 
a  wave  Length  of  528  m/u  served  to  determine  the  [Xi(o-phen)3]++  complex 
with  a  third  linear  plot.  In  each  case  the  deviations  from  linearity  reach  a 
maximum  at  the  composition  sought,  as  in  the  original  method. 

The  extended  method  of  continuous  variations  enabled  Haendler31  to 
show  that  diethylenetriamine  forms  copper(II)  and  nickel (II)  complexes 
containing  either  one  or  two  amine  molecules.  This  implication  of  a  coor- 
dination number  of  six  is  supported,  in  the  case  of  nickel,  by  Vosburgh29' 30, 
who  reports  the  existence  of  [Ni  en]++,  [Xi  en2]++,  and  [XTi  ens]4^.  As  with 
the  other  applications  of  this  method,  the  presence  of  water  molecules  in 
the  coordination  sphere  is  usually  not  detected.  Thus  the  apparent  coor- 
dination numbers  in  [Ni  en]^  and  [X^o-phen)]"^,  for  example,  are  not 
necessarily  the  true  coordination  numbers. 

Job27  has  shown  that  when  the  formula  is  known  for  a  complex  in  solu- 
tion, the  equilibrium  constant  of  its  formation  (or  its  reciprocal,  the  dissocia- 
tion constant)  may  be  found  mathematically  through  a  relation  between 
concentration  and  extinction  coefficient.  As  part  of  his  continuous  varia- 
tions studies,  Job  found  constants  for  a  number  of  complexes. 

Babko32  has  investigated  the  formation  of  copper(II)  salicylate  com- 
plexes at  various  pH  values.  A  plot  of  extinction  coefficient  against  pH 
shows  sharp  breaks  at  pH  3-5  and  pH  7-9,  indicating  the  presence  of 
[Cu (salicylate)]  and  [Cu(salicylate)2]=,  respectively.  The  same  author  has 
studied  iron(III)  thiocyanate  complexes  in  aqueous  solution33-34.  Varia- 
tions in  extinction  coefficient  with  thiocyanate  concentration  give  evidence 
for  formation  of  the  complexes  [Fe(SCXT)x]3_z,  where  x  ranges  from  1  to  G. 

31.  Haendler,  J.  Am.  Chem.  Soc,  64,  686-8  (1942). 

32.  Babko,  J.  Gen.  Chem.,  I  . >>/,'.  17,  4  13  (1947). 

33.  Babko,  J.  Gen.  Chem.,  U.SJ3.R.,  16,  33,  1549  (1946);  16,  758,  874  (1945). 

34.  Babko,  Compt.  rend.  acad.  sci.,  U.L'.S.S.,  52,  37  (1946). 


572  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Studies  on  decolorization  of  the  thiocyanate  complexes  by  addition  of  fluo- 
ride ion  have  shown  the  existence  of  such  equilibria  as 

[Fe(SCN)]++  +  nF-^±  [FeFnp-»  +  SCN". 

If  the  magnitude  of  the  extinction  at  an  absorption  peak  is  proportional 
to  the  concentration  of  the  complex  giving  rise  to  the  absorption,  the 
method  of  Moore  and  Anderson35  is  useful  in  determining  the  stability  of 
the  complex.  From  the  equilibrium 

i_L_iTl  v       [AMB]» 

raA  +  nB  ^±  AJ3n ;        K  =  , 

[A^BJ 

whence  log  [AmBn]  =  m  log  [A]  +  n  log  [B]  -  log  K. 

If  [A]  is  kept  constant  and  [B]  is  varied,  log  [Ajy  is  a  linear  function  of 
log  [B].  If  the  logarithm  of  the  optical  density,  which  is  proportional  to 
log  [AJB»],  is  plotted  against  log  [B],  the  slope  of  the  resulting  straight 
line  is  the  value  of  n.  The  value  of  m  may  be  similarly  determined,  and 
the  constant  K  may  then  be  found.  In  studying  the  system  involving 
cerium (IV),  sulfate,  and  perchlorate,  these  authors  have  concluded  from 
concordant  results  of  the  logarithmic  and  continuous  variations  methods 
that  no  colored  complex  is  formed  between  cerium  and  perchlorate  ions. 
In  solutions  having  total  ionic  concentrations  up  to  O.Olilf  the  complex 
[CeS04]++  exists.  At  higher  concentrations  the  complexes  [Ce (804)2]  and 
[Ce(S04)3]=  appear. 

Thorns  and  Gantz36  noted  the  effect  of  various  anions  on  the  absorption 
of  iron  (III)  chloride  solutions  between  350  and  750  nnx.  From  the  data, 
the  authors  ranked  the  various  anions  with  respect  to  relative  ease  of 
replacement  of  any  one  in  the  series  by  any  other:  most  stable — CN~~, 
citrate,  C204=,  C4H406=,  C2H302-,  P04a,  F~,  SCN"  B407=,  S04=,  CI",  Br-, 
I-,  N03- — least  stable.  Studies  of  this  type  have  also  been  made  by  Kossi- 
koff  and  Sickman37  on  copper(II)  nitrite  complexes;  they  concluded  that 
one,  two,  or  three  nitrite  ions  may  be  attached  to  copper,  but  each  succes- 
sive nitrite  group  is  more  difficultly  added.  Bjerrum38  has  studied  the  chloro 
complexes  of  copper (II)  and  reports  that  only  the  complex  [CuCl4]=  is 
sufficiently  stable  to  produce  absorption  measurably  different  from  that 
of  the  components. 

Numerous  investigations  have  been  made  of  the  substitution  of  chloro 
groups  for  water  molecules  in  the  hexaquocobalt(II)  ion.  Howell  and 
Jackson39  observed  maxima  in  the  plot  of  extinction  coefficient  against 

35.  Moore  and  Anderson,  /.  Am.  Chem.  Soc,  67,  168  (1945). 

36.  Thorns  and  Gantz,  Proc.  Indiana  Acad.  Sci.,  56,  130  (1946). 

37.  Kossiakoff  and  Sickman,  /.  Am.  Chem.  Soc,  68,  442  (1946). 

38.  Bjerrum,  Kgl.  Danske  Videnskab  Selskhb,  Math.-fys.  Medd.,  22,  (18),  43  (1946)' 

39.  Howell  and  Jackson,  J.  Chem.  Soc,  1268  (1936). 


PHYSICAL  METHODS  IX  COORDINATION  CHEMISTRY  573 

mole  fraction  of  added  chloride.  They  propose  the  equilibria: 

[Co(II20)6]++  +  2d-;=±  [Co(H20)4Clo]  +  2II20 
[Co(HiO)4Cl,j  +  Cl-^  [Co(H20)3Cl,]     I-  11  .» 

Gerendes40,  however,  found  it  possible  to  identify  six  separate  maxima 
with  increasing  chloride  concentration,  hydrochloric  acid  acting  as  the 

source  of  chloride.  From  this  evidence  (iereudes  concluded  thai  complete 
and  stepwise  replacement  of  water  by  chloride  takes  place,  resulting  ul- 
timately in  the  formation  of  [CoCl6]4~.  Kiss  and  his  co-workers11-  '-  have 
extended  this  study  to  nonaqueous  solvents,  noting  tendencies  toward 
solvent  coordination,  particularly  with  pyridine.  Kiss  has  also  found  that 
in  nonaqueous  solvents  there  are  frequent  exceptions  to  the  commonly 
assumed  rule  that  all  red  cobalt(II)  complexes  are  six-coordinate,  and  all 
blue  cobalt  (II)  complexes  are  four-coordinate. 

Spectral  methods  have  been  useful  in  examining  possibilities  of  the 
formation  of  unusual  oxidation  states.  Strong  spectrometric  evidence  for 
the  formation  of  the  pentavalent  molybdenum  complex  [Mo(SCX)5]  was 
found  by  Babko43.  A  sharp  extinction  maximum  corresponds  to  the  forma- 
tion of  the  complex  with  thiocyanate  concentrations  in  the  vicinity  of 
0.1  M.  Greater  concentrations  lead  to  the  formation  of  [Mo(SCN)6]~, 
whereas  dilution  produces  [Mo(SCN)2]+++  and  [Mo(SCN)]4+.  The  possible 
existence  in  solution  of  tin  (III)  and  antimony  (IV)  species  was  investigated 
spectrally  by  Whitney  and  Davidson44,  who  concluded  that  no  evidence 
suggests  the  existence  of  these  states. 

Much  information  concerning  the  mechanisms  of  reactions  of  com- 
plexes may  be  obtained  spectrophotometrically.  If  the  absorption  spectra 
of  two  complexes  are  known,  for  example,  and  one  of  them  may  undergo 
stepwise  reaction  to  form  the  other,  the  nature  of  the  intermediate  prod- 
ucts may  frequently  be  determined.  For  this  purpose  it  is  possible  to  com- 
pare the  spectra  taken  during  the  reaction  with  the  spectra  of  known 
species  thought  to  be  logical  intermediate  products.  A  second  approach 
involves  measuring  the  total  effect  of  the  reaction  on  the  position  and 
intensity  of  the  absorption  bands,  then  using  the  intermediate  spectra  as 
a  basis  for  calculated  identification  of  any  transient  species  formed.  Serf  ass 
and  Theis45  have  shown  that  sulfato  complexes  of  chromium(III)  may 
undergo  successive  replacement  of  sulfato  groups  by  hydroxy  groups.  This 

40.  Gerendes,  Magyar  Chem.  Folyoirat,  43,  31  (1937). 

41.  Kiss,  Csokan  and  Richter,  Acta  I  rniv.  Szeged.  Sect.  Set.  Nat.}  Acta  ( 'hi  m.,  Mil 

Phys.t  7,  119  (1939). 
12.  Loss  and  Csokan,  Z.  physik.  Chem.,  A186,  23!)  (1940]  . 

43.  Babko,  /.  Gen.  Chet  9LR.,  17, 642  (1947). 

44.  Whitney  and  Davidson, /.  Am.  Chem.  Soc.,  69, 2076  (1947). 

45.  Serfase  and  Theis,  J.  Am.  Leather  Chemists*  Assoc. }  43,  2()*i  (1948). 


574  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

replacement  may  be  followed  spectrophotometrically  by  observing  the 
pronounced  increase  in  extinction  at  420  ncuz,  as  well  as  a  lesser  increase 
at  580  mju,  caused  by  the  entry  of  each  hydroxy  group  into  the  complex. 
Addition  of  sulfuric  acid  reverses  the  reaction  and  reduces  these  maxima. 
Uemura  and  Hirasawa46  have  studied  the  effect  of  pH  upon  ethylenedia- 
mine  complexes  of  cobalt.  The  spectrum  of  tris(ethylenediamine)cobalt(III) 
ion  shows  little  variation  between  pH  1  and  pH  10.  By  comparison  with 
standard  curves,  however,  these  authors  noted  the  following  changes 
with  bis(ethylenediamine)  complexes: 

cis-[Co  en2  (H20)2]+++  ,    °H     \  cis-[Co  en2  (H20)  OH]++       C1~   ) 
H+ 

cis-[Co  en2  (H20)  Cl]++      H2°    ) 

trans-[Co  en2  (H20)2]+++  — 9EL^  trans-[Co  en2  H20  OH]++. 

The  complexes  [Co  en2  (H20)2]+++,  [Coen2  (H20)  Cl]++  and  [Co  en2  Cl2]+  all 
were  observed  to  be  stable  in  acid  solution;  in  basic  solution  they  are  trans- 
formed to  [Co  en2  (H20)  OH]++.  It  was  also  noted  that  the  differences  in  the 
absorption  spectra  of  the  cis  and  trans  forms  of  [Co  en2  (H20)  N02]++,  useful 
for  distinguishing  these  isomers  in  acid  solution,  are  lost  upon  the  addition 
of  base. 

The  three  isomeric  species  [Cr(H20)6]Cl3 ,  [Cr(H20)5Cl]Cl2-H20,  and 
[Cr(H20)4Cl2]Cl-2H20  were  studied  by  Datar  and  Quereski47.  It  was  found 
that  a  transition  from  the  third  complex  to  the  first  takes  place  on  standing 
in  aqueous  solution.  Irradiation  by  ultraviolet  light  weakens  the  metal- 
chlorine  bond  and  increases  the  rate  of  aquotization.  This  is  significant  in 
that  the  frequency  range  chosen  for  a  spectral  investigation  may  include 
frequencies  which  affect  the  system  under  study. 

Hagenmuller48  has  developed  a  graphical  method  for  determination  of 
complex  dissociation  constants  from  continuous  variations  data.  As  in 
Job's  original  method,  a  curve  is  drawn  to  show  the  deviations  of  a  property 
from  the  values  it  would  assume  if  no  complex  formation  took  place. 
Whereas  Job's  calculations  of  dissociation  constants  involve  application  of 
the  law  of  mass  action,  Hagenmuller's  method  permits  direct  calculation 
of  the  constants  from  the  shape  of  the  deviation  curve.  The  reader  is  re- 
ferred to  Hagenmuller's  discussion  for  mathematical  details.  For  the  equilib- 
rium, 

Hg(N02)2  +  Zn(N02)2^±  Zn[Hg(N02)4], 

46.  Uemura  and  Hirasawa,  Bull.  Chem.  Soc.  Japan,  13,  379  (1938). 

47.  Datar  and  Quereski,  J.  Osmania  Univ.,  8,  6  (1940). 

48.  Hagenmuller,  Compt.  rend.,  230,  2190  (1950). 


PHYSICAL  METHODS  IN  COORDINATION  CHEMISTRY  575 

Job's  method  of  calculation  of  K,{  for  Zn[Hg(N0i)4]  yields  the  value  0.50. 
The  graphical  method  yields  A',,  =  0.56. 
Brigando48  has  carried  out  a  spectrophotometric  continuous  variations 

study  on  solutions  of  cobalt (II)  chloride  and  bistidine.  Her  data  indicate 
formation  of  cobalt  (II I  >  complexes  containing  four  and  six  molecule-  of 
histidine  per  cobalt  (III)  ion.  These  complexes  form  slowly,  the  four-co- 
ordinate one  forming  from  30  to  180  minutes  after  mixing  the  cobalt(II) 
solution  with  histidine.  The  six-coordinate  complex  is  present  at  equilib- 
rium, attained  in  five  hours.  Although  the  complexes  form  slowly,  they  are 
sufficiently  stable  so  that  the  trivalent  cobalt  cannot  be  precipitated  by 
the  addition  of  thiocyanate  or  hydroxide  ions. 

A  large  number  of  spectral  studies  of  reactions  of  complexes  have  been 
carried  on  by  Basolo  and  his  associates50, 51.  These  studies  give  special 
emphasis  to  the  kinetics  and  mechanisms  of  reactions.  Basolo,  Hayes,  and 
Neumann50  investigated  the  mechanism  of  racemization  of  the  optically 
active  ions  tris(o-phenanthroline)nickel(II)  and  tris(2,2'-dipyridyl)- 
nickel(II).  The  rates  of  racemization  for  the  two  complexes  in  water  solu- 
tion were  compared  with  the  rates  of  dissociation  in  acid  solution,  according 
to  the  equations: 

[Ni  (o-phen)3]++  — >  [Ni(o-phen)2]++  +  o-phen 
H+  +  o-phen  — *  H  o-phen+. 

The  products  of  the  dissociation  show  different  absorption  characteristics 
from  those  of  the  reactants.  Measurement  of  the  changes  in  absorption  at 
400,  420,  440,  and  520  m^  was  sufficient  to  provide  quantitative  rate  data. 
Mathematical  analysis  shows  that  under  the  same  conditions  the  rates  of 
racemization  and  dissociation  are  equal,  within  experimental  error,  and 
that  the  activation  energies  for  the  two  processes  are  equal.  It  is  evident, 
therefore,  that  racemization  of  these  complexes  takes  place  by  a  mechanism 
of  dissociation.  This  mechanism  is  to  be  contrasted  with  the  intramolecu- 
lar rearrangement  process  which  probably  characterizes  the  racemization 
of  the  tris(oxalato)cobalt(III)  ion. 

Infrared  Spectra 

Absorption  of  radiation-  in  the  infrared  range  is  attributed  to  molecular 
vibrations  of  the  absorbing  material.  These  vibrations  comprise  motions 
of  the  atomic  masses  in  the  material  about  centers  of  vibration.  For  pur- 

49.  Brigando,  Compt.  rend,,  237,  163  (1953). 

50.  Basolo,  He  es,  :  i  r  j  « 1  Neumann,  ./.  Am.  Chem.  Soc.,  75,  5102  (1953). 

51.  Basolo,  Stone  and  Pearson,  /.  Am.  CI  76,819    1953  ;  Pearson,  Boston, 

..  75,  308  Basolo,  Stone,  Bergmann,  and 

-in.  ./.  Am.  Chem.  Soe.,  76,  3079    1964  .  Basolo,  Chen,  and  Murmann, 
J.  An,.  Chem.  Soc,  76,  9.56  (1954). 


576  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

poses  of  description,  two  atoms  which  are  covalently  bound  to  each  other 
may  be  thought  of  as  the  simplest  vibrational  system.  The  two  atomic 
masses  represent  the  bodies  which  are  displaced  during  vibration,  and  the 
strength  of  the  bond  corresponds  to  the  restoring  force.  Thus  each  such 
system  has  a  characteristic  vibrational  frequency  depending  upon  these 
factors,  and  it  absorbs  infrared  radiations  of  the  same  frequency.  In  gen- 
eral, only  vibrations  of  an  unsymmetrical  nature  are  detected  by  infrared 
absorption.  Only  completely  homopolar  bonds  are  thereby  excluded,  how- 
ever, and  even  these  must  be  isolated  from  any  other  vibrating  systems 
in  order  to  be  free  of  coupling  effects.  In  actual  practice,  the  molecular 
vibrations  in  complex  compounds  are  of  such  abundance  and  variety  that 
complete  and  precise  interpretations  of  spectra  are  usually  impossible. 
Conclusions  of  a  general  nature  are  feasible  with  respect  to  ligand  chain 
length,  presence  or  absence  of  certain  functional  groups,  multiple  bonding, 
isomerism,  free  or  bound  state  of  a  ligand,  and  degree  of  molecular  sym- 
metry. 

Duval  and  his  co-workers52- 53  have  made  many  valuable  contributions 
to  the  study  of  complexes  by  the  use  of  infrared  absorption  measurements. 
In  examining  a  large  number  of  hexacovalent  cobalt  and  chromium  am- 
mines,  Duval  found  that  nearly  all  of  them  absorbed  in  three  principal 
regions.  The  first  region,  quite  intense,  extends  between  800  and  850  cm-1 
for  the  cobalt  complexes,  and  appears  at  about  770  cm-1  for  the  chromium 
complexes.  Duval  attributes  this  absorption  to  triply  degenerate  vibration 
of  the  complex  as  a  whole,  in  the  case  of  hexammines,  and  to  doubly  de- 
generate vibration  in  the  case  of  pentammines.  A  second  prominent  region 
of  absorption,  near  1300  cm-1,  is  considered  to  be  due  to  deformation  vi- 
bration of  the  ammine  groups.  A  third  region,  extending  from  1500  to 
1600  cm-1,  shows  variable  and  generally  less  intensity.  This  absorption 
region  results  from  various  molecular  effects,  depending  upon  the  nature 
of  the  complex. 

The  work  of  Freymann54  illustrates  the  phenomenon  of  dissimulation. 
The  absorption  band  characteristic  of  a  trivalent  nitrogen  atom,  bound 
to  at  least  one  hydrogen  atom,  is  found  in  the  spectra  of  ammonia  and 
amines.  If  the  nitrogen  atom  forms  a  coordinate  bond,  thus  becoming 
quaternary,  the  band  for  the  trivalent  atom  weakens  or  disappears.  Thus 
ammine  complexes,  as  well  as  ammonium  salts,  do  not  show  the  trivalent 
absorption.  Freymann 's  measurements  of  a  number  of  ammines  of  copper, 
cobalt,  platinum,  silver,  and  rhodium  show  the  consistent  dissimulation 
of  the  trivalent  band  in  the  spectra  of  these  complexes. 

.")_'.   Duval,  Duval  and  Lecomte,  Bull.  soc.  chim.  France,  1048  (1947). 

53.  Duval,  Duval  and  Lecomte,  Compt.  rend.,  224,  1632  (1947). 

.">t.  Freymann,  Freymanii  and  Rumpf,  J.  phys.  radium,  7,  30  (1936);  Freymann, 
Ann.  chim.,  11,  40  (1939);  Freymann  and  Mathieu,  Bull.  soc.  chim.,  (5)  4,  1297 
(1937);  Freymann  and  Freymann,  Proc.  Indian  Acad.  Sci.,  8A,  301  (1938). 


PHYSICAL   METHODS  IN  COORDINATION  CHEMISTRY  577 

Duval,  Freymann,  and  Lecomte48  have  measured  the  infrared  absorption 
of  powdered  acetylacetone  derivatives  of  beryllium,  magnesium,  aluminum, 
Bcandium,  samarium,  chromium,  iron(III),  cobalt(II),  cobalt  (III),  cop- 

per(II),  and  zinc.  Whereas  in  acetylacetone  itself  both  the  keto  and  enol 
structures  are  evidenl  from  infrared  absorption,  with  the  metal  salts  only 
the  enol  form  of  the  Ligand  could  be  detected.  The  C=0  group,  which 
normally  absorbs  in  the  range  1710  to  1730  cm-1,  is  evidently  modified 
through  chelation  so  that  a  large  degree  of  single-bond  character  results, 
and  a  shift  of  electron  density  toward  the  metal  strengthens  the  coordinate 
structure. 

Infrared  evidence  was  used  by  Busch  and  Bailar56  to  confirm  the  exist- 
ence of  a  cobalt  (III)  complex  containing  ethylenediam  inetetraacetic  acid 
(EDTA")  as  a  hexadentate  ligand.  The  free  acid  shows  a  maximum  of  ab- 
sorption at  1697  cm-1,  attributable  to  the  carbonyl  structure  in  the  four 
carboxy]  groups,  which  are  normally  associated  through  hydrogen  bonding. 
The  complexes  Xa[Co(EDTA)Br]  and  Na[Co(EDTA)N02],  in  which 
EDTA  is  pentadentate,  were  found  to  exhibit  two  carboxyl  absorptions 
each,  at  1G35  and  1740  cm-1  for  the  nitro  complex,  and  1628  and  1723  cm-1 
for  the  bromo  complex.  The  lower-frequency  absorptions  may  be  ascribed 
to  the  three  complexed  carboxyl  groups,  while  the  single  free  group  is  re- 
sponsible for  the  somewhat  weaker  higher-frequency  bands.  The  barium 
salt  of  the  bromo  complex  was  ground  with  silver  oxide  to  remove  the 
bromine  and  induce  the  free  carboxyl  group  to  coordinate.  The  resulting 
hexadentate  complex  shows  only  one  carbonyl  absorption  band,  at  1G38 
cm-1,  which  may  be  assigned  to  the  four  equivalent  coordinated  carboxyl 
groups. 

A  frequent  problem  in  infrared  absorption  studies  is  the  choice  of  a  suit- 
able solvent.  Since  solvent  molecular  vibrations,  particularly  those  arising 
from  hydrogen  bonding,  may  interfere  with  the  absorption  of  the  sub- 
stance studied,  samples  are  frequently  suspended  or  emulsified  in  a  medium 
such  as  Nujol.  A  significant  development  in  the  technique  of  sample  prep- 
aration is  the  solid  disk  method  of  Stimson  and  O'DonnelF.  If  a  solid 
complex  compound  is  finely  ground,  mixed  intimately  with  potassium 
bromide  in  the  same  state,  and  subjected  to  a  high  mechanical  pressure,  a 
transparent  solid  mass  results.  This  solid  may  be  quite  conveniently  han- 
dled and  examined  spectrophotometrically. 

The  solid  disk  technique  has  been  used  to  advantage  by  Quagliano  and 
his  co-workers.  Fausi  and  Quagliano68  report  that  the  cis  and  trans  forms 
of  dinitrotetrainminecobalt(III)  chloride,  examined  as  solid  disks,  show 
different  infrared  absorptions.  The  cm  isomer  shows  a  greater  multiplicity 

55.  Duval,  Freymann,  and  Lecomte,  Bull.  soc.  ekim.  Frana  .  1952,  106. 

56.  Busch  and  Bailar.  ./.  Am.  Chem.  Soc.t  75,  1674     1863). 

57.  Stimson  and  O'Donnell,  ./.  .1///.  Chem.  8oc.,  74,  L805  (1952). 

58.  Faust  and  Quagliano,  ./.  Am.  Chem.  Soc,  76,  5346  (1954). 


578  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

of  absorption  peaks  than  does  the  trans  isomer.  This  result  is  concordant 
with  the  antisymmetric  nature  of  infrared  absorption,  inasmuch  as  the 
cis  isomer  has  a  lesser  degree  of  symmetry. 

Mizushima,  Sen,  Curran,  and  Quagliano59  have  measured  the  infrared 
absorption  characteristics  of  the  glycine  complexes  of  copper,  nickel,  and 
cobalt.  The  free  carboxyl  group  in  glycine  hydrochloride  absorbs  strongly 
at  5.85/x,  whereas  the  carboxylate  group  in  potassium  glycinate  absorbs 
strongly  at  6.35/x.  The  copper,  nickel,  and  cobalt  glycinates  absorb  strongly 
in  the  6.3-6. 5/x  region,  but  not  at  all  at  5.9/*.  The  resonance  of  the  negative 
carboxylate  is  evidently  preserved  in  the  complexes,  with  the  metal-oxygen 
bond  being  virtually  completely  electrostatic.  On  the  other  hand,  the 
nitrogen  band  in  potassium  glycinate  found  at  3.1  /z,  is  shifted  in  the  copper, 
nickel,  and  cobalt  complexes;  copper  glycinate  absorbs  at  3.22  fi,  and  cobalt 
and  nickel  glycinates  at  3.30  /x.  Evidently  the  metal-nitrogen  bonds  in  these 
complexes  are  primarily  covalent. 

Infrared  evidence  for  symmetrical  platinum-olefin  coordinate  bonds  has 
been  presented  by  Chatt  (p.  504). 

Raman  Spectra 

The  emission  spectra  resulting  from  the  Raman  effect  are  attributable 
to  molecular  vibrations  which  are  symmetrical  in  nature.  Raman  spectra 
thus  complement  infrared  spectra  as  means  of  studying  molecular  struc- 
ures.  Because  of  the  complexity  of  most  molecules  studied  by  the  Raman 
technique,  many  symmetric  effects  arise  from  coupling  of  simpler  individual 

!  vibrating  systems.  Usually,  therefore,  both  the  Raman  and  infrared  methods 

yield  significant  data  concerning  molecular  structures,  and  these  data  in 
some  cases  overlap.  Frequently  it  is  necessary  to  use  crystallographic 
methods  in  order  to  choose  among  several  structures,  each  of  which  is 
compatible  with  Raman  measurements. 

The  Raman  effect  is  produced  when  a  molecule  is  irradiated  with  a  beam 
of  monochromatic  light  of  wave  length  greater  than  the  size  of  the  mole- 
cule. The  radiation  undergoes  interaction  with  the  molecule,  loses  some  of 
its  energy,  and  then  scatters.  The  wave  length  of  the  scattered  light  is 
greater  than  that  of  the  incident  light  unless  the  molecule  is  in  an  excited 
state.  The  scattered  light  may  be  passed  through  a  spectrometer  and  re- 
ceived on  a  photographic  plate.  The  spectrum  on  the  plate  contains  a 
strong  central  line  corresponding  to  the  incident  beam,  and  removed  at 
various  distances  are  the  less  intense  Raman  lines.  The  differences  in 
energy  result  from  a  distribution  of  frequencies  among  the  various  degrees 
of  freedom  of  the  molecule. 

59.  Mizushima,  Sen,  Curran,  and  Quagliano,  Abstracts  of  Papers,  Am.  Chem.  Soc, 
124th  Meeting.  Sept.  6-11,  1953,  43R;  /.  Am.  Chem.  Soc.,  77,  211  (1955). 


PHYSICAL  METHODS  IN  COORDINATION  CHEMISTRY  579 

Frequency  shifts  of  Raman  lines  from  the  frequency  of  the  principal  line 
are  the  quantities  of  significance  in  use  of  the  method.  The  numerical 
values  of  these  shifts  an4  in  the  same  range  as  the  frequencies  of  infrared 
absorption.  It'  a  molecular  vibrational  system  is  characterized  by  symmetric 
and  antisymmetric  vibrations  of  equal  energies,  its  Raman  spectrum  shows 
a  shift  equal  in  magnitude  to  the  corresponding  absorption  frequency  in 
the  infrared  spectrum.  The  mathematical  theory  of  the  Raman  effect  shows 
that  any  Raman  emission  may  be  completely  described  by  measurement  of 
its  frequency  shift,  its  intensity,  and  a  third  coordinate,  called  degree  of 
depolarization. 

Krishnamurti60  used  the  Raman  method  to  study  the  formation  of  chloro 
romplexes  of  mercury.  A  strong  Raman  line  (frequency  shift  =  Av  =  269 
cm-1)  is  observable  with  solutions  of  mercury (II)  chloride  and  ammonium 
chloride  in  a  1:2  molar  ratio.  This  line  compares  with  the  strong  line 
(Av  =  273  cm-1)  for  solid  ammonium  tetrachloromercurate(II)  and  indi- 
cates the  formation  of  the  ion  [HgCl4]=  in  solution.  Solutions  containing 
varying  ratios  of  mercury (II)  bromide  and  alkali  bromide  show  Raman 
shifts  ascribed  to  the  formation  of  [HgBr3]~  and  [HgBr4]=.  Both  complexes 
have  been  depicted  as  tetrahedral  structures  by  Delwaulle61.  The  mercury 
ion  occupies  a  central  position  in  [HgBr4]=  and  a  vertex  in  [HgBr3]~. 

An  extensive  investigation  of  the  structures  of  complexes  has  been  carried 
out  by  Mathieu  and  Cornevin62.  These  authors  measured  the  Raman  spectra 
for  many  complexes.  It  was  found  that  complexes  of  different  metals  which 
have  similar  structures  and  bond  types  yield  similar  Raman  lines.  The 
authors  classified  the  observed  frequency  shifts  into  two  general  groups — 
those  arising  from  metal-ligand  bonds,  and  those  arising  from  the  vibrations 
of  the  coordinated  groups  themselves.  The  second  class  of  shifts  contains 
those  characteristic  of  uncoordinated  ligands,  as  well  as  those  appearing 
only  on  coordination. 

A  number  of  applications  of  the  Raman  method  have  been  made  in  the 
study  of  metal  complexes  of  unsaturated  hydrocarbons.  Nesmeyanov63  has 
reported  data  for  the  compound  CICHCH-HgCl,  proposing  both  the 
structures  [Hg(ClCH=CH)Cl]  and  [Hg(CH  =  CH)Cl]Cl.  Taufen  and  his 
co-workers64  have  suggested  that  complex  formation  between  unsaturated 
hydrocarbons  and  silver(I),  copper(I),  mercury(II),  and  platinum(II)  ions 
accounts  for  the  marked  alterations  in  the  Raman  spectra  of  the  hydro- 
carbons when  these  metal  ions  are  present.  The  hydrocarbons  used  by 

60.  Krishnamurti,  Indian  J.  Physics,  6,  7  (1931). 

61.  Delwaulle,  Francois,  and  Wiemann,  Compt.  rend.,  206,  1108  (1938);  207,  340 

(1938). 

62.  Mathieu  and  Cornevin,  J.  chim.  phys.,  36,  271  (1939). 

63.  Nesmeyanov,  Bull.  acad.  sci.,  U.R.S.S.,  class.  Set.  chim.,  239  (1945). 

64.  Taufen,  Murray,  and  Cleveland,  J.  Am.  Chem.  Soc,  63,  3500  (1941). 


580  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Taufen  with  silver(I)  ion  were  ds-2-butene,  trans-2-butene,  cyclopentene, 
cyclohexene,  ethylacetylene,  propylacetylene,  and  phenylacetylene.  The 
presence  of  the  metal  ion  lowers  the  strong  olefinic  frequency  shift  by  65 
cm-1  and  the  acetylenic  shift  by  100  cm-1. 

It  has  been  found  by  Mathieu65  that  Raman  spectra  provide  no  positive 
differentiation  between  square  and  octahedral  configurations  of  the  plati- 
num and  rhodium  ammines.  Spacu66  has  reported  different  Raman  spectra 
for  the  cis  and  trans  isomers  of  [Pt(NH3)2  py2]Cl2 ,  but  identical  spectra  for 
the  isomers  of  [Pt  py2  Cl2]  and  [Co  en2  (N02)2]N03 .  It  seems  reasonable,  in 
view  of  the  differences  of  degree  of  symmetry  of  these  cis-trans  isomers, 
that  differences  in  the  spectra  actually  exist,  although  the  distinguishing 
lines  may  be  so  weak  that  they  have  escaped  detection. 

Venkateswaran67  used  Raman  data  to  study  the  symmetry  of  a  number 
of  complexes  of  the  type  [MOn],  as  well  as  the  azide  ion.  Telluric  acid  was 
found  to  be  octahedrally  symmetrical  in  agreement  with  the  formula 
H6[TeOe].  Tetrahedral  structures  were  confirmed  for  Cr04=,  Mo04=,  W04=, 
and  I04~,  pyramidal  structures  for  C103~  and  Br03~,  and  a  linear  structure 
for  N3~~.  Raman  spectra  of  solid  NaRe04  and  KRe04 ,  studied  by  Fonteyne68, 
show  a  distorted  tetrahedral  arrangement,  changing  in  water  solution  to 
the  octahedral  [Re06]5-  complex. 

The  infrared  spectral  studies  of  Crawford  and  Cross,  and  the  Raman 
spectral  studies  of  Crawford  and  Horiwitz,  each  of  which  supports  the 
postulated  tetrahedral  structure  of  nickel  tetracarbonyl,  have  been  cited 
in  Chapter  16  (p.  519). 

Optical  Methods 
Polarimetry 

The  ability  of  a  substance  to  rotate  a  beam  of  plane  polarized  light  is  a 
function  of  molecular  or  crystalline  asymmetry.  Optical  activity  of  co- 
ordination compounds  is  almost  exclusively  due  to  molecular  asymmetry 
which  persists  in  solution. 

Rotation  of  polarized  light  is  detected  and  measured  with  the  polarime- 
ter.  Solutions  of  varying  concentrations  may  constitute  the  sample.  Greater 
concentrations  produce  a  larger  observed  rotation,  but  in  many  cases  the 
intense  colors  of  the  solutions  prevent  sufficient  transmission  of  the  po- 
larized beam  unless  very  strong  light  sources  or  solutions  of  low  concentra- 
tion are  used.  A  substance  whose  solution  rotates  polarized  light  in  a  clock- 
wise direction  is  said  to  be  dextrorotatory,  and  one  giving  counterclockwise 

65.  Mathieu,  Compt.  rend.,  204,  682  (1937). 

66.  Spacu,  Bull.  soc.  chim.  (5)  4,  364  (1937). 

67.  Venkateswaran,  Proc.  Indian  Acad.  Sci.,  7A,  144  (1938). 

68.  Fonteyne,  Natuurw.  Tijdschr.,  20,  20  (1938);  20,  112  (1938). 


PHYSICAL  METHODS  IN  COORDINATION*   CHEMISTR1  583 

rotation  is  called  Levorotatory.  The  two  optical  isomers  <>t"  the  complex 
arc  referred  to  as  the  d  and  /  forms  according  to  the  sign  of  rotation.  Dex- 
trorotation is  assigned  a  plus  value. 
It  should  be  emphasized  that  the  sign  of  rotation  cannot  1><>  used  to  find 

absolute  configurations  of  complex  substances.  Different  species  with  the 
same  sign  of  rotation  may  have  the  same  or  opposite  configurations;  indeed, 
the  sign  and  degree  of  rotation  of  any  given  complex  usually  vary  with  the 
wave  length  of  the  light  source.  This  variation  is  often  of  much  greater  use 
in  elucidating  structures  than  are  isolated  rotational  measurements  at 
single  wave  lengths. 

An  important  polarization  phenomenon  in  structural  studies  of  com- 
plexes is  the  Cotton  effect69, 70.  A  normal  rotatory  dispersion  curve,  or  plot 
of  magnitude  of  rotation  against  wave  length  of  incident  light,  is  hyperbolic 
in  form.  The  Cotton  effect  is  evidenced  by  an  abnormality  in  rotation  in 
the  vicinity  of  maximum  light  absorption  of  the  complex.  This  abnormality 
is  generally  characterized  by  a  maximum  of  rotation,  a  sharp  decrease  to 
zero  rotation,  and  an  increase  in  rotation  of  the  opposite  sign.  All  these 
variations  take  place  with  a  small  change  in  wave  length71.  Mellor72  has 
reported  a  relationship  between  the  Cotton  effect  and  the  magnetic  mo- 
ments of  several  nickel,  copper,  and  cobalt  chelates.  The  effect  evidently  is 
found  only  among  complexes  of  the  covalent  type.  Pfeiffer73  attributes  the 
Cotton  effect  in  certain  heavy  metal  tetracovalent  complexes  to  the  chro- 
mophobe nature  of  the  central  metal  atom.  Mathieu74  states  that  the  pres- 
ence of  asymmetric  carbon  atoms  in  ligands  produces  a  Cotton  effect  by 
vicinal  influence,  but  Pfeiffer's  work  shows  no  evidence  of  such  influence, 
so  long  as  the  dispersion  curves  of  the  ligands  are  normal.  A  variation  in 
vicinal  influence  with  bond  lengths  may  well  account  for  this  difference. 

The  effect  of  asymmetric  molecules,  not  necessarily  coordinated,  in  pro- 
ducing anomalous  rotations  in  solutions  of  complexes  is  termed  asym- 
metric induction.  Pfeiffer  and  Quehl75  noted  that  the  optical  rotation  of 
zinc  d-a-camphor-^-sulfonate  is  reduced  nearly  to  zero  upon  addition  of 
three  moles  of  o^/io-phenanthroline  per  mole  of  zinc.  Likewise,  the  specific 
rotation  of  zinc  c/-a-bromocamphor-7r-sulfonate  is  4.55°,  but  that  of  tris 
(o-phenanthroline)zinc  c?-a-bromocamphor-7r-sulfonate  is  8.44°.  Active  cat- 

69.  Jaeger,  "Optical  Activity  and  High  Temperature  Measurements,"  New  York, 

McGraw-Hill  Book  Co.,  1030. 

70.  Cotton,  Ann.  chim.  phys.,  8,  317  [1896). 

71.  Bruhat,  Bull.  eoc.  chim.,  17,  223  (1915). 

72.  Mellor,  ./.  Proc.  lion.  Soc.  A  .  8.  Wales,  75,  157  (1942). 

7::.  Pfeiffer,  Christeleit,  Hosse,  Pfitzner,  and  Thielert,  •/.  Prakt.  Chew.,  150,  261 

(1938). 
71.  Mathieu,  .1/'//.  phy8.t  19,  336  (1944  . 
75.  Pfeiffer  and  Quehl,  Ber.,  64B,  2667  (1931);  65B,  560  (1932). 


582  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

ions  do  not  exercise  the  inductive  effect  in  these  instances.  The  findings  of 
Pfeiffer  and  Quehl  have  been  confirmed  by  Brasted76.  Biswas77  has  observed 
a  similar  effect  of  d-tartaric  acid  in  molybdic  acid  solutions.  Dwyer78  has 
done  extensive  work  with  this  effect,  using  racemic  complexes  whose  active 
forms  are  optically  stable,  as  well  as  those  having  optically  labile  active 
forms.  Addition  of  an  asymmetric  substance  such  as  bromocamphorsul- 
fonate  to  a  racemic  complex  appears  to  affect  the  rotatory  powers  of  the 
d  and  I  forms  of  the  complex  by  different  amounts,  thus  producing  a  net 
rotation  different  from  zero.  Another  change  consists  of  a  shift  in  the  equi- 
librium of  the  isomers  away  from  the  normal  1 : 1  ratio.  This  second  change 
may  be  immediate  or  slow,  and  it  further  affects  the  observed  rotation. 
These  effects  Dwyer  attributes  to  alterations  of  the  thermodynamic  activi- 
ties of  the  isomers  in  the  presence  of  the  asymmetric  substance. 

Determinations  of  structure  from  polarimetric  data  usually  involve 
analysis  of  rotatory  dispersion  curves.  Mathieu79  has  shown  that  if  two 
complexes  of  analogous  composition  yield  curves  characterized  by  the 
Cotton  effect,  those  portions  of  the  curve  displaying  the  effect  will  have 
slopes  of  the  same  sign  if  the  complexes  have  the  same  configuration.  If 
the  configurations  are  opposite,  the  slopes  of  the  dispersion  curves  will 
have  opposite  signs  in  the  area  of  the  Cotton  effect. 

An  empirical  rule  of  Werner80  states  that  optically  active  ions  of  the  same 
configuration,  when  crystallized  with  the  same  optically  active  substance 
(e.g.,  d-tartrate),  will  have  analogous  solubilities,  either  both  less  or  both 
greater  than  the  compounds  of  their  respective  antipodes.  This  rule  has 
been  applied  by  Jaeger81  in  his  investigation  of  diamine  complexes  of  co- 
balt, rhodium,  and  chromium.  Delepine82  has  noted  that  the  active  isomers 
of  certain  complexes  may  crystallize  in  forms  which  are  different  from 
those  of  the  racemic  crystal  of  the  same  complex.  The  crystals  of  complexes 
of  the  same  chemical  type,  containing  different  metals,  sometimes  show  the 
same  differences  in  crystal  form  between  active  crystals  and  racemic  crys- 
tals. In  such  cases  the  active  forms  of  the  complex  of  the  one  metal  are 
generally  isomorphous  with  the  active  forms  containing  the  other  metal. 
Similarly,  the  racemic  crystals  are  isomorphous  with  each  other.  But  a 
crystal  may  also  be  formed  by  the  d  isomer  containing  the  first  metal, 

76.  Brasted,  Thesis,  University  of  Illinois,  1942. 

77.  Biswas,  J.  Indian  Chemical  Soc,  22,  351  (1945). 

78.  Dwyer  and  Gyarf as,  J.  Proc.  Roy.  Soc.  N.  S.  Wales,  83,  170  (1949) ;  Dwyer,  Gyar- 
fas,  and  O'Dwyer,  Nature,  167,  1036  (1951). 

79.  Mathieu,  </.  chim.  phijs.,  33,  78  (1936) ;  Bull.  soc.  chim.,  [5],  4,  687  (1937). 

80.  Werner,  Ber.,  45,  121,  1228  (1912). 

81.  Jaeger,  Proc.  Acad.  Sci.  Amsterdam,  40,  2  (1937);  Jaeger  and  Bijkerk,  Proc. 
Acad.  Sci.  Amsterdam,  40,  116  (1937). 

82.  Delepine,  Bull.  soc.  chim.,  [4],  29,  656  (1921);  [51, 1,  1256  (1934). 


PHYSICAL  METHODS  I\   COORDINATION*   CHEMISTRY 

and  the  /  isomer  containing  the  Becond  metal.  This  crystal  has  the  habit  of 
the  racemateSj  but  it  is  optically  active,  Bince  the  two  metal-  do  nol  in 
genera]  form  analogous  complexes  with  exactly  the  same  rotational  values. 
Such  crystals  are  termed  "active  racemates"  by  Delepine.  [f  one  of  the 
constituents  of  the  active  racemate  has  a  known  configuration,  the  other 
may  be  considered  to  have  tin1  opposite  configuration.  This  method  of 
determining  relative  configurations  is  clearly  limited,  since  only  complexes 
of  similar  size  and  chemical  type  are  isomorphous. 

Polarimetric  observations  enabled  Dwyer81  to  verify  bis  asymmetric 
synthesis  of  an  iron(III)  cationic  complex,  the  firsl  such  preparation  to  be 
reported.  By  oxidizing  one  of  the  isomers  of  tris(dipyTidyl)iron(II)  per- 
chlorate  with  cerium(TV  ammonium  nitrate  solution,  then  adding  sodium 
perchlorate  in  excess,  Dwyer  was  able  to  precipitate  blue  crystals  of  op- 
tically active  [Fe(dipy)j](C104V3HjO. 

Refractometry 

Refractometric  measurements  of  solutions  may  be  used  in  applying  the 
continuous  variations  method  of  Job.  The  work  of  Spacu  and  Popper84   is 

outstanding  in  this  field.  These  authors  have  reported  refractometric  evi- 
dence for  existence  of  acetato,  tartrate,  and  citrato  complexes  of  aluminum, 
as  well  as  such  complexes  as  [HgCl3]~,  [HgCl5]-,  [CdBr5]=,  [BaCl4]=,  and 
numerous  others.  Refraction  data  have  also  led  Spacu  and  Popper  to  assign 
the  nitrile  .structure  to  potassium  cyanide,  potassium  tbiocyanate,  and 
potassium  selenocyanate.  Criticism  of  the  broad  conclusions  of  Spacu  and 
Popper  has  been  advanced  by  Haldar*5,  Tahvonen86,  and  Grinberg87,  who 
dispute  the  original  authors'  use  of  additive  refraction  values  for  certain 
functional  groups.  While  the  contributions  of  constituent  groups  in  a  mole- 
cule to  the  molecular  refraction  are  roughly  additive,  care  must  be  exercised 
in  drawing  highly  specific  conclusions  from  refraction  data. 

The  nature  of  complex  ions  in  highly  concentrated  solutions  of  the  metal 
ion  and  Ligand  (as  cadmium  ion  and  cyanide  ion)  have  been  examined  by 
Brasted.  A  plot  of  direct  dipping  refractometer  readings  vs.  mole  fraction  of 
metal  ion  shows  a  maximum  at  the  point  of  .-table  complex  ion  formation. 
The  sharpness  of  this  peak  is  indicative  of  the  stability  of  the  complex. 
Addition  of  cyanide  ion  solution  to  cadmium  ion  solution  -both  at  2M 
concentration)  indicate-  by  the  sharp  maximum  the  species  [<  'd|CX)4]=. 

83.  Dwyer  ai  J4, 

-    •■■a  and  Popper,  Bull.  we.  stiinU  Civ  ,8,5    L934   ;  7,  ■   ,KolloidZ.t 

103,  19    1943   .  /  .  A180.  154    1937   ;  A182.     - 

85.  Haldar,  ./.  Indian  <  23.  206    Ifl 

96    Tahvonen,  \d.  8ci.  7  •  A49,  No.  6,  No.  7 

87.  <  -   /■      .  /'  iklad. Kkim., tl,  W 


584  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

At  such  high  concentrations  optical  or  spectrographs  methods  would  not 
in  genera]  be  applicable. 

Electrombtric  Methods 
Polarography* 

The  polarograph  received  wide  use  in  analytical  chemistry  immediately 
following  its  invention  by  Heyrovsky  and  Shikata88  in  1925,  but  not  until 
ten  years  later  did  its  usefulness  in  coordination  chemistry  become  signifi- 
cant. Among  the  important  quantitative  data  obtainable  by  polarographic 
means  are  dissociation  constants  of  complexes,  coordination  numbers  of 
metal  ions,  and  the  degree  of  stabilization  of  various  oxidation  states. 

Polarographic  studies  are  carried  out  with  an  apparatus  wThich  combines 
an  electrolytic  cell  with  a  recording  device.  The  usual  cell  is  composed  of  a 
dropping  mercury  cathode,  a  mercury  pool  anode,  and  a  solution  containing 
a  known  concentration  of  the  substance  to  be  studied  and  an  indifferent, 
or  supporting,  electrolyte.  The  recording  device  plots  current  as  ordinate 
against  a  continuously  increasing  potential  as  abscissa.  Direct  current 
sources  are  usual,  although  alternating  current  has  been  used  to  advantage. 
A  typical  analysis  may  involve  the  reduction  of  a  complex  cation  in  solution. 
As  the  electrolysis  begins,  the  potential  is  chosen  less  than  the  reduction 
potential  of  the  species  in  solution.  The  current  flowing  through  the  cell  is 
small.  So  long  as  the  cell  potential  is  less  than  the  reduction  potential  of 
the  complex  ion,  this  current  remains  practically  constant.  The  recording 
device  traces  a  nearly  horizontal  line.  Since  the  growth  and  fall  of  each 
mercury  drop  causes  a  slight  oscillation  in  the  current  value,  the  actual 

iplot  is  a  composite  of  many  waves  of  small  amplitude,  tracing  the  over-all 
horizontal  line.  When  the  reduction  potential  (decomposition  potential)  is 
reached,  a  sharp  rise  in  the  current  occurs  with  reduction,  usually  to  the 
metallic  state,  with  amalgamation  of  the  previously  complexed  metal  with 
the  cathode.  Mercury  ionizes  correspondingly  at  the  anode.  The  current 
continues  to  increase  with  increasing  potential,  but  a  limiting  value  is 
reached  in  unagitated  systems.  As  electrolysis  proceeds,  the  concentration  of 
reducible  material  falls  in  the  immediate  vicinity  of  the  cathode.  Then  more 
reducible  material  diffuses  from  the  body  of  the  solution  to  the  cathode. 
The  rate  of  diffusion  depends  upon  the  concentration  gradient  between  the 
solution  proper  and  the  reducing  area  near  the  surface  of  the  cathode.  The 
potential  eventually  reaches  a  value  corresponding  to  a  negligible  concen- 
tration next  to  the  cathode,  the  substance  being  reduced  virtually  instantty 
upon  diffusion.  Then  the  rate  of  diffusion  becomes  constant  and  essentially 

*  The  presentation  of  much  of  the  material  in  this  section  was  suggested  by  Dr. 
II.  V.  Boltzclaw  of  the  University  of  Nebraska. 
88.  Heyrovsky  and  Shikata,  Rec.  trav.  chim.,  44,  496  (1925). 


PHYSICAL  METHODS  l\  COORDINATION  CHEMISTRY  585 

independent  of  further  potential  increase,  hut  dependent  on  the  concentra- 
tion of  reducible  substance  in  the  solution  proper.  The  current  assumes  the 
limiting  value  under  these  conditions,  and  the  current  and  rate  of  diffusion 
may  be  seen  to  be  proportional  to  the  concentration  of  reducible  substance. 
Strictly  considered,  the  migration  of  ions  also  contributes  to  the  limiting 
current,  hut  in  the  presence  o\'  a  comparatively  large  amount  of  indifferent 
electrolyte,  the  limiting  current  is  due  nearly  entirely  to  diffusion;  it  is 
therefore  known  as  the  diffusion  current  (id). 

When  the  current  has  reached  a  value  one-half  that  of  the  Limiting  cur- 
rent, the  corresponding  potential  is  the  half-wave  potential  (E{).  This 
potential  is  the  characteristic  value  sought  for  the  substance  under  study, 
and  it  is  independent  of  concentration  and  type  of  electrode.  If  several 
substances  are  present  and  electroactive,  each  may  be  determined,  provided 
no  two  half-wave  potentials  are  closer  than  0.2  volts.  The  total  range  of  the 
dropping  mercury  electrode  is  taken  as  +0.6  volts  to  —2.6  volts  against 
the  standard  calomel  electrode.  In  most  solutions  the  full  range  is  not 
realizable.  The1  substance  to  be  studied  must  be  in  true  solution  and  must 
be  resistant  to  oxidation,  reduction,  and  decomposition  from  outside 
sources.  Cations  and  anions,  oxidizable  and  reducible  materials,  and  simple 
and  complex  ions  may  be  studied  by  appropriate  applications  of  the  polaro- 
graphic  method. 

A  number  of  factors  may  affect  the  electrolysis  and  alter  the  recorded 
curve.  In  this  discussion  the  most  important  factor  is  the  presence  of  com- 
plexes. Normally  a  complexed  ion  resists  the  electrolytic  reduction  more 
than  the  corresponding  uncomplexed  ion,  and  the  half-wrave  potential  is 
more  negative  for  the  complex.  The  pH  of  the  solution  may  affect  the  half- 
wave  potential  either  by  altering  the  nature  of  complexes  or  by  varying 
the  products  of  the  electrolysis.  In  the  presence  of  agar,  gelatin,  or  other 
capillary-active  substances,  undesirable  maxima  in  curves  may  often  be 
avoided;  however,  these  materials  may  alter  the  diffusive  properties  of  the 
ions  present,  thus  affecting  the  diffusion  current.  Supporting  electrolytes 
which  supply  coordinating  groups  may  deter  the  decomposition  of  complex 
ions  and  thus  bring  about  a  more  negative  half-wave  potential. 

The  polarographic  method  is  unique  among  electrometric  methods  in  that 
only  a  small  fraction  of  the  solution  is  electrolyzed.  A  further  advantage  is 
that  quite  small  concentrations  of  the  material  to  be  studied  are  sufficient. 
Among  the  favorable  features  of  the  dropping  mercury  electrode  are  its 
smooth,  reproducible,  and  renewable  surface;  ready  ascertainment  of  the 
surface  area  of  the  drops;  the  ability  of  nearly  all  metals  to  amalgamate 
with  mercury;  and  the  high  overvoltage  for  hydrogen  liberation  on  mercury, 
so  that  electrolysi.-  of  hydrogen  ions  is  minimized. 

A  thorough  treatment  of  the  methods  of  polarography  is  given  by  Kolt- 


586  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

hoff  and  Lingane89.  Pertinent  discussions  of  the  theory  and  application  of 
polarography  are  noted  in  references  89  and  90.  In  the  following  treatment 
of  applications,  no  effort  has  been  made  to  derive  mathematical  relations. 
For  convenience  Heyrovsky  and  Ilkovic91  separate  the  reduction  of  a 
metal  complex  into  two  reactions, 

MXpt"-p»  z±  M»+  +  pXb~,  (I) 

M"+  +  ne-  +  Hg  ^±  M(Hg),  (II) 

where  X  is  the  complexing  agent  and  M(Hg)  symbolizes  the  amalgam 
formed  on  the  surface  of  the  electrode.  These  reactions  may  or  may  not 
actually  occur  as  written,  but  they  serve  as  convenient  references.  The 
dissociation  constant  of  the  complex  is  given  by 

i:  K-  ~  [mxp<»-»»]  •  (m) 

This  constant  may  be  calculated  from  the  negative  shift  of  the  half-wave 
potential  upon  complexing,  as  indicated  by 


(EOc  -  (EOs  ^^\nKc-p^\n  [X»i  (IV) 

nF  n¥ 


In  this  formula  the  subscripts  c  and  s  refer  to  the  complex  and  simple 
(hydrated)  ion,  respectively.  Thus  the  difference  between  the  half-wave 
potentials  leads  to  the  determination  of  Kc ,  provided  that  p,  the  coordina- 
tion number  of  the  metal,  is  known.  The  following  formula  is  useful  in 
determining  p  from  half -wave  measurements  at  different  concentrations  of 
complexing  agent. 

A  In  [X*-]Tx  V  nF  ■ 

Usually  assumption  of  the  value  of  unity  for  the  activity  coefficient  ys 
yields  sufficient  accuracy. 

Kolthoff  and  Lingane90d  point  out  that  Equation  (IV)  is  not  a  good 
approximation  when  the  rates  of  diffusion  of  the  simple  and  complex  ions 
are  appreciably  different.  In  such  cases  the  ratio  of  the  diffusion  coefficients 
enters  the  calculation.  Sometimes  a  state  of  equilibrium  is  not  rapidly 
reached,  and  the  calculations  suffer  further  losses  in  accuracy.  Pines92, 

89.  Kolthoff  and  Lingane,  "Polarography,"  New  York,  Interscience  Publishers, 

Inc.,  (1946). 

90.  Muller,  J.  Chem.  Ed.,  18,  C5,  320  (1941) ;  Page,  Nature,  154, 199  (1944) ;  Quagliano, 

thesis,  University  of  Illinois,  1946;  Kolthoff  and  Lingane,  Chem.  Rev.,  24,  1 
(1939). 

91.  Heyrovsky  and  Ilkovic,  Collection  Czechoslov.  Chem.  Commun.,  7,  198  (1935). 

92.  Pines,  Collection  Czechoslov.  Chem.  Commun.,  1,  387  (1929). 
Pines,  Chem.  News,  139,  196  (1929). 


PHYSICAL  METHODS  l\   COORDINATION  CHEMISTRY 


587 


CIS    CATION 

rr 

h 

f 

z 

U 

/               TRANS    CATION 

m 

/  / 

cc 

1/ 

a 

)/ 

u 

J                       CIS  OR  TRANS 

i 

/   /""               CATION    OR    ANION 

/                          CIS    ANION 

m 

Z^"                          TRANS    ANION 

3r 

— 

VOLTAGE 
Fig.  IS. 2.   Limiting  currents   and   cis-tians   isomerism.    I:   with   supporting   elec- 
trolyte-diffusion  current  only.  II,  III,  IV,  V:  Without  supporting  electrolyte-diffu- 

sion  and  migration  currents. 


Brocket  and  Petit9*,  Foerster94  and  Herman95  report  delayed  equilibria 
caused  by  slow  dissociation  of  cyano  complexes  of  zinc  and  gold.  Another 
nonideality  factor  is  found  with  stable  complexes  which  reduce  directly 
without  the  dissociation  suggested  by  Equation  (I).  If  the  metal  is  well 
shielded  by  the  eomplexing  groups,  its  capture  of  electrons  from  the  cathode 
may  be  hindered.  The  extra  potential  required  for  reduction  leads  to  error 
in  the  calculated  value  of  the  constant  Kc  . 

Normally  an  excess  of  indifferent  electrolyte  suppresses  any  migration 
current  of  reducible  ions.  In  the  absence  of  an  indifferent  electrolyte,  how- 
ever, the  limiting  current  is  made  up  of  both  diffusion  and  migration  cur- 
rents. This  fact  is  useful  in  differentiating  between  cis  and  trans  forms  of 
complexes  of  the  type  [MAA]"*.  Both  forms  of  the  complex  migrate  in  the 
Bame  direction,  but  the  greater  rate  of  migration  is  shown  by  the  cis  form. 
which  has  a  dipole  moment  different  from  zero.  An  orientation  attraction 
to  the  electrode  causes  the  cis  form  to  produce  a  higher  limiting  current 
than  the  trans  form  of  both  cationic  and  anionic  complexes.  The  limiting 
current  for  either  form  of  an  anionic  complex  is  less  than  the  diffusion  cur- 
rent because  of  cathodic  repulsion  (see  Fig.  18.295). 

Lindane's  investigation  of  the  biplumbite  ionM  furnishes  a  good  example 
of  polarographic  analysis.  The  object  of  the  study  was  to  determine  the 
Dumber  of  hydroxy]  groups  coordinated  to  lead  in  the  biplumbite  complex. 
Various  concentrations  of  hydroxide  ion  were  used,  and  the  half-wave  po- 
tential corresponding  to  each  was  taken.  With  the  value  of  n  in  Equation 

Broehei  and  Petit,  Z.  Elektroehem.,  10,  900    1904). 
94.   i  ochem.,  13,  561     1907  . 

Herman,  Colit  •  '  Mt.,  6,  37    19 

96.  Lingane,  Chem.  Rev.,  29,  1  (1941;. 


588 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


(V)  taken  as  2,  the  data  are  most  nearly  satisfied  by  p  =  3.  Accordingly, 
Lingane  has  proposed  the  following  equilibria: 


Pb++  +  30H-  ^  [H3Pb03]-  ~ 


H20 


±  [HPb02]- 


+H20 

The  soluble  form  of  lead (II)  in  basic  solution  is  then  evidently  [HPbCy- 
rather  than  [Pb02]=,  which  would  be  in  equilibrium  with  the  four-coordinate 
ion  [H4Pb04]=.  Malyugina  and  his  co-workers97  have  found  a  coordination 
number  of  four  for  lead(II)  and  mercury(II)  in  the  presence  of  iodide  ion. 
The  dissociation  constants  for  [Pbl4]=  and  [Hgl4]=  are  given  as  10~7  and 
10~27,  respectively. 

A  reduction  to  a  lower  oxidation  state  but  not  to  the  metal  takes  place 
with  the  tris(oxalato)iron(III)  ion.  Stackelberg  and  Freyhold98  conclude 
that  the  iron (II)  complex  [Fe(C204)2]=  forms  with  concentrations  of  oxalate 
less  than  Q.15M  in  O.OOlilf  iron(II)  ion  solution.  With  greater  concentra- 
tions of  oxalate,  the  complex  formed  is  [Fe(C204)3]4_.  Toropova"  confirms 
the  existence  of  [Fe(C204)2]=  and  gives  dissociation  constants  for  it  and  for 
[Fe(C204)3]~.  This  reduction  of  complexed  iron  (III)  to  one  of  two  complex 
iron(II)  species  has  also  been  studied  by  Lingane100  and  by  Schaap,Laitinen, 
and  Bailar101.  Their  findings  agree  substantially  with  those  of  Toropova, 
and  of  Stackelberg  and  Freyhold,  the  most  notable  differences  being  in  the 
values  found  for  the  dissociation  constants,  summarized  below. 


Kd,  found  by 

Lingane 

Schaap 

Toropova 

[Fe11  (C204)2]=^  Fe++  +  2C204= 
[Fe11  (C204)3]4-^  Fe++  +  3C204= 
[Fe111  (C204)3]s^  Fe+++  +  3C204= 

8  X  10~6 

6.1  X  10-7 

6  X  10"20 

2.7  X  10-5 
6.1  X  10-6 

1.0  x  io-18 

2.7  X  IO"10 
1.2  X  10~24 

The  tendency  of  polymetaphosphates  to  form  complexes  has  been  studied 
polarographically  by  Caglioti  and  his  co-workers102.  Copper(II)  and  cad- 
mium (II)  ions  do  not  form  such  complexes  under  the  conditions  which 
they  used,  while  zinc(II),  manganese(II),  and  lead(II)  form  unstable  com- 
plexes, and  iron(II)  forms  a  stable  complex. 

Harris  and  Kolthoff103  have  presented  data  which  support  the  following 

'.i7.  Malyugina,  Shchemukova,  and  Korshunov,  J.  Gen.  Chem.,  U.S.S.R.,  16,  1573 
(1946). 

98.  Stackelberg  and  Freyhold,  Z.  Eleklrochem.,  46,  120  (1940). 

99.  Toropova, ./.  Gen.  Chem.,  U.S.S.R.,  11,  1211  (1941). 

100.  Lingane,  ./.  Am.  Chem.  Soc.,  68,  2448  (1946). 

101.  Schaap,  Laitinen,  and  Bailar,  J.  ,1//?.  Chem.  Soc.,  76,  5868  (1954). 
1(12.  Caglioti,  Sartori,  and  Bianchi,  Gazz.  chim.  Hal.,  72,  63  (1942). 
103.  Harris  and  Kolthoff.  J.  Am.  Chem.  Soc,  67,  1484  (1945). 


PHYSICAL   METHODS  IN  COORDINATION  CHEMISTM  589 

reaction  of  the  urany]  ion  in  0.01  to  0.2M  hydrochloric  acid 

[JO  •    <•     .   ■  I  <> 

This  reaction  suggests  that  uranium(V)  compounds  may  be  preparable  in 

acid  solution.  The  compound  UCli  is  known,  hut  its  water  solution  contains 
only  uranium(VI)  and  uranium (II). 

From  studies  of  cyano  and  thiocyanato  complexes  of  rhodium,  Willis104 
concludes  that  complexes  of  rhodium(III)  reduce  first  to  those  of  rho- 
dium(II)  and  then  to  the  metal.  There  is  some  experimental  evidence  for 
tin1  intermediate  formation  of  rhodium  (I),  but  Willis  consider-  its  existence 
questionable.  A  stability  series  for  the  cyano  complexes  of  the  Group  VIII 
metals  has  been  drawn  up  by  Willis.  Relative  shifts  in  half-wave  potentials 
indicate  that  if  the  metals  are  arranged  in  the  usual  periodic  order,  stability 
of  the  cyano  complexes  increases  downward  in  each  column. 


More 
stable 


IV  ll,,  Fe(III) 

Co(III) 

Ni(II) 

Ru(II) 

Rh(III) 

Pd(II) 

Os(II) 

Ir(III) 

Pt(II) 

Less 
stable 


Wheelwright,  Spedding,  and  Schwarzenbach105  have  found  the  polaro- 
graphic  method  useful  in  determining  formation  constants  of  the  heavier 
rare  earth  complexes  of  ethylenediaminetetraacetic  acid  (EDTA).  Meas- 
urements were  made  of  solutions  containing  the  complexing  agent  and  both 
copper(II)  and  a  rare  earth  metal  ion,  in  order  to  determine  the  amount  of 
free  copper(II)  ion  present.  These  data,  the  original  composition  of  the 
solutions,  and  the  known  dissociation  constants  of  the  ligand  and  its  copper 
complex  are  sufficient  for  the  calculation  of  the  formation  constant  Kf  in 
the  expression 

RE+++  +  EDTA<-  ^  [RE  EDTA]-;         Kf 


[RE+++][EDTA< 


All  experimental  work  was  done  at  constant  temperature  and  ionic  strength. 
A  potent  iometric  method  was  employed  as  a  check  and  found  to  be  some- 
what more  precise  for  the  lighter  rare  earths.  Comparative  values  for  the 
formation  constants  arc  listed  below. 

Metal  complex  Kf  (polarographic)  Kf    potentiometric) 

[Ce  EDTA]  15. G     ±  0.  \  L5.39  ±  0.06 

Gd  EDTA  16.6    ±  0.15  16.70  ±  0.08 

[Lu  EDTA  19.65  ±  0.12  19.06  ±  0.  I 

Frank  and  Hume""'  have  studied  the  formation  of  thiocyanate  complexes 

104.  Will.-.  ./.  .1,/,.  Chem.  Soc.,  66,  L067    1944  . 

105.  Wheelwrighl .  Bpedding,  and  Sch*  arzenbach, ./.  .1///.  ( 'Ai  m.  Soc.,  75,  1 196    1953 
inc.  Frank  and  Hume  ./ .  Am.  Chem.  Soc.,  75,  1736    1953 


590  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

of  zinc  in  solutions  containing  zinc  salts,  potassium  thiocyanate,  and  potas- 
sium nitrate.  Half-wave  potentials  of  the  zinc  ion  indicate  the  formation  of 
complexes  containing  up  to  four  thiocyanate  groups  per  zinc  ion.  The  zinc 
complexes  have  been  shown  to  be  much  less  stable  than  their  cadmium 
analogs,  but  the  gradations  of  stability  within  each  series  are  quite  similar. 
A  polarographic  distinction  between  cis  and  trans  forms  of  complexes 
containing  two  negative  groups  has  been  reported  by  Holtzclaw  and 
Sheetz107.  In  the  presence  of  potassium  chloride  as  a  supporting  electrolyte, 
the  cis  form  reduces  at  a  more  positive  potential  than  the  trans  form  for 
the  complexes  [Co(NH3)4(N02)2]+,  [Co  en2  (N02)2]+,  and  [Co  en2  (NCS) 
\<>2]+.  The  ions  [Co  en2  (NH3)N02]++,  [Co  en2  (NH3)NCS]++  and  [Co  en2 
(NH3)2]+++,  which  contain  one  or  no  negative  groups,  do  not  exhibit  this 
difference. 

Electrometric  Titrations ;  Electromotive  Force  Measurements 

Electrometric  titrations  are  generally  classified  into  three  groups:  po- 
tentiometric,  conduct ometric,  and  amperometric.  Potentiometric  titrations 
are  characterized  by  changes  in  the  potential  of  an  electrode  in  the  solution 
which  is  being  examined.  Potentials  are  referred  to  some  standard  electrode 
system.  As  a  titration  proceeds,  a  change  in  concentration  of  the  species 
studied  will  be  reflected  in  a  change  in  electrode  potential,  with  the  equiva- 
lence point  corresponding  usually  to  an  abrupt  potential  shift.  The  meas- 
urement of  pH  by  electrode  methods  is  a  special  application  of  potentio- 
metric theory.  A  hydrogen  electrode  serves  as  the  classical  electrode  for 
pH  measurements,  since  its  potential  variations  are  directly  related  to 
changes  in  hydrogen-ion  activity.  Other  electrodes,  such  as  the  quinhy- 

!  drone  electrode  and  the  glass  electrode,  are  often  more  convenient. 

The  electrode  in  a  potentiometric  titration  is  chosen  appropriately  for  a 
given  titration  reaction.  Since  it  may  be  regarded  as  a  specific  indicator  for 
the  reaction,  it  is  often  called  an  indicator  electrode.  Indicator  electrodes  for 
pH  measurement  have  been  mentioned  above.  Oxidation-reduction  titra- 
tions usually  involve  noble-metal  electrodes  such  as  platinum  wire  or 
platinum  gauze.  Silver  and  mercury  electrodes  are  often  used  in  deter- 
minations of  metal-ion  concentrations. 

Conductometric  titrations  involve  measurement  of  the  conductivity  of 
the  tested  solution  as  the  desired  reaction  proceeds.  In  potentiometric  ti- 
trations, foreign  ions  arc  permissible  so  long  as  they  do  not  affect  the  po- 
tential of  the  indicator  electrode.  In  conductometric  titrations,  however, 
all  ions  present  contribute  to  the  conductivity  and  require  consideration. 
The  equivalence  point  of  a  conductometric  titration  is  not  characterized  by 
an  abrupt  change  in  conductivity,  bul  by  a  change  in  the  slope  of  the  plot 

107.  Ilultzrhu  and  Sheetz.  ./.  Am.  Chem.  Soc,  75,  3053  (1953). 


PHYSICAL   METHODS  1^  COORDINATIOh   CHEMISTRY  591 

of  conductivity  against  volume  of  titranl  added.  It  is  <jiiite  possible  to  find 
the  equivalence  point  of  a  conductometric  titration  by  extrapolating  to 
intersection  the  lines  obtained  at  tin*  beginning  and  at  the  end  of  the  titra- 
tion. Such  a  procedure  is  valuable  when  the  reaction  product  of  ilie  titra 
tion  shows  appreciable  dissociation,  solubility,  or  tendency  toward  hydroly- 
sis. The  experimental  values  near  the  equivalence  point  in  such  cases  will 
be  in  error,  l>ut  the  intersection  of  the  two  straight-line  portions  of  the  plot 
shows  the  theoretical  values.  Conductometric  techniques  are  thus  appli- 
cable when  potent iometric  techniques  may  fail.  Generally,  however,  con- 
ductometric titrations  are  not  widely  used  because  of  the  interference  of 
foreign  ions. 

Amperometric  titrations  are  concerned  with  measurement  <»f  diffusion 
currents  at  constant  potential.  Since  the  diffusion  current  of  a  solution  at 
the  dropping  mercury  electrode  is  in  general  proportional  to  the  concentra- 
tion of  the  reducible  or  oxidizable  species,  changes  in  the  diffusion  current 
may  be  related  to  changes  in  concentration.  Either  the  material  in  solution 
or  the  titrant,  or  both,  may  produce  a  diffusion  current  at  the  potential 
chosen.  The  plot  of  an  amperometric  titration  usually  consists  of  two  inter- 
secting straight  lines,  the  coordinates  of  the  intersection  point  being  the 
equivalence  diffusion  current  and  the  equivalence  volume  of  titrant.  Inter- 
ference of  the  reaction  product  frequently  requires  extrapolation  to  the 
equivalence  point,  as  with  conductometric  titrations108. 

Jaques109  has  given  a  thorough  mathematical  treatment  of  the  deter- 
mination of  the  formula  of  a  complex  by  potentiometric  titration.  If  a 
metal  ion.  M+,  reacts  with  an  anion,  A-,  to  form  a  complex,  the  general 
equilibrium  is  given  by 

Potentiometric  measurements  are  made  for  various  concentrations  of  metal 
ion  and  anion.  The  values  for  q  and  r  may  be  found  from  the  following 
equations. 

KT       /[MgArltY'*  ... 

AAi  =  —  In  I  1      ;  (I) 

nF         yMgArfe/ 

— s-(Hr- 

A/-,'i  is  the  difference  in  potential  between  concentrations  1  and  2  at  con- 
Btant  anion  concentration,  while  AA'n  is  the  difference  between  concentra- 
tions 3  and  1  at  constant  complex  concentration. 

ins.  Kolthoff  and  Laitinen,  "pH  and  Electro  Titrations/'  2nd  ed.,  New  York,  John 
Wiley  A  Sons,  Inc.,  L941. 

109.  Jaques,  "Complex  [ona  in  Aqueous  Solution,"  Longmans  Green  and  Co.,  191  L. 


592  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Leden110  has  used  the  potentiometric  titration  method  to  demonstrate 
complex  formation  between  cadmium  ions  and  various  anions.  Cadmium 
perchlorate-sodium  perchlorate  solutions  were  titrated  with  other  sodium 
salt  solutions,  and  the  data  were  interpreted  by  Leden  to  indicate  the  for- 
mation of  [CdCl]+,  [CdClJ,  [CdClJ",  [CdBr2],  [CdBrJ",  [CdBr4]=  [Cdl]+, 
[Cdl2],  [Cdl4]=,  [Cd(SCN)2],  [Cd(SCN)8]-,  [CdN03]+,  and  [CdSOJ.  Some  of 
these  complexes  are  seen  to  be  undissociated  forms  of  normal  cadmium 
salts.  The  dinuclear  complex  [Cd2Br3]+  also  appears  to  form  in  bromide 
solutions. 

An  important  method  for  determining  complex  formation  constants  has 
been  described  by  Bjerrum111.  This  method  is  essentially  one  of  pH  titra- 
tions. The  general  equilibrium  between  a  metal  ion  M  and  ligands  A  is  writ- 
ten in  steps : 

M  +  A  ;=±  MA 

MA  +  A  ^±  MA2 

MAjv-i  +  A  ^  MAjy 

The  individual  formation  constants  are  given  by 

[MA] 


*i  = 
k2  = 


[M][A] 

[MA2] 
[MA][A] 

[MA„] 
[MA*_i][A] 


Bjerrum  defines  the  quantity  n  as  the  average  number  of  coordinated 
groups  per  metal  ion  present;  all  metal  ions  are  counted  whether  coor- 
dinated or  not. 

[MA]  +  2[MA2]  +  •    •  +  N[MAjr] 
n  = 


[M]  +  [MA]  +  [MA2]  +  •  •  •  +  [MA*] 


The  value  of  n  is  determined  experimentally  by  measurement  of  pH,  since 
removal  of  free  donor  groups  by  coordination  alters  the  pH  by  amounts 
which  may  be  used  to  calculate  the  number  of  groups  coordinated.  The 
quantity  of  ligand  added  must  be  known,  as  well  as  the  value  the  pH 
would  have  if  no  ligand  were  present.  The  difference  between  concentra- 
tion of  ligand  added  and  concentration  of  ligand  coordinated  is  the  concen- 
tration of  free  ligand,  [A].  Bjerrum  has  shown  mathematically  that  when  the 

110.  Leden,  Z.  physik.  Chem.,  A188,  160  (1941). 

111.  Bjerrum,  "Metal  Ammine  Formation  in  Aqueous  Solution,"  Copenhagen,  P. 

1  [aase  and  Son,  1941. 


PHYSICAL  METHODS  IN  COORDINATION  CHEMISTRY  593 

experimental  concentrations  air  adjusted  to  specific  values  for  ft,  the  follow- 
ing relations  hold  for  the  case  N  =  2. 


If  n  -  '.-. 

A'1_iX]- 

If  n  =  y2, 

-B- 

if  n  =  l,      Vfcifcj  =  &  =  m  • 

[A] 

The  "average  constant,"  A',  is  also  the  square  root  of  the  constant  K  of  the 

over-all  reaction 

jr 

M  +  2A  k  MAi  (JV  =  2). 

Application  of  Bjerrum's  method  is  exemplified  by  the  work  of  Calvin 
and  Melchior112  with  the  5-sulfosalicylaldehyde  complex  of  copper(II). 
These  authors  titrated  5-sulfosalicylaldehyde  with  sodium  hydroxide  and 
then  repeated  the  titration  in  the  presence  of  copper(II)  ions.  Plots  of  the 
two  titrations  were  made  on  the  same  set  of  axes,  with  the  separation  of  the 
two  curves  at  a  given  pH  value  corresponding  to  the  amount  of  hydroxide 
needed  to  neutralize  the  protons  freed  by  the  coordinating  organic  groups. 
This  amount  of  hydroxide  gives  the  quantity  of  coordinated  ligand,  and, 
when  divided  by  the  known  metal  concentration,  the  value  of  n.  The  value 
of  [A],  the  concentration  of  aldehyde  anion,  was  found  from  the  known 
concentration  of  uncoordinated  aldehyde  and  its  known  dissociation  con- 
stant. The  values  of  [A]  at  n  =  14,  n  =  1,  and  n  =  %  were  used  to  calcu- 
late log  ki ,  log  k,  and  log  k2  as  approximately  5.2,  4.5,  and  3.7,  respectively. 

A  similar  application  of  Bjerrum's  method  has  been  made  by  De,  Ghosh, 
and  Ray113,  who  studied  tris(biguanide)cobalt(III)  and  tris(phenylbi- 
guanide)cobalt(III)  complexes.  These  complexes  were  found  to  be  quite 
stable,  more  so  than  the  cobalt  ammines. 

A  number  of  workers  have  obtained  values  for  dissociation  and  forma- 
tion constants  of  complexes  by  potentiometric  means  other  than  pH  meas- 
urements. Quite  often  it  is  possible  to  calculate  standard  oxidation  poten- 
tials by  correcting  experimental  oxidation  potentials  with  activity  or 
concentration  data.  Constants  may  then  be  calculated  with  the  formula 

RT 

E°  =  — =-  In  K.  E°  is  here  the  difference  in  standard  potential  between  the 

oxidation  of  metal  to  simple  ion  and  metal  to  complex  ion.  Leden110  has 
used  this  method  to  find  an  increasing  stability  of  cyano  complexes  of 

112.  Calvin  and  Melchior,  J.  Am.  Chem.  Soc,  70,  3270  (1948). 

113.  De,  Ghosh,  and  Ray,  ./.  Indian  Chem.  Soc,  27,  403  (1950). 


594  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

cadmium  as  the  number  of  cyano  groups  increases  from  one  to  four.  Sillen 
and  Liljeqvist114  have  reported  that  halo  complexes  of  zinc  increase  in 
-lability  in  the  series  iodo  <  bromo  <  chloro.  Grinberg  and  his  co-work- 
ers"1, by  determining  the  oxidation  potential  for  the  system 

[PtX4J-  +  2X-  ^  [PtX«J-  +  2e~,  (X  =  CI",  Br~,  SCN~), 

have  found  the  stability  of  the  platinum(II)  complexes  to  increase  in  the 
series  thiocyanato  <  bromo  <  chloro.  Further  studies  by  Grinberg116  have 
established  that  the  oxidation  of  the  platinum  in  such  complexes  as 
[Pt(NH3)4][Pt(CN)4]  and  [Pt(NH8)4][PtBr4]  actually  takes  place  in  two 
steps,  with  the  ammine  platinum  being  more  easily  oxidized.  Higher  tem- 
peratures accentuate  the  difference  in  potential  between  the  two  steps,  and 
low  temperatures  frequently  eliminate  it.  The  complex  [Pt(NH3)2(CN)2] 
shows  only  one  oxidation  step. 

Potentiometric  titrations  by  Treadwell  and  Huber117  have  confirmed  the 
conclusion  of  Manchot118  that  iron(I)  is  present  in  the  nitroso  Roussin 
salts,  red  K[Fe(NO)2S]  and  black  K[Fe4(NO)7S3]-H20.  Unipositive  cobalt 
and  nickel  also  appear  to  be  present  in  the  black  salts  K3[Co(NO)2(S203)2] 
and  K3[Ni(NO)(S203)2]-2H20. 

The  cis  and  trans  isomers  of  dichlorobis(ethylenediamine)  cobalt  (III) 
and  dichlorotetramminecobalt(III)  have  been  the  subjects  of  a  number  of 
potentiometric  studies.  Mathieu119  has  made  pH  measurements  during 
aquation  of  these  complexes  and  has  postulated  the  following  steps. 

[Co  en2  Cl2]+  ->  [Co  en2  C1(H20)++  +  Cl" 

[Co  en2  ClH20]++^±  [Co  en2(H20)2l+++  +  CI". 

The  first  reaction  is  considered  to  be  complete  in  solution,  and  the  equilib- 
rium of  the  second  is  found  to  vary  with  temperature,  pH,  and  concentra- 
tion of  the  chloride  and  complex  ions.  At  elevated  pH  values  a  hydroxo 
complex  tends  to  form. 

[Co  en2  CI  (H20)]++  ;=±  [Co  en2  Cl(OH)]+  +  H+ 

The  rates  of  reaction  are  markedly  different  for  the  cis  and  trans  isomers. 
Similarly,  differences  in  rate  between  cis  and  trans  forms  have  been  noted 
by  Jensen120  and  Grinberg121  for  the  following  platinum(II)  system. 

[Pt(NH3)2(H20)2]++  ^  [Pt(NH3)2(H20)OH]+  +  11+  J±  [Pt(NH3)2(OH)2l   +  2H+. 

114.  Sillen  and  Liljeqvist,  Svensk  Kern.  Tid.,  56,  89  (1944). 

115.  Grinberg,  Ptitsyn  and  Lavrent'ev,  J.  Phys.  Chem.,  U.S.S.R.,  10,  661  (1937). 

116.  Grinberg  and  Ryabchikov,  J.  Phys.  Chem.,  U.S.S.R.,  14,  119  (1937). 

117.  Treadwell  and  Huber,  Helv.  chim.  Acta,  26,  18  (1943). 
L18.  Manchot,  Ber.,  59B,  2445  (1926). 

1 19.  Mathieu,  Bull.  soc.  chim.,  [5]  3,  2121  (1936). 

120.  Jensen,  Z.  anorg.  allgem.  Chem.,  242,  87  (1939). 

121.  Grinberg  and  Ryabchikov,  Acta  Physicochem.,  U.R.S.S.,  3,  555,  569  (1933). 


PHYSICAL  METHODS  l\  COORDIh  ITIOh   CHEMI8TR1  595 

These  authors  also  point  out  an  inequality  in  equilibrium  constant  values: 

trans 

/v',  l  .6  X  10  8  6.3  X  1<>  ■ 

/\  I  .<;  X  l<i  i  I  X  10 

Grinberg  and  Gil,dengershel1M  have  used  pi  I  titrations  to  demonstrate 
acidic  properties  of  ammine  complexes  of  platinum!  IV ).  In  one  experiment . 
a  solution  of  tris(ethylenediamine)  platinum  1 1\'  I  chloride  was  titrated  with 

sodium  hydroxide,  using  a  glass  electrode.  It  was  found  that  each  of  the 
ethylenediamine  molecules  in  turn  release.-  a  proton  from  an  amine  group 
to  render  the  complex  in  effect  a  tribasic  acid.  Equilibria  and  dissociation 
constants  as  found  by  this  study  are  given  below: 

[Pt  eni]" ;±  [Pt  en,(en  -  H)]+++  +  H+;        K,  =  3.5  X  10"« 

[Pt  en,(en  -  H)]+++^±  [Pt  en(en  -  H)2]++  +  H+;        Kt  =  1.76  X  lO"1" 

[Pt  en(en  -  H)21++  -  [Pt  (en  -  H)J+j        §*  ~  j 

X\.2  O 

Biswas123  has  combined  potentiometric  and  conductometric  titration 
techniques  to  study  the  molybdic  acid-tartaric  acid  system.  A  highly 
ionized  complex  H2[Mo03(tart)(H20)]  is  evidenced  by  peaks  in  acidity  and 
conductivity  at  a  1:1  mole  ratio. 

H,  tart  H,Mo04  H2[Mo03(tart)(H20)] 


2H+  +  tart-  +  2H+  +  Mo04  ^  2H+  +  [Mo03(tart)(H20)]~ 

Dey124  has  used  conductivity  data  to  confirm  the  existence  of  a  number 
of  copper(II)  ammine  complexes.  Mixtures  of  copper(II)  nitrate  or  cop- 
per(II)  sulfate  and  ammonium  hydroxide  show  conductivities  different  from 
the  sum  of  those  of  the  constituents.  By  plotting  the  deviations  from 
additivity  against  composition,  Dey  has  found  maxima  corresponding  to 
three,  four,  five,  and  six  moles  of  coordinated  ammonia  per  mole  of  copper 
nitrate.  The  hexammine  complex  forms  in  the  presence  of  sulfate  as  well. 

If  a  complex  ion  dissociates  negligibly  at  all  concentrations,  the  conduc- 
tivity of  its  salts  will  lie  practically  a  linear  function  of  the  Square  root  of 
the  concentration.  Swift125  has  found  the  relationship  to  be  linear  for 
K4[Fe(C,X)6].  indicating  stability  of  the  iron  complex.  On  the  other  hand. 
Brasted7'  has  reported  an  incomplete  ionization  for  tris(o-phenanthroline) 

122.  Grinberg  ;m<l  Gil'dengershel,  Izvest.  Mad.  Saul.-.  8.S.S.R.,  Otdel.  Khim.  Nauk, 

1948,  179. 

123.  Biswas,  •/.  Indian  Chem.  8oe.,  24,  345,  103    1947 

124.  Dey,  Natun  .  158,  95    1946  , 

125.  Swift.  ./.  Am  <  .  60,  728    1938 


596  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

zinc  G?-a-bromocamphor-7r-sulfonate  from  conductivity,  refractometric  and 
cryoscopic  measurements. 

[Zn(o-phen)3](CioHMOBrS03)2^  [Zn(o-phen)3]++  +  2C1oHI4OBrS03-. 

Shuttleworth126  has  made  some  interesting  qualitative  tests  of  the  stability 
of  oxalato,  tartrato,  and  citrato  chromium  (III)  complexes.  Conductometric 
titration  of  each  of  the  complexes  with  hydrochloric  acid  yields  a  straight- 
line  conductivity  plot,  indicating  that  there  is  virtually  no  replacement  of 
organic  anions  by  chloride  ions.  Similar  titrations  with  sodium  hydroxide 
show  only  slight  replacement  by  hydroxy  groups. 

A  conductometric  study127  of  the  carbonatopentamminecobalt(III)  ion 
indicates  that  solutions  of  the  ion  undergo  successive  reactions  to  form  an 
equilibrium  mixture  containing  [Co(NH3)5HC03]++,  [Co(NH3)5H20]+++, 
and  [Co(NH3)5OH]++.  This  example  serves  to  point  out  the  importance  of 
determining  the  true  compositions  of  solutions,  in  order  to  avoid  attributing 
to  pure  substances  the  measurable  properties  of  mixtures. 

A  conductometric  study  of  chromium  lactate  complexes  has  been  re- 
ported by  Shuttleworth128.  Conductometric  titration  in  very  dilute  solution 
shows  that  when  chromium  alum  is  boiled  in  the  presence  of  lactate  ion, 
protons  are  liberated  from  the  lactate,  and  coordination  takes  place,  evi- 
dently forming  H3[Cr(lactate)3].  This  complex  acid  may  be  titrated  com- 
pletely with  base  without  precipitation  of  any  of  the  chromium.  Its  char- 
acteristics are  those  of  a  fairly  strong  acid  (Ka  ^  10-2).  The  formation  of 
the  anionic  complex  is  not  complete  unless  the  protons  liberated  from  the 
lactate  are  neutralized. 

Conductance  measurements  by  Nayar  and  Pande129  on  solutions  contain- 
ing lead  nitrate  and  the  heavier  alkali  nitrates  give  evidence  of  complex 
formation.  The  existence  of  4RbN03-Pb(N03)2 ,  2RbN03-Pb(N03)2 ,  and 
IlbN03-Pb(N03)2 ,  for  example,  has  been  demonstrated  by  the  conduc- 
tance method  and  confirmed  by  viscosity  and  transference  measurements. 

Dipole  Moments 

For  the  purpose  of  this  discussion,  molecules  of  compounds  may  be  con- 
sidered as  being  composed  of  positively  and  negatively  charged  particles. 
The  number  of  positive  charges  will  numerically  equal  the  negative  charges, 
resulting  in  electronegativity  of  the  compound.  Each  molecule  has  what 
may  be  thought  of  as  centers  of  positive  and  negative  charges,  much  as 
masses  have  centers  of  gravity.  If  the  centers  of  positive  and  negative 
charge  coincide,  the  molecule  is  nonpolar.  Otherwise  it  is  polar,  and  the 
measure  of  the  degree  of  polarity  is  the  dipole  moment,  ju-  Dipole  mo- 

126.  Shuttleworth,  /.  Intern.  Soc.  Leather  Trades  Chem.,  30,  342  (1946). 

127.  Lamb  and  Stevens,  /.  Am.  Chem.  Soc.,  61,  3229  (1939). 

128.  Shuttleworth,  ./.  Am.  Leather  Chemists'  Assoc.,  45,  447  (1950). 

129.  Nayar  and  Pande, ./.  Indian  Chem.  Soc,  28,  107  (1951). 


PHYSICAL  METHODS  IN  COORDINATION  CHEMISTRY  597 

ment  is  defined  as  the  product  of  the  Del  charge  of  either  Bign  and  the 
distance  between  the  centers  of  charge.  Neither  quantity  may  be  measured 

directly,  but  the  product  may  be  obtained  in  a  Dumber  of  ways. 

All  molecules,  whether  polar  or  oonpolar,  exhibit  induced  polarity  when 
placed  in  an  electric  field.  This  induced  polarity,  symbolized  by  /'/,  ,  results 
in  a  degree  of  orientation  in  the  field.  Furthermore,  all  polar  molecules  -how 
a  permanent,  or  orientation,  polarization,  symbolized  l>y  I\  ,  which  also 
produces  orientation  in  an  applied  field.  The  total  molar  polarization  P  is 
the  sum   of  the   induced   and   orientation   polarization-;   it    may   he  found 

experimentally  because  of  its  relationship  with  the  measurable  dielectric 

constant  e. 

r-Si-f 

M  is  the  molecular  weight  of  the  substance  measured,  and  d  is  its  density. 
The  dielectric  constant  is  measured  as  the  ratio  of  capacitances  of  a 
condenser  when  filled  with  the  substance  studied  and  with  air,  respectively. 
Actually  the  constant  measures  the  force  required  to  orient  the  molecules 
in  the  field.  Debye130  has  shown  that  the  orientation  polarization  PM  is 
related  to  the  dielectric  constant  by  the  formula 


-(Hs 


where  N  is  Avogadro's  number,  /,•  is  the  Boltzmann  constant  per  molecule, 
and  T  is  the  Kelvin  temperature.  If  a  substance  is  measured  in  the  gaseous 
state,  the  average  distance  between  molecules  is  sufficient  to  render  the 
induced  polarization  PD  practically  constant.  Then 

^r  =  Pd  +  P»  =  Pd 


\VX)i~7 


6+2  d 

and  the  value  of  P,  obtainable  from  values  of  €,  is  a  linear  function  of  —  • 
Experimentally,  a  plot  is  made  of  corresponding  values  of  P  and  —  ,  and  the 

-lope  of  the  resulting  line  is  set  equal  to  the  coefficient  of  -=,  on  the  right 

side  of  Equation  (II).  This  expression  then  leads  to  the  dipole  moment. 

Most  complex  compounds  cannot  be  volatilized  without  decomposition. 
A  method  for  determining  dipole  moments  of  such  substances  involves  the 
following  relation,  which  holds  true  at  infinite  dilution  in  Bolution. 

»-SfH ™ 

130.  Debye,  "Polar  Molecules,"  New  V,.rk,.  Chemical  Catalog  Co.,  (Reinhold  Put- 

lishing  Corp.)  1929. 


598  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

The  refractive  index  n  should  be  known  for  the  far  infrared  region,  but  an 
index  for  visible  light  is  a  good  approximation  for  substances  with  fairly 
high  dipole  moments.  Experimental  values  for  total  molar  polarization, 
found  as  before,  are  extrapolated  to  infinite  dilution.  The  values  of  PD 
from  Equation  (IV)  are  subtracted,  and  the  result  is  equal  to  the  right 
side  of  Equation  (II). 

A  third  method  for  determining  dipole  moments  makes  use  of  the  Stern- 
Gerlach  molecular  beam  technique.  The  material  to  be  studied  must  be 
volatilized  and  passed  through  collimating  slits.  The  molecules  of  the  ma- 
terial are  then  subjected  to  the  deflecting  force  of  an  electrical  field  and 
condensed  onto  a  plate  so  designed  that  the  molecular  trace  may  be  ob- 
served or  photographed.  The  permanent  moments  of  polar  molecules  cause 
them  to  be  deflected  more  than  nonpolar  molecules  and  to  yield  a  broader 
trace.  A  calibration  technique  is  used  to  evaluate  the  traces  by  comparison 
with  standard  dipoles. 

Dipole  moments  have  the  dimensions  esu-cm.  Their  values  are  always  of 
the  order  of  10~18  esu-cm,  and  for  convenience  the  quantity  10-18  esu-cm 
has  been  chosen  as  the  dipole  moment  unit  and  named  the  Debye  unit 
(D.U.). 

The  measurement  of  dipole  moments  has  been  only  recently  applied  to 
structural  studies  of  complexes.  When  two  or  more  structures  for  a  mole- 
cule each  agree  substantially  with  data  from  other  physical  methods,  dipole 
studies  frequently  permit  choice  of  a  most  likely  structure.  Dipole  moment 
data  have  been  used  also  in  estimating  degrees  of  partial  ionic  character 
and  in  distinguishing  between  cis  and  trans  isomers.  Several  examples  of 
dipole  moment  studies  are  given  below. 

Martin131  reports  the  values  for  dipole  moments  of  several  halides  and 
correlates  the  values  with  the  tendency  toward  bonding  between  the  halides 
and  boron  trichloride.  The  data  are  given  in  part  in  Table  18.1.  Martin 
qoints  out  the  value  2.00  D.U.  as  an  apparent  demarcation  between  bonding 
and  nonbonding  halides.  Evidently  the  polar  character  of  the  halides  de- 
termines the  degree  of  availability  of  bonding  electrons.  Chlorine  itself, 
with  a  dipole  moment  of  zero,  forms  no  compound. 

Jensen132  investigated  the  dipole  moments  of  platinum(II)  complexes 
with  tertiary  phosphines,  arsines,  and  stibines.  The  dipole  moments  fall 
into  two  distinct  groups.  The  group  called  a  by  Jensen  is  characterized  by 
very  small  dipole  moments,  suggesting  trans  configuration.  The  dipole 
moment  of  (linitratol)is(triethyl  phosphine)platinum(II)  is  considerably 
larger  than  the  others,  presumably  because  of  unsymmetrical  coordination 
of  the  nitrate^  group.  The  dipoles  of  the  (3  group  are  quite  marked,  suggesting 

131.  Martin,  J.  Phys.  and  Colloid  Chan.,  51,  1400  (1947). 
L32    Jensen,  Z.  anorg.  allgem.  Chem.,  229,  225  (1936). 


PHYSICAL   METHODS  IN  COORDINATION*   CHEMISTRY 


599 


Table  18.1.  Dipole  Moments  oi   Cbbtain  Salides   lnd  Compound  Formation 


Willi 

BC1 

Hklide 

m(D.U.) 

Compound  1 

BC1 

1.03 

None 

CH  CI 

1.84 

None 

CAC1 

2.01 

C  11,01(3013)2 

n-C3H7Cl 

1.97 

None 

MO-CsHrCl 

2  02 

(C3H7Cl)J'.ci 

cia  forms.  Similar  results  for  analogous  trans  palladium  complexes  arc 
reported  by  Mann  and  Purdie133. 

Lamb  and  Mysels1*  report  a  thorough  study  of  carbonatotetrammine- 
cobalt(III)   and   earbonatopentamminecobalt(III)   complexes,   using  the 

method  of  dielectric  increments.  This  method  involves  measurement  of 
electrical  capacitance  of  a  substance  in  a  pulsating- electrical  field  generated 
by  an  electronic  oscillator.  The  frequency  of  the  oscillator  is  varied,  and  the 
corresponding  capacitances  are  measured.  In  order  to  calculate  the  dipole 
moment  of  the  substance,  one  must  first  determine  the  electrical  conduc- 
tance in  solution.  The  calculation  formula  involves  the  conductance,  the 
frequency  used,  the  capacitance  observed,  and  several  correction  factors. 
Resulting  values  of  the  dielectric  constant  at  several  low  frequencies  are 
compared  with  the  theoretically  obtained  value  for  infinite  frequency.  The 
average  difference,  or  dielectric  increment  for  low  frequencies,  may  be 
tised  to  find  the  dipole  moment.  Lamb  and  Mysels  show  by  this  method 
that  the  dipole  moment  of  [Co(XH3)5(C03)]+  is  sufficiently  greater  than 
that  of  [Co(XH3)4(C0.3)]+  to  warrant  postulation  of  the  structures 

O 

/   \ 
(XH,)«  Co  C=0 

\    / 

o 

and 


(XH3)5Co—  O 


i 

\ 


In  the  second  .-tincture  the  dipole  is  more  pronounced.  The  complex  be- 
-  as  it"  it  were  formed  by  loss  of  a  proton  from  the  bicarbonatopentam- 
mine  complex,  with  subsequent  localization  of  negative  charge. 

Mann  and  Purdie,  ./.  Chem.  Soe.,  1549    10311  ;  B73    L036). 
134.  Lamb  and  Mysels, ./.  Am.  Chem.  Soe.,  67,  168  (1046 


600  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Magnetic  Measurements 

While  electrical  dipoles  result  from  unbalanced  distribution  of  positive 
and  negative  charges  within  molecules  of  a  compound,  magnetic  dipoles 
result  from  unbalanced  electronic  spin  and  orbital  contributions  to  molecu- 
lar magnetism.  All  substances  display  some  sort  of  magnetic  dissymmetry, 
however,  in  contrast  to  the  existence  of  electrical  nonpolarity. 

The  intensity  of  a  magnetic  field  is  always  changed  within  a  material 
through  which  the  field  passes.  All  materials  have  in  common  a  tendency 
to  lessen  the  intensity  of  the  field  and  thus  to  be  repelled  by  it.  This  prop- 
erty, called  diamagnetism,  is  attributable  to  the  effect  of  the  field  on  elec- 
tron pairs  within  molecules.  Some  materials  also  contain  unpaired  elec- 
trons or  unbalanced  orbitals,  which  increase  the  intensity  of  the  field  within 
the  material.  This  property  is  called  paramagnetism,  and  its  magnitude  is 
so  much  greater  than  that  of  diamagnetism  that  the  latter  may  usually  be 
neglected  in  paramagnetic  materials.  A  special  case  of  paramagnetism,  in 
which  the  field  increase  within  the  material  is  of  the  order  of  a  million  times, 
is  termed  ferromagnetism.  This  phenomenon  is  exhibited  by  only  a  few 
materials,  those  which  are  capable  of  "permanent  magnetism." 

Changes  in  field  intensity  are  expressed  mathematically  by  the  relation 

B  =  H  +  4x7,  (I) 

where  B  is  the  intensity  in  oersteds  within  the  substance,  H  the  outside  field 
intensity,  and  I  the  intensity  of  magnetization.  I  has  negative  values  for 
diamagnetism  and  larger  positive  values  for  paramagnetism.  The  quantity 

K  =  —  is  termed  magnetic  susceptibility  per  unit  volume.  Susceptibility 
H 

per  unit  mass,  x,  is  obtained  as  the  quotient  of  K  and  the  density  of  the 

substance.  Molar  susceptibility,  xm  ,  is  the  product  of  x  and  the  molecular 

weight. 

Experimental  measurements  generally  determine  the  susceptibility  of  a 

substance,  but  a  quantity  of  great  theoretical  interest  is  the  magnetic 

moment,  ju.  The  relationship  between  magnetic  moment  and  susceptibility 

is  expressed  by 

Nul 
*"  =  Na  +  3^  (II) 

where  N  is  Avogadro's  number,  a  is  diamagnetic  susceptibility  per  mole- 
cule, and  k  is  the  Boltzmann  constant.  Magnetic  moments  are  expressed  in 
Bohr  magnetons.  If  the  orbital  contributions  to  magnetic  moment  are 
neglected,  the  moment  may  be  related  to  the  number  of  unpaired  electrons 
per  molecule  by  the  "spin  only"  formula. 

M  =  Vn(n  +  2)  (III) 


PHYSICAL  METHODS  IN  COORDINATION  CHEMISTRY  601 

This  theoretical  value  for  the  magnetic  moment  agrees  well  with  experi- 
mental values  for  substances  whose  orbital  contributions  are  not  shielded 
and  may  be  neutralized  by  interaction  with  surrounding  particles.  Unpaired 
electrons  of  the  rare  earth  elements  lie  in  the  1/  level  and  are  not  subject  bo 

interaction.  For  these  elements  the  "spin  only"  formula  tails  to  agree  with 
experiment,  and  refinements  musl  be  introduced  into  theoretical  calcula- 
tions. 
Comprehensive  treatments  of  magnetic  theory  are  given  by  Selwoodm, 

Klemm136,  Van  Vleck137,  and  Pauling133. 

Numerous  methods  have  been  developed  for  measurement  of  magnetic 
susceptibilities.  The  most  widely  used  method  was  developed  by  ( i<my139. 
This  method  measures  the  force  exerted  upon  a  sample  by  a  magnetic  field 
of  high  intensity  at  one  end  of  the  sample  and  nearly  negligible  intensity  at 
the  other  end.  The  force  is  measured  on  a  balance  in  terms  of  the  apparent 
added  weight  upon  application  of  the  magnetic  field.  It  is  necessary  to  cal- 
culate susceptibility  values  from  the  experimental  data. 

Other  useful  methods  have  been  worked  out  by  Quincke140,  Faraday141-  142, 
Curie  and  Cheneveau143,  Rankine144,  and  Iskenderian146. 

Measurements  of  magnetic  susceptibility  have  been  of  great  value  in 
determining  bond  types  and  structures  of  complexes.  The  various  types  of 
bonding  possible  in  a  given  complex  may  often  be  distinguished  on  the  basis 
of  the  number  of  unpaired  electrons  present  with  each  type.  If  experiment 
establishes  the  magnetic  susceptibility  and  thus  the  number  of  unpaired 
electrons,  questions  may  frequently  be  settled  concerning  orbital  hybridi- 
zation, degree  of  covalent  character,  and  probable  structure.  Theories  of 
bonding,  orbitals,  and  structure  in  coordination  chemistry  have  not  been 
thoroughly  evolved,  but  magnetic  data  constitute  a  powerful  tool  for  the 
improvement  of  current  ideas. 

Tyson  and  Adams146  have  used  magnetic  data  to  postulate  structures  for 

135.  Selwood,  "Magnetochemistry,"  New  York,  Interscience  Publishers,  Inc.,  1943. 

136.  Klemm,  "Magnetochemie,"  Leipzig,  Akademische  Verlagsgesellschaft  m.b.H., 

1936. 

137.  Van  Vleck,  "Theory  of  Electric  and  Magnetic  Susceptibilities,"  pp.  283-301, 

Oxford,  The  Clarendon  Press,  1932. 

138.  Pauling,  "The  Nature  of  the  Chemical  Bond,"  Ithaca,  N.  Y.,  Cornell  University 

Press,  1940. 

139.  Gouy,  Compl.  rend.,  109,  935  (1889). 

140.  Quincke,  Ann.  Physik.,  24,  347  (1885);  34,  401  (1888). 

141.  Stoner,  "Magnetism  and  Matter,"  London,  Methucn  and  Co.,  Ltd.,  1934. 

142.  Curie,  Ann.  chim.  phys.,  (7)  5,  289  (1895). 

143.  Cheneveau,  Phil.  Mag.,  20,  357  (1910). 

111.  Rankine,  Proc.  Phys.  Soc.  London,  46,  1,  391  (1934 

145.  Iskenderian.  P)  .     Rev.,  51,  1092     19 

146.  Tyson  and  Adams,  ./.  Am.  Chem.  8oc.,  62,  1228  (1940). 


002 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Table  18.2.  Magnetic  Moments  of  Some  Inner  Complexes  of  Copper,  Nickel, 

and  Cobalt 


MEA3URE0  NEAREST      THEO- 

MOMENT  RETICAL    VALUE 

ANC  CORRESPOND- 
ING    NUMBER  OF 
UNPAIRED  ELECTRONS 


(d) 


(e) 


H     H 

■  Nl 


CLTX-T3 


3.1 


'VO 


1.73;  i 


283;2 


Q  =  ORIGINAL    ELECTRONS 

GO  -  COORDINATION    ELECTRONS 

3d  4S  4p 

|TTm      S      [F[iin»P3.TETRAHEDIUl 


mm]  h  eed  ^^ 


]  a  o 


J      Sp3,TETRAHEDRAL 


irm  h  nmv^ 


urn  b 


spJ,TETRAHEORAL 


mrrem  s  eel>p! 


a 


H 

C=0^  tO- 

Coc 

0^       ^0  = 


"0 


3.88;  3 


0 


>p3,TETRAHEDnAL 


salicylaldehyde  and  salicylaldimine  complexes  of  divalent  copper,  nickel, 
and  cobalt  (Table  18.2).  It  is  apparent  that  magnetic  data  alone  are  not 
sufficient  to  choose  between  the  two  reasonable  structures  for  complexes 
(a)  and  (b).  Cox  and  Webster147  have  established  by  x-ray  methods  that 
both  complexes  are  planar.  The  two  inner  complexes  of  nickel  are  of  special 
interest.  Their  difference  in  structure  is  further  confirmed  by  a  pronounced 
difference  in  absorption  maxima. 

The  work  of  Mellor  and  Goldacre148  has  shown  that  a  number  of  co- 
balt(II)  nitrogen- and  oxygen-bonded  complexes  display  the  high  magnetic 
moments  characteristic  of  ionic  complexes  of  divalent  cobalt.  Most  values 
are  considerably  above  the  theoretical  three-electron  moment  of  3.88,  and 
such  values  are  to  be  expected.  The  magnetic  moments  of  [Co(NH3)6]Cl2 , 
[Co  (en)3]Cl2,  and  Na2[Co(C6H4{COOJ2)2]  are  given  as  4.96,  3.82,  and  5.35 
Bohr  magnetons,  respectively.  Orbital  magnetism  is  evidently  a  con- 
tributing factor  in  these  instances. 

A  relationship  between  complex  stability  and  magnetic  moments  has 
been  reported  by  Russel  and  his  co-workers149  for  certain  nickel (II)  and 

147.  Cox  and  Webster,  ./.  Chem.  Soc,  731  (1935). 

148.  Mellor  and  Goldacre,  J.  Proc.  Roy.  Soc.  N.  S.  Wales,  73,  233-9  (1940). 
1  ID.  Russel,  Cooper,  and  Vosburgh,  J.  Am.  Chem.  Soc.,  65,  1301  (1943). 


PHYSICAL   METHODS  I\  COORDINATION  CHEMISTRY  603 

copperl  1 1 1  complexes.  Aqueous  solutions  of  the  metal  sulfates  were  treated 
with  excesses  of  various  nitrogen-  and  oxygen-donating  groups,  two  types 
oi  donor  molecules  at  a  time.  Measurement  of  maximum  lighl  absorption 
and  comparison  with  known  values  permitted  a  conclusion  in  several  cases 
as  to  the  relative  coordinating  abilities  of  the  two  ligands  used.  Each  com- 
plex was  also  isolated  and  tested  magnetically.  A  nearly  linear  relationship 
was  discovered  between  stability  as  shown  spectrally  and  by  magnetic 
moment.  The  coordinating  groups  for  which  stability  conclusions  could  be 
drawn  are  shown  below. 

Nickel (II)  complexes:  Least  Btable-aquo  <  pyridine  <  ammine  <  ethylenediamine 

n  S  3.24 

<  o-phenanthroline-Most  stable 

M  ^  3.08 
Copper (II)  complexes:  Least  stable-aquo  <  pyridine   <  ammine  <  aminoacetate 

M  S  1.95 

<  ethylenediamine-Most  stable 

m  ^  1.85 

Srivastava,  Pande,  and  Xayar150  have  described  an  interesting  applica- 
tion of  magnetic  measurements  to  the  method  of  continuous  variations. 
Lead  nitrate  was  added  to  aqueous  solutions  of  potassium  nitrate  and 
ammonium  nitrate,  respectively.  The  magnetic  susceptibility  was  measured 
at  intervals  and  plotted  against  composition  of  the  solution.  The  results 
correspond  to  compound  formation  involving  one,  two,  and  four  molecules 
of  lead  nitrate  per  molecule  of  potassium  or  ammonium  nitrate.  The  results 
have  been  confirmed  by  a  conductometric  method. 

Apparently  anomalous  magnetic  moments  may  sometimes  be  found 
among  complexes  containing  optically  active  ligands.  French  and  his 
'dates151  have  noted  that  certain  complexes  of  nickel(II),  which  would 
be  expected  by  analogy  to  be  diamagnetic  and  planar,  are  actually  para- 
magnetic and  therefore  probably  tetrahedral.  An  example  is  bis(formyl- 
camphor)nickel(II),  [Xi(Ci0Hi4{CHO}O)2].  Both  magnetic  data  and  rota- 
tory dispersion  measurements  point  to  the  nickel  in  this  complex  as  ;i 
source  of  asymmetry  and  optical  activity  resulting  from  tetrahedral  co- 
ordination. Presumably  the  optically  active  ligand  exercises  a  kind  of 
inductive  influence. 

Mellor  and  Lockwood152  have  furnished  additional  evidence  for  the  dis- 
torting influence  of  certain  ligands.  These  investigators  found  that  coor- 
dination of  substituted  pyrromethenes  with  nickel(II)  produces  a  tetra- 

3rivastava,  Pande,  and  Nayar,  Current  Sri.,  16, 226-6  (1947  . 
151.  French,  Magee,  and  Sheffield,  J.  Am.  Chem.  Soc,  64,  1924-S  (1942  . 
162.  Mellor  and  I.ockwood,  J.  Proc.  Roy.  Soc.  A    8.  Wales,  74,  141  8  (1040);  Nai 
145,  862  (1940). 


604 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Table  18.3.  Orbital  Arrangements  for  Silver(II)  and  Silver(III)  Complexes 

4d  5s  5p 

SILVER  (n) 


SILVER  (m) 


••••X  X  XX* 

•  1   •    I    •    I    ♦    I  x  |        |  x  |        I   x   |    x 

•  I   •   I.  •    I   •   I  x  I        I  x  I        I  x  I   x   I 

•  •         •         .         X  X  XX 


dsp2,  PLANAR 


dsp2,  PLANAR 


hedral  configuration.  Bis  (3 ,3'  ,5  jS'-tetramethyM^'-dicarbethoxydipyr- 
romethene)nickel(II),  [Ni(Ci9H2304N2)2],  has  a  magnetic  moment  of  3.2 
Bohr  magnetons,  corresponding  to  two  unpaired  electrons.  The  analogous 
complex  of  palladium (II),  however,  is  diamagnetic. 

Ray153  has  used  magnetic  measurements  to  demonstrate  the  existence  of 
silver  (III)  complexes  with  ethylene  biguanide  (C4N5H9  =  big  H).  He  pre- 
pared the  salts  [Ag(big  H)2]X3 ,  where  X  may  be  nitrate,  perchlorate,  or 
hydroxide,  as  well  as  [Ag(big  11)2)2(804)3 .  All  these  salts  are  diamagnetic, 
as  would  be  expected  for  silver(III);  a  corresponding  silver(II)  salt  with 
the  same  ligand  is  paramagnetic.  See  Table  18.3. 

A  comparison  technique  has  enabled  Mellor  and  Craig154  to  support  the 
idea  that  the  diphenylmethylarsine  copper  complex,  [Cu2Cl3(Ph2MeAs)3], 
has  a  dinuclear  structure  containing  both  monovalent  and  divalent  copper. 
Two  forms  of  this  complex  may  be  isolated,  one  blue  and  the  other  brown. 
Mellor  and  Craig  determined  that  the  magnetic  moment  of  each  form  has 
a  value  in  the  neighborhood  of  1.73  Bohr  magnetons.  The  cyanoammine 
copper  complex  [Cu3(CN)4(NH3)3],  known  to  contain  one  copper(II)  atom 
per  molecule,  and  thus  one  unpaired  electron,  has  a  moment  of  1.78  Bohr 
magnetons.  The  following  structures  for  the  two  forms  of  the  arsine  complex 
are  proposed: 


CI 


\      / 


CI 


Cu1 


Cu11 


/         \ 


PhoMeAs 


CI 


AsMePh2 

AsMePh2_ 

"Ph2MeAs                CI 

\       /     \ 

Cu1              Cu11 

/       \     / 
JPh2MeAs                CI 

CI 


AsMePh2_ 
This  work  has  been  seriously  questioned  on  other  grounds  (see  p.  609). 

153.  Ray,  Nature,  151,  643  (1943). 

154.  Mellor  and  Craig,  J.  Proc.  Roy.  Soc.  N.  S.  Wales,  74,  475-94  (1941). 


PHYSICAL  METHODS  IN  COORDINATION  CHEMISTRY 


605 


A  systematic  study  of  the  relation.-  among  magnetic  moment,  color,  and 
configuration  of  complexes  has  been  made  by  Willis  and  Mellor166.  An  inter- 
esting transition  pointed  out  by  this  study  is  that  of  bis(ethylenediamine- 
formylcamphor)nickel(II)  in  pyridine  solution.  When  the  solution  is  freshly 
prepared,  the  complex  exhibits  dia magnetism  and  a  green  color,  correspond- 
ing to  a  tetracovalent  planar  structure.  Upon  standing  for  two  weeks  the 
solution  turns  brown,  and  paramagnetism  appears,  reaching  a  value  of  3.15 
Bohr  magnetons.  Evidently  the  complex  combines  with  two  pyridine  mole- 
cules per  nickel  atom  and  rearranges  to  an  octahedral  structure,  with  un- 
pairing  and  promotion  of  two  3d  electrons  to  the  4d  shell. 


h* 


0-C10HwCHzN-CH2 


O-qoH^CHrN-CK, 


GREEN 


+  2PH 


-C,0HWCH=N-CH2 


O  -C^O^  N-CH2y 


BROWN 


Consideration  of  the  completely  paired  electron  structure  of  cobalt  (III) 
complexes  showing  d2spz  hybridization  suggests  that  all  such  complexes 
should  be  diamagnetic.  That  this  is  not  the  case  has  been  demonstrated 
by  Cambi,  Ferrari,  and  Nardelli156,  who  report  magnetic  measurements  on 
a  series  of  hexanitrocobaltate(III)  complexes.  The  appreciable  paramag- 
netism of  these  compounds  suggests  contributions  from  incompletely 
quenched  orbital  magnetism. 


Complex 

Na3[Co(N02)6] 

K3[Co(N02)6J-H20 

(NTI4)2[Co(X02)6].2H20 

Tl3[Co(N02)6] 

Ba3[Co(N02)6]2-12H20 

Pb3[Co(N02)6]2-llH20 

(Me4N)2Na[Co(N02)6]-2MH20 


MB 

0.57 
0.79 
0.63 
0.52 
0.59 
0.84 
0.52 


Jonassen  and  Frey157  have  shown  that  cobalt(II)  ion  forms  a  complex 
with  tetraethylenepentamine  in  which  the  bonding  is  principally  ionic.  A 
solution  of  cobalt (II)  perchlorate  containing  tetraethylenepentamine  is 
green,  but  after  standing  for  72  hours,  it  is  red.  Thecobalt(II)  complex  which 
may  be  isolated  from  the  solution  shows  a  magnetic  susceptibility  of  4.52 

155.  Willis  and  Mdlor,  ./.  Am.  Chem.  Soc,  69,  1237-40  (1947). 

156.  Cambi,  Ferrari,  and  Nardelli,  Gazz.  chim.  Hal,  82,  816  (1952). 

157.  Jonassen  and  Frey,  ./.  Am.  Chem.  Soc,  75,  1524  (1953). 


606  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Bohr  magnetons.  This  value  is  in  the  usual  range  for  cobalt (II)  complexes 
containing  three  unpaired  electrons. 

X-Ray  and  Electron  Diffraction 
X-Ravs158,  159,  160 

The  radiations  known  as  x-rays  have  wave  lengths  of  the  same  order  as 
interatomic  distances  in  molecules  and  crystals.  For  this  reason  Laue  in 
1(.)12  suggested  that  the  regular  arrangement  of  crystal  lattices  should  act 
as  a  three-dimensional  diffraction  grating  for  x-rays.  It  remained  for 
Friedrich  and  Knipping  to  substantiate  Laue's  idea  by  passing  x-rays 
through  various  crystals  and  onto  a  photographic  plate.  The  developed 
plate  showed  a  prominent  central  area  exposed  by  undiffracted  rays,  and 
a  symmetrical  concentric  pattern  of  rings  in  diffraction  zones  outward  from 
the  center.  This  Laue  transmission  method  has  proved  to  be  of  great  value 
in  structural  analyses.  Hypothetical  crystals  having  any  arbitrary  structure 
are  analyzed  mathematically  to  determine  calculated  diffraction  patterns; 
these  patterns  are  then  compared  with  experimental  results  and  adjusted 
until  they  are  identical.  The  crystal  under  study  is  assigned  the  calculated 
structure. 

A  more  direct  and  convenient  approach  to  x-ray  analysis  is  given  by  the 
Bragg  method.  This  method  treats  the  crystal  as  a  series  of  reflecting  planes 
arranged  in  space  so  that  they  permit  reflection  and  interference  of  x-rays 
entering  at  appropriate  angles.  The  fundamental  equation  for  the  Bragg 
method  is 

nA  =  2d  sin  0,  (1) 

where  n  is  the  order  of  reflection,  d  is  the  distance  between  crystal  reflecting 
planes,  and  6  is  the  angle  at  which  the  rays  strike  the  crystal  face.  Succes- 
sive orders  of  reflection  are  spread  outward  from  the  center  of  the  reflection 
pattern,  as  well  as  weakened  in  intensity.  Knowledge  of  the  wave  length 
and  incident  angle  of  the  x-rays  permits  calculation  of  the  distance  between 
crystal  planes. 

For  practical  application  of  the  Bragg  analysis  a  crystal  is  mounted  on  a 
rotating  table.  An  x-ray  generator  is  so  arranged  that  the  rays  are  produced 
and  collimated  directly  toward  the  center  of  rotation.  After  striking  the 
crystal,  the  rays  travel  to  a  photographic  plate  or  an  ionization  chamber, 
where  their  intensities  are  measured.  A  plot  is  made  of  the  intensity  as  a 

158.  Zachariasen,   "Theory  of  X-Ray  Diffraction  in  Crystals,"  New  York,  John 

Wiley  &  Sons,  Inc.,  1945. 

159.  Roi nniut h,  J.  Chem.  Ed.,  7,  138,  860,  1313  (1930). 

160.  Pirenne,  "The  Diffraction  of  X-Rays  and  Electrons  by  Free  Molecules,"  Cam- 

I iridic,  Cambridge  University  Press,  1946. 


PHYSICAL   METHODS  IN  COORDINATIOh   CHEMISTR1  607 

function  of  angle  of  incidence;  the  most  pronounced  maximum  corresponds 
to  first-order  reflection,  and  so  on.  The  Bragg  equal  ion  serves  to  determine 
the  interplanar  distance  for  all  axes  of  crystal  rotation,  and  after  ;ill  feasible 
orientations  of  the  crystal  on  the  table  have  been  individually  tested,  the 

data  are  taken  to  he  complete. 

The  simplest  applications  of  x-ray  analysis  have  been  made  in  determin- 
ing the  lattice  structure  of  BUCh  ionic  crystals  as  the  alkali  halides.   More 

complicated  structures  are  also  amenable  to  treatment  by  the  methods  jusl 
described.  Data  from  Lane  or  Bragg  tests  are  sometime-  subjected  to  com 
plete  mathematical  analyses  of  the  Fourier  type.  The  ultimate  aim  is  con- 
struction of  an  accurate  three-dimensional  model  which  represents  com- 
pletely the  distribution  of  electron  density  in  a  crystal  and  thus  shows  the 
arrangement  and  separations  of  all  atoms  present.  This  objective  is  not 
realizable  for  structures  containing  hydrogen,  since  the  hydrogen  atom  is 
two  small  for  detection  by  x-rays.  Models  which  are  otherwise  complete 
have  been  arrived  at  for  some  systems,  but  only  with  great  difficulty  and 
tedious  calculation.  Fortunately,  such  complete  analyses  are  not  usually 
necessary  to  establish  structures. 

A  quick  and  relatively  simple  method  of  x-ray  analysis  employs  crystal- 
line powders  rather  than  a  large  crystal.  The  reflection  patterns  obtained 
by  this  method  are  not  usually  so  sharp  as  those  obtained  with  larger  par- 
ticles. Powders  are  often  readily  available,  however,  when  preparation  of 
Bingle  crystals  is  difficult.  Powder  patterns  sometimes  serve  to  identify 
unknown  substances  by  comparison  with  known  patterns.  In  such  cases 
mathematical  analyses  are  unnecessary. 

Electron  Diffraction161   162   163 

The  useful  diffractive  and  reflective  properties  of  x-rays  are  found  also 
in  rapidly  moving  beams  of  electrons.  Electron  beams  are  usually  generated 
electronically  as  cathode  rays.  A  uniform  voltage  of  the  order  of  40,000  to 
60,000  volts  per  centimeter  is  maintained.  The  beam  is  directed  toward  a 
photographic  plate,  and  vapor  of  the  substance  to  be  examined  is  interposed 
between  the  source  and  the  plate.  After  development,  the  plate  shows  a 
prominent  central  spot  and  concentric  rings,  which  may  be  analyzed  in  a 
maimer  analogous  to  x-ray  analysis.  Since  the  penetrating  power  of  elec- 
tion- is  much  lower  than  that  of  x-rays,  the  electron  diffraction  method  is 
suited  particularly  to  studies  of  gases,  while  x-ray  method-  are  besl  for 
solid  and  liquid  measurements.  Photographic  plates  may  be  made  more 
sensitive  to  electron-  than  to  x-rays,  ai  the  intensities  normally  generated 

nil.  Brockway,  /.'<      1/  8,  231     L97I 

162.  Clark  and  Wolthiua,  ./.  Chem.  Ed.,  15,  64    I 

L63.  Pauling  and  Brockway,/.  .1-/.  Chem.  Soc.,  67,2684    1935 


G08  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

in  the  laboratory.  Thus  electron  diffraction  patterns  may  be  taken  in  a 
few  seconds,  while  exposure  of  plates  to  x-rays  usually  extends  over  several 
hours.  More  rings  are  usually  produced  by  the  electron  diffraction  method; 
this  fact  is  important,  inasmuch  as  inner  rings  are  often  obscured  by  the 
central  beam. 

Applications 

X-ray  and  electron  diffraction  studies  on  complex  compounds  have  yielded 
valuable  information  concerning  properties  of  symmetry;  spatial  configura- 
tion; orientation  of  complex  ions  and  molecules  in  crystal  lattices;  differen- 
tiation between  racemates  and  optically  inactive  forms;  determination  of 
bond  angles  and  distances;  estimation  of  molecular  weights  of  complexes; 
differentiation  between  mixtures  and  single-phase  crystals;  and  identifica- 
tion of  bridging  groups. 

Electron  diffraction  studies  have  enabled  Palmer  and  Elliott164  to  propose 
a  structure  for  dimeric  aluminum  chloride  consisting  of  two  tetrahedra 
sharing  an  edge.  Chloride  ions  are  thought  to  occupy  the  corners  of  the 
tetrahedra,  with  aluminum  ions  at  the  centers.  Partial  covalent  character 
reduces  to  some  extent  the  separation  and  magnitude  of  charges  which 
purely  ionic  bonding  would  produce. 

Electron  diffraction  data  lead  to  the  conclusion  that  nickel  carbonyl  has 
a  tetrahedral  structure165.  Measured  bond  distances  for  nickel-carbon  and 
carbon-oxygen  bonds  are  1.82  A  and  1.15  A,  respectively.  These  distances 
are  in  agreement  with  Pauling's  suggestion  that  the  nickel-carbon  bonds 
should  be  considered  as  hybrids,  partaking  of  both  single-bond  and  double- 
bond  character.  The  CO  groups  in  Ni(CO)4  are  evidently  tetrahedrally 
distributed  about  the  nickel,  with  the  character  of  the  carbon-oxygen  bonds 
quite  similar  to  that  found  in  carbon  monoxide.  The  carbonyl  hydrides 
Fe(CO)4H2  and  Co(CO)4H  were  studied  by  Ewens  and  Lister166,  who  at- 
tributed tetrahedral  structures  to  both  on  the  basis  of  electron  diffraction 
patterns.  The  hydrogen  atoms  are  thought  to  be  bonded  to  oxygen,  so  that 
formulas  for  these  hydrides  may  also  be  written  Fe(CO)2(COH)2  and 
Co(CO)3(COH).  The  iron-carbon  distance  for  the  CO  groups  is  1.84  A, 
while  for  the  COH  groups  it  is  1 .79  A.  Respective  distances  for  the  cobalt 
compound  are  1.83  A  and  1.75  A.  Volatility  of  the  carbonyls  and  carbonyl 
hydrides  facilitates  their  study  by  this  method. 

Beach  and  Bauer167  have  obtained  electron  diffraction  patterns  for  the 
vapor  of  the  compound  AIB3H12 .  The  data  indicate  that  an  aluminum  atom 

Kit.  Palmer  and  Elliott,  ./.  Am.  Chem.  Soc.,  60,  1852  (1938). 

165.  Pauling,  ./.  .1///.  Chem.  Soc,  53,  1367  (1931);  64,  988  (1932). 

166.  Ewena  and  Lister,  Trans.  Faraday  Soc,  35,  681  (1939). 

167.  Beach  and  Bauer, ./.  .1///.  Chem.  Soc,  62,  3110  (1940). 


PHYSICAL   METHODS  IN  COORDINATIOh   CHEMISTRY 


609 


is  bonded  to  three  HI  1 1  groups  in  a  planar  configuration  with  the  bonds  a1 
angles  of  120°.  Bach  boron  atom  is  near  the  center  of  a  trigonal  bipyramid 
formed  by  tour  hydrogen  atoms  and  the  aluminum  atom.  The  compound 
is  electron-deficient,  and  the  authors  interpret  the  norma]  aluminum-boron 
bond  lengths  to  indicate  thai  the  deficiency  resides  in  the  boron-hydrogen 
bonding. 

Dipole  moment  studios  of  tetrachlorobis  (trimethylarsine)  palladium(II) 
suggest  three  possible  forms  for  tins  complex. 


Me  As 


\ 


Pd 


Cl  CI 

'  \  / 

Pd 


Me  As 


/     \ 


01 


Cl 


Me  As 


Cl 


Cl 


Pd 


/ 


Pd 


(I) 


Cl 

(II) 


\s\lr 


Cl 


Me3As 


Cl 


Cl 


Pd 


Pd 


Cl 

(III) 


/ 
I 

\ 


Cl 


Me3As 


X-ray  examination  in  the  solid  state  led  Mann  and  his  co-workers168  to  the 
conclusion  that  only  form  (III)  exists  as  a  solid,  although  the  other  forms 
probably  exist  in  organic  solvents  (p.  604).  Replacement  of  two  chloro 
groups  by  an  oxalato  group  in  the  analogous  tributylphosphine  complex 
raises  the  question  of  identifying  the  bridging  groups.  Chatt  and  his  asso- 
ciates169 showed  by  x-ray  investigation  that  the  separation  of  5.3  A  between 
the  palladium  atoms  corresponds  to  oxalato  bridging.  Chloro  bridges  would 
give  the  metal-metal  distance  a  value  of  3.4  A. 

Complex  metal  cyanides  have  been  the  objects  of  considerable  study  by 
x-ray  techniques.  Dothie170  has  shown  that  both  dicyanodipyridylaurate(I) 
and  dicyano-o-phenanthrolineaurate(I)  have  planar  structures,  four  ion- 
comprising  a  unit  cell. 


<> 

P 

CN               CN 

Au 


CN 


CN 


168.  Mann  and  Wells,/.  Chem  8oc.t  702    L938  ;  Mann  and  Purdie,  J.  Chem.  8oe.,  873 

1936  . 

169.  Chatt,  Mann,  and  Wells,  /.  Chem.  Soc.,  1949,2086    19 

17m    Dothie,  LleweUyn.  Wardlaw.  and  Welch,  ./.  Chem.  Soc.,  126    19 


610  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Keggin  and  Miles171  have  studied  a  number  of  cyano  complexes.  The  com- 
pound FeI1M2[FeII(CN)8],  where  M  signifies  an  alkali  metal  or  ammonium 
ion,  has  a  cubic  lattice  structure.  The  iron  atoms  occupy  corner  positions, 
and  the  cyano  groups  bridge  the  iron  atoms  along  all  edges  of  the  cubes. 
The  alkali  metal  ions  are  located  at  the  centers  of  the  cubes.  Oxidation  of 
this  compound  first  produces  alkali-containing  Prussian  blue  and  then 
Berlin  green,  Fe[Fe(CN)6]. 

7FeM2[Fe(CN)6]  -*  2Fe4[Fe(CN)6]3  +  6MCN  +  8M+  +  8e~ 
2Fe«[Fe(CN)6]3  +  6MCN  ->  7Fe[Fe(CN)6]  +  6M+  +  6e~. 

It  is  interesting  that  Weiser172  has  found  identical  x-ray  patterns  for  Prus- 
sian blue  and  Turnbull's  blue,  which  are  formally  written  as 

Fe4III[FeII(CN)6]3    and     Fe3II[FeIII(CN)6]2 , 

respectively. 

Cox  and  his  co-workers173  have  interpreted  x-ray  data  for  the  tetrachloro- 
stannate(II)  ion  to  mean  that  four-coordination  is  present  rather  than  six- 
coordination.  The  hydrated  potassium  salt  is  therefore  K2[SnCl4]-2H20, 
and  not  K2[SnCl4(H20)2].  Cox  has  also  established  the  planar  structures  of 
potassium  bis(oxalato)plumbate(II),  bis(thiourea)lead(II)  chloride,  bis- 
(salicylato)lead(II),  and  bis(benzoylacetone)lead(II). 

Beintema174  has  made  a  detailed  study  of  hexaquo  complexes  of  divalent 
metals  in  which  the  hexahydroxoantimonate(V)  anion  is  present.  Two 
crystalline  modifications  of  [Mg(H20)6][Sb(OH)6]2  are  reported.  One  is  a 
trigonal  form,  isomorphous  with  [Ni(H20)6][Sb(OH)6]2 ,  and  the  other  is 
triclinic  pseudo-monoclinic,  isomorphous  with  [Co(H20)6][Sb(OH)6]2 . 

Lambot175  has  used  x-rays  to  confirm  a  planar  structure  for  K2[Pt(N02)4]. 
The  platinum-nitrogen  distance  is  calculated  as  2.02  A,  and  the  nitrogen- 
oxygen  distance  as  1.22  A.  The  0 — N — O  angle  in  the  nitro  groups  is  127°. 

Heneghan  and  Bailar176  have  shown  that  the  cis  and  trans  isomers  of 
(lichlorobis(ethylenediamine)platinum(IV)  nitrate  yield  quite  different 
x-ray  patterns.  Formerly  all  the  preparative  methods  used  to  synthesize 
this  compound  had  produced  only  the  trans  form.  Heneghan  and  Bailar 
have  developed  a  method  of  synthesis  for  the  cis  form.  It  is  optically  re- 
solvable, and  its  x-ray  pattern  shows  clearly  that  it  is  not  the  trans  isomer. 

Moeller  and  Ramaniah177  have  used  x-ray  data  to  distinguish  between  two 

171.  Keggin  and  Miles,  Nature,  137,  577  (1936). 

172.  Weiser,  Million  and  Bates,  J.  Phys.  Chem.,  46,  99  (1942). 

173.  Cox,  Shorter,  and  Wardlaw,  Nature,  139,  71   (1937). 
171.  Beintema,  Rec.  trav.  chim.,  56,  931  (1937). 

17.").   Lambot,  Roy.  soc.  Liege,  12,  463  (1943). 

L76.  Heneghan  and  Bailar,/.  .1///.  Chem.  Soc., 75,  1840  (1953). 

177.  Moeller  and  Ramaniah,/.  .1///.  Ch em.  Soc.,  75,  3946  (1953). 


PHYSICAL  METHODS  l\  COORDINATION*   CHEMISTRY  611 

complexes  of  thorium  with  oxine  (8-hydroxyquinoline).  If  a  solution  of 
thorium(IV)  nitrate  is  treated  with  oxine  under  appropriate  conditions,  a 

product  may  be  isolated  which  contains  tour  oxinatc  anions  and  one  mole- 
cule of  oxine  per  thorium (IV)  ion.  Heating  this  product  to  120  to   L25  C 

tor  five  hours  and  then  to  L30  to  loo  for  one  hour  produces  the  normal 
inner  complex,  [Th^oxinate)||.  X-ray  diffraction  studies  show  that  the  two 
complexes  are  different,  and  that  the  I  :5  complex  IS  different  from  a  mix- 
ture of  the  1  :  I  complex  and  one  mole  of  oxine.  The  fifth  molecule  of  oxine 
is  lost  in  solution,  and  it  seems  therefore  to  he  hound  by  weak  lattice  forces. 
An  analogous  situation  occurs  with  scandium178. 

Traces  Techniques;  Exchange  Reactions 

Any  molecules,  atoms  or  ions  of  any  given  species  are  indistinguishable 
from  all  the  other  members  of  the  same  single  species  when  subjected  to 
most  physical  measurements.  This  failure  is  a  limiting  factor  in  chemical 
studies,  since  apparently  inert  chemical  combinations  may  be  in  equilibrium 
with  their  constituents  without  this  equilibrium  being  detected.  Tracer 
techniques  take  advantage  of  the  fact  that  isotopic  species  may  be  dis- 
tinguished, yet  their  presence  in  any  ratio  seldom  affects  the  course  or  rate 
of  a  reaction  by  any  measurable  amount.  It  is  theoretically  possible  to 
determine  the  distribution  in  a  reaction  of  ordinary  isotopes  of  different 
masses.  In  usual  practice,  however,  only  the  isotopes  of  hydrogen  have  a  suf- 
ficient percentage  of  mass  difference  to  permit  reasonably  accurate  measure- 
ments. The  availability  of  radioactive  isotopes  and  the  development  of 
efficient  techniques  for  measuring  radioactivity  have  been  largely  responsible 
for  the  growth  of  tracer  chemistry.  Like  isotopic  mass  difference,  radio- 
activity almost  never  alters  the  chemical  nature  of  a  system  into  which  it 
is  introduced  as  a  constituent.  A  radioactive  element  is  usually  added  to  a 
reaction  in  the  form  of  a  common  compound.  If  every  molecule  or  complex 
which  contains  this  element  is  in  rapid  equilibrium  with  its  constituent-. 
the  radioactive  substance  quickly  assumes  a  statistical  distribution  which 
is  in  proportion  to  the  distribution  of  the  ordinary  isotope.  Deviations  from 
rapid  equilibrium  are  measurable  in  terms  of  deviations  from  this  statistical 
distribution  of  radioactivity.  The  method  requires  chemical  separation  of 
the  species  present,  accurate  measurement  of  the  radioactivity,  and  ap- 
propriate calculations.  The  objective  is  a  knowledge  of  the  relative  lability, 
or  its  opposite,  the  "inertness,"  of  the  chemical  bonds  in  the  species  studied. 

Preparation  of  Radioisotopes179 •  180,  181 
Very  few  naturally  occurring  radioactive  elements  are  useful  in  tracer 

178.  Pokras,  Kilpatrick,  and  Bernays,  •/.  Am.  Chem.  Soc.,75,  L264    1953  . 

179.  Friedlander   and   Kennedy,    "Introduction    to    Radiochemistry,"    New    York. 

John  Wiley  &  Sons,  Inc.,  1949. 


612  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

chemistry.  Complexes  of  such  metals  as  uranium  and  thorium  may  be 
studied  by  application  of  natural  tracers,  but  very  careful  separations  and 
detailed  calculations  of  the  effects  of  various  isotopes  are  necessary.  Radio- 
active isotopes  also  occur  naturally  in  potassium,  rubidium,  samarium, 
lutetium,  and  rhenium.  All  these  isotopes  have  half-lives  of  the  order  of 
108-1012  years;  hence  their  activities  are  at  low  levels. 

Most  of  the  useful  tracer  elements  are  produced  artificially.  The  nuclear 
reactions  producing  the  active  isotopes  may  be  induced  by  bombardment 
with  alpha  particles,  deuterons,  protons,  neutrons,  electrons,  7-rays,  or 
x-rays.  Neutron-bombardment  reactions  produce  many  of  the  radioisotopes 
obtainable  from  the  Oak  Ridge  National  Laboratory.  Production  of  the 
radioactive  carbon  of  mass  number  14  is  illustrated  by  the  reaction 
N14(n,  p)Cu.  The  production  of  radioactive  bromine  may  also  be  effected 
by  neutron  bombardment;  in  this  case  the  reaction  takes  place  with  emis- 
sion of  7  radiation:  Br79(n,  7)Br80.  Both  the  radioactive  elements  produced 
by  these  neutron  reactions  emit  (3~  particles  at  measurable  rates.  It  should 
be  pointed  out  that  these  nuclear  reactions  are  independent  of  the  chemical 
form  of  the  target  element,  so  far  as  their  actual  occurrence  is  concerned. 
The  state  of  aggregation  and  chemical  form  do  affect  the  efficiency  of  bom- 
bardment, since  they  determine  the  number  and  position  of  atoms  within 
the  target  area. 

Since  the  actual  amounts  of  radioactive  material  produced  for  tracer  use 
are  quite  small,  ordinary  handling  procedures  are  not  applicable.  If,  how- 
ever, sufficient  quantities  of  inactive  material  of  the  same  chemical  form 
are  added,  the  active  and  inactive  portions  may  be  chemically  treated  as  a 
unit.  The  fraction  of  radioactive  material  present  may  be  found  by  meas- 
urement of  the  activity  and  weighing  of  the  entire  mass.  The  tracer  in  such 
a  case  is  contained  in  a  chemical  substance — the  "carrier" — which  holds  it 
during  manipulations  and  separations.  Carriers  with  their  radioactive 
fractions  may  be  chemically  separated  from  other  carriers  whose  chemical 
nature  is  not  objectionable,  but  whose  active  fractions  are  a  radioactive 
impurity. 

If  target  bombardment  results  in  transmutation,  so  that  the  desired  ac- 
tive product  is  not  isotopic  with  the  remainder  of  the  target,  chemical  and 
physical  means  are  useful  in  separation.  Such  common  techniques  as  ion 
exchange,  volatilization,  electrolysis,  solvent  extraction,  adsorption  on 
precipitates,  and  leaching  have  been  profitably  used.  For  example,  bom- 
bardment of  magnesium  oxide  with  neutrons  or  deuterons  produces  radio- 

180.  Wahl  and  Bonner,  "Radioactivity  Applied  to  Chemistry,"  New  York,  John 

Wiley  &  Sons,  Inc.,  1951. 

181.  Moeller,  "Inorganic  Chemistry,"  pp.  52-77,  New  York,  John  Wiley  &  Sons,  Inc., 

1952. 


PHYSICAL  METHODS  I.\   COORDINATION*   CHEMISTRY  613 

active  sodium  by  the  reactions  Mg84^,  p)NaM  and  Mg84^,  a)Na ".  rhe 
sodium  is  recovered  by  leaching  the  target  with  hot  water. 

When  the  desired  product  is  isotopic  with  the  target,  separations  are 
theoretically  possible  by  means  of  such  method-  as  gaseous  diffusion, 

thermal  diffusion,  mass  spectrography,  and  fractional  distillation"-'.  Prac- 
tically, however,  t racers  are  difficult  to  separate  from  targets  by  t hese  tech- 
niques. Szilard  and  Chalmers188  have  described  a  neutron  bombardment  of 
ethyl  iodide,  [187(n,  y)I188,  followed  by  water  extraction  of  most  of  the 
iodine  activity.  Evidently  the  energy  of  the  neutrons  is  partially  diverted 
to  break  the  carbon-iodine  bonds.  This  type  of  process  has  been  found  to 
be  applicable  to  a  number  of  radioactive  preparations.  The  necessary  char- 
acteristics of  the  process  are  rupture  of  only  those  bonds  involving  activated 
atoms,  slow"  exchange  between  the  freed  radioactive  material  and  the 
original  substance,  and  reasonable  ease  of  separation  of  the  activated  sub- 
stance in  its  new  chemical  form.  The  Szilard-Chalmers  process  has  been 
used  for  production  of  radioactivity  in  metals  by  neutron  bombardment  of 
metal  complexes.  If  the  metal  in  a  complex  does  not  undergo  appreciable 
exchange  with  uncomplexed  metal  ions  of  the  same  species,  the  radioactive 
metal  ions  produced  by  neutron  collisions  remain  free  of  complexing  during 
the  separation  process.  Successful  Szilard-Chalmers  preparations  of  radio- 
active metals  have  been  made  by  neutron  irradiation  of  salts  of  bis(ethylene- 
diamine)platinum(II),  tris(ethylenediamine)cobalt(III),  tris(ethylenedi- 
amine)iridium(III),  and  tris(ethylenediamine)rhodium(III),  as  reported  by 
Steigman184.  Mann155  has  used  bis(ethylacetoacetato)copper(II)  in  the  Szil- 
ard-Chalmers process,  and  Duffield  and  Calvin186  have  used  disalicylalde- 
hyde  o-phenylenediimine  copper(II). 

Detection  and  Measurement  of  Radioactivity 

A  typical  tracer  study  involves  introduction  of  a  tracer  of  known  activity 
and  chemical  form  into  a  system,  carrying  out  a  known  reaction  in  the  sys- 
tem, separating  the  chemical  entities,  determining  the  activity  of  each,  and 
calculating  the  deviations  from  purely  statistical  distribution.  As  an  ex- 
ample, the  work  of  Grinberg  and  Filinov187  may  be  cited.  These  authors 
prepared  radioactive  bromine  as  potassium  bromide,  KBr*,  where  the 
asterisk  denotes  the  active  element.  In  one  part  of  the  study  a  known  weighl 
of  tracer  potassium  bromide  was  added  to  a  solution  of  a  known  weighl  of 

182.  Mueller,  ibid.,  pp.  38-52. 

L83.  SzUard  and  Chalmers,  Nalun  ,  134,  462    L934 

184.  Steigman,  Ph  is.  Rev.,  59,  198  (1941). 

Mann,  Natun  ,  142,  710    1938). 
186.  Duffield  and  Calvin,  ./.  Am.  Chem.  Soc.,  68,  557,  1129  (1946). 
1^7.  Grinberg  and  Filinov.  Cnn.pi.  rend.  acad.  set.  U.R.S  >'..  23,  912    1938  ;  31,  453 
(1941). 


614  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

potassium  tetrabromoplatinate(II).  After  a  short  time  the  two  compounds 
were  separated  (e.g.,  by  precipitation  of  silver  bromide  or  [Pt(NH3)4]- 
[PtBrJ).  The  activity  of  each  was  determined  and  found  to  be  exactly  that 
dictated  by  statistical  considerations  for  the  equilibria 

K2[PtBr4]  +  KBr*  ;=±  K2[PtBr3Br*J  +  KBr; 
K2[PtBr3Br*J  +  KBr*  ^  K2[PtBr2Br2*]  +  KBr; 
K2[PtBr2Br2*J  +  KBr*  ^±  K2[PtBrBr3*]  +  KBr; 

K2[PtBrBr3*]  +  KBr*  ;=±  K2[PtBr4*]  +  KBr. 

That  is,  assuming  equimolar  amounts  of  complex  and  potassium  bromide 
four-fifths  of  the  activity  is  transferred  to  the  complex.  This  demonstrates 
rapid  exchange  between  bromide  ions  and  the  bromo  groups  of  the  complex 
and  indicates  a  lability  of  the  complex. 

The  example  just  given  points  out  the  fundamental  importance  of  ac- 
curate measurement  of  radioactivity  in  tracer  studies.  Nearly  all  common 
tracers  emit  /3~  particles,  and  some  emit  7  radiation.  Heavy,  naturally 
radioactive  elements  frequently  emit  a  particles.  All  these  types  of  radia- 
tion may  be  detected  by  the  classical  method  of  permitting  them  to  strike 
a  photographic  film,  which  on  development  shows  blackening  caused  by 
ionization  of  the  emulsion  material.  Photographic  techniques  are  useful  for 
microscopic  study  of  particle  tracks,  but  they  are  not  suitable  for  continuous 
measurement  of  radiation  rates. 

Applications 

Radioactive  tracers  have  become  increasingly  important  in  recent  years 
in  the  study  of  complexes.  Their  principal  use  has  been  in  exchange  studies, 
the  data  from  which  have  led  to  many  significant  conclusions  regarding 
bond  type.  The  example  given  above  from  the  work  of  Grinberg  and 
Filinov187  showed  rapid  exchange  between  free  bromide  ions  and  the  bromo 
groups  of  [PtBr4]=.  A  large  degree  of  ionic  character  appears  to  be  present 
in  the  platinum-bromine  bond.  The  same  series  of  studies  demonstrated 
rapid  bromide  exchange  for  the  complexes  [PtBr6]=  and  [Pt(NH3)2Br2]. 
When  radioactive  platinum  was  used,  however,  in  the  form  of  [Pt*Clc]=,  no 
metal  exchange  was  observed  with  [Pt(NH3)2Cl4],  nor  with  [Ir(py)2Cl4]  and 
[Ir*Cl6]=  or  [Ir*Cl6]=.  These  results  suggest  either  that  the  metal-chlorine 
bonds  exhibit  much  more  covalent  character  than  the  metal-bromine 
bonds,  or,  as  is  more  likely,  that  the  metal-nitrogen  bonds  in  these  platinum 
group  complexes  are  primarily  covalent.  In  the  latter  case,  regardless  of  the 
rapidity  of  the  halogen  exchange,  no  radioactive  metal  atom  could  be 
attached  to  a  nitrogen-donor  group,  since  only  the  inactive  metal  atoms 
were  originally  so  attached.  Thus  no  activity  can  appear  in  the  nitrogen- 
containing  fraction  of  the  complex  mixture. 


PHYSICAL   METHODS  IX  COORDINATION*   CHEMISTRY  615 

Polesitskii188  \\>vd  radioactive  iodine  in  his  study  of  the  tetraiodomer- 
curate(II)  complex,  formed  according  bo  the  equation 

Hgls  +  21    ;  iHglr. 

Mercury(II)  iodide  was  shaken  with  radioactive  potassium  iodide  in  one 

solution,  and  radioactive  mercury ( II  |  iodide  with  inactive  potassium  iodide 
in  another.  Silver  ion  was  added  to  precipitate  silver  iodide  and  silver 
tetraiodomercuratel  1 1 }.  Completely  statistical  distribution  of  activity  in  the 
precipitates  showed  complete  exchange  and  led  the  author  to  conclude  that 
all  four  coordination  positions  in  the  mercury(II)  complex  are  equivalent. 
Tracers  have  played  a  significant  pari  in  several  investigations  of  tris- 
(oxalato)  complexes  of  aluminum,  iron(III),  chromium(IIl),  and  co- 
balt (III).  Thomas189  suggested  that  the  resolved  form  of  the  chromium  salt 
racemizes  by  a  mechanism  whose  rate-determining  step  is 

d-  or  MCr(C204)3]s^  [Cr(C204)2]-  +  C2Or 

Thomas,  YVahl1"",  and  Burrows  and  Lauder191,  furthermore,  reported  that 
the  iron  and  aluminum  complexes  are  resolvable,  as  the  cobalt  and  chro- 
mium complexes  are  known  to  be.  Long192  and  Johnson193,  however,  were 
unable  to  confirm  these  resolutions.  In  addition,  Long  prepared  radioactive 
oxalate  by  deuteron  bombardment  of  carbon  and  successive  conversion  to 
carbon  monoxide,  carbon  dioxide,  and  oxalate.  In  solution  this  active 
oxalate  was  mixed  with  the  tris(oxalato)  complex  of  each  of  the  four  metals. 
Exchange  proved  to  be  rapid  for  iron  and  aluminum,  while  no  exchange  was 
measurable  with  cobalt  and  chromium.  These  results  indicate  predom- 
inantly ionic  bonds  in  the  iron  and  aluminum  complexes  and  predominantly 
covalent  bonds  in  the  cobalt  and  chromium  complexes.  Resolution  of  the 
first  two  complexes  therefore  seems  unlikely,  as  does  the  ionic  mechanism 
for  racemization  of  the  chromium  complex. 

An  extensive  review  of  the  use  of  tracers  in  studying  substitution  reac- 
tions in  complexes  has  been  given  by  Taube194.  The  most  important  concept 
advanced  by  Taube  is  that  the  covalent  or  ionic  character  of  metal-ligand 
bond-  is  not  the  fundamental  factor  influencing  rates  of  exchange  involving 
these  bonds.  It  is  rather  the  electron  structure  of  the  central  metal  ion 
which  exerts  a  direct  effect.  Among  the  inner  orbital  complexes,  those  hav- 
ing one  or  more  vacant  inner  d  orbitals  show  much  faster  rate-  of  substitu- 

188.  Polesitskii,  Compt    rend.  acad.  set.  U.R.S.S.,  24,  540  (193 

189.  Thomas,./.  Chem.  Soc.,  119,  1140  (1921). 

190.  W.ihl.  B<     .  60.  399  (1927). 

r.H.  Burrows  and  Lauder,  ./.  .1///.  Chem.  Soc.,  53,  3600  (1931). 

192.  Long,  •/.  .1//-.  Chem.  Soc.,  61,  570    193 

193.  Johnson.  Trans.  Faraday  Soc.,  28,  845    L932). 

194.  Taube,  Chem.  Rev.,  50,  89    r 


616  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

t ion  than  those  in  which  at  least  one  electron  occupies  each  inner  d  orbital. 
Taube  proposes  that  substitution  reactions  in  these  cases  take  place  through 
formation  of  an  intermediate  which  uses  the  vacant  orbital,  thus  increasing 
the  normal  coordination  number  by  one.  This  type  of  intermediate  can 
result  from  complexes  with  filled  d  orbitals  only  through  pairing  or  promo- 
tion of  elect  ions,  both  of  which  require  considerable  energy.  An  example  of 
the  application  of  this  concept  may  be  found  in  the  substitution  reactions 
of  vanadium  (III)  complexes,  which  have  a  vacant  d  orbital,  and  chro- 
mium(III)  complexes,  which  do  not.  The  reactions  may  be  described  in 
terms  of  electron  structure  in  the  following  manner. 

V(III)      dWdoDtSP3  -*  [dWDtSP*]  -+  dWd°D2SP3     lower-energy    intermediate; 

rapid  reaction 
Cr(III)     dWdWSP*  ->  [d*dlD*SP3]  ->  dldldlD2SP*     higher-energy      intermediate; 

slow  reaction 

Experimental  observations  confirm  the  marked  difference  in  rates  of  ex- 
change among  complexes  of  these  two  trivalent  metals. 

Complexes  of  the  outer-orbital  type,  which  are  not  subject  to  the  direct 
effect  of  d-orbital  structure,  show  a  regular  variation  in  substitution  rates 
with  charge  on  the  central  metal  ion.  Increasing  charge  corresponds  to  de- 
creasing rate  of  exchange,  and  the  secondary  effect  of  covalent  character  is 
more  important  here.  Covalent  character  likewise  accounts  for  rate  differ- 
ences in  cases  of  similar  electron  structure  among  inner-orbital  complexes, 
the  more  covalent  complexes  undergoing  slower  substitutions.  In  general, 
Taube  has  suggested  that  degree  of  covalence  is  an  index  of  substitution 
rates  when  there  is  no  significant  variation  in  electron  structure  in  the 
complexes  under  consideration,  or  when  covalent  character  has  a  direct 
influence  on  the  electron  structure.  But  covalent  character  alone  is  not  a 
reliable  guide  in  prediction  of  substitution  rates,  since  in  many  cases  its 
effects  are  opposite  to  the  determining  effects  of  electron  structure. 

Establishment  of  the  formulas  of  complexes  has  been  possible  through 
tracer  studies.  Adamson195  has  studied  the  cyano  complex  of  cobalt  (II) 
and  established  its  formula  as  [Co(CN)5]=  rather  than  [Co(CN)6]4~,  as 
previously  supposed.  The  cyano  groups  in  the  complex  show  rapid  ex- 
change (2  minutes)  with  radioactive  potassium  cyanide,  but  exchange  with 
[Co(CN)6]~  is  negligible  after  several  days.  Adamson  suggests  that  the 
cobalt (II)  complex  is  an  example  of  a  true  five-coordinate  species  in  solu- 
tion. 

Long196  has  reported  a  tracer  study  of  the  tetracyanonickelate(II)  ion, 
using  radioactive  cyanide  and  radioactive  nickel.  The  rate  of  exchange  be- 

L95.  Adamson,  ./.  Am.  Chen,.  Soc,  73,  5710  (1951). 
L96.  Long,  ./.  .1///.  Chem.  Soc.,  73,  537  (1951). 


PHYSICAL  METHODS  I\  COORDINATION  CHEMISTRY  617 

tween  the  radioactive  cyanide  and  [Ni(CN)J  ifi  Immeasurably  fast,  This 
fact  suggests  that  radioactive  nickel  ion  of  [Ni*(HiO)J++,  should  exchange 
rapidly  with  that  in  tet  racyanonickelate.  Such  is  not  the  case,  however; 
addition  of  hydra  ted  nickel  ion  to  a  solution  containing  tetracyanonickelate 

ion  results  in  the  precipitation  of  nickel  cyanide  as  a  suspension.  Then 
addition  of  dimethylglyoxime  precipitates  the  amount  of  nickel  added  ae 
[Ni*(H20)»]++,  with  no  loss  of  radioactivity.  Evidently  the  precipitated 

nickel  cyanide  actually  contains  two  unlike  kinds  of  nickel.  Long  postulates 
the  formula  \i[\i(C\u]  for  solid  nickel  cyanide. 

Johnson  and  Hall1"7  have  found  that  four-coordinate  complexes  of  nickel 
which  are  shown  by  magnetic  or  x-ray  studies  to  have  covalent  bonds  do 
not  exchange  appreciably  with  radioactive  nickel  ion.  Similarly,  the  six- 
coordinate  complexes  which  can  be  resolved  into  optical  isomers  do  not  ex- 
change, with  the  exception  of  tris(dipyridyl)  nickel(II)  ion.  This  complex 
shows  a  measurable  rate  of  exchange,  and  it  also  racemizes  measurably 
rapidly,  as  may  be  expected.  Although  bis(salicylaldoxime)  nickel  and  bis- 
(salicylaldimine)  nickel  are  diamagnetic  in  the  solid  state  and  therefore 
covalent,  both  complexes  exchange  with  radioactive  nickel  in  methyl  cello- 
Bolve  solution.  Johnson  and  Hall  interpret  this  evidence  to  signify  a  change 
of  bond  type  upon  solution. 

Using  a  tracer  method,  Cook  and  Long198  have  successfully  measured  the 
dissociation  constant  of  the  stable  complex  ion  tris(o-phenanthroline)iron 
(II),  which  is  used  analytically  as  ferroin  indicator.  Radioactive  iron  was 
used  in  preparing  the  complex.  Then  known  amounts  of  the  complex  were 
dissolved  in  known  volumes  of  water  and  treated  with  measured  quantities 
of  sulfuric  acid.  Upon  acidification  the  following  reaction  takes  place. 

[Fe(o-phen)3l++  +  3H+ ^  Fe++  +  3  H-o-phen.+ 

The  o-phenanthrolinium  ion  has  a  known  dissociation  constant,  and  the 
original  concentrations  of  complex  and  added  acid  were  known.  Xext  a 
hundred-fold  excess  of  ordinary  iron(II)  ion  was  added  to  the  solution,  and 
the  complex  was  precipitated  with  [Cdl4]=  ion.  It  was  assumed  that  precipi- 
tation was  complete  before  any  shift  in  equilibrium  took  place  and  before 
any  exchange  could  occur  between  added  iron (II)  ion  and  complexed  radio- 
active iron(II)  ion.  Both  these  assumptions  are  reasonable,  since  the  ferroin 
complex  is  quite  stable  and  slow  to  exchange.  After  precipitation,  the  total 
amount  of  radioactivity  in  the  filtrate  was  measured  and  attributed  to  the 
iron(II)  ion  originally  dissociated  from  the  complex.  The  added  excess  of 
iron (II)  ion  acted  as  a  carrier,  assuring  nearly  complete  recovery  of  the 
activity  in  solution.  The  rat  io  of  hit  rate  activity  to  original  complex  activity 

197.  Johnson  and  Hall.  ./ .  .!  .  Soc,  70,  2344     I'»48). 

198.  Cook  and  Long,  J.  Am.  Chem.  Soc,  73,  4119  (1951 


618  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

was  taken  as  the  degree  of  dissociation  of  the  complex  in  acid  solution.  All 
other  necessary  values  for  calculation  of  the  dissociation  constant  were 
known,  and  the  constant  could  then  be  found. 

[Fe(o-phen)3]++^  Fe++  +  3  o-phen 
[Fe++][o-phen]3 

^    "    L    ,     \        s      *       =     8    X    10-22 

[Fe(o-phen)3++] 

By  considering  individual  ion  activity  coefficients,  Cook  and  Long  arrived 
at  a  lower  value  of  7  X  10-22,  which  is  in  rather  good  agreement  with  the 
value  5  X  10-22  found  by  Lee,  Kolthoff,  and  Leussing199,  who  used  cell 
measurements. 

Dialysis  and  Electrolytic  Transference 

The  diffusion  of  ions  through  membranes  and  their  migration  toward 
electrodes  have  been  of  occasional  value  in  the  study  of  the  nature  of  com- 
plexes. Physical  methods  involving  these  phenomena  are  particularly  suited 
to  the  determination  of  effective  ionic  weights. 

An  ordinary  electrolyte,  when  subjected  in  solution  to  the  effect  of  an 
electric  current,  shows  the  familiar  migration  of  the  positively  charged  ion 
to  the  cathode  and  the  negative  ion  to  the  anode.  If  an  electrolytic  cell  con- 
taining such  a  system  is  divided  with  porous  walls,  or  even  imaginary 
boundaries,  into  compartments,  and  a  sample  of  solution  from  each  com- 
partment is  analyzed  after  electrolysis,  the  differences  in  concentrations  in 
the  compartments  may  be  used  to  calculate  the  fractions  of  the  current 
carried  by  each  of  the  two  kinds  of  ions  present.  These  fractions,  known  as 
transport  numbers,  are  characteristic  of  individual  ionic  species,  being  large 
for  rapidly  moving  ions  and  small  for  slow  ions. 

If  a  metal  ion  has  been  complexed  by  a  sufficient  number  of  negative 
coordinating  groups  to  render  the  overall  charge  of  the  complex  negative, 
the  electrolytic  migration  will  be  opposite  to  that  of  the  uncomplexed  metal 
ion.  Under  these  circumstances  the  formal  calculation  of  transport  numbers 
yields  a  negative  value  for  the  metal.  For  example,  the  addition  of  silver 
ion  to  an  excess  of  a  cyanide  salt,  followed  by  electrolysis,  shows  that  the 
silver  migrates  toward  the  anode  compartment.  Furthermore,  analysis  of 
the  solution  in  the  anode  compartment  shows  that  each  silver  ion  entering 
the  anode  compartment  has  been  accompanied  by  two  cyanide  ions.  These 
observations  correspond  to  the  formation  of  the  dicyanoargentate  ion. 

Ag+  +  2CN--*  [Ag(CN) si- 
ll ittorf200  has  made  transference  studies  of  several  complex  species  in 

L99.  Lee,  Kolthoff  and  Leussing,  J.  Am.  Chem.  Soc,  70,  2348  (1948). 
200.  Hittorf:  "(l>er  die  Wanderungen  der  Ionen  wahrend  der  Elektrolyse,"  Leipzig, 
W.  Engelmanh,  1912. 


PHYS/cM.   UETHODSIN  COORDINATION   CHEMISTRY  619 

solution.  His  data  for  the  tetraiodo  complex  of  cadmium,  |(MI»|  ,  show 
negative  cadmium  transport  numbers  for  concentrated  solutions,  ka  more 

water  is  added  to  the  solution,  the  cadmium  transport  number  increases  in 
value,  evidently  because  of  the  dissociation  of  the  complex  and  formation 

i>\  cat  ionic  species.  1 1  it  tort'  has  shown  thai  a  similar  dissociation  occurs  with 
the  trichloroauratel  I  I  ion.  [AuCl  ,|  . 

Electrolytic  diffusion  measurements  arc  conveniently  made  by  dialysis 
or  diffusion  of  ions  through  membranes.  Most  of  the  dialysis  studies  of 
complexes  carried  out  since  1930  are  the  work  of  Brintzinger801.  The  general 
technique  used  is  fairly  simple.  The  electrolyte  to  he  studied  is  dissolved  in 
a  solution  containing  an  excess  of  another  electrolyte  such  as  -odium  or 
potassium  chloride.  The  resulting  solution  is  placed  in  a  cup  having  a  mem- 
branous bottom.  The  cup  is  suspended  so  that  tin1  bottom  is  in  contact  with 
a  known  volume  of  solution  containing  the  foreign  electrolyte  in  the  same 
concentration  as  in  the  solution  which  also  contains  the  unknown.  Both  the 
solutions  are  stirred  for  a  known  length  of  time,  and  the  solution  in  the  cup 
is  then  analyzed.  This  procedure  is  repeated  for  the  unknown  solution,  using 
several  different  time  intervals.  Then  a  like  procedure  is  followed  for  a 
reference  electrolyte  whose  rate  of  diffusion  is  known.  The  initial  and  final 
concentrations  of  electrolyte  in  the  cup  are  used  to  calculate  the  dialytic 
constant  X  from  the  relation 

Ct  =  CV~X' 

where  Co  is  the  original  concentration  exclusive  of  foreign  electrolyte,  and 
Ct  is  the  concentration  at  time  /.  With  a  proper  choice  of  membrane  ma- 
terial, the  values  of  X  for  different  ionic  weights  obey  the  relation 

X  \/l  =  constant  . 

where  /  is  the  ionic  weight.  Thus 


i  -  fey 


where  the  subscripts  x  refer  to  the  electrolyte  to  be  determined,  and  the 
subscripts  r  indicate  the  reference  electrolyte.  This  method  is  therefore 
applicable  to  the  determination  of  ionic  weights  by  comparison  with  a 
standard. 

Brintzinger  has  reported  very  extensive  dialysis  studies  of  complex  ions 
in  the  presence  of  various  other  ions.  His  mosl  general  conclusion  is  that 
the  species  generally  regarded  as  complex,  such  as  [Co  \II:5)6]+"H"  and 
[Cot  XII    .,(  "lj~~ .  are  in  the  presence  of  other  ion-  complexed  even  further, 

201.  Brintzinger,  Z.  anorg.  allgem.  Chem.,  220,  172    1934   .  225,  221     1935  .  227,  341, 
3.51    1936  :  232.  li:»    1937  ;  256.  98    L948),  and  many  other  publications 


620  (HEM  1ST HY  OF  THE  COORDINATION  COMPOUNDS 

to  form  such  "two-shelled"  complexes  as  {[Co(NH3)6][S04]4!5_  and 
{[Co(NH3)5Cl][S04]4J6-.  The  experimentally  found  ionic  weights  for  such 
species  are  in  remarkably  good  agreement  with  those  calculated  from  the 
proposed  formulas.  There  are,  however,  certain  serious  criticisms  of  the 
method  of  dialysis.  The  most  important  of  these  is  the  fact  that  a  reference 
ion  must  be  used  in  each  experiment,  and  the  degree  of  complexing  or  hy- 
dration in  the  reference  ion  is  often  uncertain.  In  addition,  the  pore  size  of 
the  membrane  used  is  considered  by  many  workers  to  be  a  much  more  criti- 
cal variable  than  is  supposed  by  Brintzinger.  Jander202  and  Kiss203  have 
shown  to  their  satisfaction  that  slight  variations  in  pore  size  or  insufficient 
quantities  of  foreign  electrolyte  result  in  wide  variation  in  the  "dialytic 
constant."  These  criticisms  are  apparently  justified.  It  is  not  reasonable, 
however,  to  discredit  the  possibility  of  existence  of  such  two-shelled  com- 
plexes as  are  proposed  by  Brintzinger.  Laitinen,  Bailar,  Holtzclawr,  and 
Quagliano204  have  shown  that  the  half-wave  potential  of  the  hexammine 
cobalt(III)  ion  is  shifted  to  more  negative  values  in  the  presence  of  in- 
different electrolyte  anions  which  are  good  coordinating  agents,  such  as 
sulfate,  tartrate,  and  citrate.  Diffusion  rates  in  the  presence  of  these  co- 
ordinating ions  are  slower  than  with  chloride  or  nitrate.  These  findings  sug- 
gest formation  of  a  two-shelled  "super-complex"  wrhich  is  both  more  stable 
and  slower  to  diffuse  than  the  hexammine  cobalt(III)  ion.  Other  methods 
should  be  applied  to  this  problem. 

Thermal  Measurements 

The  measurement  of  temperature  has  been  useful  in  studying  partial  or 
complete  decomposition  of  coordination  compounds,  as  well  as  their  phase 
changes,  vapor  pressures,  and  other  thermodynamic  properties  such  as 
heats  of  formation,  reaction,  and  solution. 

Ephraim205  has  reported  an  extensive  series  of  studies  of  the  decomposi- 
t  ion  temperatures  of  polyhalides  and  of  ammine  complexes  of  the  transition 
elements.  His  interpretations  of  the  data  arising  from  these  studies  lead  to 
several  generalizations  concerning  thermal  stability  of  complexes. 

1 .  If  the  metal  ion  of  an  ammine  complex  may  exist  in  more  than  one 
oxidation  state,  the  higher  state  corresponds  to  the  more  stable  complex. 
This  statement  is  illustrated  by  the  much  greater  thermal  stability  of 
[Co(NH8)6]Cl3  as  compared  with  [Co(NH3)6]Cl2 . 

2.  Divalent  metals  of  small  ionic  volume  show  greater  tendencies  toward 
complex  formation  than  those  of  larger  ionic  volume,  and  their  complexes 

202.  Jander  and  Spandu,  Z.  physik.  Chem.}  A188,  65  (1941). 

203.  Kiss  and  Acs.  Z.  anorg.  allgem.  Chew.,  247,  190  (1941). 

204.  Laitinen,  Bailar,  Holtzclaw,  and  Quagliano, /.  Am.  Chew.  Soc,  70,  2999  (1948). 
206.   Ephraim,  Ber.,  36,  1177,  1815,  1912  (1903);  Z.  phys.  Chew..  81,  513,  539  (1912); 

83,  196  (1913);  84,  98  (1913) ;  Ber.,  45,  1322  (1912);  50,  1069  (1917);  Ephraim  and 
Wagner,  Ber.,  50,  1088  (1917);  Ephraim  and  Muller,  Ber.,  54B,  973  (1921). 


PHYSICAL  METHODS  IX  COORDINATION  CHEMISTRY  621 

are  more  stable.  The  hexammines  of  divalenl  manganese,  cobalt ,  nickel  and 
iron  follow  the  relationship  (FT7)  "'  =  constant,  where  V  is  the  ionic  volume 
and  T  is  the  absolute  decomposition  temperature.  Other  hexammine  com- 
plexes obey  the  relationship  only  approximately. 

3,  Hexammine  complex  salts  containing  large  anions  are  more  stable 
than  their  analogs  containing  smaller  anions.  For  example,  in  the  series 
[Ni(NH8)JXj,  the  chloride  decomposes  at  L64°C,  the  bromide  at  L95°C, 
and  the  iodide  at  221°C. 

1.  Ammine  complex  salts  containing  Large  anions  tend  to  show  an 
increased  coordination  number  in  the  cation,  so  that  the  disparity  in 
size  of  the  cation  and  anion  is  a  minimum.  For  example,  |\'i(  \'f  bOe] 
[Co(\II;;M  \()Ai|  is  difficultly  crystallized  from  solution,  bul  addition  of 
ammonia  results  in  crystallization  of  [Nil  \H:{).s|[(\>(XII:,)-j(\<  h)*]>  It  is 
questionable  whether  the  additional  ammonia  molecules  are  truly  coordi- 
nated to  the  nickel  ion;  they  are  more  likely  to  be  held  merely  by  the  re- 
quirements of  the  crystal  lattice. 

The  work  of  Hilt z-"6  is  important  among  thermal  studies  of  complexes. 
This  work  will  not  be  discussed  in  detail  here,  hut  it  should  be  mentioned 
that  Hilt z  has  collected  significant  phase  transition  data  from  studies  of 
Stepwise  dissociations  of  hexammine  complexes,  performed  at  either  con- 
stant pressure  or  constant  temperature.  Divalenl  hexammines  in  general 
decompose  directly  to  diammines,  without  intermediate  stepwise  loss  of 
coordinated  groups.  The  diammines  usually  have  a  greater  relative  thermal 
stability  than  do  the  hexammines. 

Phase-change  measurements  may  also  be  made  with  solutions  of  complex 
compounds.  Hagenmuller207  has  used  cryoscopic  measurements  of  aqueous 
solutions  of  nitrite  complexes  as  the  basis  of  continuous  variations  analyses. 
Deviations  of  freezing  points  from  additivity  indicate  the  existence  of 
[Hg(X()2)4]=  [Cd(X02)4]=  [Cd(NOf)i]-,  [Cu(N02)8]-  [PMXO^h  and 
[Pb(X()2):{]~.  Hagenmuller  assumes  that  the  trinitrite  complexes  are  singly 
hydrated  to  complete  the  coordination  sphere. 

Other  Methods 

Many  other  physical  methods  have  received  infrequent  attention  in  the 
Btudy  of  coordination  compounds.  Most  of  these  methods  are  not  suited  to 
wide  application  in  this  field;  they  are  instead  particularly  adaptable  to 
certain  unusual  types  of  problems.  Several  example-  of  the  use  of  such 

method-  arc  given  below. 

Gustavson2  '  has  carried  out  identification  and  separation  of  basic  salts 

206.  Biltz,  Z.  phyeik.  Chem.,  67,  561     L909  ;  Z.  anorg.  Chem.,  109, 132    L920 

207.  Hagenmuller,  Ann.  chim.,  6,  5    1951 

208.  Gu8tavsoi     v  Kem.Tid., 66, 14  (1944);/. Intern.  Soc  Leather  Trades  Chem., 

80,264    1946 


622  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

of  chromium  (III),  using  selective  adsorption  on  ion  exchange  columns. 
Since  the  basic  salts  consist  of  mixtures  of  complexes  of  both  negative  and 
positive  charges,  depending  upon  the  number  of  hydroxo  groups  within  the 
coordination  sphere,  both  cationic  and  anionic  exchange  treatments  are 
necessary  for  separation.  Elution  of  the  adsorbed  complexes,  followed  by 
analysis  of  the  clnatc,  determines  the  composition  of  both  the  cationic  and 
anionic  complexes  present.  Gustavson  has  used  this  method  to  study  basic 
chromium  chlorides,  sulfates,  oxalates,  and  thiocyanates. 

Mel  lor209  has  proposed  that  ion  exchange  resins  be  prepared  with  complex- 
forming  ligands  polymerized  into  their  structure  so  that  some  donor  groups 
are  left  free.  Trace  quantities  of  metal  ions  could  then  be  removed  from  a 
solution  passed  through  such  a  resin. 

Continuous  variations  studies  with  solution  surface  tension  as  a  variable 
have  been  carried  out  by  Arcay  and  Marcot210  and  by  Kazi  and  Desai211. 
Arcay  and  Marcot  report  the  formation  of  compounds  having  the  compo- 
sitions 2HgCl2-KCl,  HgCVKCl,  and  HgCl2-2KCl,  while  Kazi  and  Desai 
conclude  that  CdI2KI  and  CdI2-2KI  form  in  solution. 

Resolution  of  optically  active  complexes  in  solution  has  been  accom- 
plished in  some  instances  by  shaking  the  solution  with  finely  ground  crys- 
tals of  one  optical  isomer  of  quartz.  Columns  packed  with  the  ground  quartz 
have  also  been  used.  In  either  case  a  selective  adsorption  effect  is  responsi- 
ble. Sometimes  the  effect  seems  to  be  of  a  true  equilibrium  nature,  since  the 
time  of  contact  with  the  quartz  is  immaterial  so  long  as  it  is  sufficient  to 
bring  about  appreciable  adsorption.  In  other  cases,  however,  the  selectivity 
appears  to  take  place  kinetically,  with  one  isomer  adsorbed  more  rapidly, 
but  both  adsorbed  equally  after  a  long  period  of  time.  In  numerous  other 
instances  no  separation  has  been  achieved  by  the  use  of  this  method.  Kara- 
gunes  and  Coumoulos212  have  resolved  tris(ethylenediamine)  chromium  (III) 
chloride  with  quartz.  Tsuchida213  has  used  the  method  to  resolve  chloro- 
bis(dimethylglyoximino)-ammine-cobalt(III).  Frequent  applications  of 
quartz  resolution  have  been  made  by  Bailar  and  his  co-workers214.  Only 
partial  resolutions  have  been  achieved  by  this  method. 

Biltz  and  Stollenwerk215  have  employed  a  pressure  method  to  study  the 

209.  Mellor,  Australian  J.  Sci.,  12,  183  (1950). 

210.  Arcay  and  .Marcot.  Compt.  rend.,  209,  881  (1939). 

211.  Kazi  and  Desai,  Current  Sci.,  India,  22,  15  (1953). 

212.  Karagunea  and  ( loumoulos,  Nature,  142, 162  (1938) ;  AttiX0  Congr.  Intern.  Chim., 

2,  278  (1938). 

213.  Tsuchida,  Kobayashi,  and  Nakamura,  ./.  Chem.  S„r.  Japan,  56,1339  (1935); 

Tsuchida.  Kobayashi,    and    Nakamura,  Bull.  Chan.  Sac.  .In pun,  11  (1),  38 
1936 
21  I    Sec.  for  example,  Buscfa  and  Bailar,/.  .1///.  Chem.  Soc,  76,  4574  (1953);  Kuebler 

and  Bailar,  ibid.,  74,  3535  (1952);  Bailar  and  Peppard,  ibid.,  62,  105  (1940). 
215.   I'.ilt/.  and  Stollenwerk.  /.  anorg.  allgem.  ('Inn,.,  114,  174  (1920). 


PHYSICAL  METHODS  IN  COORDINATION  CHEMISTRY  623 

formation  of  Bilver  ammine  complexes.  These  invesl igatora  passed  ammonia 
into  an  evacuated  vessel  containing  Bilver  chloride.  The  gaseous  pres- 
sure was  observed  to  rise  steadily  until  a  reaction  took  place  between  the 

and  solid.  During  the  reaction  the  pressure  remained  nearly  constant, 
and  then  it  rose  again.  Since  the  quantity  of  ammonia  admitted  at  any  time 
was  known,  the  quantity  combined  with  the  solid  could  be  calculated  from 
the  pressure  data.  The  results  give  evidence  tor  the  formation  of  Ag(  '1  •  XI I  . 
2AgCl-3NH3)  and  A.gCl*3NH  .  The  ordinary  ammine  complex,  corre- 
sponding to  AgCl*2NH    .  doe-  not  appear  to  form  under  these  condition-. 

When  solutions  of  two  metal  salts  are  mixed  to  form  an  ideal  solution, 
the  volume  of  the  final  solution  is  equal  to  the  sum  of  the  volumes  of  the 
component  solutions.  If  there  is  complex  formation  between  the  two  -alts, 
however,  a  non-ideal  solution  results,  whose  volume  is  not  the  sum  of  the 
original  volumes.  Davis  and  Logan*1'  have  identified  reaction-  of  metal- 
pyridine  complexes  with  cyanate  and  thiocyanate  ions  by  noting  contrac- 
tions in  volume.  Among  the  metals  tested,  the  copper(II)  complexes  are 
characterized  by  the  least  contraction  upon  addition  of  cyanate  or  thio- 
cyanate solutions.  Cobalt(II)  complexes  are  intermediate,  and  nickel(II) 
complexes  <how  the  greatest  contractions.  The  addition  of  cyanate  causes 
a  greater  contraction  than  the  addition  of  thiocyanate.  Davis  and  Logan 
advance  the  hypothesis  that  the  amount  of  contraction  may  be  related  to 
the  degree  of  metal-ligand  affinity  in  these  instance-. 

Slightly  soluble  salts  are  normally  somewhat  more  soluble  in  concen- 
trated solution-  of  other  salts,  because  of  the  increased  ionic  strength  of  the 
solution  and  the  correspondingly  decreased  activity  coefficients  of  the  ions 
of  the  slightly  soluble  salt.  Sometimes,  however,  abnormal  increases  in 
solubility  indicate  complex  formation.  Hayek217  has  concluded  from  sol- 
ubility studies  that  the  increased  solubility  of  mercury(II)  iodide  and 
mercurv(II)  oxide  in  mercury  salt  solutions  is  a  result  of  complexing.  A  com- 
petition appears  to  exist  between  the  water  molecules  of  the  hydrated  mer- 
cury(II)  ion-  and  the  neutral  mercury(II)  oxide  or  mercury(II)  iodide 
molecule-.  Coordination  of  these  molecules  to  form  BUCh  8p 
[Hg(Ugh)x(H.<))v}++  and  [Hg(HgO),(HiO)y]++  accounts  for  the  increased 
solubility.  Hayek  suggests  that  the  complexes  [Hg(HgIa  <  0  and 
Hg  Hg< )  .mCK);  ,  form  in  mercury  II  |  perchlorate  solution  in  the  presence 
of  the  respective  -lightly  soluble  mercury  compounds.  This  explanation 
agrees  substantially  with  the  proposal  of  Sidgwick  and  Lewis*18  concerning 
bility  of  beryllium  oxide  in  beryllium  .-alt  solutions  through  formation 
of  complexes  of  the  type  \R{    BeO  J"*"1". 

2n;.  I)  vu   ■■.  :  Log  n,J.  .  58,  2153 

-•17.  Bayek    Z  223.  382 

218.  Sidgwick  and  Lewis,  •/.  '  .  1287    ! 


624  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Immiscible  solvent  distribution  studies  have  been  reported  by  Sinha  and 
Ray219,  who  investigated  pyridine  complexes  of  copper(II).  Pyridine  and 
benzene  were  added  to  solutions  of  copper(II)  perchlorate,  and  the  distribu- 
tion of  pyridine  between  the  aqueous  and  benzene  phases  was  measured  as 
a  function  of  the  total  quantity  of  pyridine.  The  amount  of  coordinated 
pyridine  was  calculated  from  the  known  distribution  coefficients  for  the  two 
solvents.  When  the  total  amount  of  pyridine  had  any  value  between  ten 
and  thirty  times  the  amount  of  copper  salt,  only  the  dipyridine  and  tetra- 
pyridine  complexes  were  observed  to  form.  Related  studies  by  Macdonald, 
Mitchell,  and  Mitchell220,  with  iron(III)  thiocyanate  complexes  in  an  ether- 
water  system,  indicate  that  from  one  to  six  thiocyanate  groups  may  coordi- 
nate with  the  iron  (III)  ion,  forming  all  the  complexes  in  the  series 
[Fe(SCN)]++  to  [Fe(SCN)6]= 

Complex  formation  in  solutions  containing  lead  nitrate  and  either  potas- 
sium or  ammonium  nitrate  is  indicated  by  the  compressibility  studies  of 
Venkatasubramanian221.  This  investigator  measured  ultrasonic  velocities  in 
the  solutions  and  estimated  the  compressibilities  of  the  solutions  as  a  func- 
tion of  composition.  Minima  in  the  compressibility-composition  curves 
corresponded  to  formation  of  Pb(N03)2-KN03 ,  Pb(N03)2-2KN03 , 
Pb(N03)2-4KN03 ,  Pb(N03)2-NH4N03 ,  and  Pb(N03)2-2NH4N03 . 

219.  Sinha  and  Ray,  /.  Indian  Chem.  Soc,  25,  247  (1948). 

220.  Macdonald,  Mitchell,  and  Mitchell,  /.  Chem.  Soc,  1574  (1951). 

221.  Venkatasubramanian,  Current  Sci.,  India,  20,  13  (1951). 


\/.  Coordination  Compounds  in 
Electrodeposition 

Robert  W.  Parry 

University  of  Michigan,  Ann  Arbor,  Michigan 

and 
Ernest  H.  Lyons,  Jr. 

The  Principia,  Elsah,  Illinois 

Coordination  compounds  are  widely  used  in  electrodeposition.  Deposits 
obtained  from  the  simple  salt  solutions  are  sometimes  loose,  nonadherent, 
coarsely  crystalline,  and  generally  undesirable,  while  metal  deposits  from 
appropriate  complex  salt  solutions  are  often  smooth,  adherent,  and  of  high 
protective  and  decorative'  value. 

The  methods  used  in  developing  suitable  plating  baths  are  largely  em- 
pirical; the  art  of  electrodeposition  is  far  ahead  of  its  science.  Thompson1 
suggested  that  further  progress  in  the  development  of  the  science  of  electro- 
deposition might  be  achieved  by  a  systematic  application  of  Werner's 
coordination  theory. 

The  Theory  of  Electrodeposition  from  Complex  Compounds 

The  mechanism  of  electrode  reactions,  even  for  the  so-called  simple  ions, 
is  a  subject  of  great  complexity.  As  yet  no  theory  can  adequately  explain 
all  phases  of  the  cathodic  evolution  of  hydrogen  from  dilute  acid2.  It  is  not 
surprising  that  the  much  more  complex  phenomenon  of  metal  deposition  is 
not  well  understood3.  The  most  widely  used  coordination  compounds  in 
commercial  electrodeposition  are  the  anionic  metal  cyanides,  such  as 
[Ag(CXj-j]~  and  [Cu(CN)s]".  Many  investigators  have  found  it  difficult 

1.  Thompson,  Trans. Electrochem.  Soc.,79, 417  (1941). 

2.  Bockris,  ./.  Electrochem.  Soc.,  98,  No.  11,  L63c  (1951);  Bockrie  and  Potter,  J 

Electrochem.  Soc.,  99,  169  (1962  ;  Eyring,  Glasstone,  and  Laidler,  Trans. 
Electrocht  -  76,  I  15  1939);  Sickling  and  Bait,  Trans.  Faraday  8oc.t  38, 
171    1942  . 

3.  Blum,  Beckman,  and  Meyer,  Trai  -  80,  287  (1941). 

625 


626  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

to  picture  the  reduction  of  a  negatively  charged  complex  on  a  negatively 
charged  cathode  surface. 

The  Alkali  Metal  Reduction  Hypothesis.  One  of  the  earliest  mecha- 
nisms1, usually  attributed  to  Hittorf,  suggested  that  positively  charged 
potassium  ions  are  initially  reduced  to  give  potassium  metal,  and  that  the 
discharged  potassium  metal  reduces  silver  from  the  cyanide  complex.  No 
direct  experimental  evidence  was  ever  produced.  It  is  highly  improbable 
that  alkali  metal  could  plate  out  first5  unless  the  free  energy  of  the  solid 
alkali  is  greatly  lowered  by  instantaneous  alloy  formation  on  the  electrode 
surface.  Such  alloy  formation6  may  occur  with  electrodes  such  as  mercury 
and  possibly  lead,  but  is  highly  improbable  for  other  metals.  The  hy- 
pothesis is  now  obsolete. 

A  rather  similar  hypothesis7  assumes  that  nascent  hydrogen  is  liberated 
from  the  alkaline  solution  and  reduces  the  silver  cyanide  complex  in  a 
secondary  chemical  process.  No  unequivocal  evidence  to  support  or  refute 
such  a  mechanism  is  available.  Butler8  suggests  that  such  a  mechanism  is 
apparently  operative  in  some  electrolytic  organic  reductions.  An  extension 
to  complex  compounds  is  speculative. 

The  Dissociation  of  the  Complex  to  give  "Simple"  Metal  Ions. 
This  concept  might  be  called  the  classical  picture  of  complex  ion  reduction. 
It  is  assumed  that  complex  ions  dissociate  to  give  low  concentrations  of 
simple  metal  cations  which  can  be  reduced  at  the  cathode9,  10-  n. 

[Ag(CN)2]-->  Ag+  +  2CN- 
Ag+  +  e~  -»  Ag 

The  concept  apparently  developed  from  application  of  thermodynamic  in- 
stability constants  to  the  calculation  of  electrode  potentials  in  the  presence 
of  complex  ions.  In  most  cases  experimental  differentiation  between  this 
mechanism  and  direct  reduction  of  the  complex  has  not  been  achieved; 
however,  some  evidence  to  support  the  dissociation  hypothesis  has  been 
cited.  From  very  dilute  solutions  of  silver  nitrate  or  copper  sulfate,  ranging 

4.  Classen  and  Hall,  "Quantitative  Analysis  by  Electrolysis,"  5th  ed.  p.  48,  New 

York,  John  Wiley  &  Sons,  Inc.,  1913;  Dean  and  Chang,  Chem.  Met.  Eng.,  19, 
83  (1918);  Hedges,  /.  Chem.  Soc,  1927,  1077;  Levasseur,  Technique  Moderne, 
19,  29  (1926). 

5.  Glasstone,  J.  Chem.  Soc.,  1929,  690,  702;  Sanigar,  Rec.  trav.  chim.,  44,  556  (1925). 

6.  Piontelli,  Gazz.  chim.  ital,  69,  231  (1939). 

7.  Jolibois,  Helv.  chim.  Acta,  23,  412  (1940);  Jolibois,  Compt.  rend.,  225,  1227  (1947). 

8.  Butler,  "Electrocapillarity,"  p.  199,  London,  Methuen  and  Co.  Ltd.,  1940. 

9.  Spitzer,  Z.  Elektrochem.,  11,  345;  391  (1905). 

10.  Petrocelli,  Trans.  Electrochem.  Soc,  77,  133  (1940);  Stout  and  Faust,  Trans. 

Electrochem.  Soc,  61,  341  (1932). 

11.  Levin,  ./.  Gen.  Chem.,  U.S.S.R.,  14,  31  (1944);  </.  Phys.  Chem.,  U.S.S.R.,  18, 

53  (1944);  cf.,  Chem.  Abs.,  39,  1597  (1945). 


COORDINATION  COMPOUNDS  l\   ELECTRODEPO&ITIOh  627 

in  concentration  from  10  ,;  to  10  "'.Y,  finely  crystalline,  adherent  deposits 
of  silver  or  copper  can  be  deposited  by  allowing  the  solul  ion  to  flow  rapidly 
between  charged  electrodes12,  :  .  The  size  of  crystallites  in  silver  deposits 
obtained  from  silver  nitrate  solutions  decreased  as  the  concentration  of 
silver  nitrate  was  reduced  from  10  "  to  10  '.V.  Bancroft"  stated  thai  de- 
posits become  more  finely  crystalline  as  the  potential  difference  between 
the  metal  electrode  and  the  solution  is  increased,*  bul  extension  to  the 
mechanism  of  silver  cyanide  reduction  is  certainly  open  to  question. 
Theoretical  arguments  have  been  used  against  the  hypothesis. 

From  the  equilibrium  constant  for  the  reaction  [Ag(CN)a]"  ^  \&'  -\- 
3CN~ 15,  Haber16  calculated  the  ratio  between  time  of  formation  and  time 
oi  dissociation  of  the  complex  ion.  This  ratio  is: 

Time  of  formation  of  complex 

-  =   Kenuilib.    =    1.3  X    10"22 


Time  of  dissociation  of  complex 


It  was  shown  that  if  the  time  of  formation  for  a  given  amount  of  complex 
is  10~3  or  10~4  seconds,  more  than  a  thousand  years  are  required  for  dis- 
sociation of  the  same  amount  of  complex.  Such  a  situation  precludes  electro- 
deposition  of  silver  by  dissociation  of  the  cyanide  ion. 

Alternatively,  the  time  of  dissociation  of  a  complex  ion  may  be  set  at 
10~-  seconds  or  any  other  reasonable  value  to  permit  dissociation  before 
deposition,  and  the  time  of  formation  of  the  ion  may  be  calculated.  Such 
a  calculation  shows  that  the  complex  ion  must  form  in  less  than  10-22 
seconds.  If  the  coordinating  anions  move  at  least  an  atomic  diameter  (about 
10_s  cm)  to  form  the  complex,  they  must  have  velocities  several  million 
times  greater  than  that  of  light.  The  situation  is  not  altered  by  substituting 
thedicyanide  for  the  tricyanide  of  silver.  Haber  concluded  thai  reduction  of 
silver  must  take  place  by  direct  reduction  of  the  anion  and  not  by  an  inter- 
mediate dissociation  process. 

Similar  conclusions  were  drawn  from  studies17  of  current-voltage  curves 
for  the  reduction  of  [Cu(CX)3]=. 

*  Glasstone  and  Sanigar11  have  shown  that  the  correlation  between  electrode  po 
tential  and  the  physical  properties  of  the  deposit  is  not  rigorous.  The  physical  proper- 
ties of  silver  deposited  from  argentocyanide  solutions  containing  Na+,  K    or  anions 
such  as  PO<  ,  CO    .  ><  l .  .  could  not  be  correlated  with  the  small  changes  in  elec- 
trode potential  which  accompanied  the  introduction  of  these  ions  to  the  solution. 

12.  Vahramian  and  Alemyan,  ./.  Phya.  Chem  .  [  .8  S.R.,  9,  517    1937  ;  Acta  PI 

chimica,U.S.S.R.,7, 95    1937  ; cf., Chem. Aba., 31, 6975    L937  ; 32, 2844    1938 

13.  Bancroft../.  Ph        <  9,290    1*>05). 

14.  I  and  Sanigar,  Trm  5        85,  " 

15.  Bodlander  and  Eberlin,  Z.  anorg.  Chem.,  39, 197    1904  . 

16.  Haber, Z.  I  10,  133    1904  . 

17.  Masing,  Z.  El  48,  85    L942  . 


628  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Since  such  calculations  are  based  on  questionable  assumptions,  an  ex- 
perimental answer  to  the  question  has  been  sought. 

The  Direct  Reduction  of  the  Complex  Ion.  The  direct  process15,  18>  19 
for  representative  complex  ions  is  shown  in  the  following  equations: 

[Ag(CN)2]-  +  e-  -»  Ag  +  2CN- 

[Cu(NH3)2]+  +  e-  ->  Cu  +  2NH3 

This  assumes  reduction  of  a  negatively  charged  anion  at  a  negatively 
charged  electrode5*,  which  is  reasonable  since  a  negatively  charged  cyanide 
ion  may  be  attracted  and  bound  to  a  complex  ion  which  already  bears  nega- 
tive charge: 

[Ou(ON)J-  +  CN-  ->  [Cu(CN),]- 

In  such  cases  localized  charge  distribution  may  be  of  more  importance  than 
the  over-all  ionic  charge. 

Furthermore,  certain  complex  anions  undergo  direct  cathodic  reduction.! 
In  the  reduction  of  potassium  ferricyanide  at  a  platinum  microelectrode, 
the  rate  of  reduction  is  controlled  by  the  rate  at  which  ferricyanide  ions 
diffuse  to  the  electrode  surface20.  Radioactive  iron(III)  ion  does  not  ex- 
change with  ferricyanide  ion  at  an  appreciable  rate21;  thus  no  dynamic 
equilibrium  exists  between  iron  ions  in  the  complex  and  iron  ions  in  solu- 
tion. Similar  observations  were  made  for  iron(II)  ions  and  ferrocyanide. 
Since  the  equilibrium 

[Fe(CN)6j=  ^±  Fe+++  +  6CN~ 

is  established  very  slowly,  it  cannot  be  regarded  as  essential  to  the  cathode 
reaction.  Moreover,  ferrocyanide,  as  well  as  the  corresponding  cyanides  of 
nickel  and  cobalt,  can  be  reduced  electrolytically  to  give  almost  quanti- 
tative yields  of  complex  cyanides  containing  univalent  iron,  cobalt,  or 
nickel22.  The  fact  that  ferrocyanide  is  not  in  labile  equilibrium  with  iron (II) 
ions  in  solution  makes  a  mechanism  involving  previous  dissociation  un- 
tenable. An  iron  alloy  may  be  deposited  from  a  solution  containing  iron 
only  as  K3[Fe(CX)6]10b.  Thus,  though  ferricyanide  ions  do  not  dissociate 
readily  to  produce  hydrated  iron(III)  ions  or  other  complexes,  the  entire 

t  In  using  (ho  term  "direct  cathodic  reduction"  no  definite  mechanism  for  the 
electron  transfer  is  implied. 
is.   Bodlander,  Z.  Elektrochem.,  10,  604  (1904);  Foerster,  "Electrochemie  Wasseriger 
Losungen,"  3rd  ed.,  p.  229,  footnote  1,  Leipzig,  J.  A.  Barth,  1922. 

19.  Newton  and  Furman,  Trans.  Electrochem.  Soc.,  80,  26  (1941). 

20.  Laitinen  and  Kolthoff,  J.  Am.  Chem.  Soc.,  61,  3344  (1939). 

21.  Thompson,  ./.  .1///.  Chem.  Soc.,  70,  L046  (1948). 

22.  Treadwell  and  Huber,  Helv.  chin,.  Acta,  26,  10  (1943). 


COORDINATION  COMPOUNDS  IN  ELECT  RODE  POSIT  ION  629 

anion  can  be  reduced  to  give  iron  in  the  divalent ,  moncn  alenl ,  or  zero  valenl 
state. 

The  cathodic  reduction  of  negative  ions  is  likewise  observed  with  the 
cyano  complexes  of  manganese2*,  molybdenum14,  chromium**,  tungsten16, 

and  platinum-7.   Kates  of  Substitution   reactions  with   these  ions-'  indicate 

that  they  are  not  in  mobile  equilibrium  with  the  coordinating  groups,  ;i 
conclusion  confirmed  in  Borne  instances  by  radioactive  tracer  experiments*9. 
Other  examples  are  the  electroreduction  of  citrate  complexes  of  copper10,  of 

plumbate*1,  of  stannate,  and  of  eliminate.  A  large  number  of  organic  anions 
are  also  reduced  at  the  cathode. 

A  good  metallic  deposit  of  cobalt  can  he  obtained  with  lii<2;h  current 
efficiency  from  solutions  of  [C,o(pn)2Cl2J+  and  [Co(en)3]+++ 32,  yet  Flaj 
found  no  exchange  between  simple  radioactive  cobalt (II)  ion  and  the  pro- 
pylenediamine  complex.  Since,  at  room  temperature,  racemization  of  the 
optically-active  [Co  enj+++  complex  in  water  solution  requires  several 
weeks,  equilibrium  between  the  ethylenediamine  complex  and  cobalt  ions 
in  solution  or  in  other  complexes  must  be  established  very  slowly.  A  thin 
chromium  plate  can  be  obtained  from  ammonium  trisoxalatochromi- 
um(III)34,  yet  exchange  between  the  complex  ion  and  radioactive  oxalate 
ions  in  the  solution  is  very  slow35,  showing  that  there  is  no  labile  equilibrium 
between  the  complex  and  simple  chromium(III)  ions.  Metal  deposition 
apparently  occurs  through  reduction  of  the  anion  complex. 

Deposition  from  these  compounds  probably  proceeds  through  a  lower 
valence  state.  Thus,  in  the  reduction  of  [Co(XH:5)6]+++,  [Co(XH:j)5X02]++, 
[Co(\H3)4(X02)2]+,  [Co(XH3)3(X02)3],  [Co(XH3)2(X02)4]-  and  related 
aquo  and  chloro  ammines,  the  polarographic  waves36  consist  of  two  parts. 

23.  Grube  and  Brause,  Ber.,  60,  2273  (1927). 

24.  Collenberg,  Z.  phygflc.  Chem.,  146,  81,  177  (1930);  Kolthoff  and  Tomiscek,  ./. 

Phys.  Chem.,  40,  247  (1936). 

25.  Hume  and  Kolthoff,  •/.  Am.  Chem.  Soc.,  65,  1897  (1943). 

26.  Collenberg,  Z.  phyeik.  Chem.,  109,  353  (1924). 

terre      ./.  Chem.  Soc.,  1928,  202. 
'  hem.  Rev.,  50,  69  (1952). 
_      Menken  and  Garner,/.  Am.  Chem.  Soc.,  71,  371  (I'M1' 
30.  Kalousek,  Collection  Czechcelav.  Chem.  Commune.,  11,  ~>!»2  (1939 
•  il    Glaastone  and  Hickling,   "Electrolytic  Oxidation   and   Reduction,1'   London. 
Chapman  and  Hall.  Ltd.,  1935;  Latimer,  "Oxidation  States  of  the  Elements," 
W'a  Y,,rk.  Prentice-Hall,  Inc.  1928. 
Elramer,  Swann,  and  Bailar,  Trane.  Electrochem.  Soc.  90.  55  -  L946 
Flagg,  J.Am.  Chem.  Soc.,  63,  557    L941). 
34.  Mazsucchelli  and  Baeci,  Oazz.  ehim.  ital.,  62,  7:>n    L932 
Long,  ./.  .1/-.  Chi  m.  Soc.,  61,  570    1939 

Kolthoff  and  Lingane,  "Polarography,"  p.  285,  New  York,  [nterscience  Publish 
ers,  Inc.  L941;  Laitinen,  Bailar.  Holtzclaw,  and  Quagliano,  ./.  .1///.  Chem. 
70,  -         1948  :  Willis.  Friend,  and  Mellor,  ./.  Am.  I  /<<  m.  Soc.,  67,  1680 
1945  . 


630  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

The  first,  corresponding  to  a  gain  of  one  electron,  apparently  represents  re- 
duction to  the  coball  'II)  state,  and  the  second,  corresponding  to  two  elec- 
tions, represents  reduction  to  the  metal.  The  half-wave  potential  of  the 
latter  is  always  very  nearly  that  of  the  aquated  cobalt(II)  ion,  which  is 
presumably  formed  because  cobalt(II)  ammines  are  unstable: 

[Co(NH3)6]+++  +  e~  -+  [Co(NH3)6]++ 
stable  unstable 

[Co(NH3)6]++  +  6H20  -+  [Co(H20)6l++  +  6NH3  (very  rapid) 

[Co(H20)6]++  +  2e-  -»  Co  +  6H20. 

As  explained  later,  a  two-step  process  is  likely  for  other  cobalt(III)  and 
chromium(III)  compounds,  and  possibly  for  chromate. 

From  thiosulfate  solutions,  good  deposits  of  copper  and  zinc  may  be  ob- 
tained37, but  cadmium  from  thiosulfate  contains  up  to  5  per  cent  sulfur,  and 
nickel  from  22  to  70  per  cent  sulfur.  X-ray  analysis  of  the  nickel-sulfur  de- 
posil  indicates  the  presence  of  nickel  sulfides  such  as  Ni2S3 .  The  deposition 
of  semi-crystalline  nickel  sulfide  suggests  that  dissociation  of  the  thiosulfate 
complex  does  not  precede  reduction  of  the  nickel  ions. 

Similarly,  nitrogen  has  been  detected  in  a  copper-lead  alloy  plate  from  a 
solution  containing  ethylenediamine  complexes38.  Up  to  17  per  cent  of 
halogen  has  been  found39  in  deposits  of  antimony,  cadmium,  bismuth,  cop- 
per, and  tin  obtained  from  halide  solutions  of  the  metal  ions.  Thus,  with 
stable  complexes,  reduction  appears  to  occur  directly  from  the  complex  ion. 
For  complexes  such  as  [Ag(CX)2]~,  which  is  in  labile  equilibrium  with  the 
Ag+  and  CN~  ions,  experimental  demonstration  of  the  mechanism  is  not 
conclusive;  however,  theoretical  considerations  favor  direct  reduction. 

Deposits  are  usually  not  obtained40  from  aqueous  solutions  of  complex 
ions  with  electronic  configurations  involving  hybridized  orbitals  from  the 
inner  electron  shells,  that  is,  the  "inner  orbital"  ions  of  Taube28.  From  ions 
of  "outer  orbital"  configuration,  deposits  are  generally  obtained.  This  rule 
holds  for  aquo  complexes  as  well  as  for  others,  and  suggests  that  the  com- 
plex ion  is  directly  involved. 

Reduction  of  an  Intermediate  Complex  Cation.  To  avoid  difficulty 
due  to  charge  repulsion  at  the  cathode,  Glasstone6*** 4I  suggested  that  a 
complex  cation  is  formed  from  the  complex  anion;  this  cation  then  under- 

:;;.  Gernes,  Lorenz,  and  Montillon,  Trans.  Electrochem.  Soc. 77,  177  (1940). 
38    Etoszkowski,  Hanley,  Schrenk  and  Clayton,  Tinny.  Electrochem.  Soc,  80,  235 
L941). 
-■.»ii<\  thesis,  [ndiana  University. 
Mi.  Lyons,/.  Electrochem.  Soc,  101,  363,  :>7(i,  (1964);  Lyons,  Bailar,  and  Laitinen, 

ibid,  101,  IK)  (1964). 
•11.  Glasstone,  /.  Clu  m.  Soc.,  1930,  1237. 


COORDINATION  COMPOl  NDS  l\   ELECTRODEPOSITIOh  631 

goes  cathodic  reduction: 

2[Ag(ClS  ■    v.  i  \  M  N 

[Ag,CN]+  +  e       •   \v    ■    ^gCN 

The  existence  of  cationic  complexes  in  iodide  or  cyanide  solutions  contain- 
ing an  excess  of  silver  ion  is  fairly  well  established41,  '-,  bul  the  presence  of 
appreciable  amounts  of  the  complex  cation  in  plating  solutions  containing 
a  ten-fold  excess  of  complexing  cyanide  anion  Is  open  to  question48.  Job44 
found  appreciable  amounts  of  a  cationic  cobalt  complex  (CoCl)"1  in  a  solu- 
tion containing  an  excess  of  hydrochloric  acid,  bu1  an  extrapolation  to 
silver  solutions  is  speculative. 

The  hypothesis  has  been  extended  to  the  plating  of  copper,  zinc,  cad- 
mium, and  mercury11,  and  to  silver  deposition  from  complex  iodid< 
Glazunow46  assumes  that  complex  cations  must  be  present  in  all  complex 
salt  solutions  and  reduction  of  these  cations  gives  rise  to  three  possibilities: 
(1)  new  complexes  arise  which  cannot  exist  in  the  free  state  and  decompose 
quickly  with  deposition  of  metal;  (2)  new  complexes  arise  which  give  rise 
to  insoluble  oxides,  chlorides,  etc.,  on  the  electrode  surface;  or  (3)  new 
stable  complexes  arise  which  contain  the  metal  in  a  lower  valence  state. 
The  first  possibility  is  illustrated  by  the  deposition  of  zinc  from  complex 
cyanides. 

[Zn(CN)4]-  -  [Zn(CN)]+  +  3CN 

[Zn(CN)]+  +  e-->  ZnCN 

2ZnCN->  Zn  +  Zn(CN)a 

Zn(CX),  +  2CN-  -»  [Zn(CN)4]- 

The  second  possibility  was  used480- 46(1  to  explain  the  preparationof  explosive 
antimony  by  electrolytic  reduction  of  solutions  containing  antimony 
chloride  complexes  of  the  type  [SbClJH  and  [SbCl]++: 

[SbCl,]+  +e"->  SbCli 

SbCla  +  2e-  -»  SI)  +  2C1    or  BbCla  -♦  Sb  +  CI, 

If  the  unstable  SbClo  molecule  i>  formed  more  rapidly  than  it  decomposes, 
the  unstable  neutralized  complex  Sb(  '!_■  is  included  in  the  metal  deposit .  and 

}_'.  Bellwig,  Z.  anorg.  Chem.,  25,  157    1900  , 

43.  Erdej  Gruz,  Z.  phyeik.  Chem.,  172,  157    1935  , 

14.  Job,  Ann.chim.,  [11J6,97    L936 

15.  Bchlotter,  Korpiun,  and  Bunneister,  Z    Metallkunde,  26,  L07    I 

16.  Glazunov,  Chem.  bitty,  32,  246    1938  ;  Glazunov,  Starosta,  and  Vbndrasel     / 

/.-.  Chem.,  A185,  393    1939);  i  Uazunov,  Rex .  met.,  43,  J I  l    1946);  Glazunov 
and  Lazarev,  ( 'fu  m.  Liety.,  34,  99    run  ;  ( Uazunov  ;m<l  Bchlol  ter,  First  I 
Electrod.  Conj  .    1937  ;  cf   Cfo  m     16.     31.  7766 


632  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

gives  rise  bo  explosive  antimony.  At  lower  current  density,  S0CI2  molecules 
decompose  as  fast  as  they  are  formed;  this  gives  stable  antimony. 

The  third  possibility  is  illustrated  by  reduction  of  ferricyanide  to  ferro- 
cyanide. 

Copper  has  been  deposited4613  on  thin  glass  fibers  stretched  across  the 
surface  of  a  polished  copper  cathode  in  copper  cyanide  solution.  The  pres- 
ence of  copper  on  the  nonconducting  glass  fiber  was  interpreted  as  evidence 
for  secondary  deposition;  however,  copper  on  these  fibers  might  result 
from  the  metal  lattice  growing  out  over  the  glass  fiber  in  a  primary  reduc- 
tion process. 

The  Kinetics  and  Mechanism  of  Electrodeposition  From  Complex 
Ions 

If  dissociation  takes  place  before  reduction,  any  one  of  at  least  three 
steps  may  be  rate  determining:  (1)  diffusion  of  ions  to  the  electrode  surface, 
(2)  dissociation  of  the  complex  to  give  so-called  simple  ions,  (3)  reduction 
of  the  simple  ion  and  incorporation  of  metal  atoms  into  the  lattice.  A 
number  of  investigators47  have  suggested  slow  dissociation  as  the  rate  de- 
termining step.  In  most  cases  it  is  impossible  to  distinguish  experimentally 
between  slow  dissociation  and  slow  reduction. 

Alternatively,  if  deposition  occurs  by  direct  reduction  of  the  complex 
ion,  the  process  can  be  broken  down  into  two  major  steps:  (1)  transfer  of 
ions  to  the  electrode  surface  and  (2)  reduction  of  the  ion  on  the  electrode 
surface.  Experimentally  these  processes  are  studied  by  polarization  curves. 
If  transfer  of  ions  to  the  electrode  surface  is  the  rate  controlling  factor,  the 
potential  of  the  cathode  will  rise  above  the  reversible  electrode  potential 
for  the  solution  as  a  whole,  and  the  increase  is  termed  concentration  polar- 
ization. If  the  reduction  process  is  slow  wiiile  the  transfer  process  is  rapid, 
the  potential  of  the  cathode  will  again  rise  above  the  equilibrium  electrode 
potential  before  metal  is  deposited.  The  latter  increase  in  potential  is 
termed  chemical  polarization.  Much  experimental  work  on  the  kinetics  of 
the  electrode  processes  involving  complex  ions  has  attempted  to  differ- 
entiate between  concent  rat  ion  and  chemical  polarization. 

The  Transfer  of  Ions  to  the  Electrode  as  the  Rate  Determining 
Process.  Ions  to  be  reduced  reach  the  electrode  surface  by  (1)  diffusion, 
(2)  mechanical  stirring  or  (3)  electrolytic  migration.  It  is  supposed  that 
mechanical  stirring  cannot  move  ions  directly  to  the  electrode  since  a  thin 
unstirred  liquid  layer  is  generally  considered  to  adhere  tenaciously  to  the 
metal  surface.  Ions  must  diffuse  through  this  adhering  film.  The  effects  of 

17  Dole.  Trims.  Electrochem.  Soc.,  82,  1241  (1942);  Ksin,  Acta  Plujsicochimica., 
1  I;  8  s  .  16,  L02  L942);cf.,  Chem.  Aba.,  87,  2273  1  M)43);  LeBlanc  and  Schick, 
/    Elektrochem., 9, 636  (1903);Z.  physik,  Chem., 46, 213  (1903). 


COORDINATIOA   COMPOUNDS  Ih    ELECTRODBPOSITIOh  633 

electrolytic  migration  in  the  negative  field  of  the  cathode  arc  generally  not 
of  great  importance,  and  can  be  made  negligible  by  the  presence  of  an  ex- 

-  of  an  inert  electrolyte.  An  excellent  discussion  of  ion  movement  in 
solution  is  given  by  Kolthoff  and  Lingane48.  Frequently,  diffusion  controls 
the  rate  of  ion  migration  to  the  electrode. 

If  transport  of  ions  to  the  electrode  by  diffusion  is  the  limiting  process, 
the  current  Sowing  can  be  calculated  from  Kick's  law  of  diffusion.  By  de- 
termining the  effect  of  a  change  in  conditions  of  diffusion  on  ihe  current 
Bowing  at  a  given  potential,  concentration  polarization  may  be  identified. 

The  Reduction  oi'  Ions  on  tin*  Electrode  as  the  Slow  Process.  The 
reduction  process  has  been  considered  in  three  somewhat  different  ways. 
First,  it  has  been  assumed  that  the  metal  ions  are  discharged,  then  the 
metal  atoms  find  places  in  the  metal  lattice.  Either  Btep  may  he  rate  de- 
termining. LeBlanc49  thought  that  the  slow  step  was  dehydration  or  de- 
coordination  of  the  metal  ion.  ( )ther  workers50  assume  that  free  metal  atom- 
accumulate  around  the  electrode  until  metal  crystallization  occurs.  An 
effort  has  been  made  to  correlate  the  physical  properties  of  the  metal  plate 
with  the  expected  concentration  of  metal  atoms  in  the  cathode  film.  Here 
crystallization  would  be  rate  determining. 

A  second  point  of  view  suggests  that  an  ion  must  first  find  a  suitable  place 
on  the  lattice  before  reduction  occurs61,  B2.  Two  possible  energy  harrier- 
may  he  pictured,  corresponding  to  desolvation  and  adsorpt  i<>n  of  the  ion  on 
the  electrode  surface,  and  to  transfer  of  an  electron  from  the  electrode  to 
the  adsorbed  ion.  Either  process  may  be  rate  determining.  By  applying  tin1 
theory  of  absolute  reaction  rates,  the  Nernst  equation  for  the  potent  ial  <>!  a 
reversible  electrode  is  obtained.  In  addition,  an  equation  was  developed61  to 
give  the  current  flowing  to  the  electrode  at  any  voltage  V  as  a  fund  ion  of 
the  variables  controlling  both  ion  diffusion  and  ion  reduction  on  the  elec- 
trode surface. 

The  third  hypothesis  pictures  the  adsorption  and  reduction  proa 
occurring  in  a  Bingle  Btep84.  No  attempt  is  made  to  differentiate  separate 

tv  Kolthoff  and  Lingane,  "Polarography,"  Chapt.   II,  New  York,   [nterscience 

Publishers,  [nc,  L941. 
r».  LeBlanc,  Trans.  Faraday  Sue.  9,  251  l  191  I  . 

50.  An-ii  rind  Boerlage,  Rec.  trav.  chim.,  39,  7_'i>    1920  ;  Brandes,  '/.    physik.  Chun., 

142,  !•:    1929  ;  Fink.  ./.  Phys.  Chem.,  46,  7<i    1942  ;  Hughes,  Dept.  oj  Scientific 
and  Ind.  Research  Bull.,  No.  6,    1922) ;  Hunt,  Tram    I  hem.  Soc.,  65,  113 

1934   ,  Hum.  ./.  Phys.  Chem.,  36,  1006,  2259    L! 

51.  I  ru2  and  Volmer,  /  physik  Chem.,  A157,  165  (1931). 

52.  Glasstone,  Laidler,  and  Eyring,  "Theory  of  Rate  Processes, "  pp   575  81,  New 

5Tork,  McGraw  Hill  Book  Co..  1941. 
( rlasstone,  Laidler,  and   I .;  ring,  "Theoi 
York.  McGraw  Hill  Book  Co.,  1941. 
■~>\.  Blum  and  Rawdon.  Trai  14,  :;'i7    191 


634 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


(A)  (B) 

Fig.  19.1.  Potential  energy  of  a  metal  ion  at  the  surface  of  the  metallic  lattice  (A), 
and  in  the  complexed  state  (B).  Distance  of  separation  great. 

steps  in  the  process.  Because  of  inherent  simplifications  in  this  mechanism 
it  may  readily  be  applied  to  the  reduction  of  complex  ions55-59. 

A  metal  may  be  pictured  as  metal  ions  surrounded  by  mobile,  loosely 
held  electrons58- 59.  The  variation  of  the  energy  of  a  metal  ion  near  the  sur- 
face of  the  metal  is  represented  by  the  potential  energy  diagram  in  Fig. 
L9.1A.  The  energy  of  an  isolated  ion  in  vacuo  is  represented  by  the  hori- 
zontal line  .1 ;  the  ion  loses  energy  Um  when  it  is  bound  to  the  metal  surface 
and  comes  to  rest  at  an  equilibrium  distance  "d"  from  the  bulk  of  the  metal. 
The  first  horizontal  line,  BB,  represents  the  ground  energy  level  of  the  ion 
and  the  other  lines  represent  higher  energy  levels.  As  the  temperature  of 
the  metal  increases  there  is  greater  probability  that  higher  energy  levels 
will  be  occupied. 

Similarly,  Fig.  19. IB  is  a  potential  energy  diagram  for  a  metal  ion  in  the 
vicinity  of  a  water  molecule,  group  of  water  molecules,  or  other  coordi- 
nating groups.  Ua  is  the  energy  of  hydration  or  energy  of  coordination  and 
solvation  for  the  ion.  If  a  solvated  ion  from  the  solution  approaches  the 
metal  surface,  the  two  curves  may  overlap  and  combine  to  give  a  curve  of 
the  type  shown  in  Fig.  19.2  (AorB).  Now  we  have  two  equilibrium  positions 
for  the  ion,  separated  by  an  energy  barrier  C.  The  height  of  this  barrier  is 
determined  by  how  close  the  ion  may  approach  to  the  metal  surface.  In 
-Dine  cases  the  potential  hill  may  completely  vanish  at  the  moment  of 
impact  and  reappear  immediately  as  the  ion  rebounds.  At  the  present  time 
we  have  little  information  concerning  such  energy  barriers. 

55.  Butler,  Trans.  Fannin  a  Soc,  19,  729  (1924). 

56.  Gurney,  Proc.  Roy.  S<><-.  London,  A136,  378  (1032). 

57.  Fowler,  Proc.  Roy.  Soc.  London,  A136,  391  (1932). 

58.  Butler,  "Electrocapillarity,"  pp.  30  34,  London,  Methuen  ;uul  Co.  Ltd.,  1940. 
Gurney,  "Ions  in  Solution,"  Chapt.  IV,  London.  Cambridge  University  Press, 


COORDINATION  COMPOUNDS  l\    ELECT RODEPOSITIOh 


635 


ETAL 


ENERGY    OF     ISOLATED     ION 
IN       VACUUM 


w 


METAL  - 


ENERGY    OF    ISOLATED 

ION 

1 

f 

.    d     „ 

\- 

U-r. 

• 

w 

Fig.  19.2.  New  potential  energ}r  relationship  associated  with  approach  of  solvated 
ion  to  electrode. 


If  the  potential  valleys  are  of  equal  depth,  there  will  be  no  tendency  for 
transfer  of  ions  from  one  side  to  the  other,  but  if  energy  Levels  in  the  metal 
are  available  below  the  levels  of  the  ion  in  solution  (Fig.  19.2A),  spon- 
taneous transfer  of  ions  will  take  place  from  the  solution  to  the  metal 
surface,  providing  the  ions  can  get  over  the  energy  barrier  in  the  middle. 
For  many  cases  this  hump  may  be  negligible,  as  for  readily  reversible  elec- 
trodes, but  in  other  cases  the  rate  of  the  transfer  may  be  limited  by  this 
barrier.  The  height  of  the  barrier  determines  an  activation  energj  for  the 
proees.-.  If  the  number  of  positive  ions  being  deposited  initially  exceed-  the 
number  of  ions  Leaving  the  metal  surface,  the  metal  will  acquire  a  positive 
charge,  which  retards  and  finally  stops  further  deposition  of  positive  ions 

On  the  metal  .-uilace.  In  elect  rodepo-it  ion  an  extraneous  negative  potential 

is  imposed  on  the  electrode  to  prevent  this  accumulation.  The  imposed 
E.M.F.  maintain-  the  energy  levels  for  positive  ions  in  the  metal  below 
those  in  the  solution. 

The  reverse  situation,  illustrated  in  Fig.  L9.2B,  comes  about  when  ions  on 
the  metal  surface  have  higher  potential  energy  than  solvated  or  coordi 
ions.  Positive  ions  arc  then  transferred  spontaneously  from  the  metal  to 


636  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

the  solution,  a  negative  charge  builds  up  on  the  electrode  and  a  positive 
charge  in  the  solution  until  the  energy  levels  of  ions  on  the  electrode  and  in 
the  solution  are  equal.  If  metal  is  to  be  deposited  from  solution,  a  larger 
external  negative  potential  must  be  imposed  on  the  cathode  until  energy 
levels  in  the  metal  are  below  those  of  the  ions  in  the  solution.  The  first 
situation,  Fig.  19.2A,  might  be  represented  by  a  noble  metal  such  as  silver 
while  the  second  situation,  Fig.  19.2B,  would  represent  a  less  noble  metal 
such  as  zinc.  In  general,  the  effect  of  complex  formation  is  to  lower  the  po- 
tential energy  of  ions  in  solution  relative  to  the  potential  energy  of  "simple" 
hydrated  ions.  As  a  result,  the  dips  on  the  right  in  Fig.  19.2A  and  19.2B  will 
usually  be  deeper  for  the  complex  ions  than  for  the  simple  hydrated  ions. 
This  means,  for  instance,  that  the  potential  for  the  reaction: 

[Ag(CN)2J-  +  e~  ->  Ag  +  2CN- 

will  be  more  negative  (reaction  has  less  tendency  to  go)  than  the  potential 
for  the  corresponding  reaction  involving  the  simple  hydrated  ion  of  silver. 

[Ag(H20)2]+  +  e-  ->  Ag  +  2H20 

This  treatment  does  not  require  dissociation  of  the  complex  into  simple 
ions,  but  rather  assumes  that  the  complex  is  in  direct  equilibrium  with  the 
electrode  surface.  The  possibility  that  the  reduction  process  is  sometimes 
slow  is  suggested  by  the  energy  barrier  in  Fig.  19.2. 

Rate  Determining  Steps  in  the  Reduction  of  a  Number  of  Com- 
plex Ions.  Electrode  polarization  has  been  used  as  a  criterion  for  identifying 
the  slow  process  in  electrode  reactions.  Conclusions  are  generally  based  on 
the  shape  of  experimentally  determined  current  voltage  curves  or  upon  the 
variation  of  such  curves  with  changes  in  experimental  conditions.  The  study 
of  such  curves  is  subject  to  a  number  of  experimental  errors2*  >  50a- 60.  Further, 
detailed  interpretation  of  the  data  varies,  depending  upon  the  assumptions 
used.  It  is  possible,  however,  in  some  cases  to  distinguish  between  diffusion 
and  retarded  reduction  as  the  rate  controlling  process. 

In  the  deposition  of  silver  from  solutions  of  the  complex  ions  [Ag(NH3)2]+ 
and  [Ag(CN)2]~~  5a'  61,  the  maximum  current  density  which  gives  100  per 
cent  cathode  efficiency  for  metal  deposition  is  determined  by  the  rate  at 
which  complex  ions  can  diffuse  to  the  surface  of  the  cathode.  With  am- 
nion in,  tliiocyanate,  and  iodide  complexes  of  silver,  the  rate  of  diffusion  of 

60.  Butler,   "Electrocapillarity,"  p.  167,  London,  Methuen  and  Co.  Ltd.,  1940; 

Glasstone,  J.  Chem.  Soc,  127,  1824  (1925);  Kohlschutter  and  Torricelli,  Z. 
Elektrochem.,  38,  213  (1932);  Smartsev,  Compt.  Rend.  Acad.  Sci.,  U.S.S.R.,  2, 
178  (1935);  Khim.  Referat  Zhur.,  4,  no.  5,  119  (1941);  Acta  Physicochim.  U.R. 
8.S.,  16,  206  (1942);  Mathers  and  Johnson,  Trans.  Electrochem.  Soc,  81,  267 
(1942). 

61 .  Glasstone,  J.  Chem.  Soc,  1932,  2849. 


\RDINATIOh   COMPOl  \  D8  I  \   ELECTRODBPOSITIOA  637 

ions  to  the  cathode  determines  cathode  potential  while  the  diffusion  <>!'  ions 
from  the  anode  determines  anode  potential  (concentration  polarization  '  , 
Erdey-Gruz  ami  Volmer48, M  concluded  from  current-voltage  curves  that 
under  conditions  such  that  concentration  polarization  is  minimized, 
metal  discharge  is  the  rate-controlling  step  in  deposition  from  ammoniacal 
solutions  o\  silver  bromide  or  chloride.  For  an  ammoniacal  solution  of  silver 
oxide,  as  well  as  for  solutions  of  [Agk]    [AgBrJ  ,  [Ag(CN)s]  ,  and  [AgCli]  , 

the   rate   appears   to   be  determined    by  the  orientation  of  the  ions  in  the 

lattice  before  reduction.   Equations  were  derived  for  the  curves  under 
different  circumstances  of  lattice  formation. 

These  methods  (see  also  Ref.  63)  have  been  applied  to  other  systems*, 
bul  are  subject  to  errors  in  measuring  the  active  electrode  surface  and  ex- 

ave   concentration    polarization   around    small    active   areas  of   crystal 
growth*5. 

In  the  deposition  of  copper  from  solutions  containing  pyrophosphate, 
oxalic  acid,  or  thiocyanate,  concentration  polarization  was  observed64,  M. 
With  ammonia,  ammonium  oxalate,  and  thiosulfate  as  complexing  agents 
the  slow  process  was  attributed  to  ion  discharge.  LeBlanc  and  Schick'7, 
believe  that  the  rate  of  copper  deposition  from  potassium  cyanide  solution 
is  limited  by  a  slow  dissociation  of  the  [Cu(CX);<]=  complex.  This  idea  has 
been  used67  to  explain  deposition  of  copper-gold  alloys  from  cyanide  solu- 
tion. The  rate  of  deposition  of  gold,  but  not  that  of  copper,  was  that  calcu- 
lated from  diffusion  theory.  It  was  concluded  that  the  rate  of  discharge  of 
gold  cyanide  is  probably  determined  by  the  rate  of  diffusion  of  the  complex 
ions  to  the  electrode,  but  the  rate  of  discharge  of  copper  cyanide  ion-  is 
probably  determined  both  by  diffusion  and  by  rate  of  dissociation  (or  rate 
of  reduction)  of  the  complex  at  the  electrode  surface.  However,  ( rlasstoi 
found  that  the  potential  of  a  copper  electrode  in  a  copper  cyanide  solution 
Is  dependent  upon  the  concentration  of  cyanide.  Relatively  small  increases 
in  cyanide  content  bring  about  considerable  increase  in  potential  required 
for  copper  deposition.  If  the  cyanide  concentration  is  large  or  becomes 
large  due  to  accumulation  of  cyanide  around  the  cathode,  hydrogen  may  be 

evolved  along  with  copper.  He  concluded  that   polarization  of  the  cathode 

is  due  to  depletion  of  complex  copper  cyanide  ions  and  accumulation  of 

simple  cyanide  ions.  This  suggests  diffusion  as  the  rate  controlling  pro© 
as  is  indicated  by  current-voltage  data' 

62.  Levin,  ./.  Phys.  Chem.,  U.S.S.R.,  17,  247  (1943  j  19,  365    1946);  cf.  Chi 

38,  1960  (1944);40,  1738  (1946). 

63.  Butler,  "Electrocapillarity,"  p.  169,  London,  Methuen  and  Co.  Ltd.,  L940. 

64.  Levin,  /.  Phys.  Cfo  I.S.R.,  16,  948  (1941);  cf.  Chen     16      86,6087    L942 

65.  Vahramian,  Acta  Physicochimica,  19.  L48,  159    1944  . 

66.  Levin  and  Btonikova,  •/.  Gen.  I  3  B.,  13,  667    Lfl 

I.mu  and  Alfimova,  •/.  /  9  R  .  8,  L37    16  .  31, 

1706  (1947). 


038  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

The  rate  of  deposition  of  zinc  and  cadmium  from  solutions  of  metal  am- 
mines  or  metal  cyanides  is  controlled  by  the  diffusion  of  ions  to  the  elec- 
trode41- 81« 67- G9, 7,).  Both  diffusion  and  retarded  discharge  play  a  part  in  the 
reduction  of  zinc  from  zincate  solutions70. 

Only  concentration  polarization  has  been  found61- 71  in  the  deposition  of 
mercury  from  [Hg(CN)4]=,  though  both  diffusion  and  slow  reduction  are 
important  in  the  deposition  of  mercury  from  Hg(CN)2  (or  perhaps 
[Hg(CN)2(H20)2]). 

In  the  deposition  of  bismuth  from  hydrochloric  or  nitric  acid  solutions, 
concentration  polarization  predominates72,  while  chemical  polarization  due 
to  slow  discharge  is  important  in  the  deposition  of  bismuth  from  sulfuric 
acid  solutions.  Similarly,  deposition  of  antimony73  from  hydrochloric  acid 
solution  is  limited  by  ion  diffusion,  while  ion  discharge  is  important  in  the 
deposition  from  sulfuric  acid  solution. 

In  these  experiments,  concentration  of  the  solution,  current  density,  and 
temperature  and  other  factors,  play  such  large  roles  in  determining  the 
identity  of  the  rate  determining  step  that  a  distinct  and  unambiguous 
answer  is  obtainable  only  for  certain  ions  under  specific  conditions. 

Extensive  investigations  on  electrode  kinetics  are  summarized  by  Dela- 
hay73a.  The  most  notable  result  is  the  determination  of  reaction  rate  con- 
stants for  metal  deposition.  In  some  instances,  it  appears  that  the  complex 
involved  in  the  deposition  mechanism  has  a  lower  coordination  number 
than  that  of  the  predominant  species  in  the  solution73B. 

Electronic  Configuration  and  Deposition  Mechanism 

The  electronic  configurations  of  the  ions  to  be  deposited  exercise  a  con- 
trolling influence40.  For  example,  the  electronic  structure  of  the  aquated 
iron (II)  ion  is  represented: 

3d  4S  4p  4d 

[Fe|H20)J++     ,sW2p'3sV      [TUM    0    H    f^TFTH 

68.  Esin  and  Mantansev,  J.  chim.  phys.,  33,  631  (1936). 

69.  Levin,  ./.  Gen.  Chem.,  U.S.S.R.,  14,  795  (1944);  cf.,  Chem.  Abs.,  39,  3736  (1945). 
7D.  Esin  and  Beklemysheva,  •/.  Phys.  ('hem.,  U.S.S.R.,  10,  145  (1937);  cf.,  Chem. 

Aba.,  32,  430  (1938);  J.  Gen.  Chem.,  U.S.S.R.,  6,  1602  (1936). 

71 .  Esin  and  Alfimova,  •/.  Gen.  Chem.,  U.S.S.R.,  7,  2030  (1937);  Esin  and  Malarzev, 

Z.  physik.  Chen..  A174,  384  (1935). 

72.  Esin,  Lashkarev,  Levitina,  and  Rusanova,  ./.  Applied  Chem.,  U.S.S.R.,  13,  56 

(1940);  17,  111  (1944). 

73.  Esin,/.  Applied  Chem.,  U.  S.S.R.,  17,  111  (1944);  cf.,  Chem.  Abs.,  89, 1359  (1945). 

Pelahay,  W\\    Instrumental  Methods  in  Electrochemistry,  New  York,  Inter- 
3ci(  nee  Publishers,  Inc.,  1954. 
73b.  Gerischer,  Z.  Electrochem.,  57,  604  (1953). 


COORDINATION   COMPOUNDS  IN  ELECTRODEPOSITIOh  639 

in  which  the  Crosses  represent  electrons  donated  by  water  molecules  to  the 

n/>V-  hybridized  orbitals.  The4  presence  of  four  unpaired  electrons  is  indi- 
cated by  magnetic  data.  In  the  hexacyano  ion,  however,  the  Bingle  electrons 
become  paired,  and  the  hybridization  is  cPsp*t  involving  3d  levels  as  well  as 

\s  and   \p: 


[r»M.]  4 


Is  2s  2p  3s  3p 


3C 

4S 

4p 

• 

• 

• 

X 

X 

K 

■ 

■ 

• 

• 

• 

X 

«. 

y 

• 

» 

' 

in  which  the  crosses  represent  electrons  from  the  cyano  groups.  The  ion  is 
diamagnetic,  indicating  that  no  unpaired  electron.-  are  present. 

Iron  is  readily  deposited  from  the  aquated  ion,  but  not  from  the  cyano 
ion,  (except  as  an  alloy  under  special  conditions)101',  [n  aqueous  solutions, 
deposition  generally  does  not  occur  where  hybridization  involves  an  inner 
orbital.  Sucha  configuration  may  represent  unusual  stability,  and  apparent  ly 
less  energy  is  required  to  reduce  hydrogen  ion  than  to  break  up  hybridiza- 
tion. Consequently,  hydrogen  rather  than  metal  is  discharged. 

Inner  orbital  complexes  react  slowly  or  not  at  all  in  substitution  reac- 
tions28 except  when  half  filled  orbitals  are  present ;  a  similar  situation  seems 
to  hold  for  electron  transfer  reactions40.  These  observations  suggest  that 
.at ion  of  a  coordinated  group  from  an  inner  orbital  complex.  AX„  — > 
AX  ,  +  X,  occurs  only  with  difficulty.  Since  the  configuration  is  also 
unfavorable  for  elect rodeposition,  it  is  inferred  that  difficulty  of  dissocia- 
tion is  reflected  in  the  deposition  reaction,  and  that  an  intermediate  of  the 
type  AX     !  is  important781*. 

Reduction  of  ferricyanide  ion  to  ferrocyanide  is  reversible.  Evidently 
little  activation  energy  is  needed  to  transfer  an  electron  to  the  complex. 
Reduction  to  iron,  however,  does  not  generally  occur.  Since  there  appear- 
to  be  no  difficulty  in  transferring  a  single  electron  to  the  iron!  Ill'  complex, 
it  has  been  suggested40  that  the  obstacle  lies  in  the  stripping  of  the  coordi- 
nated group-.  1  association  would  be  the  first  step  in  this  process.  The  diffi- 
culty of  dissociating  an  inner  orbital  complex  would  be  shown  by  very 
large  potential  energy  humps  in  Figs.  l(.).l  and  19.2. 

With  the  aquated  iron(Il  I  ion.  on  the  other  hand,  substitution  and  elec- 
tron transfer  studies  indicate  that  dissociation  occurs.  Likewise  the  metal 
may  be  deposited.  The  necessary  electrons  are  relatively  easy  to  add,  and 

loSS  of  water  gTOUOS  takes  place  readily. 

"Flash"  deposits  are  sometimes  obtained  from  inner  orbital  complexes. 
In  s<»me  instance-,  the  deposits  appear  to  be  the  result  of  codeposition  of 
impurities,  and  in  other-,  the  nature  of  the  basis  metal  may  permit  de- 
position until  it  is  completely  coated.  In  either  case,  deposition  so 

To  account    for  the  attachment    of  the  metal   loll   to  1  he  cathode  surface, 

n  ha-  been  suggested40  that  the  dissociated  ion,  A \     :  .  replaces  the  lost 


640  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

coordinated  group  with  a  molecule  on  the  aquated  cathode  surface.  Sub- 
sequently this  water  bridge  is  eliminated,  perhaps  because  of  the  elec- 
t  lost  at  ic  attraction  of  the  cathode  for  the  positive  metal  ion,  and  a  metallic 
bond  is  established.  As  other  metal  atoms  are  deposited  in  neighboring 
positions,  the  remaining  coordinate  bonds  are  replaced  by  metallic  bonds. 
Transfer  of  electrons  to  depositing  ions  is  needed  only  to  maintain  the 
average  electrical  potential  of  the  cathode.  Details,  in  terms  of  Pauling's 
theory  of  the  metallic  state,  are  given  in  reference  40,  and  provide  an  ex- 
planation for  the  nature  of  inclusions  in  deposits.  There  is  evidence  that 
non  metal  inclusions  consist  largely  of  residual  coordinated  groups. 

Reversibility  in  the  deposition  of  metal  ions  is  found  only  when  no  re- 
arrangement of  the  electronic  configuration  of  the  ion  is  necessary  to  attain 
the  configuration  of  the  metallic  atoms.  Among  transition  elements,  re- 
arrangement of  electrons  is  associated  with  deposition;  this  requires  ex- 
penditure of  energy  and  is  responsible  for  the  observed  irreversibility.  In 
Figs.  19.1  and  19.2,  this  would  correspond  to  potential  humps  higher  than 
those  for  such  metals  as  zinc  and  lead,  but  not  quite  as  high  as  that  for 
hydrogen,  which  is  commonly  codeposited  with  these  metals. 

Another  cause  of  irreversibility  is  the  tendency  of  such  metals  as  tin, 
bismuth,  and  gallium  to  form  multinuclear  aquo  or  hydroxo  complexes 
which  are  slow  to  dissociate.  The  effect  of  chloride  ions  in  reducing  the  ir- 
reversibility is  presumably  to  be  attributed  to  formation  of  mononuclear 
chloro  complexes. 

Coordination  Compounds  as  Important  Factors  in 
Electrodeposition 

It  is  well  known  that  metal  deposits  obtained  from  solutions  of  complex 
ions  frequently  have  better  physical  properties  than  those  from  simple  salt 
solutions.  Further,  small  quantities  of  addition  agents  produce  truly  re- 
markable changes  in  the  physical  properties  of  the  deposited  metal.  The 
causes  of  these  phenomena  are  not  understood,  though  both  are  of  sub- 
stantial technological  importance. 

Crystal  Structures  of  Electrodeposits 

Metal  deposits  obtained  from  solutions  of  complex  salts  are  made  up  of 
submicroscopic  crystals8,  16a,  but  it  is  not  true  that  the  crystals  must  be 
smaller  than  the  wave  length  of  light  to  produce  bright  deposits.  Bright 
and  dull  deposits  of  chromium  contain  crystals  of  comparable  size74,  but  in 
1  >right  deposits,  crystals  show  regular  orientation.  Blum4' 75  emphasized  the 
importance  of  crystal  oriental  ion  and  suggested  that  copper  deposited  from 

71.  Wood,  Trans.  Faraday  Soc.,  31,  1248  (1935). 

75    Blum.  Beckman.  and  Meyer.  Trans.  Electrochem.  Soc,  80,  249,  288,  254  (1941). 


ORDINATION    COMPOl  \l>s/\    ELECT  RODEPOSITIOh  641 

cyanide  complexes  is  dull,  not  because  of  crystal  size,  bul  because  of  random 
orientation.  Recenl  investigations79,  however,  indicate  that  neither  crystal 
size  nor  orientation  is  directly  related  to  brightness.  Ii  can  only  be  asserted 
that  the  surface  must  be  smooth  enough  for  Bpecular  reflection,  regardless 

of  the  structure  beneath. 

It  has  been  suggested11  '■  that  the  increased  deposition  potential  on  the 
cathode  a>  a  result  of  complex  formation  is  responsible  for  small-grained, 
and  sometimes  oriented,  deposits;  however,  this  does  not  explain  the  actual 
function  of  the  complex  ion,  hut  rather  emphasizes  a  nonrigorous  corre- 
lation14 between  electrode  potential  and  character  of  deposited  metal. 

Kohlschutter77  suggested  that  insoluble  cyanides  deposited  on  the  elec- 
trode surface  prevent  the  growth  of  large  crystals,  and  attention  has  been 
directed41,  fi:>  toward  the  possible  adsorption  of  complexing  ions  on  the  elec- 
trode. Microscopic  studies43,  50a- 60c' 60d- 65- 7s  show  that  from  perchlorate  or 
nitrate  solutions  silver  is  not  deposited  uniformly  over  the  face  «»!  ;i  Bilver 
crystal  hut  only  on  a  number  of  active  centers  on  the  crystal  face.  The  num- 
ber of  such  active  centers  on  the  crystal  surface  is  increased  by  a  decreas* 
in  the  concentration  of  the  silver  salt  in  the  solution12.  If  the  current  is 
interrupted  for  a  short  time,  the  old  crystal  surfaces  will  not  develop  again, 
but  when  electrolysis  is  resumed,  new  localized  sites  become  active  and 
crystallites  grow  from  the  new  sites12* 80d.  In  silver  nitrate  solution  from 
which  all  organic  matter  had  been  removed,  localized  passivation  and  acti- 
vation of  the  silver  crystal  face  did  not  develop65.  Addition  of  0.2  per  cent 
dextrin  solution  brought  about  a  strong  passivation,  suggesting  that 
passivation  is  due  to  adsorption  of  surface-active  organic  impuritii 

In  contrast  to  the  behavior  for  simple  salts,  an  entire  face  of  the  crystal 
may  develop  in  solutions  of  complexes  such  as  cyanide.  In  general,  the 
materials  present  in  the  solution  determine  which  crystal  face  develops*  43. 
The  absence  of  passivation  in  the  electrodeposition  of  silver  from  cyanide 
solutions  is  accounted  for  by  the  high  adsorption  of  the  cyanide-silver  com- 
plex, which  prevents  adsorption  of  surface-active  impurities 

A  Study*1  of  the  deposition  of  cobalt  and  nickel  from  a  wide  variety  of 
complex  i  -  sted  that  the  nature  of  the  coordinating  group  as  well  as 

*  It  is  Interesting  in  this  connection  that  the  crystalline  form  «»t"  an  electrode 
posited  metal  is  dependent  upon  the  bath  from  which  it  is  obtained.  For  instance, 
body-centered  cubic  chromium  is  tonne. l  in  the  essential  absence  of  trivalent  chro 
mium.  whereas  deposition  of  the  hexagonal  form  depends  upon  the  presence  of  tri 
valenl  chromium751 . 

Clark  and  Simonsen,  •/.  EUctroi  >  ■       S        98,  1  in    195]  » ;  Denise  and  Leidfa 
ibid.,  100, 490    19S 

77.  Kohlschutter,  /   Eleki  ■    •  ■        19.  181     1911 

78.  Vahramian.  Compt.  i:<  nd.  Am, I.  &    .  1  .R  8  8    22.  _  I  im. 

U.R.SJS.,  7.  ftg 


642  (IIEMIST/IY  OF  THE  COORDINATION  COMPOUNDS 

the  thermodynamic  stability  of  the  complex  ions  is  important  in  determin- 
ing whether  good  plates  will  be  formed.  In  general,  large  coordinating 
groups  or  those  containing  aromatic  ring  systems  gave  poor  plates.  It  was 
also  observed  that  complexes  which  are  reduced  either  with  great  difficulty 
or  too  easily  gave  poor  plates.  Complexes  in  an  intermediate  range  of  sta- 
bility (i.  e.,  [Co(en)3]+++)  gave  good  plates. 

The  Effect  of  Brighteners.  Mathers79 f  suggested  that  brighteners  and 
addition  agents  may  owe  their  action  to  ability  to  form  complexes  with  the 
metal  ions  in  solution.  Mathers  used  the  terms  "complex  ion"  and  "com- 
plex compound"  very  broadly  and  implied  that  all  ions  present  in  the  ionic 
atmosphere  are  part  of  the  complex. 

However,  it  does  not  seem  justifiable  to  postulate  that  all  addition  agents 
form  Werner  type  coordination  compounds  with  metal  ions  in  solution.  A 
survey  of  over  one  hundred  organic  addition  agents  used  in  the  plating  of 
nickel  failed  to  reveal  any  relation  between  structure  of  the  compounds  and 
efficacy  as  brighteners  or  polarizers81.  In  the  deposition  of  silver  and  copper, 
on  the  other  hand,  various  substances  such  as  glycine,  tartaric  acid,  citric 
acid,  and  metaphosphoric  acid  can  improve  the  quality  of  the  deposit  even 
when  the  addition  agent  is  present  in  very  small  concentrations  (.013/  in 
1  M  A<i\()3)82.  The  addition  agents  were  found  in  the  deposits  in  small 
amounts,  and  it  was  established  by  transference  studies  that  each  of  the 
agents  was  able  to  form  complex  cations  with  silver  or  copper  ions.  From 
this,  a  close  correlation  between  the  efficacy  of  an  addition  agent  and  its 
ability  to  form  complex  compounds  was  suggested.  However,  no  single 
simple  explanation  will  correlate  all  of  the  observed  facts  with  the  struc- 
tures of  the  wide  variety  of  addition  agents  now  in  use. 

An  addition  agent  is  usually  a  substance  added  in  relatively  small 
amounts  to  modify  physical  properties  of  the  deposit.  Addition  agents  are 
often  used  to  produce  bright  deposits,  to  reduce  or  "level"  surface  ir- 
regularities on  the  cathode,  or  to  alter  stresses  in  the  deposits. 

Addition  agents  may  be  grouped  in  three  classes:  (a)  Grain  refining 
agents,  such  as  gelatin  in  copper  sulfate  and  many  other  baths,  reduce  the 

t  Mutscheller80  suggested  earlier  that   gelatin  forms  complexes  with  the  anions 
in  solutions  of  CuS()4  and  AgNOs  .  bul  his  definition  of  complex  was  much  broader 
than  that  used  for  the  metal  complexes  now  under  consideration. 
79.  Mathers,  Proc.  .\»i.  Electro  platers  Soc,  June,  134  (1939);  Mathers  and  Kuebler, 

Trans.  Am.  Electrochem.  Soc,  29,  117  (1916);  36,  234  (1919);  38,  133  (1920). 
BO.  Mutscheller,  Met.  and  Chem.  Eng.,  13,  353  (1915). 

81.  Raub  and  Wittum,  Metal  Tnd.}    V.  r\),  38,  206,  315,  429  (1940). 

82.  Fuseya  and  Maurata,  Trans.  Am.  Electrochem.  Soc,  50,  235  (1926);  Fuseya  and 

Nagano,  Trans.  A m .  Elect rod,,  m .  So,-..  52,  249  (1927);  Fuseya,  Murata,  and 
Yunuito.  Tech.  Il< "ports  Tohobu  Imp.  (nir.,9,  do.  1,33  (1929);  cf.,  Chem.  Abs.t 
24,  3446    1930 


ORDINATION  COMPOUNDS  l\   ELECT RODEPOSITIOh  643 

grain  size  of  the  deposit,  and  often  diminish  the  tendency  of  the  depoail 
to  fonn  "trees"  and  nodules;  (b)  active  agents,  including  brighteners  such 
as  zinc,  cadmium,  sulfonated  aryl  aldehydes,  safranines,  etc.,  in  nickel 
baths81,  which  modify  the  surface  of  the  deposit,  and  usually  the  structure 
as  well,  and  often  produce  the  desired  effects  only  over  a  narrow  range  of 
current  density,  temperature,  pH,  and  other  conditions;  and  (c)  carrier 
agents,  such  as  naphthalene  disulfonic  acids  or  p-toluenesulfonamide  in 
nickel  baths8*,  which  greatly  extend  the  effective  operating  range  of  the 
active  brightener,  imparl  greater  tolerance  towards  impurities,  and  in 
some  instance-  enhance  brightness. 

Bright  deposits  ordinarily  have  a  handed  structure,  the  cause  of  which 
is  unknown.  They  are  almost  invariably  more  brittle  than  typical  deposits 
made  in  the  absence  of  the  brightener.  This  is  usually  attributed  to  the 
inclusion  of  the  brightener,  or  its  decomposition  product,  in  the  deposit, 
resulting  in  a  strained  or  distorted  metal  lattice. 

Brightening  is  only  one  result  of  the  action  of  addition  agents.  Far  more 
frequently,  addition  agents  cause  the  formation  of  spongy  deposits;  this, 
of  course,  is  not  desirable  for  electroplating.  Other  results  are  "wrinkled" 
deposits,  discolorations,  and  roughness  resembling  that  of  sand  paper. 
Studies  have  usually  been  directed  toward  brightness;  a  thorough  study  of 
addition  agents  seems  not  to  have  been  made. 

Grain  refining  agents  are  generally  colloidal.  Most  active  agents  are  elec- 
tron  donors  and  a  tendency  toward  coordination  is  to  be  expected.  Appar- 
ently at  least  two  donor  pairs  are  required84.  Where  the  coordination  is  very 
strong,  as  between  glycine  and  many  metals,  spongy  deposits  are  produced. 
This  is  attributed  to  failure  to  convert  coordinate  linkages  to  metallic 
bonds  on  the  cathode,  so  that  the  agent  is  included  in  the  deposit,  making 
it  impossible  to  build  up  a  normal  metallic  lattice.  It  is  suggested  that  an 
effective  brightener  must  have  sufficiently  strong  coordinating  tendency 
to  modify  the  cathode  surface  by  preventing  formation  of  protruding 
crystal  edges,  and  yet  not  so  strong  that  it  cannol  readily  be  decoordinated 
to  form  metallic  bonds.  Probably  a  few  residual  coordinated  groups  remain 
in  the  deposit — enough  to  produce  the  characteristic  banded  structure 
well  as  the  desired  smooth,  bright  surface. 

A  possible  explanation  of  smoothing  action  lies  in  the  tendency  .,t  deco- 
ordinated groups  t<>  remain  a1  the  cathode  surface  and  form  new  linkages 
with  metal  ions  as  they  diffuse  toward  the  cathode.  By  this  action,  atoms 
may  be  ••\\'i\"  into  the  proper  level,  and  the  build-up  of  crystals  may  be 

prevented.  This  is  closely  related  t<»  adsorption  pro  iggested  by 

other-'1.  The  function  of  the  carrier  type  of  addition  agent    is  not   under- 

83.  Pinner,  Boderberg,  and  Bakei  -•<     80.  699    I'M! 

34    Rdth  and  Leidheiaer,  J   Elet  I  100,  190    I 

85.  Henrick>  Electrochem.  Sor.,  82,  237    1942). 


(ill  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

stood;  in  some  instances  the  carrier  may  coordinate  with  trace  impurities 
and  prevent  them  from  influencing  the  deposit. 

It  is  usually  relatively  difficult  to  find  brighteners  for  metals  which  are 
deposited  reversibly  or  very  nearly  so,  such  as  tin  and  lead.  On  the  other 
hand,  many  brighteners  are  known  for  metals  which  are  deposited  irre- 
versibly, such  as  nickel.  In  fact,  irreversibility  is  generally  accompanied  by  a 
tendency  to  smooth,  fine-grained,  semi-bright  deposits  even  in  the  absence 
of  specific  addition  agents.  It  is  suggested  that  in  these  instances,  the  co- 
ordinating  tendency  is  so  strong  that  even  water  functions  to  some  extent 
as  an  addition  agent. 

Complexes  and  Throwing  Power 

In  the  practice  of  electroplating,  an  important  consideration  is  ability  to 
deposit  coatings  of  relatively  uniform  thickness  on  articles  of  irregular 
shape,  even  though  the  current  distribution  is  far  from  uniform,  as  on  pro- 
truding edges  or  in  recesses.  This  ability,  known  as  "throwing  power," 
represents  the  net  result  of  several  characteristics  of  the  bath  and  also  of 
the  geometry  of  the  plating  cell.  Polarization,  conductivity,  and  variation 
of  current  efficiency  with  current  density  are  important.  "Throwing  power 
is  not  a  single  measurable  property  of  a  solution"75;  a  definitive  discussion 
has  not  been  given  and  is  perhaps  impossible. 

In  a  general  way,  throwing  power  seems  to  parallel  the  stability  of  the 
complex  ions  in  the  baths.  Thus,  in  industrial  practice,  silver  and  copper 
cyanide  baths  have  the  highest  throwing  powers,  the  cadmium  bath  is 
somewhat  inferior,  and  the  cyanide  zinc  bath  is  still  poorer.  This  is  exactty 
parallel  to  the  stability  constants  of  the  cyano  complexes.  All  of  these  baths, 
as  well  as  the  the  stannate  bath,  have  much  better  throwing  power  than 
the  corresponding  sulfate  baths  or  the  silver  nitrate  bath. 

Furthermore,  the  throwing  power  of  cyanide  baths  may  be  improved  by 
increasing  the  concentration  of  cyanide,  although  at  the  expense  of  cathode 
efficiency.  The  improvement  of  throwing  power  by  complexing  has  been 
considered  to  be  the  result  of  diminution  of  "free"  metal  ions  in  the  bath. 
However,  since  deposition  appears  to  occur  directly  from  complex  ions, 
this  explanation  is  unsatisfactory.  Neither  can  the  influence  on  cathode 
efficiency  explain  the  results,  since  efficiencies  in  the  silver  bath  are  close 
to  100  per  cent. 

It  seems  likely  that  concentration  effects  at  the  cathode  surface  are  im- 
portant. Glasstone6*  observed  that  small  changes  in  cyanide  concentration 
have  large  effects  on  electrode  potential.  However,  ordinary  polarization 
measurements  do  not  parallel  throwing  power  very  closely. 

All  commercial  baths  with  good  throwing  power  are  alkaline.  It  is  not 
known  whether  this  rule  applies  to  other  baths.  Metals  remain  in  alkaline 
solution  only  by  forming  complexes,  and  hence  good  throwing  power  is  to 


COORDlNATIOh   COMPOl  ND8  l\   ELECT RODEPOSlTIOh  645 

be  expected;  it  is  unlikely  thai  alkalinity  exerts  any  direct  influence.  There 
•11  to  be  do  data  on  the  throwing  power  of  highly  stable  complexes  in 
acid  solution.  That  of  the  chromium  bath  is  very  poor  but  this  bath  is  ex- 
ceptional in  many  ways. 

The  Plating  of  Specific    Metals  prom  Aqueous  Soli  pions 
of  Complex  Eons 

Metals  which  can  be  deposited  from  aqueous  solution  in  nearly  pure  form 

(i.e.,  not  as  amalgams  or  alloys)  arc  located  in  one  area  of  the  periodic 
table.  Furthermore,  it'  the  metals  are  classified  according  to  the  inner  or 
outer  orbital  configuration  of  their  complexes,  they  fall  into  four  fairly  well 
defined  regions  (see  Fig.  19.3).  The  plating  of  pure  zirconium,  columbium, 
molybdenum,  tungsten,  and  tantalum  is  still  classed  as  doubtful81  though 
several  alloys  of  the  latter  group  of  metals  can  be  plated  from  aqueous 
solution. 

If  one  considers  the  hydrated  ion  a  complex,  complex  ions  are  involved 
in  all  cases  of  elect rodeposition  from  aqueous  solution;  however,  in  agree- 
ment with  genera]  practice,  solutions  containing  the  hydrated  ions  will  be 
classed  as  SOlul  ions  of  the  simple  salts  unless  hydrate  isomerism  is  observed, 
as  in  the  case  of  chromium.  In  only  a  tew  cases  have  the  complex  ions  pres- 
ent in  specific  plating  solutions  been  identified.  Even  isolation  of  a  specific 
solid  complex  Buch  as  one  of  the  cyanides  of  copper  gives  no  assurance  thai 
the  particular  complex  is  present  as  such  in  solution. 

.Metals  forming  cyanide  anions  with  low  coordination  numbers  tend  to 
deposit  readily.  Dicyanide  is  very  favorable,  tetracyanide  intermediate, 
and  hexacyanide  and  octocyanide  are  very  unfavorable  for  deposition1. 
Since  metals  of  ( iroup  I B  form  the  dicyanide  while  members  of  ( rroup  VIII 
form  the  hexacyanide  and  molybdenum  and  tungsten  form  the  octocyanide, 
this  generalization  also  emphasizes  periodic  relationships.  In  solutions  of 
copper  cyanide,  increase  in  cyanide  concentration  reduces  cathode  effici- 
ency, since  copper  complexes  of  higher  coordination  number,  [Cu(CN)s]~ 
and  [Cu(CN)J  ,  are  formed.  Likewise  in  the  cadmium  bath,  which  contains 

largely  [Cd(CN)j]  ,  an  increase  in  cyanide  ion  concentration  lower.-  the 
current  efficiency.  Pure  zinc  cyanide  bath-  contain  chiefly  [Zn(CN 
and  show  such  low  current  efficiencies  that  cyanide  and  zincate  solutions 
arc  mixed  t<»  produce  commercial  baths.  Mercury  deposits  readily  from 
cyanide  solution  and  the  deposition  is  not  affected  by  excess  cyanide.  The 
solution  appears  to  contain  [Hg(CN  •.  with  traces  of  [Hg(CN  .  Deposi- 
tion comes  largely  from  the  tetracyano  ion,  which  is  scarcely  affected  by 
excess  cyanide'*.  These  observations  are  in  accord  with  the  hypothesis  that 
one  of  the  cyanide  groups  is  lost  by  dissocial  ion  ;i-  the  first  step  in  the  depo- 
sition pro. 

86.  Blum,  Monthly  Rev.  Am.  Blectroplater'i  Soc.,  27, 923    1940 


646 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


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648  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

The  deposition  of  metals  other  than  those  of  Groups  IB  and  IIB  from 
cyanide  solutions  may  oecasionally  be  promoted  by  the  presence  of  a  second 
complex-forming  ion  such  as  tartrate  and  by  the  deposition  of  certain  alloys 
rather  than  that  of  the  pure  metal.  The  effect  of  the  complex  forming  ion 
is  not  understood,  but  the  function  of  the  alloy  seems  to  be  related  to  the 
reduction  of  the  free  energy  of  the  metal  in  the  deposit. 

Group  VIII  metals  are  deposited  from  ammines  or  nitroso  complexes  and 
not  from  cyanides.  Both  cobalt  and  nickel  may  be  plated  from  ethylene- 
diamine  complexes;  ammonia  complexes  are  useful  in  deposition  of  platinum 
and  palladium,  while  ruthenium  may  be  plated  from  nitroso  ammine  com- 
plexes. Ammine  complexes  are  not  suitable  for  the  technical  plating  of  silver 
and  gold  in  Group  IB  because  of  low  anode  corrosion. 

The  elements  near  each  end  of  the  plating  groups  are  more  difficult  to 
deposit.  Oxyanions  such  as  Cr04=,  Re04=,  Ge044_,  Sn044_,  As03=,  Se03=, 
and  Te03=  may  be  used.  Except  for  chromium,  the  baths  range  from  moder- 
ately to  strongly  alkaline. 

Lead,  tin,  and  bismuth  appear  to  be  deposited  from  the  simple  hydrated 
ions  better  than  from  "complex"  ions.  Solutions  containing  complex  ions  of 
low  coordinating  ability,  such  as  BF4~,  N03~,  and  C104~,  are  suitable. 
In  the  absence  of  addition  agents,  deposits  of  these  metals  are  frequently 
coarsely  crystalline. 

Other  elements  can  be  deposited  about  equally  well  from  simple  or  com- 
plex solutions.  This  group  includes  cobalt,  copper,  iron,  gallium,  manganese, 
nickel,  rhenium,  thallium,  bismuth,  and  zinc  (see  Fig.  19.4). 

Deposition  of  Pure  Metals  from  Aqueous  Solution 

Arsenic,  Antimony,  and  Bismuth.  Arsenic  is  deposited  from  solutions 
of  arsenite  or  thioarsenite  ions,  preferably  with  small  amounts  of  cyanide87 
or  chloride88.  At  the  dropping  mercury  electrode,  it  is  deposited  from  acid 
solutions,  if  chlorides  are  present89. 

Chlorides  are  also  necessary  for  the  polarographic  reduction  of  antimony- 
(V)90.  Reduction  of  antimony (II I)  to  the  metal  apparently  does  not  require 
halides.  The  so-called  "explosive"  antimony  is  deposited  from  chloride  solu- 
tion at  current  densities  so  high  that  appreciable  amounts  of  chloride  are 
included  in  the  deposit  (see  p.  031).  Other  coordinating  substances  invest  i- 

87.  Hammett  and  Lorch, ./.  Am.  Chan.  S<><-.,  55,  71  (1933). 

88.  Rodionov,  Russian  Patent  27,546  (1927);  Torrance,  Analyst,  63,  104  (1938). 

89.  Khlopin,  Zkur.  Obschei  Kkim.,  18,  264  (1948);  Kolthoff  and  Lingane,  "Polarog- 

raphy,"  p.  261,  New  York,  [nterscience  Publishers,  Inc.,  1941. 

90.  Lingane  and  Nichida,  •/.  .1///.  Chem.  Soc.,  69,  530  (1947). 


)RDINATION  COMPOUNDS  l\   BLECTRODBPOSITIOA  f»l(.» 

gated  include  fluoride91,  sulfate9*,  tartrate91-1  w,  oxalate914,  and  sulfide91.  The 
fluoride  bath  is  preferred.  Deposits  are  also  obtained  from  antimony  poly- 
sulfide99. 

Bismuth  is  deposited  from  solutions  of  the  fluosilicate,  fluoborate,  per- 

chlorate94  and  nitrate"''.  Since  these  anions  have  little  tendency  to  form 
complexes,  it  is  probably  theaquated  ion  which  is  reduced.  Chloride  com- 
plexes, as  NaBiCU,  have  also  been  used-"1''  ". 

Cadmium.  Cadmium  forms  only  outer  orbital  complexes,  and  accord- 
ingly it  appeal's  to  be  deposited  from  all  of  its  water-soluble  compounds. 
Commercial  cadmium  plating  is  conducted  from  cyanide  baths  containing 
addition  agents97.  Cadmium  sulfate  baths  used  in  electrowinning  give  rough, 
crystalline,  "treed"  deposits  unless  an  addition  agent  such  as  gelatin  is 
used'71'.  The  cyanide  bath  requires  addition  agents  to  give  smooth,  bright 
deposits.  Organic  agents  such  as  sulfonic  acids,  resins,  aldehydes,  and  lic- 
orice extract,  and  inorganic  agents  such  as  nickel  or  cobalt  salts,  are  used, 
often  simultaneously. 

Both  Xa2[Cd(CX)4]  and  Na[Cd(CN)8]  exist  in  solution*1'98  and  also  in 
crystalline  form1.  Excess  cyanide  lowers  the  current  efficiency",  probably 
by  increasing  the  proportion  of  the  tetracyano  complex.  On  the  other  hand, 
conductivity,  anode  corrosion,  and  throwing  power  are  improved. 

Poor  deposits  are  obtained  from  the  ammine,  [Cd(XH:j)4]++10U,  as  well 

91.  Betts,  Trans.  Am.  Elect rochem.  Soc,  8, 186  (1905);  Bloom.  British  Patents 567794 

(1945)  and  559164  (1944);  U.  S.  Patent  2389131  (1945);  British  Patent  2941  13 
(1927);  Mathers,  Mental  Cleaning  and  Finishing,  7,  339  (1935);  Mathers  and 
Means,  Trans.  Am.  Electrochem.  Soc.,  31,  289  (1917);  Mathers.  Means,  and 
Richards,  Trans.  Am.  Electrochem.  Soc,  31,  293  (1917). 

92.  Piontelli  and  Tremolada,  Met.  ital.,  32,  417  (1940);  cf.,  Chem.  Abs.}  37,  1336 

(1943). 

93.  Salmoni,  AttiX  congr.  intern  cAtm.,3,614  (1939);  cf.  Chem.  Abs.,9S,  8504    1 

94.  Harbaugh  and  Mathers,  Trans.  Electrochem.  Soc,  64,  293  (1933);  Kern  and  Joi 

Trans.  Am.  Electrochem.  Soc,  57,  255  (1930);  Piontelli,  Atti  X  cot 
caim.,  3,  609  (1939);  cf  ,  Chem.  Abs.,  33,  9148  (1939). 

95.  Vozduizhenskii,  Kamaletdinov,  and  Khusianov,   Trans.   Bulk    0     / 

Tech.,  Kazan,  Xo.  1,  102  (1934);  cf.  Chem.  Abs.,  29,  391S  (1935). 

96.  Levin,  J.  Applied  Chem.,  U.S.S.R.,n,W>    1944  ; cf.  Chem.  Abs.,  40, 2075    1946 
Hall  and  Hogaboom,  "Plating  and  Finishing  Guidebook,"  1  1th  ed.,  p.  15, 

York,  The  Metal  Industry  Publishing  Co.,  1915;  Russell  and  Wbolrich,  British 
Patent  12526  (1849);  Soderberg  and  Weetbrook,  /  at      I  < ..  80, 

492  (1941). 
98.  Britton  and  Dodd,  ./.  Chi  m.  Soc,  1932,  1940. 

.  //,   .  .i .     /.  eel  ■       $oc  ,  30,  603    194 

100.  Brand.  Z.  anal.  Chem.,  28,  581  (1889,-  Clark.  Bet  .  11,  1409    1879  ;  Davison, 
./.  Am.  Chem.  8oe.t  27,  1275    1905  ;  Yvei    Bull.  soc.  chim.t  Pa\  -   34.  18    1880 


050  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

as  from  the  acetate,  formate,  lactate,  succinate,  and  oxalate101.  Deposits 
from  the  thiosulfate  bath  may  contain  5  per  cent  sulfur37,  and  those  from 
complex  halide  solutions  contain  halides79a.  The  sulfamate102,  fluoborate103, 
and  ethylenediamine104  baths  have  also  been  studied. 

Chromium.  In  chromium(III)  baths,  the  formation  of  chromium(II) 
at  the  cathode  is  vital105,  since  chromium(III)  complexes  are  inner  orbital. 
The  deposits  are  poor,  although  acceptable  for  electrowinning106,  and  cath- 
ode efficiencies  are  low.  Additions  of  acetate,  tartrate107,  benzoate,  and 
salicylate108  are  not  beneficial,  but  oxalates  are  helpful.  It  is  reported34-  109 
that  better  results  are  obtained  with  the  blue  ammonium  trisoxalatochro- 
mate(III)  than  with  the  red  ammonium  diaquobisoxalatochromate(III); 
the  presence  of  ammonium  ion  is  essential.  Deposits  are  also  obtained  from 
citrate  complexes110.  Contrary  to  some  reports,  there  is  no  significant  dif- 
ference between  plating  from  the  violet  hexaquochromium(III)  sulfate  or 
chloride,  and  from  the  green  chloraquo  or  sulfatoaquo  isomers105.  Deposits 
are  not  obtained  from  [Cr(NH3)6]Cl3 ,  [Cr(NH3)5Cl]Cl2 ,  [Cr(en)8]Cl3  , 
[Cr(en)3](CNS)3 ,  [Cr(urea)6]Cl3 ,  or  K3[Cr(ox)3]-3H20105. 

Commercially,  chromium  is  plated  from  chromic  acid  solutions  contain- 
ing sulfate  ion  in  the  proportion  of  1  part  to  100  parts  Cr03m.  Fluosilicic 
acid,  fluorides,  and  fluoborates  may  replace  a  portion  or  all  of  the  sulfate. 
The  cathode  efficiency  is  low-rarely  greater  than  15  per  cent. 

The  mechanism  of  reduction  is  not  understood.  Although  chromium  (III) 
ion  is  produced  in  the  operation,  radioactive  trivalent  chromium,  when 
added  to  the  bath,  does  not  enter  the  deposit112.  A  divalent  complex  is 
probably  involved.  At  the  dropping  mercury  electrode,  both  trivalent  and 
divalent  states  are  recognized  in  the  reduction113. 

Cobalt.  Cobalt  is  generally  deposited  from  sulfate  baths;  ammonium 
salts,  boric  acid,  and  sodium  fluoride  or  chloride  may  be  added114.  The 

101.  Mathers  and  Marble,  Trans.  Am.  Electrochem.  Soc,  25,  297  (1914). 

102.  Piontelli  and  Giulotto,  Chimica  e  industria,  Italy,  21,  278  (1939);  Piontelli, 

Korrasion  u.  Metallschvlz.,  19,  110  (1943);  cf.,  Chem.  Abs.,  38,  2571  (1944). 

103.  Anantharaman  and  Balachandra,  J.  Electrochem.  Soc,  100,  232  (1953);  Narcus, 

Metal  Finishing,  43,  188  (1945). 

104.  Harford,  U.  S.  Patent  2377228  (1945) ,  2377229  (1945). 

105.  Parry,  Swann,  and  Bailar,  Trans.  Electrochem.  Soc,  92,  507  (1947). 

106.  Lloyd,  Rawles,  and  Feenej  ,  Trans.  Electrochem.  Soc,  89,  443  (1946). 
K)7.  Britton  and  Wescott,  Trans.  Faraday  Soc,  28,  627  (1932). 

108.  LeBlanc,  Trans.  Am.  Electrochem.  Soc,  9,  315  (1906). 
L09.  Mazzucchelli,  Atti  acad.  Lincei.,  12,  587  (1930). 

110.  Kasper, ./.  Research., Nat.  Bur.  Standards,  11,  515  (1933);  Yn tenia, ,/.  Am.  Chem. 

Soc,  54,  3775  (1932). 

1 1 1 .  I  lubpernell,  Trans.  Electrochem.  Soc,  80,  589  (1941). 

I L2.  I  >gburo  and  Brenner,  Trans.  Electrochem.  Soc,  96,  347  (1949). 
113.  Lingane  and  Kolthoff,  •/.  Am.  Chem.  Soc,  62,  852  (1940). 

11  I.  Soderberg,  Pinner,  and  Baker,  Trans.  Electrochem.  Soc,  80,  579  (1941);  Watts, 
Trans.  Am.  Electrochem.  Soc,  23,  99  (1913). 


OKDINATIOh   COMPOl  \l>s/\   ELECT RODE  POSIT IOh  651 

aquated  ion  is  readily,  though  irreversibly,  reduced.  Poor  results  are  ob 
tained  from  thiocyanate  solution11',  but  bright  plates  are  reported  from 
a  triethanolamine  hath11". 

Cobalt(III)  complexes  show  varying  results12.  These  inner  orbital  com 
plex  ions  are  reduced  to  outer  orbital  cobalt  MI)  complexes  prior  to  depo 
sition,  as  is  clearly  shown  at  the  dropping  mercury  electrode16,  in  accord- 
ance with  the  discussion  on  p.  629.  No  deposit  is  obtained  from  inner  orbital 
cobalt(I]  |  complexes. 

Copper.  Commercial  copper  deposition117  is  carried  out  from  sulfate 
baths118,  used  chiefly  for electrorefining  and  electrotyping,  and  from  cyanide 
baths1,119,  used  largely  for  electroplating. 

In  sulfate  baths,  copper  is  present  mainly  as  the  tetraquocopperl  II)  ion. 
It  has  been  supposed  that  its  planar  configuration  indicates  inner  orbital 
configuration,  but  recently  the  existence  of  two  series  of  copper(I]  I  com- 
plexes has  been  demonstrated40,  12°,  one  of  which  docs  not  permit  electro- 
deposition  and  is  presumably  inner  orbital,  whereas  the  other  gives  electro- 
deposits  and  is  outer  orbital120.  The  aquated  ion  belongs  to  the  latter  series. 

Added  tartrates  form  a  complex  with  iron  which  accumulates  in  the  bath 
and  prevents  contamination  of  the  deposit  from  tin-  source121.  Urea  and 
thiourea  produce  bright  plates122,  but  it  has  not  been  shown  that  they  form 
complexes  in  the  bath. 

Since  copper  is  univalent  and  diamagnetic  in  cyanide  baths,  it  has  only 
outer  orbital  configuration.  The  tricyano  complex  is  tin  principal  con- 
stituent, and  it  is  in  dynamic  equilibrium  with  di-  and  tetracyano  ions'. 
It  i-  assumed  that  deposition  occurs  from  the  dicyano  ion,  the  supply  of 
which  is  replenished  by  rapid  dissociation  of  other  complexes.  Thus,  factors 
promoting  a  shift  in  equilibrium  toward  higher  coordination  numbers,  such 
as  increase  in  cyanide-copper  ratio,  or  reduction  in  temperature,  decrease 
the  cathode  efficiency.  If  the  cyanide-copper  ratio  is  sufficiently  low.  and 
a  large  amount  of  sodium  or  potassium  hydroxide  is  added,  a  "high  Bpeed" 
copper  bath  is  obtained"9*,  which  at  high  temperatures  has  anode  and 
cathode  efficiencies  approaching  100  per  cent,  even  at  high  current  densi- 
ties. High  temperature-  favor  the  dicyano  ion.  The  bath  is  vigorously 
stirred  so  as  to  reduce  concentration  polarization. 

115.  Mathers  and  Johnson,  Trans.  Electrocht  74,  _'_"•    I 

116.  Broekman  and Nowlen,  Trans.  Ele<  69,  v>:;    I 

117.  Bandes,  Trans.  Electroch*  88,  263    I 

118.  Winkler,  Trans.  Elect  80.  :>-'!    1941   . 

119.  Bennei  and  Wernlund,  T  ant    Electrochei     S        80,  355    1941   ;  Graham  and 

Read,  Trans.  I  s        80,  :;il    1941). 

120.  Ray  and  Sen,  J.  Indian  Chem.  Sac.  26,  17:;    1948  ;  Sen,  Miznshima,  Curran,  and 

Quagliano,  •/    Am.  Chem.  Soc.,  77.  -Ml     1965 

121.  Rasumovinkov  and  Maslenikov,  ./.  Inst.  Hetals,  I:  ■      an,  42.  500    192* 

Caem.46«.,24,344; 
Keller,  l\  S.  Patent  2462870. 


652  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

In  the  conventional  cyanide  plating  bath,  both  anode  and  cathode  effi- 
ciencies are  low.  Under  some  conditions  the  anode  efficiency  may  fall  to 
zero  unless  Rochelle  salt  (potassium  sodium  tartrate)  is  added123.  Graham 
and  Read119b  suggest  that  the  tartrate  forms  temporary  complexes  with 
electrolysis  products  in  the  anode  film.  Citrate  has  also  been  used124. 
Sodium  sulfite  and  thiosulfate  are  recommended  as  addition  agents119b. 

Strangely  enough,  both  anode  and  cathode  efficiencies  are  improved  by 
increasing  the  total  concentration  of  the  tricyano  ion.  This  unexpected  effect 
on  anode  corrosion  is  attributed10"1  to  depolarization  as  follows: 
2[Cu(CN)3]=  +  Cu  ->  3[Cu(CN)2]-  +  e~. 

Since  cyanide  baths  are  extremely  toxic  and  have  other  defects125, 
many  other  complexes  have  been  investigated,  but  no  bath  equivalent  in 
all  respects  to  the  cyanide  solution  has  been  developed.  Pyrophosphate 
baths  have  had  some  application126.  Copper(II)  complexes  which  give  ac- 
ceptable results  include  the  ammine19-  127,  oxalate125-  127c,  formate128,  ethylene- 
diamine129,  diethylenetriamine129b-  13°,  thiosulfate37-  131,  thiourea115, 132*,  thio- 
cyanate133,  and  the  sulfamate102.  Monoethanolamine45, 52,  diethanolamine46, 
and  triethanolamine134  give  poor  deposits  unless  oxalate  is  added,  possibly 
forming  mixed  oxalato-amine  complexes.  Good  deposits  are  obtained  from 
baths  containing  copper  (I)  chloride  complexes  and  gelatin135. 

123.  McCullough  and  Gilchrist,  U.  S.  Patent  1863869. 

124.  Smith  and  Munton,  Metal  Finishing.,  39,  415  (1941). 

125.  Fink  and  Wong,  Trans.  Electrochem.  Soc,  63,  65  (1933). 

126.  Coyle,  Proc.  Am.  Electroplater's  Soc,  p.  113  (1941);  Gamov  and  Fomenko,  Rus- 

sian Patent  54546  (Feb.  28,  1939) ;  cf .,  Chem.  Abs.,  35,  2800  (1941) ;  Gershevich 
and  Gamburg,  Korroziya  i  Borba  s  Nei.,  6,  no.  2,  46  (1940) ;  cf .,  Chem.  Abs.,  36, 
4031  (1942);  Stareck,  U.  S.  Patent  2250556;  British  Patent  509650;  Canadian 
Patent  379802;  German  Patent  680304. 

127.  Hansel,  German  Patent  688696  (1940);  Kudra  and  Kleibs,  Zapiski  Inst.  Khim., 

Akad.  Nauk.,  U.S.S.R.,  6,  No.  3-4, 203  (1940), -of.  Chem.  Abs.,  35,2796  (1941); 
Levin,  J.  Applied  Chem.,  U.S.S.R.,  13,  686  (1940);  14,  68  (1941);  cf.  Chem. 
Abs.,  35,  3536  (1941) ;  36,  972  (1942). 

128.  Stareck  and  Passal,  U.  S.  Patent  2383895  (1945). 

129.  Brockman  and  Mote,  Trans.  Electrochem.  Soc,  73,  371  (1938);  Greenspan,  U.  S. 

Patent  2195454;  Trans.  Electrochem.  Soc,  78,  303  (1940);  Wilson,  U.  S.  Patent 
2111671  (Nov.  26,  1946). 

130.  Brockman,  Trans.  Electrochem.  Soc,  71,  255  (1937). 

L31.  Govaerts  and  Wenmaekers,  German  Patent  406360  (1924) ;  384250  (1923) ;  Thomp- 
son. Chem.  A-  Met.  Eng.,  10,  -458  (1912). 

132.  Gockel,  /.  Elektrochem.,  40,  302  (1934). 

♦Thiourea  shows  a  strong  tendency  to  stabilize  univalent  copper.  It  is  likely  that 
the  solution  contains  appreciable  amounts  of  the  copper(I)  complex. 

133.  Schlotter,  Oberjtacheniech.,  12,  45  (1935). 

131.  Brockman  and  Brewer,  Trans.  Electrochem.  Soc,  69,  535  (1936);  Brockman  and 

Tebeau,   Trans.  Electrochem.  Soc,  73,  365  (1938);  Schweig,  British  Patent 
503095  (March  31.  1939). 
136.  Dievand  I  ashkarev,/.  Applied  Chem.,  U. 8. 8.R.,  12, 686  (1939);  cf. ,  Chem.  Abs., 


I  OORDINATION  COMPOl  NDS  l\  ELECTRODEPOSITIOh  653 

Two-step  reduction  processes  are  observed  with  copper(II)  complexes  of 
ammonia19' 89b* m,  bromide  and  chloride116,  thiocyanate  '  Bb,  thiourea, 
and  pyridine1***'  l**b.  These  agents  stabilize  the  copper(I)  state  sufficiently 
for  it  to  be  observed  in  the  deposition  process.  Satisfactory  deposits  arc 

obtained  from  COpper(I)  thiosulfate*1  and  thiocyanate101  hath-. 

Gallium  and  Germanium.  Gallium  is  deposited  from  sulfate  or 
alkaline  gallate  solutions1*7.  The  process  is  irreversible,  presumably  because 

the  metal  ion  is  hound  in  a  colloidal  sol  by  hydrolysis   ". 

Deposits  of  germanium  are  obtained  from  both  sulfate  and  germanate 
solutions1**.  Oxalate,  tan  rate,  carbonate,  and  phosphate  additions  have 
been  suggested140;  it  is  not  known  whether  complexes  are  formed. 

Inner  orbital  complexes  are  not  formed  by  these  metals. 

Gold.  Although  deposition  from  many  gold  complexes  has  been  investi- 
gated, only  the  cyanide  and  chloride  baths  have  found  extensive  applica- 
tion141. The  former  contains  the  outer  orbital  dicyanoaurate(I)  ion.  In  early 
days  it  was  prepared  from  the  ferrocyanide,  which  was  available  in  higher 
purity  than  the  cyanide.  The  suggestion  that  the  gold(III)  complex  is 
formed141  is  doubtless  in  error.  The  ferrocyanide  is  still  employed  in  the 
••-alt  water"  process1411'. 

The  tetrachloroaurate(III)  complex  is  used  mainly  in  electrorefining. 
It  has  square  planar  configuration,  and  therefore  is  presumably  of  inner 
orbital  dsp2  type.  At  the  cathode,  it  is  reduced  to  the  unstable  dichloro- 
aurate(I)143,  which  has  outer  orbital  configuration.  The  bromide  bath  be- 
haves similarly.  Iodide  baths144  contain  gold  as  the  monovalent  complex, 
[Aulo]-.  Thiourea146,  thiocyanate,  thiosulfate,  polysulfide,  phosphate,  and 

33,  s.504  (1939);  Kameyama  and  Makishima,  ./.  Soc.  Chem.  I  ml..  ./<i/,<iri,  34, 
462  (1932);  36,  365  (1933);  38,  18  (1935). 

136.  Kolthoff  and  Lingane,  "Polarc-graphy,"  p.   17<>.  279,  New  York,  [nterscience 

Puhlishers,  Inc.,  1941;  Verdieck,  Ksychki,  and  Yntema,  Trans.  Electrochem. 
Soc.,  80,  n    1941). 

137.  Fogg,  Trans.  Electrochem.  Soc.,  66,  107  (1934);  Sebba  and  Pugh,  ./.  Chem.  S 

1937,  1371. 

138.  Moeller  and  King,  J.Am.  Chem.  Soc.,  74,  L355    1952  . 

139.  Alimarin  and  [vanov-Emin,  ./.  Applied  Chem.,  V  S  S  R  .  17.  _'ni     l1.  it  .  link 

and  Doki  Electrochem.  Soc.,  93,  80    1949  ;  Hall  and  Koenig,  Trans. 

Electrochi  m.  Soc.,  65,  215    1934 
140    Schwartz,  Heinrich,  and  Hollstein,  Z.  anorg.  aUgem.  Chem.f  229,  164    19 
141.  Frary,  Trans.  Am.  Electrochem.  Soc.,  23.  25,   19    L913  ;  Weisberg  and  Graham, 

Trans.  Elect  ochen    Soc  .  80,  5Q9    1941   . 
L42.  Beutel,  Z.  angew.  Chem  ,86,995    1912). 

Bjerrum,  Bull.  soc.  chim.  Beiges,  57,  132    1948 
144.  Schlotter,  V   8    Patent  1857664    M-c   I"    1932) ;  German  Patent  608268    Jan.  19, 

1935  . 
L45    Schonmann,  German  Patent  731043    Dec    24,  1942  . 


65  1  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

sulfite  baths141*-  li"'  have  been  described.  Kushner  summarized  noncyanide 
baths148,  and  general  commercial  practice147. 

Indium.  Indium,  plated  from  simple  sulfate  or  from  cyanide  solutions148, 
has  recently  found  rather  extensive  use  as  a  wear  and  corrosion  resistant 
coating  for  bearing  surfaces.  It  is  the  only  trivalent  metal  known  to  be 
deposited  readily  from  a  cyanide  bath149.  The  complex  cyanides  are  not  well 
understood;  Thompson1  states  that  only  the  tetracyanide,  [In(CN)4]~, 
is  well  known,  although  the  experimental  basis  for  this  statement  is  not 
given.  The  coordination  number  of  six,  attributed  to  indium150,  is  observed 
in  certain  compounds151.  Regardless  of  its  formula,  the  cyano  complex  is 
unstable  and  slowly  precipitates  the  hydroxide  from  water.  Stability  is 
improved  by  the  presence  of  a  large  excess  of  alkali  cyanide  together  with 
other  substances,  such  as  glucose,  tartrates,  and  glycine149. 

Deposits  from  sulfate  baths  containing  formate152,  citrate153,  fluoride154, 
hydroxylamine,  or  pyridine152  are  good,  but  oxalate  or  acetate  gives  poor 
results. 

Indium  forms  only  outer  orbital  complexes.  The  low  current  efficiencies 
in  both  the  sulfate  and  cyanide  baths155  and  the  corresponding  polarographic 
irreversibility156  are  probably  the  results  of  hydrolysis157.  In  the  presence 
of  chloride  ions,  the  reduction  becomes  reversible158,  presumably  because 
the  chloro  complex  is  less  readily  hydrolyzed. 

Iron.  Iron  is  electroplated  from  sulfate  or  chloride  baths159.  The  pres- 
ence of  iron(III)  ions  is  undesirable.  Chloro  or  sulfato  complexes  probably 
exist  in  solution  along  with  aquated  iron (II)  ions.  The  chloride  bath  gives 
better  deposits  at  high  temperatures;  the  sulfate,  at  low  temperatures. 

146.  Kushner,  Products  Finishing,  6,  no.  3,  22  (1941). 

147.  Kushner,  Products  Finishing,  4,  No.  12,  30  (1940),  5,  Nos.  1-12  (1940-41). 

148.  Hall  and  Hogaboom,  "Plating  and  Finishing  Guidebook,"  14th  ed.,  p.  61,  New 

York,  The  Metal  Industry  Publishing  Co.,  1945. 

149.  Cray,  Trans.  Electrochem.  Soc,  65,  377  (1934). 

150.  Mueller,  ./.  Am.  Chem.  Soc,  62,  2444  (1940);  64,  2234  (1942). 

151.  Ensslin  and  Dreyer,  Z.  anorg.  allgem.  Chem.,  249,  119  (1P42);  Klemm  and  Kilian, 

Z.  anorg.  allgem.  Chem.,  241,  93  (1939). 
L52.   Dennis  and  Geer,  ./.  Am.  Chem.  Soc,  26,  437  (1904). 
L53.  Westbrook,  Trans.  Am.  Electrochem.  Soc,  57,  289  (1930). 
i:,l.  Bartz,  British  Patent  564053  (Sept.  11.  1944 

155.  Linford,  Trans.  Electrochem.  Soc.  ,79, 443  (1941),  Whitehead,  Metal  Finishing,  42, 

105  (1944). 

156.  Kolthoff  and  Lingane,  "Polarography,"  p.  274,  New  York.  Interseienee  Publish- 

I  in-.,  11)  11 . 

[57.  Hattoxand  DeVi ies,  ./.  .1///.  Chem.  Soc.,  58,  2126  (1936);  Takagi,  J.  Chem.  Soc.t 
1928,  301. 

158.  Kolthoff  and  Lingane,  "Polarography,"  p.  263,  New  York,  Interseienee  Pub- 
lishers, Inc.  1941. 

l.V.i.  Thomas.  Trans.  Electrochem.  Soc.,  80,  499  (1941). 


COORDINATION  COMPOUNDS  TN  ELECTRODEPOSITIOh  655 

Sulfamate108  and  fluoroborate1811*'  m  baths  have  been  suggested.    Eron   is 
presenl  probably  as  the  aquatod  ion.  Deposit  ion  from  an  alkaline  bath  con 
taining  ethylenediaminetetracetic  acid  and  triethanolamine  has  recently 
been  reported160.  In  this  bath  iron  is  undoubtedly  presenl  as  a  complex  ion, 

but  its  nature  has  not  been  established. 

Iron  forms  both  inner  and  outer  orbital  complex  ions.  Deposition  is  pos- 
sible from  the  outer  orbital  aqUO,  chloro,  and  SUlfatO  complexes,  hut    not 

from  the  inner  orbital  cyano,  o-phenanthroline,  and  a,a'-dipyridyl  com- 
plexes, although  certain  alloys  may  be  deposited  from  the  cyanide  com- 
plex ions,  as  discussed  on  page  667. 

Lead.  The  best  lead  deposits  are  obtained  from  solutions  containing 
anions  of  low  complexity  power.  Lead  nitrate,  per  chlorate,  and  Baits  of 
fluoro  acids,  especially  fluoroantimonate,  fluorost annate,  fluoroborate,  and 
fluorosilicate  have  been  tried161.  The  last  two  have  found  commercial  ap- 
plication162. The  sulfamate  bath  also  gives  good  deposits102,  163.  Lead  is 
present  probably  as  the  aquated  lead(II)  ion.  The  deposition  is  reversible 
both  at  lead  and  at  mercury  cathodes164,  as  would  be  expected  from  the 
outer  orbital  configuration  of  the  ion. 

In  alkaline  solutions,  the  acetate165  gives  poor  deposits161.  A  bath  con- 
taining potassium  bisoxalatoplumbate(II)  with  excess  potassium  oxalate 
gives  good  deposits,  but  the  corresponding  ammonium  bath  gives  spongy 
metal166.  Lead  tartrate  in  the  presence  of  sodium  acetate  gives  compart 
deposits161-  167. 

Manganese.  Manganese  is  usually  deposited  from  a  sulfate  solution 
containing  excess  ammonium  sulfate168,  although  the  corresponding  chlorides 
may  be  used.  Because  of  the  strong  tendency  of  manganese  to  form  coordi- 
nation compounds169,  it  is  probable  that  deposition  occurs  from  outer  orbital 
sulfate  or  chloro  complexes.  Deposits  are  not  obtained  from  the  inner  orbital 
hexacyanomanganate(II),  except  at  a  mercury  cathode,  a1  which  the  high 
hydrogen  overvoltage  and  the  free  energy  of  amalgam  formation  allow 

160.  Foley,  Linford,  and  Meyer,  Plating,  40,  887  (1953). 

161.  Mathers.  Trans.  Am.  Electrochem.  So,-.,  23,  153  (1913). 

162.  Gray  and  Blum,  Trans.  Electrochem.  Soc.,  80,  645  (1941). 
L63.  Mathers  and  Forney.  Trans.  Electrochem.  Soc.,  76,  371  (19! 

L64.  Kolthoff  and  Lingane,  "Polarography,"  p.  267,  New  York,  [nterscience  Pub- 
lishers, Inc.,  1941  . 

L66.  Friend,  "A  Textbook  of  Inorganic  Chemist ry,"  vol.  5,  p.  433,  London, C. Griffin 
Co.,  1921. 

166.  Classen,  Ber.,  15,  1096  (1882). 

167.  Glazunov  and  Jenicek,  Korrosion   u.   Metallschutz,  17,  384     1941);  cf.  Chi 

Abs.,  36,  5095    L942). 

168.  Bradt  and  Taylor,  T  ans.  Electrochem.  Soc.,  73,  327    1938 

169.  Morgan  and  Buratall,  "Inorganic  Chemistry     A  survey  of  Modern  Develop 

ments,"  p.  195,  Cambridge,  England,  W.  Beffer  A  Bona    18 


656  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

deposition  to  proceed170.  The  addition  of  excess  ammonium  thiocyanate  has 
been  recommended168  for  the  sulfate  bath.  Manganese (II)  fluoroborate, 
benzoate,  acetate,  and  citrate  solutions  all  give  deposits,  as  do  sodium  ci- 
trate solutions  of  manganese(II)  dithionate,  tartrate,  formate,  acetate,  and 
fluorosilicate168,  m.  Complexes  with  amines,  such  as  mono-,  di-,  or  trietha- 
nolamines,  also  give  deposits172. 

Mercury.  Mercury  is  readily  deposited  from  the  complex  cyanide  bath; 
the  tetracyanomercury(II)  ion  predominates41,98,  although  there  may  be 
small  amounts  of  the  tricyano  ion1.  Little  activation  is  needed,  even  with 
the  divalent  ion.  Reduction  probably  proceeds  through  the  univalent  state, 
which  forms  only  outer  orbital  complexes.  Acetate  solutions  have  also  been 
studied173. 

Nickel.  Commercial  nickel  plating  baths  contain  nickel  sulfate  and 
chloride83,  usually  with  boric  acid.  An  all-chloride  bath  is  also  used174. 
Chloride  is  necessary  to  dissolve  the  nickel  anode  under  operating  condi- 
tions175, probably  through  forming  a  chloroaquo  complex.  Deposition  oc- 
curs from  both  aquo  and  chloro  complexes.  According  to  magnetic  data28, 
these  ions  have  two  unpaired  electrons,  indicating  outer  orbital  spzd2  hy- 
bridization. The  cyano  complex  has  no  unpaired  electrons,  so  that  the 
hybridization  is  inner  orbital  dsp2.  Deposits  from  cyanide  baths176  appear 
to  be  only  flash  deposits  and  plating  soon  ceases84.  The  deposition  of  nickel 
alloys  from  cyanide  baths  is  discussed  on  page  667. 

Ammoniacal  solutions  of  a  number  of  nickel  salts114b  contain  the  tetram- 
mine  complex177,  and  give  good  deposits.  Dark  deposits  of  so-called  black 
nickel  which  contain  sulfur  are  obtained  from  baths  prepared  b}^  dissolving 
nickel  carbonate  in  concentrated  solutions  of  potassium  thiocyanate114b, 
probably  giving  [Ni(SCN)4]=. 

Plating  solutions  containing  such  complex-forming  substances  as  oxa- 
late95,    citrate178,    pyrophosphate95,    tartrate179,    lactate178*-  178b,    thiocya- 

170.  KolthofT  and  Lingane,  "Polarography,"  p.  254,  New  York,  Interscience  Pub- 

lishers, Inc.,  1941. 

171.  Bradt  and  Oaks,  Trans.  Electrochem.  Soc,  71,  279  (1937);  69,  567  (1936);  U.  S. 

Patent  2398614  (Apr.  16,  1946). 

172.  Dean,  U.  S.  Patent  2317153  (Apr.  20,  1943);  cf.,  Chem.  Abs.,  37,  5663  (1943). 

173.  Malkin,  Ber.  Inst,  physik.  Chem.,  Akad.  Wiss.  I'kr.S.S.R.,  11,  109  (1938);  cf., 

Chem.  Abs.,  34,  2261  (1940). 

174.  Wesley  and  Carey.  Trans.  Electrochem.  Soc,  75,  209  (1939). 

175.  Dorrance  and  Gardiner,  Trans.  Am.  Electrochem.  Soc,  54,  303  (1928). 

176.  Bennett,  Rose,  and  Tinkler,  Trans.  Am.  Electrochem.  Soc,  28,  339  (1915);  Watts, 

Trans.  Am.  Electrochem.  Soc,  27,  141  (1915). 

177.  Kato,  •/.  Chem.  Soc,  Japan,  58,  1146  (1937). 

178.  Ballay,  Compt.  rend.,  198,  1494  (1934);  Franssen,  Oberfiachenteck.,  14,  174  (1937); 

cf.,  Chem.  Abe.,  31,  8387  (1947);  Nichols,  Trans.  Electrochem.  Soc,  64,  265 
L933). 

179.  Mathers,  Webb,  and  SchafT,  M vial  Cleaning  and  Finishing,  6,  412,  148  (1934). 


COORDINATION   COMPOl  NDB  IN  ELBi  TRODEPOSITIOh  667 

Qatelu' 180,  fluoride181,  triethanolamine11',  and  sulfamic  acid"-'  have  been 
studied.  In  general  the  depoeitfl  are  fairly  good,  but  the  baths  offer  no 
advantages  over  tin1  chloride  or  sulfate  bath.*  Fluoroborate  and  sulfamate 
baths  are  occasionally  used. 

In  tin4  presence  of  excess  thiosulfate,  the  deposits  are  smooth,  adherent, 
and  metallic,  bul  contain  from  22  to  70  per  cenl  sulfur11.  Ni  s.  wa&  identi- 
fied by  means  of  x-ray  diffraction,  The  presence  of  sulfide  may  be  taken 
to  indicate  that  coordinate  bonds  are  not  always  easily  converted  to 
metallic  bonds. 

A  Btudy*  of  elect rodeposits  from  nickel  complexes  showed  that  the 
smaller  the  coordinating  group,  the  better  the  form  of  the  deposit.  Thus 
the  tris(ethylenediamine)  complex  gives  better  plates  than  the  correspond- 
ing propvlenediamine  compound,  which  in  turn  is  better  than  the  butylene- 
diamine  ion.  It  is  possible  that  the  larger  groups  prevent  close  approach 
of  the  nickel  ion  to  the  cathode  so  that  conversion  of  coordinate  bonds  to 
metallic  bonds  is  more  difficult  than  with  the  smaller  groups. 

Metals  which  are  irreversibly  reduced,  such  as  nickel,  tend  to  be  de- 
ported more  smoothly  than  those  which  are  deposited  reversibly,  perhaps 
because  the  hindrance  to  deposition  precludes  the  formation  of  large  crys- 
tals. Accordingly,  nickel  deposits  are  particularly  susceptible  to  the  in- 
fluence of  addition  agents.  Nevertheless,  the  formulation  of  a  nickel  bath 
t<>  yield  bright  deposits  under  the  conditions  encountered  in  industry  is 
difficult.  Two  classes  of  addition  agents  are  recognized,  the  active  agent- 
and  the  carriersS3  (see  discussion,  page  643).  Although  the  mechanism  by 
which  these  function  is  unknown,  there  is  probably  a  better  empirical 
knowledge  of  nickel  brighteners  than  of  those  for  other  metals. 

The  Platinum  Group  Metals:  Ruthenium,  Rhodium,  Palladium, 
Osmium,  Iridium,  Platinum.  The  water-soluble  compounds  of  the  plat- 
inum metals  all  seem  to  be  inner  orbital  complexes.  Nevertheless,  depo 
have  been  reported.  Lyons  suggests  that  this  may  result  from  the  extreme 
stability  of  the  metallic  state, f  so  that  the  energy  required  to  break  the 
inner  orbital  hybridization  doe-  not  greatly  exceed  that  needed  to  discharge 

*  Triethanolamine  and  ammoniacal  citrate  baths  permit  direct  plating  on  zinc. 
Ordinary  baths  plate  nickel  on  zinc  by  displacement  and  such  deposits  are  BDOngy 
and  do  not  afford  a  satisfactory  base  for  subsequent  electrodepoflits.  The  deposition 
potential  of  nickel  in  these  special  I  pparently  raised  to  thai  of  zinc  (see 

alloy  plating,  page  666).  A  nickel  sulfate  bath  containing  substantia]  amounts  of 
sodium  sulfate  has  also  been  used;  a  sulfato complex  was  probably  formed. 

t  The  heat  of  sublimation  of  platinum  is  1.86  electron-volts11*. 

180.  Schone,  Metal  Finishing,  41,  77     ' 

181.  house,  Can.  Patent  101154;  Spiro  and  Wohlgemuth,  British  Patent  584877   Jan. 

28,  19 

182.  Jv-lley,  "Heats  of  Fusion  of  Inorganic  Compoundi  "'     s  itesBull., 

393_(1936). 


658  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

the  hydrogen  ion40.  The  current  efficiencies  are  quite  low,  and  the  deposition 
reactions  are  irreversible.  The  deposition  of  heavy  coatings  seems  generally 
to  be  difficult,  and  most  investigators  have  been  satisfied  with  "flash" 
deposits.  Information  on  the  plating  of  osmium,  iridium,  and  ruthenium 
is  scanty,  and  it  may  be  that  only  "flash"  deposits  are  obtained.  With 
rhodium,  platinum,  and  palladium,  heavier  deposits  are  obtained,  although 
with  difficulty.  Cyanide  complexes  give  no  deposits  of  the  platinum  met- 
als183. 

Electrodeposition  of  these  metals  is  not  well  developed,  owing  largely  to 
these  difficulties  and  to  the  expense  of  the  metals.  Rhodium  plating  has 
received  attention  because  of  the  high  reflectivity,  corrosion  resistance, 
and  hardness  of  the  deposit.  It  appears  to  be  the  easiest  of  the  group  to 
electrodeposit. 

Rhodium  is  generally  plated  from  acid  electrolytes184.  The  most  common 
baths  are:  (1)  a  solution  of  rhodium  sulfate  in  sulfuric  acid;  (2)  a  solution 
of  rhodium  phosphate  in  phosphoric  acid;  or  (3)  a  mixture  of  the  two. 
Additional  alkali  sulfates  or  phosphates  may  be  added185.  The  solutions  are 
undoubtedly  complex,  and  may  contain  ions  of  the  type  [Rh(S04)3]=  or 
[Rh(P04)2]=186.  No  simple  solid  phosphates  of  rhodium  have  been  isolated; 
only  complex  phosphates  of  variable  composition  have  been  produced. 

Addition  agents  suggested  for  the  sulfate  bath  include  the  complex 
forming  substances,  di-  and  trimethyl-  and  ethylamines  and  tartaric  and 
lactic  acids187. 

Complexes  recommended  for  rhodium  plating  include  chlorides,  as 
Xa3[RhCl6],  K3[RhCl8],  (NH4)3[RhCl6],  and  H3[RhCl6]188;  and  nitrites,  as 
(NH4)3[Rh(X02)6]189  and  [Rh(NH3)4(N02)2]N02190.  Good  deposits  of  rho- 
dium have  been  reported  from  solutions  prepared  by  dissolving  rhodium 
hydroxide  in  sulfamic  acid102a,  nitric  acid191,  fluoroboric  acid,  and  perchloric 
acid192. 

Platinum  black  is  a  typical  powdery  deposit,  obtained  from  the  hexa- 

183.  Grube  and  Reinhardt,  Z.  Elektrochem.,  37,  316  (1931). 

184.  Schumpelt,  Trans.  Electrochem.  Soc,  80,  489  (1941). 

185.  Fink  and  Lamhros,  Trans.  Electrochem.  Soc,  63,  181  (1933). 

186.  Yamamato,  Rept.  Chem.  Research,  Prefectiual  Inst.  Advancement  Ind.,  Tokyo., 

no.  2,2-12  (1940) ;  ci.,  Che?n.  Abs.,  35,  7840  (1941). 
L87.  Spies,  German  Patent  692122  (May  16,  1940). 
L88.  Weisberg,  Metal  Finishing, .38, 687  (1940). 

189.  Keitel,  U.  S.  Patent  2067534  (Jan.  12,  1937);  Can.  Patent  365965  (May  11,  1937); 

Zimmermann,  U.  S.  Patent  2067747  (June  12,  1937). 

190.  Keitel,  T.  S.  Patent  1779436  (Oct.  28,  1930);  Zschiegner,  U.  S.  Patent  1779457 

Oct.  28,  L930). 
I'M.  British  Patent  480145  (Feb.  17,  1938). 

192.  link  and  Deren,  Trans.  EUctrochem.  Soc,  66,  471  (1934);  Grube  and  Resting, 
/.  Elektrochem..  39,  951  (1933). 


\RDINATIOh   COMPOUNDS  IN  ELECTRODEPOSITJOh  659 

chloroplatinate  IV  ion;  the  reduction  proceeds  through  the  tetrachloro 
platinate(II)  ion  to  the  free  metal181.  Although  inner  orbital  in  configuration, 
the  latter  ion  is  thermodynamically  unstable1,7b  and  disproportionates  to 
metal  and  the  tetravalenl  ion.  Sometimes  this  results  in  colloidal  metal  in 
the  plating  bath191.  The  instability  of  this  inner  orbital  complex  probably 
reflects  the  high  stability  of  the  metal. 

Bright  platinum18*  is  generally  plated  from  a  bath  prepared  by  boiling 
potassium  hexachloroplatinate(IV)  with  a  solution  of  disodium  and  di- 
ammonium  phosphates.  A  color  change  during  boiling  and  the  dissolving 
of  the  precipitate  of  (NH^itPtCle]  suggesi  formation  of  an  ammine-phoe 
phato  complex,  hut  it  has  not  been  isolated.  Thick  deposit-  cannot  be 
obtained,  the  current  efficiency  is  low,  and  the  hath  deteriorates  in  use, 
since  metal  must  he  replenished  by  adding  more  complex,  and  thus  phos- 
phates and  chlorides  accumulate.  Deterioration  is  less  marked  if  accumula- 
tion of  chloride  is  avoided  by  replenishing  with  dinitrodiammineplati- 
num(II). 

A  somewhat  superior  hath  is  prepared  from  ammonium  nitrate,  ammo- 
nium hydroxide,  sodium  nitrite,  and  [Pt(NH  lj ■  ■■  X ( )-j » -j | .  The  complex  ex- 
pected under  these  conditions  is  tetrammineplatinum(II)  nitrite, 
[Pt(NHj)4](NOj)j184.  The  bath  is  replenished  with  the  dinit  rodiammine- 
platinum(II),  and  excessive  accumulation  of  salts  is  avoided  by  decompo- 
sition of  ammonium  nitrite. 

Although  rhodium  is  deposited  at  the  dropping  mercury  electrode194, 
platinum  is  not  deposited  but  catalyzes  hydrogen  evolution87, 194» 1M. 
(irube196,  however,  reported  the  reduction  of  platinum  from  the  tetracyano 
ion  on  a  mercury  cathode. 

Palladium  is  similar  to  platinum.  A  solution  containing  palladium(II) 
chloride,  disodium  and  diammonium  phosphate.-,  and  benzoic  acid  has 
been  used184.  Solutions  containing  dinitrodiamminepalladium(II), 


[Pd(NH,)2(NO, 


have  also  been  recommended197.  Baths  prepared  with  ammonium  tetra- 
chloropalladate(II)  give  good  deposits,  but   corresponding  potassium  or 

sodium  -alt-  give  no  deposit1"''.  Since  the  tetrachloropalladate  II     ion  is 
-aid  "  to  be  rapidly  reduced  by  hydrogen  in  the  cold,  easy  electrodeposition 

193.  McCaughey,  Trans.  Electrochem.  Soc.,  15.  623    1909  ;  McCaughey  and  Pat  ton, 

T  an*.  Ele*  I  ocht  ■■  .  Soe.t  63,  181  L910 

194.  Willis,  ./.  .  66.  1067  1944 

195.  Latinen  and  Onstott,  J.  A       <  ■  72,  1565    L9f 

196.  Grube  and  Beiacher,  Z.  Elel  30.    - 

107.  Klochko  and  Medvedeva,  /.  Applied  Chen   .'    8.S.R    15,25 
L96.  [patiev,  and  Tronev,  J.  Gen.  Chi  s   3LR.,  5.  643    IS 


660  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

would  be  expected.  Unlike  platinum,  palladium  is  deposited  at  the  drop- 
ping mercury  electrode  from  the  tetrachloro  complex194. 

Ruthenium  may  be  deposited  from  a  solution  prepared  by  dissolving 
the  nitrosochloride,  [Ru(NO)Cl3]-H20,  in  dilute  sulfuric,  phosphoric,  hydro- 
chloric, or  oxalic  acid.  Since  the  normal  coordination  number  of  ruthenium 
is  six,  water  or  sulfate  may  be  coordinated  in  the  remaining  positions. 
Nitrosoammine  complexes  of  unspecified  composition  have  also  been 
recommended199. 

Little  is  known  of  the  deposition  of  osmium  and  iridium,  though  baths 
containing  chloro  complexes  have  been  described200.  Ions  of  the  type 
[IrCl6]~  would  be  expected201  in  these  baths.  Ruthenium,  osmium,  and 
iridium  are  not  deposited  at  the  dropping  mercury  cathode194. 

Polonium.  Polonium,  or  radium  F,  has  not  been  available  in  sufficient 
quantities  to  permit  study  of  its  complex  compounds  on  a  macro  scale; 
however,  certain  of  them  are  known  to  be  isomorphous  with  complexes  of 
lead,  tellurium,  and  tin.  By  assuming  that  they  have  similar  formulas, 
compounds  such  as  (NH^PPoCle]  and  (NH4)3[PoCl6]  have  been  suggested202. 
Haissinsky203  states  that  polonium  forms  complexes  with  a  large  number  of 
ions  such  as  sulfate,  acetate,  oxalate,  and  even  ions  of  low  complexing 
tendency  such  as  nitrate.  Polonium  is  readily  deposited  from  solutions  of 
such  complexes,  which  are,  of  course,  outer  orbital  in  type.  A  summary  of 
the  electrochemistry  of  polonium  is  given  by  Haissinsky203. 

Rhenium.  Electroplated  rhenium  is  bright  and  hard,  resistant  to  hydro- 
chloric acid139b,  but  readily  attacked  by  nitric  acid  or  moist  air204.  Baths 
are  prepared  by  dissolving  potassium  perrhenate,  KReCU  ,  in  dilute  solu- 
tions of  sulfuric13913 •  204,  phosphoric13915,  oxalic139b,  and  hydrofluoric  acids205. 
Dilute  nitric  and  hydrochloric  acids  are  unsatisfactory204.  Perrhenate  baths 
somewhat  resemble  chromate  baths.  A  solution  of  the  chloride  complex, 
K2[ReCl6],  gives  only  traces  of  metal  on  a  platinum  cathode,  even  at  high 
current  density206.  With  a  mercury  cathode  an  amalgam  of  rhenium  is 
formed. 

Selenium.  Selenium  is  semimetallic  in  nature  and  forms  few  coordina- 
te. Zimmerman  and  Zschiegner,  U.  S.  Patent  2057638;  French  Patent  799251 ;  British 
Patent  466126;  German  Patent  647334  (1936). 

200.  Rossman,  Metal  Tnd.}  {N.  )'.),  29,  245  (1931). 

201.  Morgan  and   Hnrstall,   "Inorganic  Chemistry — A  survey  of  Modern  Develop- 

ments/' p.  233.  Cambridge,  England,  W.  Heffer  &  Sons,  1936. 

202.  Emeleua  and    Anderson,    "Modern    Aspects  of  Inorganic  Chemistry,"  p.  371, 

New  York.  I).  Van  Nostrand  Co.,  Inc.,  1938. 

203.  Baisaineky,  Trans.  Electrochem.  Soc.,  70,  343  (1936). 

204.  I.nndell  and  Knowlee,  ./.  Research  Natl.  Bur.  Standards,  18,  629  (1937). 
20.").  Holemann,  '/..  anorg.  allgem.  Chan.,  235,  1  (1937). 

206.  Holemann,  Z.  anorg.  allgem.  Chem., 211, 195  (1933). 


COORDI.XM  I<>\   COMPOl  VDB  l\   ELBCTRODBPOSITIOh  661 

tion  compounds.  It  is  deposited  in  alloys  with  such  metals  as  copper,  bis- 
muth, or  nickel  from  an  acid  solution  containing  SeOi"  and  various  addition 

agents  such  as  oxalic  acid-"7.  These  alloys  probably  resemble  the  nickel- 
sulfur  deposits  mentioned  above.  Pure  selenium  may  be  plated  on  the 

anode  by  electrolysis  of  solutions  of  selenides,  such  as  \a-_.Si 

Silver.  Univalent  silver  forms  only  outer  orbital  ions,  from  which  it 
deposits  bo  readily  that  it  tends  to  form  coarse  crystals.  No  addition  agent 
lias  been  found  which  will  give  compact,  smooth  deposits  from  theaquated 

silver  ion  in  nitrate,  perchlorate,  or  fluorohorate  baths. 

The  Bole  bath  of  commercial  importance  is  the  cyanide-"',  which  has  been 
used  with  only  minor  modification  since  its  introduction  in  1838.  The 
principal  complex  ion  is  the  dicyano,  [Ag(CX)2]~;  the  existence  of  tri-  or 
tetracyano  ions  is  negligible  under  most  conditions1.  Correspondingly,  the 
cathode  efficiency  is  not  much  affected  by  changes  in  cyanide  ion  concen- 
tration or  in  temperature;  it  is  substantially  100  per  cent  under  most  con- 
ditions. Evidently  the  dicyano  ion  is  well  suited  to  the  deposition  mecha- 
nism. The  ferrocyanide  used  in  early  baths141a  was  undoubtedly  converted 
to  the  dicyano  ion. 

A  number  of  complexing  agents  have  been  proposed  to  replace  the  toxic 
cyanide.  Chloride,  [AgCl2]~,  and  iodide,  [Agl2]~,  were  suggested  early14111. 
Plates  comparable  to  those  from  cyanide  solution  have  been  obtained  from 
iodide  baths1-7'  144;  the  addition  of  citric  acid210  has  also  been  recom- 
mended. Thiosulfate  complexes,  probably  [Ag(S203)2]~  169, 201,  m,  give  good 
plates141** *u,  but  such  deposits  are  adherent  only  when  very  thin210.  Al- 
though the  thiourea  complex  gives  good  results115,  132,  the  bath  docs  not 
compare  favorably  with  the  cyanide  solution213. 

A  variety  of  ammines  has  been  tested.  Baths  containing  [Ag(XH3)2]+ 214 
or  the  ethylenediamine  ion,  [Ag(en)]+,  give  good  plates,  but  anode  effi- 
ciency is  poor.  Cood  deposits  are  obtained  from  baths  containing  AgCN 
dissolved  in  various  amines215;  guanidine  hydrocyanide  and  ethylenedi- 
amine hydrocyanide  give  plates  equal  to  those  from  the  cyanide  bath. 
Possibly  deposition  occurs  from  the  cyano  complex.  Plates  from  the  tri- 

207.  Jilek  and  Luk  /./.<//,  21,  576  (1927),  Mougey  and  Wirshing,  U.  S.  Patent 

_     352  I    .1   q.  1944). 

208.  Bloom, U.S.  Patent  2414438  (Jan.  1947);  cf.,  Chem.  Abe.,  41,  3383  (1947). 
_'"•..  Promise]  and  Wood,  Trans.  Electrochem.  8oe.,  80,  159    1941  . 

210.  Fleetwood  and  Yntema,  //"/.  Eng.  <')><<>,..  27,  340    r< 

211.  Morgan  and  Buretall,  "Inorganic  Chemistry      \  survey  of  .Modem  Develop- 

ate,"  p.  04,  00.  Cambridge,  England,  W.  Heffer  A:  Bone,  I 

212.  Yuzhnyi.  Khiti    Refi  "'  Zl        1,  no.  11   12,104  (1938  , cf.,  Chem.  Abe.t  33,  8506 

213.  Walter,  A. Her.  and  Riemer,  MonaUch.,  65.  .V    ! 

-Ml.  Hughes,  and  Withrow,  J.  Am.  Chem.  8oc.,  32.  1571     1910 

215.  Gilberteon  and  Mathers  foe.,  79,  139    L941). 


662  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

ethanolamine  bath  are  good,  but  those  from  guanidine  and  cyclohexyl- 
amine  solutions  are  unsatisfactory. 

Silver  salt  solutions  containing  complex-forming  organic  acids,  such  as 
tartaric,  acetic,  oxalic,  and  citric,  are  inferior  to  the  cyanide  bath141a. 
Fairly  good  deposits  of  silver  are  obtained  from  a  solution  of  silver  sulfa- 
mate  containing  a  small  amount  of  tartaric  acid102b. 

Tellurium.  Tellurium  resembles  selenium.  It  may  be  deposited  on  steel 
as  an  adherent  metal  plate  from  a  strongly  alkaline  solution  of  an  alkali 
metal  tellurite216,  or  from  a  solution  of  tellurium  dioxide  in  a  mixture  of 
sulfuric  and  hydrofluoric  acids217.  Nitric  and  hydrochloric  acids  give  inferior 
deposits.  Tellurium  may  be  separated  from  selenium  by  electrolysis  in  a 
mixture  of  hydrofluoric  and  sulfuric  acids,  in  which  the  existence  of  fluoride 
complexes,  [TeF5(H20)]~~  and  [TeF6]=,  is  probably  important. 

Tin.  Inasmuch  as  its  d  orbitals  are  full,  tin  does  not  form  inner  orbital 
complexes.  Correspondingly,  electrodeposits  appear  to  be  obtained  from 
all  water-soluble  compounds.  Reversible  deposition  would  therefore  be 
expected,  with  the  formation  of  coarse  crystals,  as  observed.  However, 
the  tendency  of  tin  to  hydrolyze  is  apparently  responsible  for  a  small  de- 
gree of  irreversibility  in  the  absence  of  halides.  Thus,  at  the  dropping 
mercury  electrode,  the  reduction  is  irreversible  unless  chloride  is  present158. 
The  fact  that  the  sulfate  bath  responds  more  readily  to  addition  agents  than 
the  chloride  bath  is  doubtless  due  to  this  irreversibility,  and  the  scarcity 
of  effective  addition  agents  even  for  the  sulfate  bath  indicates  that  the 
deposition  is  not  far  from  being  reversible.  It  may  be  assumed  that  in  the 
latter  bath,  the  tin  is  present  as  partly  hydrolyzed,  aquated  ions,  while  in 
the  chloride  bath,  the  hexachloro  ion  or  a  mixed  chloroaquo  complex, 
which  is  not  readily  hydrolyzed,  is  present.  Effective  addition  agents  for  a 
mixed  fluoride-chloride  bath  have  been  found218. 

The  commercial  sulfate  bath219  contains  tin(II)  sulfate,  sulfuric  acid, 
and  various  addition  agents.  Sometimes  sulfate  is  replaced  wholly  or  par- 
tially  by  phenolsulfonate  or  other  organic  sulfonates,  or  by  fluoroborate, 
but  this  does  not  appear  to  influence  the  cathode  reaction.  During  operation, 
the  tin (II)  ion  appears  to  hydrolyze  slowly,  probably  with  oxidation,  and 
to  precipitate.  In  conformity  with  the  near  reversibility  of  the  reduction, 
the  currenl  efficiency  approximates  100  per  cent. 

Various  complexing  agents  such  as  fluoride,  oxalate,  tartrate,  citrate, 
pyrophosphate,  cyanide,  thiosulfate,  and  hydroxylamine  have  been  investi- 
gated2*0. Good  deposits  are  obtained  over  a  narrow  currenl  density  range 

216.  Well  and  Gore,  U.  S.  Patenl  2258963  (Oct.  14,  1941). 

217.  Mathers  and  Tinner.  Trans.  .1///.  Electrochem.  Sac,  54,  293  (1928). 
218    British  Patenl  592442  (Sept.  L947). 

219.  Pine,  Trans.  Electrochem.  So,-.,  80,  631  (1941  l. 

220.  Kern.  Trans.  Am.  Electrochem.  Soc,  23,  193  (1913). 


COORDINATIOh   COMPOUNDS  I  \   ELECT RODEPOQITIOh  663 

from  tin(II)  oxalate  in  oxalic  acid80**811.  Oxidation  of  the  complex  causes 
deterioration  of  the  bath.  Similarly,  polarographic  irreversibility  is  ob- 
Berved  with  tartrate  solutions111,  which  i>  probably  due  to  hydrolysis. 
Irreversibility  in  oxidation  to  the  tin(IV)  complex  results  from  the  required 
change  in  configuration. 

Electrodeposition  from  an  alkaline  stannite  bath  Bhows  high  efficiency  as 
expected,  t>ut  the  deposits  are  spongy  or  powdery  because  the  ion  dispro- 
portionates  spontaneously  into  the  metal  and  Btannate  ion-":;.  Better  de- 
poedts  are  obtained  from  the  hexahydroxystannate  IV  ion.  This  Btannate 
hath--5  has  come  into  extensive  use  since  it  was  found  t  hal  it'  the  tin  anodes 
are  mated  by  preliminary  electrolysis  at  high  current  density,  the  products 
of  subsequent  anodic  dissolution  are  in  the  quadrivalent  state  exclusively. 
Without  this  pretreatment,  stannite  soon  appears  in  the  hath,  and  the 
deposits  become  spongy.  If  stannite  is  present,  it  may  he  oxidized  by  hydro- 
gen peroxide.  The  hath  consists  of  sodium  or  potassium  stannate  with  an 
excess  of  the  corresponding  hydroxide.  Sodium  acetate  is  sometime-  added. 
Addition  agents  are  generally  omitted,  since  none  of  them  are  very  effec- 
tive. 

Deposition  from  stannate  ion  doubtless  passes  through  the  divalent 
state,  hut  since  stannite  ions  are  reduced  and  deposited  as  fast  as  they  are 
formed,  they  do  not  accumulate  in  the  bath  with  consequent  risk  of  spongy 
deposits.  Reduction  from  tin(IY)  to  tin(II)  is  irreversible,  possibly  because 
it  requires  a  change  in  configuration;  this  may  account  for  the  rather  low 
•  athode  efficiencies.  The  two-step  reduction  is  observed  polarographic- 
ally--4A;  the  first  step  is  irreversible,  and  the  second  nearly  reversible.  The 
ineffectiveness  of  addition  agentsis  to  be  associated  with  the  reversible  na- 
ture of  the  second  step. 

Good  deposits  are  obtained  from  the  tin  (I I)  polysulfide  complex,  hut 
the  hath  is  difficult  to  maintain. 

Thallium.  As  its  electronic  structure  permits  only  outer  orbital  ions, 
thallium  is  elect rodeposited  with  very  little  activation--'"'.  Solution-  of 
thallium  (I)  sulfate,  carbonate,  or  pen-hlorate  have  been   used29       n§.  The 

221.  Qotheraall  and  Brajdah&w,  J*  Electrodepositora  Tech.  Soc. ,16,  19    1939  ;  Mathers 

and  Cockrum,  Trans.  Am.  Electrochen    S        29.  til    1916  . 

222.  Lingane,/.  An  65,  866    19 

Bidgwick,  'The  Chemical  Elements  and  Their  Compounds,"  p.  621 ,  Oxford,  1 
Oplinger  and  Bauchf  Trc       Eleei      •■       fi       80.  ♦»  1 7    1941   ;  Sternfels  and  Lou 
enheim  82.  77    1942  .  84.  I'm    194 

224a.  Lingane,  •/   A  -       67,  919    I'M.") 

ni  and  McGlynn,  Tram     I       /  53.  .;.">i  thoffand 

Lingane,  "Polarography,"  i».  260,  New  York,  [nterscience  Publishers,  [nc  , 
1941. 
226.  Sopkina,  "Chapters  in  the  Chemistry  ■  >!  the  Leaf  Familiar  Element*,"  Cham- 
paign,  111..  Stipes  Publishing  ('<>     19 


664  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

univalent  ion  has  weak  coordinating  tendencies,  but  a  few  complexes,  such 
as  KT1(CN)2,  have  been  established22511' 227.  Deposition  of  a  silver-thallium 
alloy  from  a  cyanide  bath  has  been  reported228.  The  trivalent  ion  readily 
complexes,  but  reduction  undoubtedly  proceeds  through  the  univalent 
state. 

Zinc.  As  with  other  metals  which  form  only  outer  orbital  complexes, 
zinc  appears  to  be  deposited  from  all  of  its  soluble  compounds,  although  the 
deposits  are  not  always  compact.  Deposition  from  the  aquated  ion,  as  in 
sulfate229,  chloride-acetate230,  or  fluoroborate  baths231,  is  very  nearly  reversi- 
ble. Accordingly  cathode  efficiencies  are  substantially  100  per  cent,  and 
effective  addition  agents  are  relatively  scarce. 

A  solution  of  zinc  cyanide  in  sodium  cyanide  contains  the  tetracyano 
ion,  [Zn(CN)4]=  and,  perhaps,  traces  of  the  tricyano  ion  [Zn(CN)3]_. 
White  matte  deposits  of  zinc  are  obtained  at  current  efficiencies  usually 
less  than  15  per  cent.  On  the  other  hand,  from  a  solution  of  sodium  zincate, 
in  which  zinc  probably  exists  as  [Zn(OH)4]=,  efficiencies  are  as  high  as  90 
per  cent,  but  the  deposits  are  spongy  and  poor232.  A  suitable  mixture  of  the 
two  baths  gives  excellent  deposits  at  current  efficiencies  of  80  to  90  per 
cent.  Cyanohydroxo  complexes  may  be  present. 

Deposition  from  the  cyanozincate  bath  is  irreversible.  The  reasons  for 
this  are  not  clear,  although  it  is  possible  that  the  zincate,  at  least,  may 
form  polynuclear  complexes  which  are  slow  in  de coordinating.  The  highest 
cathode  efficiencies  obtained  with  cyanide  baths  are  observed  with  gold 
and  silver,  in  which  the  ions  are  dicovalent;  the  removal  of  four  cyanide 
groups,  as  with  zinc,  may  require  a  longer  time  and  produce  a  hin- 
drance. 

Addition  agents  are  generally  employed  in  alkaline  baths.  However, 
bright  deposits  are  produced  without  addition  agents  by  "bright  dipping" 
the  plated  surface  in  dilute  nitric  or  chromic  acid,  provided  no  traces  of 
heavy  metals,  especially  lead  and  copper,  are  present  in  the  bath.  Many 
so-called  brighteners  act  only  to  remove  these  metals  as  sulfides;  soluble 
sulfides,  polysulfides,  thiosulfates,  and  thiocyanates,  have  been  used.  Op- 
eration in  the  presence  of  a  suspended  precipitate  of  zinc  sulfide  is  common. 
The  function  of  the  bright  dip  is  not  understood;  zinc  is  dissolved,  but 
without  the  usual  characteristics  of  bright  dipping233.  Often  it  appears  that 

227.  Bassett  and  Corbet,  J.  Chem.  Soc,  125,  1660  (1924). 

228.  Hensel,  Am.  Inst.  Mining  Met.  Engrs.,  Inst.  Metals  Div.,  Tech.  Pub.  No.  1930 

(1945);  cf.,  Chem.  Abs.,  40,  307  (1946). 

229.  Lyons,  Trans.  Elcctrochcm.  Soc,  80,  387  (1941). 

230.  Hogaboom,  U.  S.  Patent  2421265  (May,  1947). 

231.  Anantharaman  and  Balachandra,  ./.  Electrochcm.  Soc.,  100,  237  (1953). 

232.  Hull  and  Wernlund,  Trans.  Elcctrochcm.  Soc,  80,  407  (1941). 
Soderberg,  Trans.  Electrochem.  Soc,  88,  115  (1945). 


COORDINATION   COMPOUNDS  Xh  ELEi  TRODBPOSITION         866 

a  powder  or  61m  is  removed  from  the  surface.  Bright  dips  arc  also  effective 

on  deposits  from  sulfate  baths.  It  is  claimed  that   bright  dipping  1ms  a 
passivating  effect  on  the  surface. 

True  brighteners  produce  bright  deposits  without  :i  brighl  dip.  Com- 
pounds which  have  been  used  are  various  organic  resins214,  ketone* 
molybdic  oxide28*,  piperonal  or  vanillin287,  or  thiourea  with  various  met- 
als286,288.  Most  of  these  may  form  complexes  with  the  zinc,  hut  the  exist- 
ence of  the  complexes  in  the  baths  has  not  been  demonstrated.  For  proper 
operation,  the  cyanide-metal  ratio  must  he  carefully  controlled,  suggesting 
that  certain  cyano  complexes  are  necessary  for  brightening. 

Deposits  from  ammoniacal  solution,  in  which  the  tetrammine  ion, 
[Zn(XH3)4]++,  predominate-,  are  similar  to  those  from  the  sulfate  bath. 
Zinc  sulfate  in  a  hath  containing  ammonium  thiocyanate  and  ammonium 
chloride  gives  deposits  covered  with  gray  powder,  which  can  he  huffed  to 
a  bright  coat37.  Triethylamine  and  polyamines  such  as  ethylenediamine 
have  been  recommended  as  brighteners239. 

Metals  Whose  Deposition  from  Aqueous  Solution  in  Pure  Form  is 
Doubtful — Tungsten,  Molybdenum,  Tantalum,  Zirconium,  and 
Colu  mbium 

Although  thin  metallic  plates  of  tungsten  and  molybdenum  have  been 
reported460'  110bi  239-243J  it  appears  that  the  deposits  are  alloys  and  deposition 
s  as  soon  as  the  codepositing  metal  impurity  is  exhausted. 

Results  with  tantalum,  zirconium,  and  columbium  are  similar86;  deposi- 
tion of  these  metals  in  the  pure  form  has  yet  to  be  demonstrated. 

Failure  to  obtain  deposits  of  these  metals  as  well  as  others  of  the  va- 
nadium and  titanium  groups  is  not  surprising  since  all  of  their  complexes 
are  inner  orbital.  The  oxygen  complexes  in  particular  are  very  stable,  form- 
ing such  ions  as  vanadyl  and  zirconyl,  which  strongly  resist  dissociation. 
Such  complexes  will  form  in  water  solution  unless  a  still  stronger  coordi- 

234.  Henricks,  U.  S.  Patent  2101580;  2101581. 

235.  Mattacotti,  U.  S.  Patent  2109887. 

236.  Westbrook,  U.  S.  Patent  2080520. 

237.  Westbrook,  l".  S.  Patent  2218734;  2233600. 

238.  Hoff,  I".  >.  Patent  2080479;  Hull,  (".  S.  Patent  2080423. 

Bray  and  Boward,  U.  8.  Patent  2393741    .Ian.  1946  Aba.,  40.  2395 

(1946); Harford, U.S.  Patent  2384300  (Sept.  1946  ;  U.B.  Patent  2384301  (1946). 

240.  Fink  and  Jones,  Trans.  Electron  59,   161     1931    ;  Bolt,  T  ana.  El 

Soc.,  71,  301  (1937). 

241.  Glazunov  and  Jolkin,  Atti  X  cong.  intern,  chim.,  4,  363    1939  ;  cf.,  Chem.  Aba., 

34,  3184  (1940). 

242.  Price  and  Brown,  Trana.  E  .  70,  \s\    1936). 

243.  Hokhshtein,  •/.  Gen.  Chem.,  UJS.S.R.,  7,  2486    19 ...  32,2434 

1938). 


666  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Dating  agent  is  present.  In  either  case,  the  discharge  of  hydrogen  ion  will 
probably  require  less  energy  than  that  required  to  discharge  the  metal 
from  the  complex. 

The  Electrolytic  Separation  of  Metals  and  the  Deposition  of  Metal 
Alloys  from  Solutions  of  Complex  Compounds 

The  deposition  potential  of  a  metal  can  be  markedly  altered  by  complex- 
ing  the  metal  ions  in  solution  (Chapter  11).  As  is  seen  in  Table  19.1,  the 
magnitude  of  the  shift  brought  about  by  a  given  complexing  agent  is  dif- 
ferent for  different  metals;  thus  it  is  frequently  possible  to  separate  by 
complex  formation  the  deposition  potentials  of  two  metals  whose  deposi- 
tion potentials  in  simple  aqueous  solution  are  very  close  together.  This 
permits  selective  deposition  of  the  metals,  as  in  the  purification  of  metals 
and  in  quantitative  electrometric  analysis244.  Copper  or  bismuth  cannot  be 
selectively  deposited  from  a  solution  containing  the  simple  salts,  but  pure 
copper  may  be  selectively  deposited  leaving  bismuth  in  solution  if  tartrate 
is  added.  Similarly,  antimony  may  be  separated  from  many  metals  such  as 
copper,  bismuth,  or  lead,  since  antimony  deposits  with  great  difficulty  from 
aqueous  alkaline  solutions  containing  tartrate  or  fluoride.  Zinc  may  be 
separated  from  iron  by  using  a  cyanide  bath  in  which  the  iron  forms  very 
stable  inner  orbital  complexes. 

The  plating  of  metal  alloys  may  be  considered  from  an  analogous  view- 
point. The  deposition  potentials  for  each  component  of  the  alloy  should 
have  nearly  the  same  value.  Table  19.1  shows  that  the  deposition  potentials 
for  copper  and  zinc  may  be  brought  together  through  complex  formation 
with  cyanide  in  the  brass  plating  bath. 

Thermodynamically,  alloy  deposition  is  more  complicated  than  that  of 

pure  metals.  The  Nernst  equation  for  the  potential  of  a  metal  in  an  alloy 

must  include  a  term  for  its  activity  in  the  alloy  as  well  as  for  the  activity 

of  the  metal  ions  in  solution.* 

RT         (activity  metal  ions  in  solution) 

E  =  E0  —  — -  In    ; — — — ■ 

nF  (activity  metal  in  alloy) 

The  activity  of  a  metal,  "A,"  in  the  alloy  is  dependent  upon  the  type  of 
alloy  formed.  If,  in  a  binary  alloy,  the  solid  is  a  two  phase  mixture  of  crys- 
tals of  two  metals,  the  activity  of  the  metal  "A"  in  the  alloy  is  the  same  as 
that  of  the  pure  metal  alone.  If,  however,  the  alloy  is  a  single  phase  solid 
solution  of  metal  "A"  in  a  second  metal  "B,"  the  activity  of  metal  "A"  in 
the  alloy  may  be  reduced  appreciably.  Under  these  conditions  metal  may 

*  The  Nernst  equation  is  used  only  for  thermodynamic  calculations  and  does  not 
imply  any  definite  mechanism  for  the  reduction  process.  The  actual  plating  poten- 
tial will  include  another  term  which  represents  the  excess  potential  necessary  to 
keep  i  be  deposit  ion  process  going  at  an  appreciable  rate. 

JH.  Band,  "Electrochemistry  and  Electrochemical  Analysis,"  Chapt.  IV,  Vol.  II, 
London,  Blackie  and  Son,  Limited.  1940. 


COORDINATION  COMPOUNDS  IN  BLBCTRODBPOSITIOA  867 

Tabll  19.1.  The  Variation  oi  thi  Elbctbodi  Potbntialb  roi  Zinc 

\m>   COPPBB   A8    POTASSIl  M   CtANIDI    IS   ADDED    K)    mi.   BOL1   n«>\ 

KU-ctrolytr 
(i.l  Mole  Metal  Cyanide  Plus 
0  1  Mole  Metal  Sulfate 
Metal  per  Liter  0.2Jf  KCN  0.4JT  KCN  l.u.l/  KCN 

Zinc  -0.816  -1.033  -1.182  -1.231 

Copper  0.292  -0.G11  -0.964  -1.169 

be  deposited  as  an  alloy  from  a  complex  ion  which  doea  not  permit  deposi- 
tion of  the  pure  metal.  Stout  and  Faust1"1,  were  able  to  deposit  ternary  al- 
loys of  copper,  iron,  and  nickel  from  a  solution  containing  iron  as  the  com- 
plex ferricyanide,  [Fe(CX)6]-?  although  iron  will  not  deposit  as  the  pure 
metal  from  ferricyanide  solution.  Similarly,  a  tungsten-iron  alloy  has  been 
reported  from  a  solution  containing  iron  as  ferrocyanide245.  This  explanation 
applies  also  to  the  deposition  of  other  tungsten  alloys,  since  pure  tungsten 
cannot  he  deposited  from  aqueous  solution24015. 

If  the  two  metals  ''A"  and  "B"  form  an  intermetallic  compound  of  the 
type  AB:  ,  the  deposit  may  consist  of  the  pure  compound,  or  of  solid  solu- 
tions of  A  or  B  in  ABj ,  or  of  two  or  three  phase  mixtures246.  To  shift  elec- 
trode potentials,  a  single  agent  which  forms  complexes  with  all  metal  ions 
to  he  deposited  may  he  used,  as  in  the  silver-cadmium  cyanide  bath247;  or 
two  complexing  agents  may  be  selected  so  that  the  metals  are  present  in 
different  complex  ions,  as  in  a  bath  which  contains  silver  as  the  cyanide 
and  lead  as  the  complex  tartrate248.  The  latter  type  of  bath  permits  more 
or  less  independent  control  of  the  activities  of  the  two  metal  ion  compo- 
nents. For  example,  in  the  silver-lead  bath,  excess  cyanide  reduces  silver 
ion  activity  with  relatively  small  effect  on  the  activity  of  lead  ions.  Simi- 
larly, in  a  copper-tin  alloy-plating  bath,  copper  is  present  as  the  complex 
cyanide,  and  tin  as  the  stannate  ion.  This  permits  control  of  copper  activity 
by  adjustment  of  cyanide  concentration,  and  of  tin  activity  by  adjustment 
of  alkali  concentration.* 

Alloy  plating  baths  are  usually  of  a  type  known  to  be  suitable  for  deposi- 
tion of  at  least  one  of  the  alloy  components.!  ( Cyanide  is  the  most  common 

*  Addition  of  alkali  has  a  secondary  effect  on  the  activity  of  copper,  hut  the  effect 
is  small  in  comparison  to  that  on  the  activity  of  tin. 

II       •  v.-r,  an  alloy  of  nickel  and  iron  may  ho  deposited  at  a  current  efficiency  of 
nearlvGO  percent  from  a  hath  containing  iron  as  potassium  ferrocyaiiide,  I\.[Fc<  ( !N)«], 
and  nickel  as  potassium  nickelocy aside,  KJXi  CN   ,,"-'''  though  neither  pure  iron 
nor  pure  nickel  can  be  plated  readily  from  cyanide  solution. 
245.  Berghaus,  Germai  Apr.  14, 1939  em.  .1/-* .,  St,  4886    LI) 

Allmand  and  Ellingham,  "The  Principles  of  Applied  Electrochemistry,"  p.  128, 
New  York.  Longmans,  Green  and  Co.,  1   - 
247.  Faust,  Henry,  and  France.  Trans,  1  72,  179    1 

Joe.,  75,  186    r« 
Stout  and  Carol,  Trans.  Am.  Electrochem.  8a     68,357    !■ 


<.<is  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

complexing  agent;  however,  tartrate,  oxalate,  thiocyanate,  amines,  etc. 
have  been  used.  The  literature  on  alloy  plating  is  summarized  by  Faust250. 

Cyanide  Solutions  for  Alloy  Plating.  Many  binary  and  ternary  alloys 
have  been  deposited  from  cyanide  solutions.  Binary  alloys  include:  alloys 
of  copper  with  zinc251,  nickel,  iron,  cadmium,  tin,  gold,  and  other  metals; 
alloys  of  silver  with  metals  such  as  cadmium,  indium,  palladium,  nickel, 
lead,  and  thallium;  and  alloys  of  gold  with  nickel,  tin,  etc.250b.  Ternary 
alloys  include  combinations  such  as:  cadmium-zinc-mercury252,  copper- 
cobalt-tin253,  copper-nickel-zinc254,  cadmium-zinc-antimony255,  copper-cad- 
mium-zinc256 and  copper-tin-zinc229.  Certain  of  these  metals  do  not  form 
cyanide  complexes  and  are  present  in  other  forms. 

Attempts  to  identify  ions  present  in  alloy  plating  solutions  have  not 
given  convincing  results  (for  example  see257).  However,  it  is  probable  that 
ions  such  as  [Cu(CN)3]=  and  [Zn(CN)4]=,  which  exist  in  the  copper  and 
zinc  baths,  are  present  also  in  the  brass  baths.  In  some  cyanide  solutions, 
other  complex  forming  ions  are  essential.  For  instance,  potassium  tartrate 
is  necessary  for  the  satisfactory  deposition  of  iron  from  the  iron-nickel-cop- 
per alloy  bathllb,  and  additions  of  sodium  acetate  are  recommended  for 
zinc-cadmium  alloy  plating258.  Ammonia  is  considered  by  some  to  be 
valuable  in  improving  brass  deposits. 

Solutions  for  Alloy  Plating  Which  Contain  Complexing  Agents 
Other  than  Cyanide.  A  number  of  alloys  have  been  deposited  from 
solutions  containing  complex  forming  salts  of  the  organic  acids.  A  copper- 
tin  alloy  may  be  deposited  from  a  bath  containing  the  oxalates259.  A  copper- 
zinc  alloy  is  plated  from  a  basic  solution  of  the  sulfates  and  sodium 
tartrate260.  Silver-nickel,  and  silver-lead  alloy  deposits  have  been  obtained 
from  solutions  of  tartrates,  citrates,  or  oxalates115.  An  alloy  of  tungsten  and 
nickel  is  deposited261  from  a  bath  of  sodium  tungstate,  citric  acid,  nickel 
sulfate,  and  ammonium  hydroxide  containing  ions  such  as  the  complex 

250.  Faust,  Trans.  Electrochem.  Soc,  80,  301  (1941);  78,  383  (1940). 

251.  Coates,  Trans.  Electrochem.  Soc,  80,  445  (1941). 

252.  Roberts,  U.  S.  Patent  2250842  (1941). 

253.  Sklarew  and  Cinamon,  U.  S.  Patent  2216605  (Oct.  1940). 

254.  Faust  and  Montillon,  Trans.  Electrochem.  Soc,  65,  361  (1934);  67,  281  (1935). 

255.  Stout  and  Goldstein,  Trans.  Electrochem.  Soc,  63,  99  (1933). 

256.  Ernst  and  Mann,  Trans.  Electrochem.  Soc,  61,  363  (1932). 

257.  Pan.  Trans.  Electrochem. Soc. ,82,63  (1932). 

258.  Belyaev  and  Agababov,  Korroziya  i  Borba  s.  Net,  5,  137  (1939);  cf.,  Chem.  Abs., 
36,  347  H942). 

2.7».  Bechard,  ./.  Electrodepoaiters'  Tech.  Soc,  11,  15  (1936). 

260.  Sukhodski,  Kheifetz,  and  Chapurskii,  Repts.  Central  Inst.  Metals,  Leningrad, 
no.  17,  209  (1934);  cf.,  Chem.  Abs.,  29,  5357  (1945). 

261.  Vaaler  and  Holt,  Trans.  Electrochem.  Soc,  90,  43  (1946). 


COORDINA  TION  COM  1 '01  NDS  IN  ELECTRODEPOSITION  669 

nickel  citrate,  nickel  tetrammine  ami  complex  nickel  tungstate.  Alloys  of 
tungsten  with  iron,  manganese,  and  silver  have  been  obtained  from  similar 
baths.  Molybdenum-cobalt  and  molybdenum-iron  alloys  are  deposited262 
from  solutions  containing  tartrates,  glycols,  glycerol,  and  sugars,  which 
supposedly  form  complexes  with  iron  and  cobalt.  A  solution  containing 
both  citrate  ions  and  fluoride  ions  has  been  specified  for  plating  alloys  of 
tungsten  with  nickel,  iron,  cobalt,  and  antimony245- 263.  A  nickel-iron  alloy 
may  be  deposited  from  a  formate-sulfate  bath264. 

Alloys  of  copper-zinc,  nickel-cobalt,  nickel-iron,  cadmium-zinc,  karat 
gold,  cadmium-silver,  copper-tin,  and  silver-lead  are  deposited  on  the 
commercial  scale.  Certain  alloy  deposits  may  be  obtained  from  solutions 
containing  the  aquated  ions.  Thus,  zinc-cadmium  alloys  are  deposited265 
from  solutions  of  the  sulfates,  and  lead-tin  alloys  from  mixtures  of  the 
fluoroborates266. 

Deposits  from  nickel  and  tin  chloride-fluoride  solutions  contain  50  atom- 
per  cent  of  each  metal  over  a  considerable  range  of  nickel-tin  ratios  in  the 
bath267.  This  has  been  interpreted  to  indicate  that  deposition  occurs  from 
a  double  fluoride  complex  containing  an  atom  of  each  metal;  the  existence 
of  such  a  complex  has  been  demonstrated  by  application  of  the  method  of 
continuous  variations  (p.  569)267A.  Similar  deposits  are  said  to  have  resulted 
from  an  acetate  bath.  The  concept  is  interesting,  but  constant  composition 
deposits  have  not  as  yet  been  reported  for  other  baths. 

ELECTRODEPOSITION    FROM   NONAQUEOUS   SOLUTIONS 

The  use  of  nonaqueous  solvents  in  metal  deposition  was  reviewed  by 
Audrieth  and  Xelson268  in  1931.  Anhydrous  liquid  ammonia,  the  nitrogen 
prototype  of  water,  has  been  studied  as  a  solvent  by  a  number  of  work- 
ers26s,  269  Copper,  silver,  gold,  beryllium,  zinc,  cadmium,  mercury,  thallium, 
tin,  lead,  arsenic,  chromium,  manganese,  iron,  nickel,  cobalt,  palladium 
and  platinum,  can  be  plated  from  liquid  ammonia  solution,  but  attempts 
to  deposit  aluminum,  thorium,  bismuth,    antimony,  molybdenum,  and 

262.  Yntema,  U.  S.  Patent  2428404  (Oct.  7, 1947). 

263.  Berghaus,  German  Patent  674430  (Apr.  14, 1939);  cf.,  Chem.  Abs.,  33,  4886  (1939). 

264.  Kersten  and  Young,  Ind.  Eng.  Chem.,  28,  1176  (1936). 

265.  Fink  and  Young,  Trans.  Electrochem.  Soc,  67,  131  (1935). 

266.  Blum  and  Haring,  Trans.  Electrochem.  Soc,  40,  147  (1921);  Carlson  and  Kane, 

Monthly  Rev.  Am.  Electro-platers'  Soc,  33,  255  (1946). 

267.  Cuthbertson,  Parkinson,  and  Rooksby,  J.  Electrochem.  Soc,  100,  107  (1953); 

Rooksby,  ./.  Electrodepositors'  Tech.  Soc,  27,  129  (1951). 
267A.  Ran,  thesis,  University  of  Illinois,  1955. 

268.  Audrieth  and  Xelson,  Chem.  Revs.,  8,  335  (1931). 

269.  Audrieth  and  Yntema,  /.  Phys.  Chem.,  34,  929  (1930);  Booth  and  M<rlub-Sobel, 

./.  Phys.  Chem.,  35,  3303  (1931);  Taft  and  Barham,  ./.  Phys.  Chem.,  34,  929 
(1930). 


(170  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

tungsten  were  unsuccessful.  Beryllium  is  of  particular  interest  since  it 
cannot  be  plated  from  aqueous  solution. 

From  a  solution  of  their  salts  in  formamide  or  acetamide,  lead,  copper, 
zinc,  tin,  thallium,  cadmium,  nickel,  and  cobalt  have  been  deposited270, 271. 
Iron  and  metals  above  zinc  in  the  electromotive  series  could  not  be  de- 
posited. An  alloy  of  aluminum  and  iron  has  been  plated  from  formamide272. 
Cathode  current-voltage  curves  for  metal  deposition  in  formamide  and 
pyridine  have  been  reported66.  Pyridine  as  a  solvent  permits  the  deposition 
of  silver,  magnesium,  calcium,  zinc,  copper,  iron,  potassium,  sodium,  and 
lithium.  No  plate  was  obtained  with  beryllium268. 

Miscellaneous  organic  solvents  from  which  metals  have  been  deposited 
include  glacial  acetic  acid273,  acetone268,  ether127b- 274,  ethyl  bromide  and 
benzene  mixture272b,  substituted  benzenes275,  and  phosphorous  oxychlo- 
ride276.  Of  particular  interest  is  an  ethyl  bromide-benzene  bath  containing 
dissolved  metallic  aluminum  and  a  small  amount  of  aluminum  bromide. 
Aluminum  was  deposited  from  this  bath  at  a  current  efficiency  of  60  per 
cent272b. 

In  general,  the  salts  which  are  most  soluble  in  a  variety  of  solvents  are 
the  nitrates,  bromides,  iodides,  thiocyanates,  and  cyanides268.  Nonaqueous 
baths  resemble  aqueous  baths  in  that  small  amounts  of  addition  agents, 
temperature,  and  current  density  are  of  major  importance  in  determining 
the  type  of  plate  obtained. 

Solvents  such  as  liquid  ammonia,  liquid  hydrogen  cyanide,  glacial  acetic 
acid,  anhydrous  amines,  ether,  and  acetone  are  of  particular  importance 
in  studying  the  electrochemistry  of  coordination  compounds.  Since  ions  in 
solution  are  always  solvated,*  the  so-called  simple  ions  in  water  are  com- 
plexes of  the  type  [M(H20)x]H~;  in  liquid  ammonia  the  "simple"  ions  are 
metal  ammines  of  the  form  [M(NH3)J1/+;  and  in  liquid  hydrogen  cyanide 
they  are  probably  complexes  of  the  type  [M(CN)J2-.  An  aquated  ion  in 
liquid  ammonia  is  a  complex  ion  just  as  metal  ammines  are  complexes  in 
water.  Thus,  the  distinction  between  simple  and  complex  ions  is  entirely 

*  Energy  considerations  do  not  permit  the  existence  of  the  unsolvated  simple  ion 
M+  in  the  body  of  the  solution,  though  it  may  be  adsorbed  on  the  electrode  surface. 

270.  Rohler,  Z.  Elcktrochem.,  16,  419  (1910). 

271.  Yntema  and  Audrieth,  J.  Am.  ('hem.  Soc,  52,  2693  (1930). 

272.  Blue  and  Mathers,  Trans.  Electrochem.  Soc,  63,  231  (1933);  65,  339  (1934). 

273.  Stillwell  and  Audrieth,  ./.  .1///.  Chem.  Soc,  54,  472  (1932). 

274.  Kudra  and  Klcil.s,./.  Phye.  Chem.,  U.S.S.R.,  16,  228  (1941) ;  cf.,  Chem.  Abs.,  36, 

6417  (1942). 
275    Gorenbein,    ./.    Gen    Chem.   U.S.S.R.,   8,  233    (1938);  cf.,  Chem.  Abs.  32,  5310 
L938);  Plotnikov  and  Gorenbein,  Mem.  Inst.  Chetii.,  Acad.  Sci.  Ukrain.S.S.R. 

4,  Xo.  3,249  (1937);  cf.,  Chem,  Aba.,  32, 6310  (1938). 
276.  Cady  and  Taft, ./.  Phys.  Chem.,  29,  1057,  1068  (1925). 


COOIWIXATIOX  COMI'OIXDS  IX  ELECT  HODEPOSITIOh  671 

arbitrary-'77.  Metals  have  been  deposited  from  a  variety  of  nonaqueous 

solvents  such  as  ammonia,  formamide,  sulfur  dioxide  and  acetone,  and 
little  distinction  need  be  made  between  reduction  of  solvated  and  other 
complexes.  The  potential  energy  treatment  of  Gurney  and  Fowler  (see 
page  1)34)  can  be  applied  equally  well  to  all  situations. 

Since1  the  more  reactive  metals  cannot  be  deposited  from  water  solution, 
it  is  necessary  to  use  other  solvents  to  obtain  metallic  deposits.  The  deposi- 
tion potential  of  the  metal  must  not  exceed  the  reduction  potential  of  the 
solvent.  A  number  of  the  more  reactive  metals  can  be  deposited  from 
nonaqueous  solvents  which  are  very  weak  Bronsted-Lowry  acids.  For  ex- 
ample, beryllium  is  deposited  from  solutions  of  its  salts  in  anhydrous  liquid 
ammonia8Wb,  and  aluminum  alloys  can  be  plated  from  solutions  containing 
aluminum  chloride  and  another  appropriate  metal  salt  such  as  iron(III) 
chloride  in  anhydrous  formamide272.  There  is  no  general  correlation  between 
metal  activity  and  the  deposition  of  the  metal  from  basic  solvents.  Beryl- 
lium can  be  deposited  from  liquid  ammonia,  and  the  active  alkali  metals 
can  be  reduced  to  give  their  characteristic  blue  solution  in  liquid  ammonia, 
but  much  more  noble  metals  such  as  aluminum,  magnesium,  antimony 
and  bismuth  cannot  be  deposited  or  reduced  in  liquid  ammonia  solution. 

277.  Densham,  Trans.  Faraday  Soc,  33,  1513  (1937). 


2\J.  The  Use  of  Coordination  Compounds 
in  Analytical  Chemistry 

James  V.  Quagliano 

Notre  Dame  University,  Notre  Dame,  Indiana 

and 

Donald  H.  Wilkins 

University  of  Illinois,  Urbana,  Illinois 

When  a  metal  ion  becomes  part  of  a  complex,  it  achieves  new  properties 
which  may  be  strikingly  different  from  those  of  the  original  ion.  Such 
changes  include  those  in  color,  stability  toward  oxidation  or  reduction, 


magnitude  of  ionic  charge  (frequently  even  a  change  in  sign),  and  solubili- 
ties and  crystalline  form  of  the  salts.  These  new  properties  used  in  the  iden- 
tification or  determination  of  either  the  metallic  ion  or  the  coordinating 
agent  illustrate  applications  of  complexing  to  analytical  chemistry.  Such 
applications  to  qualitative  analysis  are  found  in  the  dissolution  of  silver 
chloride  in  ammonium  hydroxide  and  in  the  generation  of  a  red  color 
when  iron  (III)  ion  is  treated  with  thiocyanate.  In  quantitative  analysis, 
coordination  compounds  are  widely  used  in  gravimetric,  volumetric  and  col- 
orimetric  determinations,  as  well  as  in  polarimetry  and  microscopy.  In  the 
broadest  sense,  any  analysis  carried  out  in  solution  might  be  considered 
to  involve  coordination,  for  "the  chemistry  of  solutions  is  the  chemistry  of 
complexes." 

Applications  to  Precipitation  Methods 

Insoluble  Inner  Complexes 

Inner  complexes  often  have  properties  useful  in  analysis  and  remarkably 
different  from  the  ions  from  which  they  are  generated.  These  complexes 
were  formerly  called  "inner  complex  salts,"  but  the  term  is  a  misnomer,  for 
they  arc  qoI  sails;  their  usefulness  in  analytical  chemistry  depends  largely 
upon  their  nonsalt-like  character.  An  inner  complex  is  a  completely  chelated, 
nonionic  structure,  formed,  usually,  by  the  union  of  a  metal  ion  with  a 
bidentate  group  which  has  a  charge  of  minus  one.  Obviously,  for  such  a 

672 


COORDINATION  COMPOUNDS  IN  ANALYTICAL  CHEMISTRY     673 

group  to  form  an  inner  complex,  the  coordination  number  of  the  metal  ion 
must  be  twice  its  ionic  charge;  this  is  frequently,  but  not  always,  the  case. 

Inner  complexes  containing  beryllium,  aluminum,  cobalt  (III),  iron(III), 
and  chromium(III)  are  common;  those  containing  cobalt(II)  or  iron(II) 
are  rare  because  the  usual  coordination  number  of  these  ions  is  six;  they 
could  form  inner  complexes  by  union  with  a  tridentate  ligand  of  uninegative 
charge. 

'1  ne  value  of  inner  complexes  in  analytical  chemistry  rests  largely  upon 
three  properties: 

(1)  Many  of  them  are  insoluble  in  aqueous  media,  but  maybe  extracted 
into  organic  solvents  immiscible  with  water,  thereby  permitting  a  separa- 
tion of  certain  ions  from  a  large  volume  of  aqueous  solution  into  a  small 
volume  of  organic  solvent.  The  extractability  is  often  a  function  of  the  pH 
of  the  aqueous  phase,  so  that  selective  extraction  and  subsequent  return 
to  a  new  aqueous  phase  are  possible(p.  44). 

Solubility  characteristics  of  inner  complexes  may  be  quite  different  if  the 
organic  coordinating  agent  contains  a  functional  group  of  such  a  nature, 
or  in  such  a  position,  that  it  cannot  take  part  in  coordination.  For  example, 
the  zinc  derivative  of  8-hydroxyquinoline  is  quantitatively  insoluble  in 
water,  whereas  the  zinc  compound  of  5-sulfo-8-hydroxyquinoline  is  readily 
soluble. 


Strictly  speaking,  such  substances  are  not  true  inner  complexes,  for  they 
give  ions  in  aqueous  solution;  however,  they  are  often  referred  to  as  inner 
complexes  because  the  coordinate  bonds  about  the  metal  ion  are  the  same 
as  in  the  derivatives  of  the  ligands  which  do  not  contain  solubilizing  groups. 

(2)  The  formation  of  inner  complexes  is  sometimes  accompanied  by  pro- 
nounced color  changes  which  permit  colorimetric  measurements.  This 
development  of  color  is  striking,  but  it  is  by  no  means  as  general  as  many 
chemists  suppose. 

(3)  The  metal  ion  to  be  determined  is  often  a  part  of  a  complex  of  high 
molecular  weight;  this  gives  a  favorable  conversion  factor. 

Correlation  of  structures  of  organic  coordinating  agents  with  structures 
pecific  metal  ions  with  which  they  react  is  largely  empirical.  The  rela- 
tionships are  doubtless  very  complex,  involving  not  only  the  varying  nature 
of  the  bond  between  the  metal  and  the  ligand,  but  also  steric  factors  and 
Bolubilities. 

Dioximes.  The  use  of  inner  complexes  in  analytical  chemistry  began  with 


674 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Tschugaeff1,  who  discovered  that  biacetyldioxime  (dimethylglyoxime)  re- 
acts with  nickel  ion  to  give  an  insoluble  red  compound.  This  reaction  has 
been  extensively  studied,  for  it  furnishes  a  very  sensitive  and  specific 
method  for  the  determination  of  nickel  by  direct  weighing. 
In  general,  compounds  designated  by  the  formula 

R—  C=NOH 
R— C=NOH 

(in  which  R  represents  an  aliphatic,  aromatic,  or  heterocyclic  group)  pre- 
cipitate comparable  red  compounds,  so  the  functional  group 

— C=NOH 


— C=NOH 

is  apparently  responsible  for  the  reaction.  The  dimethyl  compound  is  the 
best  known  and  most  widely  used  glyoxime,  but  several  other  members  of 
the  series  possess  distinct  advantages  over  it. 

Furildioxime2,     1 ,2-cyclohexanedionedioxime3,     and    1 , 2-cycloheptane- 
dionedioxime4 


ou. 


H    H 


NOH 


0"c 


=  NOH 


are  all  more  soluble  in  water  than  dimethylglyoxime  and  give  more  favor- 
able conversion  factors.  The  nickel  derivative  of  the  cycloheptane  com- 
pound, moreover,  may  be  precipitated  from  slightly  acid  solution. 

Diaminoglyoxime5,  H2N — C=NOH,  can  be  used  in  place  of  the  dimethyl 

I 
H2N— C=NOH 

analog,  but  replacement  of  the  NH2  group  by  NH2CONH —  so  increases 

the  acidity  of  the  molecule  that  it  acts  as  a  dibasic  acid6,  and  in  ammoniacal 

solution  produces  a  precipitate  of  the  formula 


NH2— CO— NH— C=NO 


NH2— CO— NH— C=NO 


/ 


NH3 


Xi 


NH:i 


1.  Tschugaeff,  Z.  anorg.  Chem.,4A,  144  (1905). 

2.  Soule,  •/.  .1///.  Chem.  Soc,  47,  981  (1925) 
:;.  Wallack,  Ann.,  437,  148,  175  (1924). 

I.  Voter  and  Banks,  Anal  Chem.,  21,  1320  (1949). 

5.  Chatterjee,  •/.  Indian  Chem.  Soc.,  16,  608  (1938). 

6.  Fei^l  and  Christian]  Kronwald,  Z.  anal.  Chem.,  65,  341  (1924). 


COORDINATION  COMPOUNDS  IN  ANALYTICAL  CHEMISTRY     675 

The  dioximes  also  yield  (yellow)  precipitates  with  palladium  Baits,  hut 
not  with  the  ions  of  any  other  metals.  The  palladium(II )  derivative  of 
dimethylglyoxime  La  insoluble  in  dilute  mineral  acid  solutions,  whereas  the 
nickel  compound  must  be  precipitated  in  a  buffered  acetate  or  amnion iacal 
medium;  the  ortho-dioxime  group  may  thus  be  considered  specific  for  both 
palladium(II)  and  nickel (II)  ions.  Palladium (II)  dimethylglyoxime,  unlike 
the  nickel(II)  derivative,  is  soluble  without  decomposition  in  solutions  <>i 
alkali  hydroxide7  to  form  the  ion 


O        O 


CH3— C=N 


/ 


Pd 


N=C— CH5 


/ 
/ 
CH3— C=X  X=C— CH3 

\  / 

o      o 


The  specificity  of  the  ortho-dioxime  group  toward  nickel  and  palladium 
vanishes  when  the  oxime  groups  are  attached  to  an  unsaturated  ring.  Thus, 
a,j8-naphthquinonedioxime  and  orthoquinonedioxime  act  as  dibasic  acids 
and  precipitate  many  metal  ions  from  neutral  solutions8. 

The  symmetrical  dioximes  exist  in  three  isomeric  forms: 


R— C— C— R 

II       II 
X     X 

/     \ 

0                     OH 

11- 

-c— 

II 
X 

1 

OH 

—  C- 

II 
X 

1 
HO 

-R 

R— C— C— R 

II       II 
X     N 

/         1 
HO         HO 

Ami 

Syn- 

Amphi- 

Of  the  three  isomers  of  biacetyldioxime,  only  the  anti-isomer  forms  the 
characteristic  red,  insoluble  nickel(II)  compound;  the  syn-isomer  is  in- 
capable of  reacting  with  metallic  salts,  and  the  amphi-isomer  gives  a  yellow 
or  green-yellow  compound  in  which  one  molecule  of  dioxime  is  combined 
with  one  nickel  ion,  the  hydrogens  of  both  oxime  groups  being  replaced  by 
the  metal6-  9. 
Following  the  demonstration  of  the  existence  of  two  tautomeric  form-  of 


Feigl  and  Suter,  J    <  .,  1948,  378. 

Feigl,  Ind.  Eng.  Chem.,  Ann!.  Ed.,  8,  401  (1936). 

Tschugaeff,  Ber.,  39,  3382  (1906);  41,  1678,  2219  (1908);  ./.  Chem.  Soc.,  105,  2187 
1914);  Atack,  ./.  Chem.  Soe.,  103,  1317  (1913);  Pfeiffer,  Ber.,  63,  1811  (1930  j 

Hieber  and  Leutert,  Ber.,  60,  2296,  2310  (1927);  Tschugaeff  and  Lebedinski, 

Z.  anorg.  Chem.,  83,  1  (1913). 


G76 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


1  he  oxime  group1 


and 


OH 


/ 


II 


Pfeiffer9e-  n  proposed  that  the  nitrone  form  is  involved  in  the  formation  of 
the  nickel  derivative,  which  then  contains  nickel-nitrogen  bonds  in  five- 
membered  rings. 

OH  O 


R— C=N 


N=C—  R 


Ni 


R— C=N  N=C— R 

I  I 

O  OH 


From  the  facts  that  the  anti-isomer  of  a-benzilmonoxime  (I),  the  mono- 
ethers  of  a-benzildioxime  (II),  and  a-benzilmonoximeimine  (III)  form  red 
precipitates  with  nickel (II),  Pfeiffer  inferred  that  the  nickel  ion  is  bonded 
to  the  nitrogen  atom  of  the  dioxime  group  rather  than  to  the  oxygen  atom. 


— C=0 


<~> 


-C=NOR 


-C=NOH 


-C=NH 


C=NOH 


C=NOH 


(I) 


(ID 


(III) 


Brady  and  Meurs12  have  proposed  the  following  formula  for  the  nickel 
derivative  of  biacetyldioxime : 


H3C- 


C C- 

II         II 
N        N, 


CH. 


0'  \    /    O 

/  Ni 

H         /     \        /H 

O-N       N-0 

it        II 

H3C-C— C-CH. 


The  postulated  hydrogen  bonding  eliminates  the  possibility  of  cis-trans 

K).  Brady  and  Mehta,  ./.  Chem.  Soc,  125,  2297  (1924). 
1 1 .  Pfeiffer  and  Richarz,  Ber.,  61,  103  (1928). 
1l\  Brady  and  Meurs,  J.  Chem.  Soc,  1930,  1599. 


COORDIX ATIOX  COMPOUNDS  IN  ANALYTICAL  CHEMISTRY     677 

isomerism  and  also  explains  the  lark  of  reactivity  of  the  hydroxyl  group. 

The  nitrogen-nickel  bonds  are  eovalent  and  planar  and  two  isomeric  nickel 
derivatives  of  unsymmetrically  substituted  dioximes  correspond  to  cis  and 
trans  configuration-'  . 


II  II  2^   5 

N  N 

o'  W   xo 

<        /\       > 

II         II 

H3C-C C-CH^CgHs 

CIS— 


C^I5H2C_C  —  C  — CH, 
II        II 
N        N 

P       Ni        O 
H         /\  H 

N0-N         N-O'' 
II  II 

H^C-  C         C— ChUCgHe 

TRANS  — 


Similar  isomerism  exists  in  the  case  of  the  palladium  derivative11.  Isomerism 
of  several  sorts  may  be  found  in  complexes,  and  since  the  isomers  may  differ 
in  color  and  in  solubility,  their  existence  is  of  great  analytical  interest. 

In  the  determination  of  nickel  with  dimethylglyoxime,  the  precipitate 
may  be  dried  and  weighed,  or  redissolved  and  titrated.  In  acid  solution,  it 
hydrolyzes  to  hydroxylamine,  which  can  be  titrated  with  a  bromate- 
bromide  mixture,  or  oxidized  by  iron(III)  ion,  the  resulting  iron(II)  ion 
being  titrated14. 

An  interesting  application  of  the  reaction  between  biacetyldioxime  and 
nickel  ion  is  found  in  the  determination  of  biacetyl,  (CH3CO)2 ,  in  butter 
and  other  natural  products.  The  biacetyl  is  converted  to  the  oxime  and 
precipitated15. 

8-IIydroxyquinoline  and  Derivatives.  In  8-hydroxyquinoline  and  its 
derivatives,  the  hydroxyl  and  heterocyclic  nitrogen  combine  with  metal 
ions  to  form  chelate  rings. 


8-Hydroxyquinoline  has  been  used  in  the  determination  and  detection  of 
over  thirty  elements16.  Attempts  have  been  made  to  increase  the  selectivity 

-  igden,  J.  CI         9foc.,  1932,  246;  Cavell  and  Sugden, ./.  Chew.  80c.,  1935,  621. 

14.  Tougarinoff,  Ann.  soc.  sci.  Bruxelles,  54B,  314  (1934  . 

16.  Barnicoat,  Analyst,  90,  053    I 

16.  Berg,  "Die  Chemische  Analyse/1  2nd  ed.,  Vol.  34,  Enke,  Stuttgart,  1938;  "Or- 
ganic Reagents  for  Metals."  lib  ed.,  London.  Sopkin  and  Williams,  1943; 
Y<h-  and  Sarver:  "Organic  Analytical  Reagents,"  New  York,  John  Wiley  & 
Sons,  Inc.,  1945. 


678  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

and  sensitivity  of  these  reactions  by  the  use  of  derivatives17,  and  by  varia- 
tions in  the  pH  of  the  solutions18. 

8-Hydroxyquinaldine19  (2-methyl-8-hydroxyquinoline)  is  a  useful 


CH3 
OH 

derivative  of  8-hydroxyquinoline,  but  the  methyl  group  in  the  2-position 
appears  to  limit  the  number  of  ions  with  which  it  will  react.  In  particular, 
8-hydroxyquinaldine  does  not  precipitate  aluminum  from  acetic  acid  solu- 
tions buffered  with  acetate,  whereas  8-hydroxyquinoline  gives  quantitative 
precipitation20. 

Many  techniques  have  been  devised  for  the  termination  of  analyses  in- 
volving 8-hydroxyquinoline  and  its  derivatives.  The  usual  methods  involve 
weighing  the  precipitate  directly  or  igniting  it  to  the  oxide,  but  sometimes 
it  is  more  convenient  to  redissolve  the  precipitate  and  titrate.  8-Hydroxy- 
quinoline  precipitates  are  conveniently  titrated  either  by  oxidation  or  by 
bromination.  For  example,  the  8-hydroxyquinoline  may  be  oxidized  by  an 
excess  of  hexanitratocerate(IV),  the  excess  being  back  titrated  with  ox- 
alate21. The  reaction  is  not  strictly  stoichiometric,  but  a  reproducible  em- 
pirical factor  may  be  determined.  The  bromination  technique,  using  stand- 
ard br ornate,  is  extremely  sensitive. 

Hydroxyoximes.  The  hydroxy oxime  grouping  is  found  in  salicylal- 
doxime,  2-hydroxy-4-methoxyacetophenoneoxime,  2-hydroxy-5-methoxy- 
acetophenoneoxime,  and  o-vanillinoxime.  With  copper,  it  forms  salts  in 

17.  Holland,  Compt.  rend.,  210,  144  (1940);  Fresenins,  Fischbach,  and  Frommes, 

Z.  anal.  Chem.,  96,  433  (1934);  Berg,  Z.  anorg.  allgem.  Chew.,  204,  208  (1932), 
Boyd,  Degering,  and  Shreve,  Ind.  Eng.  Chem.,  Anal.  Ed.,  10,  606  (1938);  Wen- 
ger,  Duckart,  and  Rieth,  Helv.  chim.  Acta,  25,  406  (1942);  Gutzeit  and 
Monnier,  Helv.  chim.  Acta,  16,  478,  485  (1933). 

18.  Moyer  and  Remington,  Ind.  Eng.  Chem.,  Anal.  Ed.,  10,  212  (1938);  Soto,  J.  Chem. 

Soc,  Japan,  54,  725  (1933);  56,  314  (1935);  Fleck  and  Ward,  Analyst  58,  388 
(1933);  62,  378  (1937) ;  Marsson  and  Hasee,  Chem.  Ztg.,  52,  993  (1928);  Halber- 
stadt,  Compt.  rend.,  205,  987  (1937). 

19.  Doebner  and  Miller,  Ber.,  17,  1698  (1884);  Merritt  and  Walker,  Ind.  Eng.  Chen.. 

Anal.  Ed.,  16,  387  (1944). 

20.  Merritt,  Record  Chem.  Progr.,  10,  No.  2,  59  (1949). 

21.  Nielson,  Ind.  Eng.  Chem.,  Anal.  Ed.,  11,  649  (1939);  Gerber,  Claassen,  andBoruff, 

Ind.  Eng.  Chem.,  Anal.  Ed.,  14,  658  (1942). 


COORDINATION  COMPOUNDS  IN  ANALYTICAL  CHEMISTRY     679 

which  the  phenolic  hydrogen  is  assumed  to  be  replaced  with  the  formation 
of  an  inner  complex. 

VC  =  N'0H   0-c' 

-c'        V       > 

C—  O  N  =  C 

HO 
The  reactions  of  the  isomeric  methyl  ethers  of  salicylaldoxime  support  this. 


H 
0=N— 0- 

-CH3 

and 

/\ 

H 
— C=NOH 

OH 

— 0— CH3 

The  compound  containing  the  free  phenolic  hydroxyl  group  reacts  with 
copper(II)  ion,  whereas  the  isomeric  phenolic  ether  does  not. 

The  functional  group  must  be  a  part  of  an  aromatic  system  to  react  with 
metal  ions.  Thus,  acetonylcarbinol  and  chloralacetophenone  contain  the 
characteristic  group  of  atoms  but  do  not  form  complexes  with  copper (II). 
Apparently,  an  acidic  hydrogen,  such  as  is  present  in  phenols,  is  necessary. 
Other  reagents  containing  this  functional  group  do  not  offer  any  special 
advantages  over  the  more  readily  available  salicylaldoxime.  However,  in 
some  cases,  the  metal  derivatives  are  more  intensely  colored22. 

The  acyloin  oxime  group  is  found  in  a  number  of  compounds  which 
possess  valuable  analytical  properties.  It  acts  as  a  dibasic  acid,  with  the 
oxime  group  tautomerizing  to  the  nitrone  form  under  the  influence  of  alkali: 

I 
R— C C— R' 

II  I 

O— N  O 

\     / 
Cu 

The  nature  of  the  R  and  R'  groups  has  little  effect  on  the  water-insolubility 
or  the  color  of  the  copper(II)  salt,  but  has  a  marked  effect  on  the  solubility 
in  excess  ammonia.  Feigl  believes  that,  if  the  R  and  R'  radicals  are  capable 
of  coordinating  with  the  copper  ion,  the  inner  complex  formed  is  incapable 
of  adding  ammonia  and  is  insoluble  in  aqueous  ammonia23. 

a-Benzoinoxime  exhibits  a  selective  action  in  precipitating  only  copper 
ion  from  ammoniacal  solutions.  In  acidic  solutions,  the  reagent  is  useful  for 

agg  and  Furrnan,  Ind.  Eng.  Chem.,  Anal.  Ed.,  12,  529  (1940 

Feigl  and  Bondi,  Ber.,  64,  2819  (1931). 


080 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


the  determination  of  molybdenum24  and  tungsten25,  even  though  the  pre- 
cipitate is  of  indefinite  composition  and  must  be  converted  to  some  other 
form  for  weighing. 

Nitroso  Hydroxylamines.  The  use  of  the  nitroso  hydroxylamine  group  in 
analytical  procedures  is  best  represented  by  the  extensive  use  of  cupferron 
and  neocupferron  (phenylnitrosohydroxylamine  and  naphthylnitrosohy- 
droxylamine).  Both  reagents  react  with  the  ions  of  a  large  number  of  heavy- 
metals,  forming  inner  complexes  insoluble  in  acid  solutions: 


Those  obtained  from  the  naphthyl  compound  are  less  soluble  than  the  de- 
rivatives of  the  phenyl  compound ;  this  illustrates  the  general  rule  that  an 
increase  in  molecular  weight  lowers  the  water  solubility  of  inner  complexes. 

Nitrosophenols.  Several  nitrosophenols  find  use  in  analytical  chemistry. 
It  is  interesting  that  2-nitroso-l-naphthol  is  eight  times  as  sensitive  as  the 
isomeric  l-nitroso-2-naphthol  in  the  precipitation  of  cobalt26.  Since  the 
nitrosophenol  precipitates  often  carry  reagent  with  them,  and  are  not  suffi- 
ciently stable  to  be  dried,  they  must  be  ignited  and  weighed  as  the  oxide. 
Nitrosonaphthols  are  used  primarily  for  the  determination  of  cobalt27,  but 
have  also  been  used  to  determine  iron28,  palladium29,  and  copper30.  Potas- 
sium has  been  determined  indirectly  by  precipitation  of  potassium  hexani- 
trocobaltate(III),  and  by  the  subsequent  determination  of  the  cobalt  in  the 
precipitate  with  l-nitroso-2-naphthol. 

Amino  Acids.  The  amino  acids  are  useful  reagents,  especially  for  di- 
valent elements  of  the  transition  series.  The  solubilities  of  the  metal  com- 

24.  Sterling  and  Spuhr,  Ind.  Eng.  Chem.,  Anal.  Ed.,  12,  33  (1940);  Arrington  and 

Rice,  U.  S.  Bur.  Mines,  liept.  Inv.,  1939,  3441 ;  Knowles,  /.  Research  Natl.  Bur. 
Standards,  9,  1  (1932);  Taylor-Austin,  Analyst,  62,  107  (1937);  Thompson  and 
Stott,  Foundry  Trade  J.,  123  (Aug.  23,  1934). 

25.  Steele,  Iron  Steel  Ind.,  11,  267  (1938);  Baker,  Chemist- Analyst,  30,  No.  2,  31 

I'.Ul );  Esibasi,/.  Chem.Soc. ,  Japan,  61, 125  (1940);  Yagoda  and  Falos,./.  Am, 
Chem.  Soc.,  60,  640  (1938). 

26.  Giua  and  Cherchi,  Gazz.  chim.  ita.,  49,  284  (1919). 

27.  Mayr  and  Feigl,  Z.  anal.  Chem.,  90,  15  (1932);  Clennell,  Mining  Mag.,  36,  270 

(1927);  Philippot,  Bull.  soc.  chim.,  Belg.,  44,  140  (1935);  Eder,  Chem.  Zig.,  46, 
430  (1922);  Craig  and  Cudroff,  Chemist-Analyst,  24,  No.  4,  10  (1927);  Hoffman, 
./.  Research  Natl.  Bur.  Standards,  8,  659  (1932). 

28.  Ilinski  and  Knorre,  Bcr.,  18,  2728  (1885);  Knorre,  Z.  angew.  Chem.,  1904,  641,  676; 

Jolles,Z. anal. Chem., 88,  149  (1897); Mathers,/.  Am. Chem. Soc., 30, 209  (1908) 

29.  Schmidt,  Z.  anorg.  Chem.,  80,  335  (1913). 

30.  Burgase,  Z.  angew.  Chem.,  1896,  596;  Knorre,  Ber.,  20,  283  (1887). 


COORDINATION  COMPOUNDS  I\  ANALYTICAL  CHEMISTRY      681 

plexes  are  pll  dependent ,  and  useful  separations  may  be  accomplished  by  an 
adjustment  of  the  pH  of  the  precipitating  medium.  Aromatic  liganda  are 
ordinarily  much  weaker  coordinating  agents  than  their  aliphatic  analogs; 
however,  the  favorable  disposition  of  coordinating  groups  in  anthranilic 
acid  makes  it  a  reasonably  good  complexing  agent,  and  it  forms  complexes 
suitable  for  analytical  procedures  with  cadmium81,  cobalt82,  copper82,  w,  and 

zmc31c.  34 

There  are  numerous  sulfur  compounds  applicable  to  the  formation  of 
inner  complexes  not  listed  under  the  functional  groups  mentioned.  Among 
these  are  2-benzothiazolethiol35,  2,5-dimercapto-l  ,3,4-thiodiazole36,  ru- 
beanic  acid37,  and  thiocarbanilide37e • 38. 

Complex  Ions  as  Precipitants 

Many  complex  ions  are  stable  enough  to  be  used  as  precipitants  of  ions 
to  be  detected  or  quantitatively  determined.  The  precipitates  may  often 
be  dried  and  weighed;  in  other  cases,  they  are  ignited  to  oxide,  or  redis- 
solved  and  then  determined  colorimetrically  or  by  titration.  The  ammines 
of  cobalt  and  chromium  have  received  the  most  study  as  precipitants,  but 
even  with  these,  the  field  has  hardly  been  touched. 

Complex  Cations  as  Precipitants.  When  hexamminecobalt(III)  ion  is 
added  to  neutral,  basic,  or  acidic  solutions  of  metavanadate  ion,  the  insolu- 
ble compounds,  [Co(NH8)«]  (V03)3 ,  [Co(NH3)6]4  (V207)3 ,  and  [Co(NH3)6]4- 
(V60n)3  are  formed,  respectively39.  The  yellow  precipitate  formed  in  acid 
solution  separates  vanadium  quantitatively  from  phosphate,  arsenate, 
iron(III),  copper(II),  and  calcium  ions.  Hexamminecobalt(III)  ion  may 

31.  Funk,  Z.  anal.  Chem.,  123,  241  (1942) ;  Wenger  and  Masset,  Helv.  chim.  Acta,  23, 

34  (1940);  Funk  and  Ditt,  Z.  anal.  Chem.,  91,  332  (1933). 

32.  Funk  and  Ditt,  Z.  anal.  Chem.,  93,  241  (1933) ;  Wenger,  Cimerman,  and  Corbaz, 

Mikrochemie,  27,  85  (1939). 

33.  Wenger  and  Besso,  Mikrochemie,  29,  240  (1941). 

34.  Anderson,  Ind.  Eng.  Chem.,  Anal.  Ed.,  13,  367   (1941);  Caldwell  and  Mover, 

./.  Am.  Chem.  Soc,  57,  2372  (1935);  Cimerman  and  Wenger,  Mikrochemie,  18, 
53  (1935) ;  Wenger,  Helv.  chim.  Acta,  25,  1499  (1942);  Mayr,  Z.  anal.  Chem.,  92, 
166  (1933). 

35.  Spacu  and  Kuras,  /.  prakt.  ('hem.,  144,  106  (1935);  Dubsky,  Mikrochemie,  28,  145 

(1940);  Spacu  and  Kuras,  Z.  anal.  Chem.,  102,  24     1935). 

36.  Dubsky,  Okac,  and  Trtilek.  Mikrochemie,  17,  332  (1935);  Ray  and  Gupta,  ./. 

Indian  Chem.  Soc,  12,  308  (1935). 

37.  Ray,  Z.  anal.  Chem..  79,  94  (1929),  Wolbling  and  Steiger,  Mikrochemie,  15,  295 

(1934);  Feigl  and  Kapulitzas,  Microchemie,  8,  239  (1930);  Center  and  Macin- 
tosh, Ind.  Eng.  Chem.,  Anal.  JM.,17,239    1930  ; Wolbling,  Ber.,e7,773  (1934 
Wohler  and  Mets,  Z.  anorg.  allgem.  Chem.,  138,  368    L924  I;  Singleton,  Ind.  Chem 
Ut,  3,  121  (1927 
39.  Parks  and  Prebluda. ./.  Am.  Chem. Soc., 07 ,  1676  (1935). 


682  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

also  be  used  as  a  precipitant  for  the  quantitative  determination  of  ferro- 
cyanide  ion40  in  the  absence  of  chromate,  dichromate,  and  vanadate  ions. 

The  nitratopentamminecobalt(III)  ion,  [Co(NH3)5N03]++,  has  been  em- 
ployed in  the  determination  of  semi-micro  quantities  of  phosphates41.  The 
insoluble,  high  molecular  weight  complex  compound,  [Co(NH3)5N03]- 
[H3PM012O41],  has  the  advantage  of  a  favorable  conversion  factor  and 
avoids  the  post-precipitation  and  occlusion  phenomena  which  are  so  trouble- 
some with  ammonium  molybdate.  It  is  interesting  to  note  that  the  complex 
cations,  [Co(NH3)6]+++  and  [Co(NH3)5Cl]++  failed  to  give  satisfactory  pre- 
cipitates in  this  procedure.  Inconsistent  results  were  obtained  in  attempts 
to  use  the  complex  ammines  in  the  determination  of  germanates  and  ar- 
senates. 

Frequently,  metal  cations  can  be  converted  to  anions  and  precipitated 
by  the  addition  of  complex  cations.  Thus,  after  the  addition  of  excess 
iodide,  bismuth  may  be  precipitated  as  the  orange-yellow  complex,  trans- 
[Co  en2  (SCN)2][BiI4]42.  Bismuth  can  be  determined  also  by  precipitation  of 
[Cr(NH3)6][BiCl6],  which  is  then  analysed  by  ammonia  distillation43.  Simi- 
larly, antimony44  can  be  precipitated  and  weighed  as  the  stable  and  very 
insoluble  chromium  compound  [Cr  en3][SbS4].  For  the  determination  of 
semi-micro  quantities  of  antimony,  the  method  is  more  rapid  and  conveni- 
ent than  the  usual  method  of  weighing  as  antimony(III)  sulfide.  Feigl  and 
Miranda45  used  the  tris(o:,a,-dipyridyl)iron(II)  ion  for  the  detection  of 
complex  anions  which  have  large  atomic  volumes,  such  as  [Cdl4]=,  [Hgl4]=, 
[Co(CN)6]s,  and  [Ni(CN)4]=.  The  similar  tris(orthophenanthroline)iron(II) 
ion  is  also  useful  for  the  precipitation  of  these  anions. 

Complex  Anions  as  Precipitants.  Complex  anions  can,  of  course,  be 
used  as  precipitants,  too,  as  is  illustrated  by  the  well-known  determination 
of  ammonium  and  potassium  ions46  by  the  precipitation  of  their  chloroplati- 
nates.  Potassium  is  also  determined  by  precipitation  of  K2Na[Co(N02)e]47 
or  the  still  less  soluble  salt  K2Ag[Co(N02)6]48. 

Cadmium  can  be  separated  from  zinc  and  determined  quantitatively 
by    precipitation    as    the    insoluble    thiourea    complex    [Cd(thiourea)2] 

40.  Hynes,  Malko,  and  Yanowski,  Ind.  Eng.  Chem.,  Anal.  Ed.,  8,  356  (1936). 

41.  Furman  and  State,  Ind.  Eng.  Chem.,  Anal.  Ed.,  8,  420  (1936). 

42.  Spacu  and  Spacu,  Z.  anal.  Chem.,  93,  260  (1933). 

43.  Mahr,  Z.  anal  Chem.,  93,  433  (1933). 

44.  Spacu  and  Pop,  Z.  anal.  Chem.,  Ill,  254  (1938) ;  Spacu  and  Pop,  Mikrochemie  ver. 

Mikrochim.  Acta,  3,  27  (1938). 

45.  Feigl  and  Miranda,  Ind.  Eng.  Chem.,  Anal.  Ed.,  16,  141  (1944). 

46.  Tenery  and  Anderson,  /.  Biol.  Chem.,  135,  659  (1940) ;  Salit,  J.  Biol.  Chem.,  136, 

191  (1940). 

47.  Snell  and  Snell,  "Colorimetric  Methods  of  Analysis,"  New  York,  D.  Van  Nos- 

trand,  1936. 

48.  Burgess  and  Kamm,  ./.  Am.  Chem.  Soc.,  34,  652  (1912). 


COORDINATION  COMPOX  VDS  l\     INALYTICAL  CHEMISTRY     683 

[Cr(SCN)4]249.  The  greal  variety  of  coordinating  agents  thai  can  be  used  to 
alter  the  properties  of  the  ion  to  be  determined  and  the  tremendous  array 
of  complex  ions  that  can  be  used  as  precipitants  makes  the  oumber  of  com- 
binations almost  without  limit. 

Applications  to  Volumetric  Analysis 

The   phenomena   of   coordination   find   wide  application    in    volumetric 

analysis,  both  in  the  use  of  complexing  ligands  and  in  the  use  of  preformed 
complex  ions.  Coordinating  agents  are  used  to  "sequester"  or  "mask" 
interfering  ions  or  to  discharge  their  colors,  to  change  oxidation-reduction 
potentials,  and  to  alter  or  intensify  the  colors  of  ions  to  be  determined. 
Thus,  citrates,  tartrates,  malates,  and  other  organic  hydroxy  anions,  which 
can  form  five-  of  six-membered  chelate  rings,  are  used  to  prevent  the  pre- 
cipitation of  metallic  hydroxides  in  alkaline  solution50.  Fluoride  ion  forms 
such  stable  complexes  with  many  metallic  ions  that  the  usual  characteristic 
reactions  of  the  simple  ions  no  longer  appear;  e.g.,  the  reaction  of  the  fluoride 
ion  with  iron(III)  ion  forms  the  colorless,  soluble  hexafluoroferrate(III) 
ion,  which  is  so  stable  that  copper  can  be  determined  iodometrically  in  its 
presence51.  The  addition  of  excess  fluoride  ion  to  a  solution  of  an  iron  salt 
lowers  the  oxidation  potential  of  the  iron(II)-iron(III)  system  sufficiently 
to  make  possible  the  use  of  diphenylamine  as  an  indicator  in  the  titration 
of  iron(II)  with  dichromate52.  Phosphate  ion  also  reacts  with  iron(III)  ion 
to  form  a  colorless,  soluble  complex  and  is  used  frequently  instead  of  fluo- 
ride53, as  in  the  well-known  iron-permanganate  titration. 

Titration  of  Liberated  Hydrogen  Ion 

In  general,  the  formation  of  an  inner  complex  from  a  salt  and  an  organic 
substance  liberates  an  equivalent  quantity  of  hydrogen  ion;  the  metal  can 
be  determined  by  titration  of  the  liberated  hydrogen  ion.  This  is  illustrated 
in  the  volumetric  determination  of  nickel54.  Obviously,  such  a  method  can 
be  used  only  with  organic  substances  which  do  not  themselves  liberate 
protons,  except  when  coordinated  with  metal  ions. 

Fit  ration  of  Metal  Ions  with  a  Complexing  Agent 

When  the  complex  ion,  formed  between  a  metal  ion  and  a  donor  mole- 
cule, is  sufficiently  stable,  i.e.,  Kd  is  a  small  Dumber,  it  may  be  possible,  by 
use  of  a  suitable  indicator  system,  to  titrate  the  metal  ion  with  the  complex- 

19.  M.-.hr  and  Ohle,  Z.  anal.  Chem.,  109,  1  (1937). 

GO.  Willanl  and  Young,  •/.  .1///.  Chem.  Soc.,  50,  1322,  1334,  1368    1928 

51.  Park,  Ind.  Eng.  Chem.,  Anal.  Ed.,  3,  77  (1931 

52.  SzebeUedy,  Z.  anal.  Chem.,  81, 97    1930 
khollenberger,  ./.  .1/.-.  Chem.  Soc.,  53,  88    L931 

54.  Bolluta,  Monaiseh.,  40,  281  (1919 


684  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

ing  agent.  Chelating  agents,  especially  those  containing  enough  donor  atoms 
within  one  molecule  to  saturate  the  coordination  sphere  of  the  metal  ion, 
are  more  generally  useful  in  this  technique,  since  monodentate  donor  species 
commonly  undergo  stepwise  reaction  with  the  metal  ion,  with  the  result 
that  a  plot  of  concentration  of  uncombined  metal  ion  against  moles  of  com- 
plexing  agent  gives  no  sharp  break55.  However,  a  number  of  well-known 
determinations  are  based  on  titration  of  metal  ions  with  monodentate 
donors.  The  cyanometric  titrations  of  nickel  and  cobalt  ions  serve  to  illus- 
trate this  point. 

Indicator  Systems.  Although  a  number  of  indicator  systems  can  be  devised 
for  determinations  of  this  type,  only  two  have  found  extensive  use.  The 
first  is  the  pH  indicator.  Since  complexing  agents  are  basic  substances 
(amines,  anions  of  weak  acids,  and  the  like)  the  first  addition  of  excess 
complexing  agent  is  accompanied  by  a  rapid  rise  in  pH.  This  principle  has 
been  applied  to  the  determination  of  a  wide  variety  of  metal  ions  by  titra- 
tion with  the  anions  of  ammonia  triacetic  acid56,  uramil  diacetic  acid57,  and 
ethylenediaminetetraacetic  acid58. 

The  second  technique  involves  tying  up  the  metal  ion  in  a  colored  com- 
plex of  lesser  stability  than  that  formed  between  the  metal  ion  and  the 
complexing  agent  which  serves  as  the  titrant.  The  success  of  this  method 
depends  on  a  sharp  color  change  accompanying  the  destruction  of  the 
indicator  complex.  Since  donor  molecules  or  ions  which  undergo  color 
change  upon  reaction  with  a  metal  ion  are  also  color  sensitive  toward  hy- 
drogen ion,  titrations  of  this  type  are  carried  out  in  buffered  solutions. 
Schwarzenbach  and  coworkers59  have  employed  purpureate  ion  (murexide) 
in  the  formation  of  indicator  complexes  with  calcium,  magnesium,  cadmium, 
zinc,  and  copper. 

H  /       •  °Ns  H 

/N_cx  /C_N 

0  =  C  C  =  N — C  C=0 

H       %         -o'       H 

MUREXIDE 

These  investigators60  have  also  used  0,0 '-dihydroxyazo  dyes  in  indicator 

55.  Schwarzenbach,  Chimin,  3,  1  (1949);  Anal.  chim.  Acta,  7,  141  (1952). 

56.  Schwarzenbach  and  Biedermann,  Helv.  chim.  Acta,  31,  331  (1948). 

57.  Schwarzenbach  and  Biedermann,  Ihlv.  chim.  Acta,  31,  456  (1948). 

58.  Schwarzenbach  and  Biedermann,  Helv.  chim.  Acta,  31,  459  (1948). 

59.  Schwarzenbach,  Biedermann,  and  Bangerter,  Helv.  chim.  Acta,  29,  811  (1946); 

Schwarzenbach  and  Gyeling,  Helv.  chim   Acta.,  32,  1108,  1314  (1949). 

60.  Schwarzenbach  and  Biedermann,  Helv.  chim.  Acta, 31, 678  (1948);  Schwarzenbach 

and  Biedermann,  chimin.  2.  56  (1948). 


I  OORDINATWX  coMI'oi  \DS  IX  AXALYTICAL  CHEMISTRl       685 

complexes  in  the  titration  of  magnesium,  calcium,  zinc,  and  cadmium  with 
disodium  ethylenediaminetetraacetic  acid.  This  indicator  technique  lias 
been  used  most  extensively  in  the  determination  of  water  hardness 

Polydentate  complexing  agents  have  also  been  utilized  in  procedures 
based  on  amperometric  titrations551'"61,  polarimetric  titrations'1',  poten- 
tiometric  titrations61,  and  spectrophotometric  titrations84. 

Applications  of  the  Technique.  The  most  widely  employed  complex- 
ing agent,  ethylenediaminetetraacetic  acid,  is  quite  nonspecific  in  its  action, 
and  may  he  applied  to  the  analysis  of  the  alkaline  earth  ions,  almost  all  of 
the  dipositive  and  tripositive  transition  element  ions,  and  the  metallic  ions 
of  periodic  groups  IB,  IIB,  and  IIIB,  as  well  as  to  the  analysis  of  lead  and 
bismuth.  Specificity  of  the  reagent  for  the  alkaline  earth  ions  or  for  lead  or 
bismuth  ions  may  be  attained  by  masking  the  transition  metal  ions  with 
cyanide65.  This  masking  technique  has  been  applied  to  the  determination 
of  the  calcium  ion  content  of  mineral  waters  containing  large  amounts  of 
copper,  cobalt,  zinc,  nickel,  and  iron  salts66. 

It  is  also  possible  to  determine  a  particular  ion  selectively  by  the  proper 
choice  of  an  indicator  complex  and  by  adjusting  the  pH  of  the  medium. 
Thus,  it  is  possible  to  determine  the  total  hardness  of  water  (magnesium 
and  calcium  ion)  by  titrating  an  aliquot  with  ethylenediaminetetraacetic 
acid,  using  Eriochrome  Black  T  as  the  indicator,  at  a  pH  of  about  1059a-  60b, 


ERIOCHROME    BLACK     T 

and  then  to  determine  the  calcium  ion  independently  by  titrating  a  second 
aliquot  of  the  sample  with  the  same  reagent  in  strongly  alkaline  solution, 
using  murexide  as  the  indicator5911  ■  67. 
This  technique  has  also  been  applied  to  the  determination  of  magnesium 

61.  Pribil  and  Matyska,  Collection  Czechoslov.  Chem.  Communs.,  16,  139  (1051). 

62.  Pribil  and  Matyska,  Chem.  Listy,  44,  305    1950  . 

Halm.  .1//.//.  chim.  Acta,  4,  583    I960);  Pribil,  Koudela,  and  Mat -ska.  Collection 
Czechoslov.  Chem.  Commune.,  16,  BO     L951);  Pribil  and  Malicky,  Collection 
Czechoslov.  Chem.  Commune.,  14,   H3    L949  ;  Pribil  and  Horacek,  Collection 
1  h,  m.  Commune.,  14,  626    1949  . 
64.  Sweetster  and  Bricker,  Anal.  Chem.,  25,  253    L953 
I  laschka  and  Huditz,  Z.  anal.  Chem.,  137,  172    1952 
Botha  and  Webb,  ./.  Inst.  Watet  Engrs.,  6,  159    1942 
Cheng,  Kurtz,  and  Bray,  Anal.  Chem.,2it  1640    1955 


G8() 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


ill  plant  materials68,  the  estimation  of  the  effectiveness  of  polyphosphates 
in  sequestering  calcium  ions69,  and  to  a  number  of  microdeterminations70. 

Complex  Ions  as  Oxidation -Reduction  Indicators 

If  a  complexing  agent  gives  stable,  highly  colored  complexes  with  a 
metal  in  two  different  oxidation  states,  and  if  the  oxidation-reduction  po- 
tential of  the  resulting  couple  is  suitable,  the  couple  can  be  used  as  an  oxi- 
dation-reduction indicator.  Such  cases  are  rare,  but  the  1 ,  10-phenanthro- 
line-iron  and  ruthenium  complexes  furnish  interesting  and  important 
examples.  Tris-(orthophenanthroline)iron(II)  ion  (ferroin) 

+  + 


is  an  intense  red  color  and  the  corresponding  iron(III)  ion  is  a  faint  blue. 
The  reaction*  is  reversible;  both  complexes  are  stable  in  acid  media,  and 
the  system  has  a  high  oxidation-reduction  potential  (1.10  volts  in  OAF 


[Fe(C12H8N2): 


+  e" 


[Fe(C12H8N2): 


acid)71.  The  potential  may  be  varied  to  suit  the  requirements  of  the  analysis 
by  placing  substituents  in  various  positions  in  the  organic  rings.  The  change 
in  potential  brought  about  by  the  substitution  of  methyl  groups  for  hydro- 
gen atoms  has  been  found  to  be  an  additive  function,  so  if  the  oxidation 
potentials  for  the  complexes  with  methyl  groups  in  the  3,  4,  or  5  positions 
are  known,  the  potential  for  any  combination  of  methyl  substitutions  can 
be  calculated71.  As  Table  20.1  shows,  methyl  groups  in  the  3  or  8  positions 
lower  the  value  by  0.03  volt;  in  the  5  or  6  positions,  lower  it  by  0.04  volt; 
and  in  the  4  or  7  positions,  lower  it  by  0.11  volt.  Substitution  of  a  nitro 
group  in  position  5  of  1 ,  10-phenanthroline  changes  the  oxidation-reduction 
potential  of  the  couple  to  1.25  volts,  which  makes  this  couple  an  excellent 
indicator  for  cerate  oxidimetry72. 

Sec  footnote,  page  399,  Chapter  11  for  discussion  of  sign  conventions  used;  the 
convention  adopted  by  polarographers  is  used  in  this  chapter. 

68.  Forster,  Analyst,  78,  17!)  (1953). 

69.  Kurias,  Textil  Rundschau,  5,  224  (1950);  cf.  Chem.  Abs.,  44,  8824e  (1950). 

70.  Flaschka,  Mikrochemie  ver.  Mikrochim.,  Acta,  39,  38  (1952);  Debney,  Nature, 

169,  ll()t  (1952);  Flaschka,  Mikrochemie  vcr  Mikrochim.  Acta,  39,  315  (1952). 

71.  Brandt  and  Smith.  Anal  Chem., 21,  1313  (1949). 

72.  Salomon,  Gabrio,  ;md  Smith,  Arch.  Biochem.,  11,  433  (1946);  Smith  and  Frit/. 

Anal  chem.,  20,  S74  (1946). 


COORDINATION*   COMPOl  NDS  IN  ANALYTICAL  CHEMISTRY     687 


Table  20.1.  Effect  of  hhb   l\TK"iiicrin\  01  Methyl  Gboups  om  the  Redox 
Potentials  of  the  1,10  Phenanthboline  [bom  Couple 


Methyl  Substituted  Derivative 

unsubstituted 

3 
4 

5 
3,  1 

::.  8 
4,5 

\,  (i 

*j  " 

5,6 
3,  4,  6 
3,  4.  7 
3,5 
3,  5,  8 
3,  4,  6,  7 
3,  4,  6,  8 
3,  4,  7,  8 
3,  4,  6,  8 


Redoi  Potential,  Found 

1.10 

1.07 


1.06 
0.97 
1.03 
0.95 
0.95 
0.88 
1.00 
0.92 
0.88 
0.93 
0.99 
0.84 
0.89 
0.85 
0.93 


Volts  in  0.1/'' acid,  Calc. 


0.99 

0.96 
1.04 
0.95 
0.95 
0.88 
1.02 
0.92 
0.85 
0.92 
0.99 
0.81 
0.89 
0.82 
0.96 


Several  modifications  of  the  1 ,  10-phenanthroline-ruthenium  structure 
have  been  studied,  but  none  has  come  into  use  as  an  indicator.  Dwyer73  in- 
vestigated the  ruthenium(II)-ruthenium(III)  couple  and  found  it  to  have 
an  oxidation-reduction  potential  of  — 1.29  volts  in  IN  nitric  acid.  Cagle  and 
Smith  studied  the  use  of  tris(a,a'-dipyridyl)iron(II)  ion  and  its  methyl 
derivatives  and  found  them  to  be  suitable  as  indicators  in  the  determina- 
tion of  iron74. 


Stiegman  and  his  co-workers  found  the  oxidation  potential  of  the  ruthe- 
nium(II)-ruthenium(III)  dipyridyl  system  to  be  1.33  volts  in  IN  nitric 
acid75.  Brant  and  Smith,  however,  report  that  this  value  is  1.25  volts71. 
The  change  in  the  redox  potential  of  metallic  couples  by  coordination  is 
well  known,  and  i>  discussed  in  Chapter  11.  Many  applications  of  the 
phenomenon  have  been  made  in  analytical  chemist  ry. 

'      Dwyer,  Humpoletz,  and  Xyholin,  ./.  Proc.  Roy.  Soc.  N.  S.  Wales,  80,  212  (1946). 
:\.  Cagle  and  Smith,./.  Am.  Chem.  .w.,69,  I860  (1947);  Ind.  Eng.  Chem.,  Anal.  Ed., 

19,384    I 
75.  Steigman,  Biernbaum,  and  Edmonds,  Ind.  En  /.  Chi  m.t  Anal.  E*l  ,  14,  30    1942). 


688  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

The    Application    of    Coordination     Compounds     to    Colorimetric 
Methods  of  Analysis 

The  display  of  colors  shown  by  coordination  compounds  of  the  transition 
metals  is  utilized  in  the  manufacture  of  pigments  (Chapter  22)  and  in  analy- 
sis by  colorimetric  methods.  The  familiar  qualitative  tests  for  iron76  and 
cobalt77  with  thiocyanate  depend  upon  this  property;  the  thiocyanate  group 
probably  coordinates  through  the  nitrogen  atom. 

At  low  concentrations  of  thiocyanate,  iron  forms  the  deep  red  complex 
[Fe(NCS)]++.  At  higher  concentrations,  other  red  complexes  of  the  type 
[Fe(NCS)n]3-"  are  formed,  where  n  may  be  any  interger  from  one  to  six78. 
Partition  studies  with  the  solvents  ether  and  water79,  thermometric  titra- 
tions80, and  spectrophotometric  studies78  indicate  that  these  species  are  in 
stepwise  equilibrium.  The  complexes  are  stable  in  strongly  acidic  solutions, 
and  so  have  a  definite  advantage  over  other,  more  sensitive  reagents. 

Cobalt(II)  ion  reacts  with  thiocyanate  in  aqueous  solutions  containing 
alcohol  or  acetone,  producing  complex  species  which  may  be  extracted  into 
ether-alcohol  solutions770.  Spectrophotometric  studies  indicate  that  the 
cobalt  ion  reacts  stepwise,  forming  a  series  of  complexes  of  the  formula 
[Co(NCS)  J  2_n,  where  n  is  an  interger  between  one  and  four,  inclusive81. 
The  intense  blue  coloration  developed  in  ether-alcohol  solutions  has  been 
variously  attributed  to  dehydration  effects  of  the  solvent82  and  to  a  change 
in  the  coordination  number  of  the  cobalt  ion81b. 

The  use  of  complexing  agents  in  quantitative  colorimetric  analyses  is 
well  illustrated  by  the  application  of  1 ,  10-phenthroline,  a,a:'-dipyridyl, 
and  aja/ja/'-tripyridyl. 


NN"/     \N"/     \N/ 
a,a',o:'/-tripypidyl 

1 ,  10-Phenanthroline  has  been  applied  to  the  colorimetric  determination  of 

76.  Frank  and  Oswalt,  J .  Am.  Chem.  Soc,  69,  1321  (1947);  Woods  and  Mellon,  Ind. 

Eng.  Chem.,  Anal.  Ed.,  13,  551  (1941);  Peters  and  French,  Ind.  Eng.  Chem., 
Anal.  Ed.,  13,  604  (1941). 

77.  Tomula,  Z.  anal.  Chem.,  83,  6  (1931);  Uri,  Analyst,  72,  478  (1947);  Young  and 

Hall,  Ind.  Eng.  Chem.,  Anal.  Ed.,  18,  264  (1946). 

78.  Babko,  /.  Gen.  Chem.,  U.S.S.R.,1B,1549  (1946);  cf.  Chem.  Abs.,  41,  4732e  (1947). 

79.  Macdonald,  Mitchell,  and  Mitchell,  J.  Chem.  Soc.,  1961,  157  1 

80.  Chatterjee,  Science  and  Culture,  15,  209  (1949). 

81.  Katzin  and  Gebert,  /.  Am.  Chem.  Soc,  72,  5659  (1950);  Lehne,  Bull.  soc.  chi?n., 

France,  1951,  76;  Babko  and  Drako,  ./.  Gen.  Chem.  U.S.S.R.,  19,  1809  (1949); 
cf.  Chem.  Abe.,  44,  1355  (1950);  Babko  and  Drako,  Zavodskaya  Lab.,  16,  1162 
(1960);  cf.  Chem.  Abe.,  47,  3175  (1951). 

82.  West  and  De  Vries,  Anal.  Chem.,  23,  334  (1951). 


COORDIXATIUX  COMI  -  IN  ANALYTICAL  CHEMISTRY     I  - 

iron  in  fruit  and  wine*,  in  leather**,  and  in  biological  materials*5.  A  large 
number  of  modified  1 .  10-phenanthroline  derivatives  have  been  studied  in 
recent  years.  The  substitution  of  methyl  groups  for  hydrogen  ate:   -    s 
additive  function  with  regard  to  the  wave  length  of  maximum  absorption 
and  molecular  extinction  coefficient.  This  relationship  was  discovered  for 
iron^IP  complexes  by  Brandt  and  Smith71  and  for  copper  by  McCurdb 
A  most  sensitive  colorimetric  reagent  for  iron  is  4. 7-diphenyl-l .  10-phe- 
nanthroline: the  iron vIT^  complex  has  a  molecular  extinction  coefficier. * 
22, 400.  In  addition  to  iron  II  .   abaft  II  .  molybdenum^,  ruthenium  II  . 
and  copper  I    g         olored  solutions  with  this  reagent.  However,  the 
per(D  complex  does  not  form  at  a  pH  less  than  7^f.  The  iron  II    complex 
may  be  extracted  into  solvents  such  as  isoamyl  alcohol  over  a  pH  range  of 
'2  to  9,  whereas  the  cobalt  (JD  complex  is  not  extractable. 

ivl  has  found  use  in  colorimetric  methods  oi  analysis,  the 
complexes  being  very  similar  to  those  of  1 . 10-phenanthroline.  but  not  so 
stable-7.  Several  substituted  dipyridyls  do  not  give  colored  complexes  with 
ironvll)  ion*  Perhaps  this  can  be  explained  on  steric  grounds  in  the  follow  - 

-    B 


N  N 


0OO 


but  in  the  case  of 


HOOC  COOH 

it  must  be  attributed  to  a  lessening  of  tl.        -       y  oi  the  nitrogen  at 
:vl  has  been  applied  to  the  spectrophotometry 

'.!  and  Cruess.  I  ml.  :  I    - 

S      Smith  and  I  tt,  195  (194? 

K     Willard  and  Hummel.  7  10,13(1938 

Curdy,  thesis,  University  of  Ulinou 
•- 

B7    Blau.  MomaU         19,      ~      9W  .d  Griffin.  Can.  J  171 

H    Willink  and  Wibsttt  M,  271 


690  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

mination  of  iron89  and  cobalt(II)90.  The  reagent  is  not  particularly  sensitive 
for  cobalt91  but  it  may  be  used  over  a  wide  range  of  concentrations  (0.5  to 
50  ppm)  and  the  cobalt  complex  is  stable  over  the  pH  range  2  to  10.  Sur- 
prisingly, copper(I)  ion  does  not  form  a  colored  complex  with  tripyridyl, 
although  it  does  so  with  a,a'-dipyridyl92  and  with  1 ,  10-phenanthroline93. 
The  stereochemistry  of  tripyridyl  is  evidently  such  that  it  is  not  as  strong 
a  coordinating  agent  as  its  bidentate  analogs.  Morgan  and  Burstall94  have 
isolated  and  characterized  a  number  of  complexes  of  a,a:'-dipyridyl, 
a,a',a"-tripyridyl,  and  a:,a:',a!",a:'"-tetrapyridyl;  tripyridyl  occupies  three 
coordination  positions  in  compounds  of  the  type,  [Pt  tripy  Cl3]Cl94e'  90b. 
It  is  probably  significant  in  the  chemistries  of  these  higher  polypyridyls 
that  they  tend  to  form  bridged  structures  and  in  so  doing  enter  into  the 
coordination  spheres  of  two  metal  atoms  simultaneously. 

Hoste95  pointed  out  the  specificity  of  6, 6 '-substituted  dipyridyls  for 
copper(I),  stating  that  2,2'-biquinoline  forms  the  most  stable  complexes 
of  this  series.  Indeed,  copper(I)  is  almost  unique  among  the  metal  ions  in 
forming  colored  complexes  with  compounds  such  as  2,2/-biquinoline96, 


\/\N/     \N/\/ 


OjG'-dimethyl^^-bipyridine,  and  2, 9-dimethyl-l,  10-phenanthroline71.  Of 
the  many  substituted  derivatives  available,  2 , 9-dimethyl-4 , 7-diphenyl- 
1 , 10-phenanthroline  is  the  most  sensitive  reagent  for  copper  now  avail- 
able97. Copper(I)  ion  reacts  with  biquinoline,  whereas  iron(II)  ion  does  not. 
8-Hydroxyquinoline  is  also  used  for  the  colorimetric  determination  of 
many  metallic  ions.  Alten,  Weiland,  and  Loofman98  coupled  the  hydroxy- 
quinolate  of  aluminum,  in  the  precipitate,  with  a  diazo  compound  to  obtain 
a  strongly  colored  dye,  Avhich  was  then  compared  with  a  standard. 

89.  Moss  and  Mellon,  Ind.  Eng.  Chem.,  Anal.  Ed.,  14,  862  (1942). 

90.  Moss  and  Mellon,  Ind.  Eng.  Chem.,  Anal.  Ed.,  14,  931  (1942) ;  Morgan  and  Bur- 

stall,  /.  Chem.  Soc,  140,  1649  (1937);  135,  20  (1932). 

91.  Moss  and  Mellon,  Ind.  Eng.  Chem.,  Anal.  Ed.,  15,  74  (1943). 

92.  Ignatieff,  J.  Soc.  Chem.  Ind.,  56,  407t  (1937);  Gerber,  Claassen,  and  Boruff,  Ind. 

Eng.  Chem.,  Anal.  Ed.,  14,  364  (1942). 

93.  Tartarini,  Gazz.  chim.  ital.,  63,  597  (1933) ;  Wenger  andDuckert,  Helv.  chim.  Acta, 

27,  757  (1944) 

94.  Morgan  and  Burstall,  J.  Chem.  Soc,  1937,  1649;  1938,  1672,  1675,  1662;  1934, 

965,  1498. 

95.  Hoste,  Anal.  chim.  Acta,  4,  23  (1950). 

96.  Breckenridge,  Lewis,  and  Quick,  Can.  J.  Research,  B17,  258  (1939). 

97.  Smith  and  Wilkins,  Anal.  Chem.,  25,  510  (1953). 

98.  Alten,  Weiland,  and  Loofmann,  Angew.  Chem.,  46,  668  (1933). 


COORDINATION  COMPOUNDS  IN  ANALYTICAL  CHEMISTRY     691 

Magnesium  and  some  other  quinolates  give  a  green  color  when  dissolved 
in  dilute  acid  and  treated  with  iron(III)  chloride911,  or  they  can  be  con- 
verted toiron(III)  quinolate,  which  is  dissolved  in  alcohol  to  give  a  green- 
black  color1"".  Alternatively,  the  quinolate  precipitates  may  be  dissolved  in 
dilute  hydrochloric  acid  and  the  absorption  of  the  solution  measured; 
8-hydroxyquinoline  absorbs  strongly  at  252  m^u.  Aluminum,  gallium,  and 
indium  hydroxyquinolates  fluoresce  strongly  in  chloroform.  Lacroix  has 
given  a  comprehensive  theoretical  treatment  of  the  equilibria  involved  in  the 
extraction  of  some  hydroxyquinolates101. 

Dithizone  (diphenylthiocarbazone) 

NHNHC6H5 

/ 

S=C 

\ 

N=N— C6H5 

is  used  primarily  in  qualitative  analysis,  particularly  in  trace  analysis102.  It 
forms  highly  colored  inner  complexes  with  a  great  number  of  metallic  ions, 
doubtless  through  chelation.  Most  of  these  inner  complexes  are  extractable 
into  carbon  tetrachloride,  but  under  proper  conditions,  separations  of  indi- 
vidual elements  can  be  made.  The  ions  which  give  colored  complexes  may 
be  divided  into  five  groups102,  103: 

(1)  Copper,  silver,  gold,  mercury,  and  palladium  ions — extractable  from 
dilute  mineral  acid  solutions. 

(2)  Zinc,  cobalt,  nickel,  palladium,  and  rather  large  quantities  of  cad- 
mium and  tin  ions — extractable  from  acetic  acid  solutions. 

(3)  Silver,  mercury,  copper,  gold,  palladium,  cobalt,  nickel,  cadmium, 
and  large  amounts  of  zinc  ions — extractable  from  sodium  hydroxide  solu- 
tion.-. 

(4)  Tin(II),  thallium(I),  bismuth,  and  lead  ions — extractable  from 
slightly  alkaline  solutions  containing  cyanide. 

(5)  Cobalt,  nickel,  and  cadmium  ions — extractable  from  strongly  alkaline 
solutions  containing  tartrate. 

Two  types  of  complexes  may  be  formed  when  a  metal  ion  combines  with 
dithizone,  a  complex  containing  the  bidentate  keto  form  of  dithizone,  in 
which  one  hydrogen  has  been  displaced  from  the  imido  group  (I),  and  a 

99.  Gerber,  Claassen,  and  Boruff,  Ind.  Eng.  Chem.,  Anal.  Ed.,  14,  658  (1942);  Weeks 
and  Todd,  Ind  Eng.  Chem.,  Anal.  Ed.,  15,297  (1943);  Wolff,  Compt.  rend.  toe. 
biol.,  127,  1445    19 

100.  Lavollay,  Bull.  sac.  chim.  Biol.,  17,  432  (1936). 

101.  Lacroix,  Anal.  chim.  Acta.,  1,  200    L947). 

I'i2.  Fischer,  11  ■  ntlich,  Si>  rru  nt  Konz<  r,  4,  158  (1925). 

103.  Fischer,  /  angew.  Chem.,  47,  685    1034  ;  Fischer  and  Leopoldi,  Z.  anal.  Chi 
107,  24]  '1936). 


692 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


complex  containing  a  tridentate  enol  form  of  the  ligand,  which  structure 
envisions  replacement  of  both  hydrogen  ions  from  the  hydrazide  function 

(II). 


s  =  c 


It  is  significant  that  those  metal  ions  which  are  good  sulfur  coordinators 
(groups  1  and  3  above)  show  the  greatest  tendency  to  form  the  "enol" 
type  of  complex  (II).  Formation  of  the  "enol"  species  takes  place  only  in 
basic  solution. 

The  selectivity  of  the  dithizone  is  generally  increased  by: 

(1)  Addition  of  complexing  reagents  to  remove  interfering  metal  ions; 
this  is  exemplified  in  group  4  above. 

(2)  Control  of  pH  of  the  solution  to  be  extracted;  compare  groups  1,  2, 
3,  and  5. 

(3)  Oxidation  or  reduction  of  interfering  metals;  platinum(II)  follows  the 
same  pattern  as  does  palladium;  however  platinum  (IV)  does  not  react  with 
dithizone104. 

Diphenylcarbazide  and  diphenylcarbazone  react  with  the  ions  of 


NH— NH— C6H£ 


0=C 


\ 


o=c 


NH— NH— C6H5 


NH— NH— C6H5 


N=N— C6H; 


heavy  metals  to  form  inner  complexes105,  which  are  extractable  into  organic 
solvents  such  as  benzene  and  chloroform.  Unipositive  copper  and  silver 
give  complexes  in  which  the  ratio  of  metal  ion  to  ligand  is  one  to  one  and  in 
which  the  ratio  is  two  to  one.  Since  these  reagents  do  not  contain  a  co- 
ordinating sulfur  atom,  they  react  with  an  entirely  different  group  of 
metallic  ions  than  does  diphenylthiocarbazone.  They  are  useful  for  the 
determination  of  chromium,  which  forms  a  soluble  red- violet  compound  in 
dilute  mineral  acid  solution,  and,  by  an  indirect  procedure,  for  the  deter- 
mination of  lead,  through  the  precipitation  of  the  chromate106. 

Colored  lakes,  even  though  they  be  insoluble  in  water,  can  often  be  used 
in  colorimetric  work  by  extracting  them  into  organic  solvents.  Chloroform 

104.  Sandell,  "Colorimetric  Determination  of  Traces  of  Metals,"  New  York,  Inter- 

science  Publishers,  Inc.,  1944. 

105.  Feigl  and  Lederer,  Monatsch.,  45,  63, 115  (1924). 

106.  Letonoff  and  Reinhold.  Tnd.  Eng.  Chem.,  Anal.  Ed.,  12,  280  (1940). 


COORDIXATI<>\    COMl'OCXDS  IX  AXALYTICAL  CHEMISTRY      ()93 


is  more  generally  useful  for  inner  complexes  than  other  organic  solvents1"7. 
By  the  introduction  of  solubilizing  groups  into  organic  molecules  which 
normally  give  insoluble  inner  complexes,  it  is  often  possible  to  obtain  ma- 
terials which  are  water  soluble  and  suitable  for  colorimetrie  determinations. 
Thus,  alizarin  sulfonic  acid  gives  a  soluble  aluminum  complex,  whereas 
alizarin  itself  gives  an  insoluble  one10S.  The  structural  formula  for  the  sul- 
fonic dye  is  probably 


so3  oh 


Similarly,  the  cobalt  (III)  compound  of  l-nitroso-2-naphthol  is  insoluble, 
but  the  disulfonic  derivative  is  soluble109. 


-,  6- 


In  these  cases,  as  in  many  others,  the  introduction  of  solubilizing  groups 
does  not  greatly  change  the  coordinating  ability. 

Quinalizarin  (1 ,2,5,8-tetrahydroxyanthroquinone)  has  been  used  for  the 


—OH 


OH   O 


colorimetrie  detection  of  germanium  and  the  rare  earths  and  for  the  deter- 
mination of  beryllium,  gallium,  magnesium,  aluminum,  and  boron11". 
Willard  and  Fogg111  have  developed  a  quantitative  procedure  for  the  deter- 

KiT.  Feigl,  Anal.  Chem.,  21,  1298    1949 

108.  Atack,  J.  Soc.  Chun.  /ml.,  34,  936    1915). 

109.  van  Klooster,  ./.  Am.  Chem.  Soc  .  43,  746    1921 

110.  Komarowsky  and  Poluektov,  Mikrochemie,  18,  66    I" 

111.  Willard  and  Fogg,  -I    Am.  Chem.  So,-..  59,  to    1937 


694  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

mination  of  gallium  based  on  the  pink  to  amethyst  color  of  the  gallium 
quinalizarin  compound. 

Boron,  a  powerful  oxygen  coordinator,  forms  an  inner  complex  with  hy- 
droxyanthraquinone  in  concentrated  sulfuric  acid  solution112,  and  mag- 
nesium, scandium,  the  rare  earths,  nickel,  cobalt,  and  beryllium  ions  give 
sensitive  reactions  with  this  reagent  in  sodium  hydroxide  solution113. 

Thiourea  and  its  derivatives  have  been  used  for  the  detection  and  deter- 
mination of  a  number  of  ions  which  are  good  sulfur  coordinators114.  Bismuth 
ion115  forms  a  yellow  compound  upon  the  addition  of  thiourea.  Mahr116  has 
proposed  a  method  for  the  determination  of  cadmium,  chromium,  and 
mercury  by  precipitation  as  the  slightly  soluble  compounds  [M  (thiourea)  2] 
[Cr(NH3)2(SCN)4]2  •  These  red  compounds  are  soluble  in  organic  ketones 
and  are  suitable  for  colorimetric  determinations. 

Storfer  has  used  thiourea  for  the  detection  of  ferricyanide  through  the 
formation  of  the  red-violet  compound  [Cu(thiourea)3]3[Fe(CN)6]  -2H20117. 
The  reagent  will  detect  0.48  mg  of  ferricyanide  at  a  dilution  of  1 :  100,000. 

Thiourea  forms  colored  compounds  which  are  suitable  for  colorimetric 
determination  of  osmium118,  ruthenium114'  115a-  119  and  other  platinum 
metals120.  The  osmium  compound  has  the  composition  [Os(NH2CSNH2)6] 
Cl3-H20.  Thiocarbanilide  reacts  with  salts  of  osmium  and  ruthenium,  both 
of  which  are  good  nitrogen  and  sulfur  coordinators119,  121.  The  resulting 
highly  colored  complexes  can  be  extracted  into  ether,  which  increases  the 
sensitivity  of  the  test  and  suggests  that  the  compounds  are  probably  inner 
complexes. 
I 

Utilization  of  Fluorescence  in   the  Application  of  Complexes   to 

Analytical  Chemistry 

The  intense  green  fluorescence  produced  by  the  addition  of  morin 
(S^V^'^'-pentahydroxyflavone)  to  a  solution  of  an  aluminum  salt122 

112.  Smith,  Analyst,  60,  735  (1935). 

113.  Fischer  and  Wernet,  Angew.  Chem.,  A60,  729  (1948). 

114.  Yoe  and  Overholser,  Ind.  Eng.  Chem.,  Anal.  Ed.,  14,  435  (1942). 

115.  Mahr,  Z.  anal.  Chem.,  94,  161  (1933);  97,  96  (1934). 

116.  Mahr,  Angew.  Chem.,  53,  257  (1940). 

117.  Storfer,  Mikrochemie,  17,  170  (1935). 

118.  Tschugaev,  Compt.  rend.,  167,  235  (1918);  Gilcrist,  J.  Research,  Natl.  Bur.  Stand- 

ards, 6,  421  (1931). 

119.  Singleton,  Ind.  Chemist,  3,  121  (1927);  Wohler  and  Metz,  Z.  anorg.  allgem.  Chem., 

138,  368  (1924). 

120.  Whitmore  and  Schneider,  Mikrochemie,  17,  279  (1935). 

121.  Wolbling,  Ber.,  67,  773  (1934). 

122.  Goppelsroder,  ./.  prakt.  Chem.,  101,  408  (1867). 


COORDINATION  COMPOUNDS  1\   ANALYTICAL  CHEMISTRY     695 

OH 

HO— r//\/0x— <  >—  OH 


oil 


O 


can  he  used  for  the  determination  of  small  quantities  of  aluminum1*1.  Good 
oxygen  coordinators,  such  as  beryllium,  "allium,  indium,  and  scandium, 
also  form  complexes  which  show  a  strong  green  fluorescence47.  However, 
these  ions  are  easily  separated  from  aluminum  ion. 

l-Amino-4-hydroxyanthro(jtiinone  gives  an  intense  red  fluorescence  with 

O      MI. 


O       OH 

beryllium  ion  in  alkaline  solution  and  with  thorium  in  acid  solution.  The 
reagent  is  less  sensitive,  but  more  specific,  than  morin124. 
Benzoin 


II 


C6H; 


"C — CeH^ 


OH       O 


has  been  suggested  as  a  qualitative  reagent  for  the  fluorometric  determina- 
tion of  zinc1'23.  In  the  presence  of  magnesium  hydroxide  as  an  adsorbing 
agent,  the  reagent  is  highly  specific,  only  beryllium,  boron,  and  antimony 
interfering.  The  stability  of  the  fluorescent  material  suggests  that  the  zinc 
replaces  the  hydroxy]  hydrogen  and  forms  a  five-membered  ring  with  the 
oxygen  atoms. 
White  and  Lowe'-'  used  the  fluorescence  of  the  sodium  Ball  of  4-sulfo-2- 

hydro\y-a-naphthalene-azo-d-naphthol    (Pontachrome    Blue   Black    I(»    bli- 
the quantitative  determination  of  aluminum.  Although  not  as  sensitive  as 

_      White  and  Neustadt,  Ind.  Eng.  Chem.,  Ann!.  Ed.,  15.  599    1943 
124.  Fletcher,  White,  and  Sheftel,  Ind.  Eng.  Chem.,  Anal.  Ed.,  18,  179    1946 
125    White  and  Lowe,  Tnd.  Eng.  Chem.,  Anal.  Ed.,  9,  130    I 


696  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

morin,  this  reagent  gives  a  direct  chemical  differentiation  between  alum- 
inum and  beryllium. 

The  Role  of  Complex  Formation  in  Polarographic  Analysis 

Where  applicable,  the  polarographic  method  is  convenient  and  rapid. 
This  is  especially  true  where  several  metals  must  be  determined  simultane- 
ously or  where  simple  precipitation  or  titration  procedures  are  not  avail- 
able126. 

Complex  formation  may  be  utilized  to  provide  a  system  especially  suited 
for  the  determination  of  a  particular  substance  and  to  mask  the  effect  of 
interfering  ions.  A  third  function  of  complex  formation  is  dependent  upon 
the  stabilization  of  valence  states  of  metal  ions  through  coordination 
(Chapter  11).  This  latter  effect  often  makes  possible  the  simultaneous  po- 
larographic determination  of  two  metal  ions  whose  uncomplexed  forms 
(aquated  ions)  normally  reduce  at  potentials  which  are  too  nearly  the  same 
to  give  distinct  polarographic  curves. 

The  first  of  the  functions  of  complex  formation  mentioned  above  may  be 
illustrated  by  the  relationships  found  among  the  complex  ions  of  rhodium127. 
The  chloro  complex  of  rhodium(III)  is  reduced  to  free  metal  upon  contact 
with  elemental  mercury  while  the  rhodium(III)  complexes  with  nitrite, 
oxalate,  tartrate  and  ethylenediamine  do  not  give  polarographic  reduction 
waves,  apparently  because  of  their  great  stability.  However,  rhodium  may 
be  determined  when  present  in  the  tripositive  state  in  a  complex  ion  of 
intermediate  stability,  such  as  [Rh(NH3)5Cl]++  or  [Rh(CNS)6]= 

The  polarographic  determination  of  manganese  in  the  presence  of  copper, 
chromium,  zinc,  cobalt,  nickel,  and  iron  provides  an  excellent  example  of 
the  masking  effect  exerted  by  particular  complexing  agents  on  the  ease  of 
reduction  of  metal  ions.  If  the  sample  is  contaminated  with  iron,  copper, 
chromium,  or  zinc,  the  addition  of  cyanide128  facilitates  the  determination 
of  manganese  (so  long  as  the  iron  is  in  the  dipositive  state)  since  the  cyano 
complexes  of  these  other  metals  are  not  reduced  at  the  dropping  mercury 
elect  rode.  Similarly,  the  addition  of  tartrate  eliminates  the  reduction  waves 
of  cobalt(II),  nickel(II),  and  iron(III)129. 

The  separation  of  very  similar  half-wave  potentials  of  metals  is,  of  course, 
a  less  extreme  case  of  the  phenomenon  of  masking.  Perhaps  the  most 
interesting  examples  are  found  among  the  complexes  of  cobalt  and  nickel. 
Although  the  polarographic  reduction  curves  for  hexaquocobalt(II)  and 
hexaquoniekel(II)  ions  overlap,  the  two  metals  may  be  estimated  from  a 

126    Kolthoff  and  Lingane,  "Polarography,"  2nd  Ed.,  Vol.  II,  p.  582,  Now  York,  In- 
terscience  Publishers,  Inc.,  1952. 

127.  Reference  126,  p.  490. 

128.  Verdier,  Collection  Czechoelov.  Chem.  Commune  .  11,  238  (1939). 
L29.  Verdier,  Collection  Czechoelov.  Chem.  Commune.,  11,  233  (1939). 


COORDINATION  COMPOUNDS  IN  ANALYTICAL  CHEMISTRY     697 

single  polarogram  by  the  use  of  a  pyridine-pyridinium  chloride  supporting 
electrolyte110.  A  determination  of  cobalt  in  the  presence  of  nickel  also  in- 
vokes the  oxidation  of  the  cohalt(II)  to  the  tripositive  state  with  perborate 
in  an  ammonia-ammonium  chloride  solution131.  The  resulting  hexammine- 
eobalt(III)  ion  is  reduced  at  -.53  volts  vs.  S.C.E.,  while  the  nickel(II)- 
ammonia  complex  is  reduced  at  a  much  more  negative  potential.  A  similar 
technique  is  used  tor  the  oxidation  of  cobalt(II)  to  cobalt(III)  in  the  pres- 
ence of  ethylenediaminetetraacetic  acid132. 

Estimation  of  the  several  metals  in  an  alloy  is  among  the  most  practical 
applications  of  the  polarographic  method126.  Here,  the  ease  and  rapidity  of 
routine  analyses  are  of  considerable  value.  The  following  scheme  for  the 
determination  of  copper,  zinc,  and  nickel  will  serve  to  illustrate.  After 
dissolution  and  preliminary  treatment,  the  polarogram  obtained  from  an 
ammonia-ammonium  carbonate  solution  of  the  mixed  salts  gives  one  wave 
attributable  to  the  copper  and  a  second  which  arises  from  both  the  nickel 
and  the  zinc.  The  nickel  may  then  be  determined  from  a  second  polarogram 
run  on  a  cyanide  medium.  Neutralization  and  addition  of  cyanide  ion  to  an 
aliquot  of  the  test  solution  leads  to  the  formation  of  very  stable  zinc  and 
copper  complexes  which  are  not  reduced  polarographically.  However,  the 
nickel  cyanide  complex  gives  a  well-defined  wave. 

A  very  clever  application  of  complex  formation  in  the  polarographic 
determination  of  metal  ions  was  reported  by  Willard  and  Dean133.  The 
o  ,o'-dihydroxyazo  dye,  5-sulfo-2-hydroxy-a-benzene-azo-/3-naphthol 

OH  HO 


— X=X 


I 

SO    \;i 


is  more  difficultly  reduced  in  the  presence  of  aluminum  ion  than  in  its 
normal  state.  Apparently,  stabilization  of  the  dye  in  a  complex  of  the  type 
[Al(dye)2]  results  in  two  waves  in  the  reduction  curve  of  the  dye,  and  the 
second  wave  is  proportional,  in  height,  to  the  concentration  of  aluminum 
ion. 

The  application  of  the  polarographic  method  to  the  study  of  complex- 
ions is  further  discussed  in  Chapter  18. 

190.  Lingane  and  Kerlinger,  //"/.  Eng.  Chem.,  Anal.  Ed.,  13,  77  (1941). 

131.  Wattera  and  Kolthoff,  Anal.  Chem.,  21,  1466  (1949  , 

132.  Souchay  and  Faucherre,  Anal.  ehim.  Acta.  3,  2.52  (1949 

133.  Willard  and  Dean.  Anal.  Chem.,2*,  1264   (1950   . 


jL\.    Coordination  Compounds  in 
Natural  Products 


Gunther  L.  Eichhom 

Louis/ana  State  University,  Baton  Rouge,  Louisiana 
and  the  National  Institutes  of  Health,  Bethesda,  Maryland 

As  a  consequence  of  the  ability  of  coordinated  metal  ions  to  influence 
many  of  the  complex  reactions  upon  which  the  vital  processes  of  living 
organisms  depend,  coordination  compounds  of  many  varieties  are  found 
widely  distributed  in  nature.  A  comprehensive  coverage  of  so  vast  a  subject 
in  a  short  chapter  is  impossible;  instead,  it  is  our  purpose  to  demonstrate 
how  the  versatility  of  coordination  reactions  has  been  incorporated  into 
nature's  pattern,  to  record  some  of  the  progress  that  has  been  made  in  the 
elucidation  of  this  pattern,  and  to  illustrate  how  a  knowledge  of  coordina- 
tion chemistry  can  yield  clues  to  the  mechanisms  of  biochemical  processes 
and  thus  serve  as  a  valuable  tool  in  biochemical  research. 

Excellent  treatises  are  available  on  some  of  the  naturally  occurring  co- 
ordination compounds1;  particular  emphasis  has  therefore  been  placed 
upon  topics  not  covered  elsewhere  from  the  point  of  view  of  coordination 
chemistry,  and  an  attempt  has  been  made  to  set  the  whole  subject  matter 
into  a  context  that  provides  the  maximum  possible  opportunity  for  an 
understanding  of  the  dynamic  relationships  that  exist  between  the  natural 
coordination  compounds. 

The  Detection  of  Coordination  Compounds  in  Natural  Products 

Clues  to  the  existence  of  complex  compounds  in  nature  range  from  those 
that  offer  conclusive  proof  to  those  that  provide  circumstantial  evidence; 
they  have  been  classified  into  four  categories,  in  the  order  of  decreasing 
amount  of  available  knowledge  concerning  the  nature  of  the  compound: 

(1)  The  isolal  ion  and  determination  of  structure  of  the  coordination  com- 
pound, with  metal  ion  and  donor  molecule  intact. 

1.  Lemberg  and  Legge,  "Hematin  Compounds  and  Bile  Pigments,"  New  York, 
Intcrsciciicr  Publishers,  Inc.,  1949;  Martell,  and  Calvin,  "Chemistry  of  the 
Metal  Chelate  Compounds,"  Chapt.  8,  New  York,  Prentice-Hall,  Inc.,  1952. 


COORDIXM  I<>\   COMPOl  VDS  IN  NATURAL  PRODUCTS 


(ill!) 


700  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

(2)  The  activation  of  a  specific  biochemical  process  by  a  metal  ion.  Fre- 
quently such  a  metal  is  part  of  an  enzyme  system,  and  it  is  possible  to 
deduce  some  of  the  donor-acceptor  relationships  from  a  knowledge  of  the 
structure  of  the  coenzyme,  the  reactants,  and  the  products. 

(3)  Observations  on  the  mineral  nutritional  requirements  of  organisms. 
When  these  include,  even  in  trace  amounts,  metals  with  high  coordinating 
ability,  the  existence  of  complexes  maj^  be  suspected  and  tested  by  radio- 
active tracer  techniques  or  through  feeding  of  competing  coordinating 
agents  to  the  organism. 

(4)  Detection  of  organic  metabolic  intermediates  that  are  good  coordinat- 
ing agents  (e.g.,  compounds  containing  two  donor  groups  separated  by  two 
or  three  carbon  atoms).  Relationships  between  such  molecules  may  suggest 
the  participation  of  metal  ions  in  the  form  of  complexes. 

Functions  of  Complex  Compounds 

Some  of  the  reactions  that  are  known  or  believed  to  occur  in  plant  or 
animal  metabolism  have  been  outlined  in  Fig.  21.1;  the  names  of  coordina- 
tion compounds  have  been  capitalized,  so  that  their  omnipresence  be- 
comes evident  from  an  inspection  of  the  chart.  Dotted  lines  have  been 
used  to  represent  functional  relationships  between  two  compounds. 


For  the  benefit  of  the  inorganic  chemist  unversed  in  the  biochemical  literature  a 
brief  explanation  of  this  chart  will  be  presented. 

It  may  be  observed  that  many  of  the  capitalized  compounds  are  catalytic  agents. 
Biochemical  catalysts  are  called  enzymes;  they  are  generally  involved  in  the  chemical 
transformation  of  a  specific  compound  or  group  of  compounds.  The  latter  are  known 
as  the  substrates  of  the  enzyme.  An  enzyme  is  frequently  named  by  addition  of  the 
suffix  "ase"  to  the  name  of  its  substrate.  Enzymes  generally  consist  of  a  protein  por- 
tion that  accounts  for  the  bulk  of  the  weight  of  the  molecule,  and  of  a  non-protein 
part,  the  "prosthetic  group"  of  the  enzyme.  Coenzymes  are  relatively  simple  organic 
molecules  in  whose  absence  the  enzyme  cannot  function.  The  distinction  between  co- 
enzymes and  prosthetic  groups  is  not  clear-cut;  the  most  important  difference  prob- 
ably lies  in  the  greater  ease  with  which  the  former  may  be  detached  from  the  protein 
component  of  the  enzyme.  Man}'  enzymes  are  coordination  compounds;  frequently 
the  donor  groups  are  contributed  by  the  prosthetic  group  or  by  the  coenzyme. 

The  reactions  outlined  in  the  chart  include  processes  that  occur  in  plant  or  animal 
metabolism;  a  large  number  of  them  are  found  in  both  types  of  organism.  Whereas 
plants  are  capable  of  producing  the  matter  required  for  their  structure  and  mainte- 
nance from  simple  compounds  through  photosynthesis,  animals  ingest  rather  com- 
plicated "food"  materials:  the  proteins,  fats,  and  carbohydrates,  which  are  con- 
densation products  of  amino  acids,  fatty  acids,  and  monosaccharides,  respectively. 

Proteins,  fats,  and  carbohydrates  are  related  to  each  other  in  both  plant  and 
animal  metabolism,  because  they  can  be  broken  down  into  simple  substances,  which 
may,  in  turn,  be  condensed  into  the  appropriate  large  molecules  as  required  by  the 
organism. 


COORDINATION  COMPOUNDS  IN  NATURAL  PRODUCTS  701 

To  illustrate  this  relationship  let  us  consider  the  molecule  of  pyruvic  acid,  which 
i-  centrally  located  on  the  chart .  This  molecule  may  be  produced  by  "transamination" 

(page  712)   from  alanine,  one  of  the  amino  acids  formed  by  the  degradation  of  pro 
teins.  The  chain  of  events  by  w  hich  proteins  (upper  left  I  arc  decomposed  is  init  iated 

by  the  so-called  endopeptidases  (page  703),  which  split  the  protein  into  relatively 

large  fragments,  the  polypept  ides.  The  degradat  ion  process  is  1  hen  I  aken  over  by  t  be 

exopeptidases  which  have  been  so  named  because  they  remove  the  terminal  amino 

acids,  the  acids  on  the  "outside"  of  the  polypeptide  chain.  Some  of  t  hese  amino  acids 
are  molecules  of  alanine,  which  may  then  be  converted  into  pyruvic  acid. 

Pyruvic  acid  may  result  from  the  metabolism  of  fats  (lower  left)  in  the  following 
manner:  Fat  degradation  ultimately  yields  acetic  acid,  or  acetate  ion,  winch  may  be 
converted  into  an  extremely  reactive  form,  acetyl  coenzyme  A.  When  the  acetyl  group 
is  so  combined,  it  may  react  with  carbon  dioxide  to  produce  pyruvic  acid,  or  it  may 
pass  through  the  tricarboxylic  acid  cycle  (see  below)  to  oxaloacetic  acid,  which  is 
converted  to  pyruvic  acid  through  decarboxylation. 

The  formation  of  pyruvic  acid  from  carbohydrates  may  be  followed  in  the  lower 
right  section  of  the  chart. 

A  group  of  substances  that  play  a  central  role  in  biochemistry  are  the  compounds 
of  the  tricarboxylic  acid  cycle.  The  cycle  consists  of  the  removal  of  citric  acid  in  ;i 
-cries  of  reactions,  in  which  two  of  its  carbon  atoms  are  lost  by  decarboxylation  (in 
the  presence  of  carboxylase  enzymes),  and  the  replenishment  of  citric  acid  through 
the  condensation  of  the  enol  form  of  oxaloacetic  acid  with  acetyl  coenzyme  A  in  the 
piesence  of  "condensing  enzyme"  (page  711).  The  cycle  may  be  followed  in  the  center 
of  the  chart . 

In  the  upper  right  section  are  outlined  some  of  the  enzymatic  reactions  which  re- 
sult in  the  oxidation  of  the  decomposition  products  of  the  amino  acids  into  quinoid 
compounds.  Here  also  are  summarized  some  of  the  reactions  which  provide  the 
energy  for  muscular  activity  through  the  splitting  of  phosphate  bonds;  e.g.,  the  con- 
version of  adenosine  triphosphate  (ATP)  (page  710)  into  adenosine  diphosphate 
ADP). 

The  left  central  section  of  the  chart  includes  relationships  between  the  compounds 
involved  in  the  metabolism  of  iron,  e.g.,  the  oxygen  carrying  molecule  hemoglobin, 
and  some  of  the  oxidizing  enzymes  of  the  "cytochrome  system".  The  latter  is  engaged 
in  numerous  oxidation-reduction  reactions;  the  dotted  lines  that  could  be  drawn  be- 
tween the  cytochromes  and  many  of  the  compounds  on  the  chart  have  been  omitted 
for  the  sake  of  simplicity. 

It  must  be  remembered  that  the  comprehension  of  the  coordination 
aspects  of  biochemical  processes  is  limited  in  the  same  manner  as  is  the  pure 
organic  chemistry  involved.  The  structures  of  relatively  small  molecules 
that  arise  a-  intermediates  or  through  degradative  action  may  be  com- 
pletely known,  but  the  structures  of  larger  aggregate-,  such  as  the  proteins 
and  the  nucleic  acids,  cannot  be  completely  described  in  terms  of  the  rela- 
tive positions  of  all  the  atoms  with  respeel  to  each  other. 

The  participation  of  complex  compounds  in  nearly  every  phase  of  bio- 
logical activity  may  be  classified  tinder  the  following-  general  headings: 

(1)  Bond  formation  and  cleavage. 

(2)  Exchange  of  functional  groups. 


702  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

(3)  Blocking  of  functional  groups. 

(4)  Influence  upon  stereochemical  configuration. 

(5)  Oxidation-reduction  reactions. 

(6)  Storage  and  transfer. 

(7)  Transmission  of  energy. 

The  first  four  of  these  functions  prescribe  that  a  complex  must  be  pro- 
duced as  an  intermediate  in  a  reaction,  the  completion  of  which  depends 
upon  the  decomposition  of  the  complex.  These  intermediates  are  labile 
coordination  compounds,  and  generally  involve  metal  ions,  such  as  mag- 
nesium, that  are  not  exceedingly  strong  electron  acceptors.  To  perform  the 
last  three  functions  the  complexes  must  remain  more  or  less  intact;  as  a 
result,  the  coordination  compounds  are  more  inert  than  those  in  the  first 
four  groups,  and  they  include  metals  like  copper  and  iron,  or  else  very  strong 
coordinating  agents. 

Bond  Formation  and  Bond  Cleavage 

Metal  ions  play  an  important  role  in  many  of  the  bond-making  and 
bond-breaking  reactions  of  natural  processes.  In  the  catalysis  of  bond 
formation  the  metal  ion  can  serve  as  a  point  of  attachment  for  the  two  donor 
atoms  between  which  reaction  is  to  take  place.  The  acceleration  of  bond 
cleavage  as  a  result  of  coordination  may  be  attributed  to  the  polarization 
of  electrons  toward  the  metal2,  and  therefore  away  from  the  organic  mole- 
cule; the  activation  energy  necessary  for  the  severence  of  the  weakest  link 
in  such  a  molecule  may  thus  be  considerably  lowered.  Many  bond -forming 
and  bond-breaking  reactions  are  reversible,  and  catalyzed  by  similar  metal- 
enzyme  compounds. 

Cleavage  of  Peptide  Bonds 

The  metabolic  decomposition  of  proteins  into  amino  acids  occurs  through 
a  complicated  series  of  reactions,  in  which  the  large  molecules  are  first  frag- 
mented into  polypeptides  by  the  endopeptidases,  and  the  resulting  poly- 
peptides are  further  degraded  by  aminopeptidases  and  carboxypeptidases, 
which  act,  respectively,  upon  the  amino  and  carboxyl  terminals  of  the 
peptide,  thus  splitting  off  amino  acids  one  by  one  from  each  end.  When  only 

2    Smith,  Proe.  Natl.  Acad.  Sci.  U.  S.,  35,  80  (1949);  Kroll,  /.  .4m.  Chem.  Soc,  74, 
2063  (1952);  Eichhorn  and  Bailar, ./.  Am.  Chem.  Soc,  75,  2905  (1953). 


COORDINATION  COMPOUNDS  IN  NATURAL  PRODUCTS  703 

two  amino  acids  remain,  the  dipeptide  is  susceptible  to  the  action  of  a  di- 
peptidase. 

Ml       (III!      CD-NH-  CHR— CO— NH CHR— CO-J-NHn 

!  I 

:- 

aminopeptidase  endopepl  Ldase 

r J 

I  ; 

I  ; 

LCHR— CO-f-NH— CHR— COOH 

exopeptidases  >  carboxj  peptidase 

I 

dipeptidase 
N 1 1  —CHR— CO^NH— CHR— COOH 

Kiulopcptidases.  Not  much  information  is  available  about  the  par- 
ticipation of  metal  ions  in  the  action  of  the  endopeptidases.  It  is  known, 
however,  that  the  enzyme  enterokinase,  which  is  involved  in  the  removal 
of  a  protective  polypeptide  from  tripsinogen,  producing  the  active  enzyme 
trypsin,  is  a  calcium  protein3.  Recently,  it  has  been  demonstrated  that  the 
activity  of  trypsin  may  be  enhanced  by  the  addition  of  a  variety  of  metal 
ions  and  chymotrypsin  by  calcium4.  It  is  possible  that  the  metal  ions  in 
these  reactions  are  coordinated  in  a  fashion  similar  to  the  linkage  in  the 
exopeptidase  complexes. 

Exopeptidases.  That  many  exopeptidases  are  metal  complexes  has  been 
amply  demonstrated  in  a  variety  of  activation  and  inhibition  experiments5. 
Many  of  the  enzymes  lose  their  activity  if  the  metal  is  removed,  and  regain 
it  upon  readdition  of  the  ion;  inhibition  by  powerful  complexing  agents. 
such  as  cysteine,  cyanide,  and  sulfide,  also  constitutes  evidence  that  metal 
ions  are  involved.  It  has  been  demonstrated  in  the  case  of  leucine  amino- 
peptidase  that  the  initial  rate  of  the  enzyme  catalyzed  reactions  is  much 
higher  if  the  enzyme  is  treated  with  metal  ion  prior  to  the  addition  of  sub- 
Btrate,  rather  than  if  the  metal  and  substrate  are  added  at  the  same  time. 
This  leads  to  the  conclusion  that  the  formation  of  bonds  between  the  metal 
ion  and  the  protein  portion  of  the  enzyme  is  a  time  consuming  proo 
indicating  thai  these  bonds  are  of  essentially  covalenl  character*. 

■  '-.  McDonald  and  Kunitz,  J.  Gen.  Physiol., 2&,  53  <  1041;. 

1.  Green,  Gladner,  and  Cunningham, /.  Am.  <'fi<  m .  Soc.,  74,  2122    1952 

•V  Smith,  "Enzymes  and  Enzyme  Systems,"  Edsall,  pp    19  T « v .  Cambridge,  Harvard 

University  Pre--.  1951 . 
5.  Smith  and  Bergmann,  J.  Biol.Ckem.,  188, 789,    1941   .  153,  r,_>7    l"!i  . 


704  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Dipeptidases.  Glycylglycine  Dipeptidase.  The  structure  of  the  dipepti- 
dase-substrate  intermediates  has  been  formulated  with  the  metal  coordi- 
nated to  the  amino  group,  the  peptide  nitrogen,  and  the  carboxyl  group  of 
the  substrate,  the  remaining  covalences  of  the  metal  being  satisfied  by 
positions  on  the  enzyme  protein.  Thus  the  cobalt(II)  enzyme  glycylglycine 
dipeptidase7  may  be  depicted113: 


ENZYME 
PROTEIN 


Coordination  to  the  amino  and  peptide  nitrogens  is  suggested  by  the  fail- 
ure of  the  enzyme  to  act  upon  the  dipeptide  having  two  methyl  groups  sub- 
stituted on  the  amino  nitrogen  or  one  methyl  group  on  the  peptide  nitro- 
gen8. Glycylglycine  dipeptidase  has  no  effect  upon  glycyl glycylglycine. 
Smith  has  shown  that  the  susceptibility  of  a  molecule  to  cleavage  by 
glycylglycine  dipeptidase  may  be  correlated  with  the  intensity  of  absorp- 
tion of  its  cobalt (II)  complex7:  the  absorption  of  the  glycylglycine  complex 
is  much  higher  than  that  of  either  the  glycine  or  glycylglycylglycine  com- 
plexes. Similar  data  have  been  obtained  by  Klotz9,  who  has  shown  that  the 
spectra  of  the  copper(II)  complexes  of  peptides  containing  even  numbers  of 
glycine  molecules  are  more  intense  than  those  of  the  odd-numbered  glycine 
peptide  complexes.  The  generalization  that  intensity  of  absorption  can  be 
used  as  a  qualitative  indication  of  the  stability  of  a  complex,  and  so  provide 
evidence  for  the  correlation  of  stability  with  enzyme  susceptibility10  ap- 
pears to  have  been  misleading  in  this  instance,  since  the  stabilities  of  glycyl- 
glycine complexes  have  since  been  determined  quantitatively  and  found  to 
be  lower  than  those  of  glycine11.  The  spectra  may  be  explained  by  the 
assumption  that  coordination  of  a  polypeptide  always  requires  the  partici- 
pation of  either  an  amino  group  or  a  carboxyl  group;  this  postulate  leads 
to  a  structure  for  the  triglycine  complex  (A)  that  resembles  the  glycine 
complex,  and  to  structures  for  di-  and  tetraglycine  complexes  (B)  that 

7.  Smith,  ibid.,  173,  571  (1948). 

s.  Smith,  ibid.,  176,  21  (1948). 

9.  Kh.tz,  Feller,  and  Urqhart, ./.  Phys.  Coll.  ('hem.,  54,  18  (1950). 
10.  Smith,   "The   Enzymes,"  Sumner  and  Myrbaeck,   Vol.   I,  p.  817,  New  York, 

Academic  Press,  Inc.,  1951. 
1  1 .  Monk,  Trans.  Faraday  Soc,  47,  297  (1951). 


COORDINATION  COMPOX  ND8  IN  NATURAL  PRODI  <  is  705 

contain  fused  ring  systems: 

)i  NH2  C=0 

CH2  NH  — CH2-C  — NH  C  =  0  —  M NH 

NH? M—  — M+:^0  O.  ^CH2 

A  B 

The  spectra  of  complexes  of  higher  polypeptides  of  glycine  show  no  sharp 
differences  depending  upon  the  presence  of  odd  or  even  numbers  of  glycine 
molecules,  since  all  of  the  complexes  probably  contain  the  condensed  ring 
structure  (B).  The  inability  of  glycylglycine  dipeptidase  to  act  upon  tri- 
glycine  can  be  interpreted  on  the  basis  of  this  structure  if  it  is  assumed  that 
structure  B  is  essential  for  enzyme  activity. 

Other  l)i peptidases.  Dipeptidases  in  general  are  specific  in  their  action 
only  upon  one  dipeptide,  and  in  their  requirement  of  a  particular  metal 
ion.  However,  the  same  substrate  may  be  acted  upon  by  different  enzymes 
in  different  tissues,  and  these  enzymes  sometimes  require  different  metals 
(e.g.,  zinc  or  magnesium  for  various  glycyl-L-leueine  dipeptidases)1*2.  Ap- 
parently the  metal  specificity  in  these  cases  depends  not  so  much  upon  the 
donors  in  the  substrate  as  it  does  upon  the  donors  in  the  enzyme  protein. 

Aminopeptidases  and  Carboxypeptidases.  The  complex  intermedi- 
ates in  the  action  of  these  enzymes5,  13  may  be  formulated  like  those  for 
the  dipeptidases;  possibly  coordination  with  the  substrate  involves  only 
two  rather  than  three  donors,  the  amino  and  peptide  groups  in  the  amino- 
peptidases,  and  the  carboxyl  and  peptide  groups  in  the  carboxypeptidases. 
Klotz •  l  has  suggested  that  the  metal  may  be  coordinated  to  the  substrate 
at  only  one  position;  he  postulates  that  the  metal  stabilizes  a  complex  be- 
tween the  peptide  bond  and  hydroxyl  ion: 

OH 

I 
R— C— NH— R' 

I 

I 

M   -■ 

I 
—Protein — 

12.  Smith,  J.  Biol.  Chem.,  176,  9  (1948). 

3mith  and  Hanson,  ibid'.,  176,  997    L948  ;  179,  902  [1949  ;  Smith,  in  'The  En 
Byrnes,"  pp.  838-40. 

Klotz,  "The  Mechanism  of  Enzyme  Action,"  McElroy  and  Glass,  pp.  267  285, 
Baltimore,  Johns  Hopkins  Press,  1964. 


706  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

The  metal  has  a  twofold  purpose  in  this  scheme:  to  attract  hydroxyl  ions 
to  the  cleavage  site,  and  to  stabilize  the  transition  state. 

Carboxylation  and  Decarboxylation  Reactions 

The  addition  and  removal  of  carbon  dioxide  are  also  widely  occurring  re- 
versible processes  which  are  catalyzed  by  metal  ions14  through  the  forma- 
tion of  complex  intermediates.  Some  of  these  reactions,  such  as  the  decar- 
boxylation of  pyruvic  acid,  may  be  accompanied  by  oxidation  or  reduction15, 
whereas  others,  such  as  the  decarboxylation  of  a-ketoglutaric  acid,  are  not16. 
The  metal  ion  is  generally  magnesium,  and  sometimes  manganese,  although 
these  may  be  replaced  by  other  metal  ions17;  in  addition,  some  carboxylase 
reactions  require  the  presence  of  diphosphothiamine,  Vitamin  Bx  pyrophos- 
phate, as  coenzyme. 

Metal-containing  carboxylase  enzymes  catalyze  the  conversion  of  oxalo- 
succinic  acid  to  a-ketoglutaric  acid20,  and  of  a-ketoglutaric  to  succinic 
acid18. 

CH2 CH— CO— COOH  CH2— CH2— CO— COOH 

||  ->    |  +  CO, 

COOH     COOH  COOH 

oxalosuccinic  acid  a-ketoglutaric  acid 

VzO-i 
4 
CH2 CH2 

I  !         +  co2 

COOH     COOH 

The  first  of  these  reactions,  as  well  as  the  decarboxylation  of  oxaloacetic 
acid16b' 17  proceeds  through  the  influence  of  metal  ions  even  in  the  absence 
of  enzyme  protein20.  Since  many  of  these  acids  are  polyfunctional,  a  number 
of  different  structures  have  been  assigned  to  the  complex  intermediates, 
among  them  the  formulation  of  the  complex  as  a  six-membered  chelate 

14.  Green,  Herbert,  and  Subrahmanyan,  /.  Biol.  Chem.,  138,  327  (1941);  Kossel,  Z. 

Physiol.  Chem.,  276,  251  (1942);  Veenesland,  Ref.  10,  Vol.  II,  pp.  183-215;  Ochoa, 
ibid.,  Vol.  II,  929-1023;  Physiol.  Rev.  31,  56  (1951). 

15.  Lipmann,  Enzymologia,  4,  65,  (1937);  Lipmann,  J.  Biol.  Chem.,  155,   55   (1944); 

Kolnitsky  and  Werkman,  Arch.  Biochem.,  2,  113  (1943);  Utter  and  Werkman, 
ibid.,  2,  491  (1943);  Koepsell  and  Johnson,  /.  Biol.  Chem.,  145,  379  (1942); 
Koepsell  and  Johnson  and  Meek,  ibid.,  154,  535,  (1944);  Stumpf,  ibid.,  159,  529 
(1945). 

16.  Krampitz  and  Werkman,  Biochem.  J.,  35,  595  (1941);  Speck,  ./.  Biol.  Chem.,  178, 

315  (1949);  Veenesland,  /.  Biol.  Chem.,  178,  591  (1949). 

17.  Krebs,  Biochem.  J.,  36,  303  (1942). 

18.  Green,  Westerfeld,  Veenesland,  and  Knox,  J.  Biol.  Chem.,  145,  69  (1942). 
20.  Kornberg,  Ochoa,  and  Mehler,  ibid.,  174,  159  (1948). 


COORDINATION  COMPOUNDS  IN  NATURAL  PRODI  <  707 

involving  the  carbonyl  and  the  carboxyl  groups  in  0-positions  to  each 
other80: 

R 
0=C  C-COOH 


O  O 

(A) 

Martell  and  Calvinlb  have  pointed  out  that  acetoacetic  acid, 

CH3— CO— CH2— COOH, 

should  be  capable  of  this  t}rpe  of  chelation,  but  its  decarboxylation  is  not 
affected  by  metal  ions.  Moreover,  esterification  of  the  a-carboxyl  group  of 
oxaloacetic  acid  prevents  the  metal  ion  catalysis  of  the  decarboxylation21, 
thus  implicating  this  group  in  the  process,  a  circumstance  not  explainable 
on  the  basis  of  the  above  formulation;  consequently  a  chelate  between  the 
keto  and  the  a-carboxyl  groups  has  been  proposed,  and  the  mechanism  of 
the  decarboxylation  has  been  formulated  as  follows113: 

R 
0=C C— CH-COHD  0=C C=CHR    +    C02 


0~  O  *  0~  0~ 

(B) 

Such  a  mechanism  suggests  that  there  may  be  no  fundamental  difference 
bet  ween  these  so-called  decarboxylations  of  /3-ketoacids  and  the  decarboxyl- 
ations of  a-ketoacids,  such  as  pyruvic  and  a-ketoglutaric,  which  may  form 
complexes  of  type  B,  but  not  of  type  A. 

Recently  Westheimer  and  Graham23  have  studied  the  iron(UI)  complexes 
of  dimethyl  oxaloacetic  acid.  The  initially  formed  yellow  complex  ifl  COn- 

-  einbergerand  Westheimer,/.  Am.  Ckem.  Soc.,73,  429  (1951). 
22   Westheimer  and  Graham,   Paper  at    conference  on   Coordination    Chemistry, 
Indiana  Univerait 


708  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

verted  to  a  blue  substance  as  a  result  of  the  decarboxylation: 

CH3    O  CH3 

I        II  I 

o=c C— C C— O"  o=c — c=c        -f        co2 

I    II  I       „    I    I  I 

cr     p   ch3  *■         cr     cr  CH, 

\  /  \/ 

Fe  Fe 

YELLOW  BLUE 

When  the  /3-carboxyl  group  is  esterified,  the  production  of  the  iron  complex 
is  not  hindered,  but  esterification  of  the  a-carboxyl  group  prevents  the  for- 
mation of  any  yellow  color,  providing  further  evidence  that  the  alpha,  and 
not  the  beta,  carboxyl  group  is  involved  in  the  chelation. 

Since  decarboxylation  reactions  are  catalyzed  by  metal  ions  in  the  ab- 
sence of  protein,  the  purpose  of  the  latter  becomes  problematic ;  it  is  reason- 
able to  suppose  that,  in  addition  to  its  rather  marked  influence  upon  the 
rate  of  the  reaction,  the  protein  is  responsible  for  rendering  an  enzyme 
specific  for  one,  and  only  one  substrate,  a  specificity  of  which  the  simple 
metal  ion  is  quite  incapable.  The  function  of  diphosphothiamine  may  be  to 
increase  the  stability  of  the  complex  between  substrate,  metal,  and  protein. 
Perhaps  the  amino  group  of  thiamine  combines  with  the  carbonyl  group  of 
the  keto  acid  to  produce  a  Schiff 's  base,  which  then  constitutes  the  active 
substrate  for  decarboxylation. 

Carbonic  Anhydrase.  This  enzyme23  catalyzes  the  reaction  between 
water  and  carbon  dioxide  to  produce  carbonic  acid,  and  for  that  reason  is 
very  important  in  the  regulation  of  pH.  The  enzyme,  which  contains  zinc, 
obviously  cannot  function  through  the  mechanism  that  has  been  postu- 
lated for  the  keto  acids;  it  may  possibly  involve  an  intermediate  zinc- 
carbonato  complex. 

Phosphorylation 

Many  biological  bond-forming  and  bond-breaking  processes,  especially 
those  connected  with  carbohydrate  and  nucleoprotein  metabolism,  are 
accompanied  by  the  synthesis  or  destruction  of  phosphate  bonds;  indeed, 
the  energy  required  for  many  biochemical  reactions  is  derived  from  the 
cleavage  of  phosphate  bonds,  especially  the  conversion  of  adenosine  tri- 
phosphate (ATP)  to  adenosine  diphosphate  (ADP).  Many  of  the  enzymes 
associated  with  these  reactions,  the  phosphorylases  that  catalyze  the  phos- 
phorolytic  degradation  of  organic  molecules,  and  the  phosphatases  that  are 
23.  Vallee  and  Altschule,  Physiol  Rev.  29,  370  (1949). 


COORDINATION  COMPOUNDS  IN  NATURAL  PRODUCTS         709 

engaged  in  the  cleavage  of  phosphate  bonds,  have  metal  ion  constituents24, 
usually  magnesium,  and  are  inhibited  by  competing  complexing  agents. 

Since  magnesium  forms  relatively  strong  bonds  with  phosphates,  the 
presence  of  this  ion  in  phosphorylating  enzymes  points  to  the  formation  of 
complex  intermediates  in  which  the  donor  and  the  recipient  of  the  phos- 
phate are  brought  together  through  complex  formation  with  the  metal  ion. 
'Inns  a  possible  intermediate  in  the  phosphorylation  of  glucose  by  ATP 
under  the  influence  of  magnesium-containing  hexokinase  may  he  formu- 
lated as  follows: 

0  o  o 

1  I  I 

ADENOSINE-0 — P — O — P  —  O  —  P — O H 

O  O  O — Mg O  —  CH2 


H- 


O 


The  existence  of  such  an  intermediate  would  indicate  that  the  magnesium 
can  perform  a  dual  function  by  bringing  about  contact  between  the  re- 
acting molecules,  and  by  labilizing  the  phosphorus-oxygen  bond.  More 
work  on  the  magnesium  complexes  of  carbohydrates  and  of  ATP  should 
prove  of  great  value  in  the  further  elucidation  of  these  reactions. 

Insulin.  A  very  important  biochemical  coordination  compound  possibly 
related  to  carbohydrate  phosphorylation  reactions  is  the  zinc  protein, 
insulin-6.  Removal  of  zinc  greatly  decreases  the  stability  of  this  molecule, 
suggesting  that  coordination  stabilization  may  be  one  of  the  functions  of 
the  metal.  Insulin  reacts  readily  with  other  divalent  metals,  such  as  cad- 
mium, cobalt,  and  nickel2627. 

Little  is  known  concerning  the  exact  mechanism  of  the  metabolic  function 
of  insulin,  but  it  has  been  proposed  that  insulin  stimulates  carbohydrate 
metabolism  through  its  antagonism  toward  a  "diabetogenic  hormone" 
(pituitary  factor)28,  which,  in  turn,  inhibits  the  phosphorylation  of  glucose. 
These  relationships  might  be  interpreted  by  postulating  that  the  hormone 

24.  Roche,  Hef.  10,  Vol.1,  473-510;  Frisell  and  Hellermann.  .1////.  Rev.  Biochem.,  20, 

24  (1951);  Humphrey  and  Humphrey,  Biochem.  ./.,  47,  238   (1950);  Meyerhof 

and  Lohmann,  Biochem.  Z.  271,  102  (1934);  Warburg   and   Christian,    ibid., 

311,    209  (1942);  314,    149,    (1943);  Jenner  and    Kay.  ./.  Biol.   Chun.,  93,   733 

1931  ;  Roche,  Nguyen-von-Thoai,  and  Danzas,  Bull.  Soc.  Chim.  Biol.,  26,  411 

\'*\\  ;Massaii  and  Vandendriessche,  Naturwis.,  28,  L43  (1940) ; Nguyen-von- 

Thoai,  Roche,  and  Roger,  Biochim.  and  Biophi/s.  Attn,  1,  61  (1947). 

26.  Scott  and  Fischer,  Biochem.  ./..  29,  1048  '1935). 

-;>iga  and  deBarbieri,  Boll.  Soc.  Ital.  Biol.  Sper.,  21,  64  (1946). 
28.  Bjering,  Acta  Med.  Scand.,  94,  483  (1936);  Baldwin,  "Dynamic  Aspects  of  Bio- 
chemistry," p.  413,  Cambridge  University  Press,  1952. 


10 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


can  coordinate  either  with  the  magnesium  of  the  phosphorylase,  or  with 
the  zinc  of  insulin;  whenever  the  former  occurs,  phosphorylation  is  barred, 
but  when  the  hormone  is  tied  to  insulin,  magnesium  is  again  free  to  catalyze 
the  phosphorylation  reaction. 


Insulin  Zn  diabetogenic  hormone 
(phosphorylation  occurs) 


diabetogenic  hormone-Mg-phosphorylase 
(no  phosphorylation) 


Actomyosin.  The  mechanical  energy  of  muscle  contraction,  like  the 
chemical  energy  of  carbohydrate  metabolism,  is  derived  from  the  decom- 
position of  ATP  into  ADP,  and  is  brought  about  through  the  contractile 
protein  actomyosin.  Although  it  has  been  claimed  that  magnesium  ions 
actually  inhibit  the  contraction  reaction29,  predominating  opinion  holds 
that  actomyosin  is  a  metal-protein  complex  and  that  the  metal  plays  an 
active  role  in  the  contraction240  • 30.  A  possible  explanation  of  this  role  is 
that  contraction  involves  coordination  of  magnesium  with  the  ATP,  and 
the  consequent  cleavage  of  the  phosphate  radical: 


°\/< 

->    OH 

P-0 

/ 

0       o 

\ 

c 

)— p-o- 

N 


NH2 
Mg— -Acto- 
^  Myosin 


The  existence  of  a  metal-actomyosin-ATP  complex  has  been  discussed  by 
Walaas31.  The  proposed  structure  of  a  complex  intermediate  bears  much 
resemblance  to  a  structure  recently  proposed  for  the  Vitamin  BJ2  molecule. 

Other  Condensation  and  Cleavage  Reactions 

Several  other  bond-forming  and  bond-breaking  enzymes  are  metallopro- 
teins  that  do  not  fit  into  any  of  the  categories  that  have  already  been  dis- 
cussed. 

29.  Braverman  and  Morgulis,  J.  Gen.  Physiol.,  31,  411  (1948);  Mommaerts,  Science, 

104,  605   (1946);  Watanabe,  Yago,  Sugekawa,  and  Tonomura,  J.  Chem.  Soc. 
Japan,  73,  761  (1952). 

30.  Bpicer  and,  Bowen,  ./.  Biol.  Chem.,  188,  741  (1951);  Perry,  Biochem.  J.,  47,  xxxviii 

I960);  Swaneon,  ./.  Biol.  Chem.,  191,  577  (1951);  Szent-Gyoergyi,  "Chomistry 
of  Muscular  Contraction,"  2nd  101.,  New  York,  Academic  Press,  1951;  Port- 
E6hl,  Z.  Naturforsch.,™,  1  (1952). 

31.  Walaas,  Nord.  Med.,  43,  1047  (1950). 


COORDIXATIOX  COMl'OCX  1)S  fX  A   1/77/1/,   I'h'ODlCTS 


711 


Condensing  Enzyme.  This  enzyme  is  the  catalyst  for  the  condensation 
of  acetate  in  the  form  of  acetyl  coenzyme-A  (produced  by  decarboxylation 
of  pyruvic  acid  or  the  degradation  of  fatty  acids)  with  oxaloacetic  acid  enol 
to  form  citric  acid.  The  latter  then  passes  through  the  tricarboxylic  acid 

cycle,  losing  in  the  process  two  carbon  atoms  and  forming  again  oxaloacetate 
which  may  undergo  another  condensation.  This  condensation  is  therefore 

of  fundamental  importance;  the  so-called  condensing  enzyme  requires  mag- 
nesium, calcium,  or  manganese  for  its  activity32*;  these  metals  probably 
function  by  exercising  their  ability  to  bring  the  condensing  molecules  into 
contact: 


CvH    H  °    COENZYME   A 


CH, 


HOOC-C 


\  / 


\ 


\\ 


O—  M 
H  \ 


HOOC-C 


/ 


CH; 


COENZYME 

I 


■Mg 


PROTEIN 


\ 


PROTEIN 


Enolase.  Another  very  important  enzyme  is  the  magnesium-containing 
enolase,  which  catalyzes  the  dehydration  of  D-2-phosphoglyceric  acid  to 
phosphoenolpyruvate33.  The  natural  enzyme  apparently  contains  mag- 
nesium, although  manganese  or  zinc  may  be  substituted.  The  following  type 
of  intermediate  may  be  postulated  for  such  a  reaction: 


OH 

/  o         o 

CH2  \    / 

\  P 

°\       / 

Mg 

-O  X  PROTEIN 


0=C- 


CH2 

\\ 


o=c- 


V 

Mg 
0         ^PROTEIN 


H20 


According  to  this  mechanism,  the  protein  magnesium  complexes  with  the 
phosphate  and  carboxyl  groups,  producing  a  five-membered  chelate  ring. 
Coordination  of  the  oxygen  on  the  central  carbon  atom  labilizes  the  bond 

*  Note  added  in  proof:  It  now  appears  that  the  metal  requirement  is  for  the  pro 
duct  ion  of  acetyl  coenzyme  A.  The  metal  may  be  omitted  when  preformed  acetyl 
CoA  is  used.  There  is  no  evidence  at  present  that  metal  ions  are  involved  in  the  con- 
densation. See  Ochoa  in  "Methods  in  Enzymology,"  Colowicb  and  Kaplan,  Vol.  I, 
p.  685.  Academic  Press,  Inc.,  1955. 

32.  Stern  and  Ochoa,  J.  Biol.  Chem.,  191,  161  (1951). 

33.  Kun,  Proc.  Soc.  Exptl.  Biol.  Med.,  75,  68  (1950);  Warburg  and  Christian,  Bio- 

chem.  Z.  310,  384  (1941). 


712  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

between  the  carbon  atom  and  the  proton.  The  latter  is  consequently  re- 
leased, thus  placing  a  negative  charge  on  the  carbon.  The  molecule  then 
regains  neutrality  by  the  loss  of  a  hydroxyl  ion  and  the  formation  of  a 
double  bond;  the  net  result  of  the  loss  of  a  proton  and  a  hydroxide  ion  is 
the  dehydration  of  the  molecule. 

Exchange  of  Functional  Groups — Transamination 
Closely  related  to  the  bond-breaking  and  bond-forming  reactions  are  the 
group  transfer  reactions,  in  which  metal  ions  may  participate  because  (a) 
they  are  able  to  bring  reacting  molecules  together  to  form  an  activated 
complex,  (b)  they  can  serve  in  the  cleavage  of  bonds  that  occurs  prior  to 
the  transfer,  and  (c)  the  relative  stabilities  of  the  complexes  of  the  reaction 
products  may  exceed  the  stabilities  of  the  complexes  of  the  reacting  sub- 
stances. 

Probably  the  most  important  exchange  reaction  of  this  sort  is  trans- 
amination34, which  provides  a  link  between  carbohydrate  and  protein 
metabolism  through  the  transfer  of  amino  groups  from  amino  acids  to  keto 
acids.  An  example  of  a  transamination  whose  natural  occurrence  has  been 
demonstrated  is  the  reaction  of  glutamic  acid  with  pyruvic  or  oxaloacetic 
acid34  • 35  to  produce  a-ketoglutaric  acid  and  alanine  or  aspartic  acid 

O 

II 
HOOC— CH2— CH2— CH— COOH  +  CH3— C— COOH  -> 

I 
NH2 

glutamic  acid  pyruvic  acid 

HOOC— CH2CH2—C— COOH  +  CH3— CH— COOH 

II  I 

O  NH2 

a-ketoglutaric  acid  alanine 

These  reactions  are  catalyzed  by  transaminase  enzymes,  the  coenzyme  of 
which  has  been  firmly  established  as  pyridoxal36  (vitamin  B6)  or  pyridox- 
amine36a  phosphate. 

CH2OP03H 

OCH 

HO 

34.  Cohen,  J.  Biol.  Chem.,  136,  565  (1940);  Cohen,  Ref.  10,  Vol.  I,  p.  1040. 

36.  Nisonoff  and  Barnes, ./.  Biol.  Chan.,  199,  699  (1952);  Green,  Leloir,  and  Nocito, 

,7m/.,  161,  559  (1945). 
36.  Lichstein,  Gunsalus,  and  Umbreit,  ./.  Biol.  Chem.,  161,  311  (1945). 
36a.  Meister,  Sober,  and  Peterson,  J.  Am.  Chem.  Soc.,  74,  2385  (1952). 


COORDINATION  COMPOUNDS  IN  NATURAL  PRODUCTS         713 

The  requirement  of  the  vitamin  has  led  to  the  speculation37  that  the  amino 
acid  initially  forms  a  Schiff's  base  with  the  pyridoxal  (see  equation  below), 
that  subsequently  the  double  bond  shifts  to  the  amino  acid  carbon  atom, 
while  a  hydrogen  is  transferred  from  the  latter  to  the  pyridoxal,  and  that 
finally  the  newly  created  double  bond  is  cleaved,  yielding  a  keto  acid  and 
pyridoxamine.  The  latter  is  then  supposed  to  transfer  the  amino  group  that 
it  has  just  picked  up  to  a  keto  acid,  the  overall  effect  being  the  transfer  of 
the  amino  group  from  the  amino  acid  to  the  keto  acid,  with  pyridoxal  acting 
as  catalyst. 

The  nonenzymatic  transfer  of  amino  groups  from  a  large  number  of 
amino  acids  to  pyridoxal,  and  from  pyridoxamine  to  a-ketoglutaric  acid 
at  100°  has  been  thoroughly  investigated38.  It  has  been  discovered  that 
the  reaction  is  inhibited  by  ethylenediaminetetraacetic  acid  and  catalyzed 
by  copper(II),  aluminum(III),  iron(II),  and  iron(III)  (in  order  of  decreas- 
ing activity).  It  has  been  postulated  that  the  intermediates  in  these  trans- 
amination reactions  are  metal  complexes  of  the  Schiff's  bases  described 
above;  the  mechanism  may  be  depicted  as  follows: 


CH20H  CH,OH 


/  \ 

R-CH-NH2  +    OCH-(  N 

COOH                              I  I 

HO  CH3 


B 

Coordination  with  the  metal  ion  stabilizes  these  Schiff's  bases  because  the 
presence  of  the  carboxyl  group  makes  possible  the  formation  of  a  second 
ring.  (When  the  production  of  such  a  fused  ring  system  is  prevented  by  the 
absence  of  an  additional  donor  group,  metal  ion  coordination  decreases  the 
stability  of  the  Schiff's  base.)  Confirmatory  evidence  for  the  existence  of 
the  postulated  Schiff's  base  complexes  at  room  temperature  has  been  ob- 

37.  Schlenk  and  Fischer,  Arch.  Biochem.,  12,  60  (1(J47). 

38.  Snell,  ./.  Am.  Chem.  Soc,  74,  979  (1952);  Snell,  ibid.,  67,  194   (1945). 


714  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

tained  recently  through  spectrophotometric  investigations  of  systems  con- 
taining copper (II)  and  nickel (II)  ions  in  solution  together  with  pyridoxal 
and  alanine,  or  with  pyridoxamine  and  pyruvic  acid39.  Solutions  of  the 
metals  in  the  presence  of  two  such  reactants  exhibit  completely  different  ab- 
sorption phenomena  from  those  of  the  complexes  of  pyruvic  acid  or 
alanine  alone,  or  of  the  vitamin  alone,  thus  indicating  Schiff's  base  complex 
formation.  Moreover,  the  spectra  of  the  pyruvic  acid-pyridoxamine  complex 
solutions  gradually  change  upon  standing  until  they  have  become  identical 
with  those  of  the  pyridoxal-alanine  complexes,  indicating  that  under  the 
experimental  conditions  employed,  the  equilibrium  favors  Schiff's  base  A, 
which  is  produced  by  a  double  bond  shift  from  the  initially  formed  B. 

Although  the  formation  of  metal-Schiff 's  base  complexes  as  intermediates 
in  nonenzymatic  transaminations  appears  thus  to  have  been  well  estab- 
lished, the  enzymatic  reaction  does  not  necessarily  follow  the  same  course. 
Indeed,  the  presence  of  metal  ions  in  transaminase  itself  has  not  been  es- 
tablished; only  a  trace  of  metal  would  be  required,  however,  in  the  catalytic 
process  that  has  been  described. 

Other  Vitamin  B6  Catalyzed  Reactions.  Closely  related  to  the  trans- 
amination reactions  are  the  deamination41  and  decarboxylation42  of  amino 
acids,  the  deamidation  of  amino  acid  amides,  and  the  racemization  of  amino 
acids43,  all  of  which  are  catalyzed  by  vitamin  B6  in  the  presence  of  metal 
ions,  and  probably  involve  the  same  type  of  Schiff's  base  complex  inter- 
mediates. Thus  it  has  been  observed  that  L-alanine  undergoes  extensive 
racemization  in  the  presence  of  both  aluminum  ion  and  pyridoxal,  although 
it  is  quite  stable  in  the  presence  of  aluminum  ion  alone.  The  racemiza- 

!tion  can  be  explained  in  terms  of  an  equilibrium  between  structures  A 
and  B ;  reformation  of  A  from  B  and  subsequent  hydrolysis  results  in  the 
production  of  the  racemate. 

Blocking  of  Functional  Groups 

Many  biochemical  processes  involve  reactions  of  polyfunctional  molecules 
at  one  specific  point  with  reagents  that  could  presumably  attack  elsewhere. 
A  possible  function  of  coordination  with  a  metal,  therefore,  is  to  block  those 
groups  whose  participation  in  the  reaction  is  to  be  avoided. 

39.  Eichhorn  and  Dawes,  /.  Am.  Chem.  Soc,  76,  5663  (1954). 

41.  Metzler  and  Snell,  /.  Biol.  Chem.,  198,  363  (1952). 

42.  Schales,  Ref.  10,  Vol.rII  p.  246;  Gunsalus,   Bellamy,  and   Umbreit,   /.  Biol. 

Chem.,  155,  685  (1944). 

43.  Olivard,  Metzler,  and  Snell,  ibid.,  199,  669  (1952). 


COORDIXAT/o.X  COMPOl  NDS  /\    NATURAL  PRODI  CTS 


71 


Arginase 

The  degradation  of  arginine  to  ornithine  is  an  illustration  of  a  reaction 
in  which  coordination  blocking  may  be  involved. 

IIOOC— CH— (CIM       NH— C— NHS 


Nil 


Ml 


arginine 


HOOC— CH— (CH2)3—  NH2  +  NH2— CO      N  1 1 


ornithine 


urea 


The  catalyst  for  the  reaction  is  the  enzyme  arginase44,  which  in  its  natural 
form  apparently  contains  manganese,  but  it  may  also  become  activated  by 
divalent  cobalt,  nickel,  and  iron45.  The  products  of  the  reaction  are  urea 
and  ornithine;  the  latter,  but  not  the  former,  inhibits  the  decomposition 
reaction.  Moreover,  other  amino  acids  besides  ornithine  are  inhibitors46, 
although  the  inhibiting  capacity  appears  to  depend  upon  the  structural 
similarity  of  the  amino  acid  to  ornithine  (lysine  is  next  in  line  after  orni- 
thine). It  is  probable  therefore,  that  the  enzyme  metal  is  coordinated  with 
the  ornithine  portion,  rather  than  the  guanidine  part,  of  the  arginine;  the 
mechanism  of  the  reaction  is  then  illustrated  by  the  following  equation: 


Nc—  or 

/        \ 

O  NH2 

X 

PROTEIN 


/NH2 

BX     /Si 

NH        NNH 


ARGININE 


Y— CH 

O  NH, 


PROTEIN 


NH- 


ORNITHINE 


NH- 


/ 

NH; 


C=0 


The  inhibition  by  amino  acids  is  probably  due  to  their  ability  to  compete 
with  arginine  for  the  metal  ion.  The  inclusion  of  manganese  in  a  second, 
seven-membered  or  larger,  chelate  ring  involving  one  of  the  guanidine 
nitrogens  is  not  out  of  the  question,  and,  if  it  occurs,  may  explain  the 
exercise  by  the  metal  of  its  bond-breaking  capacity;  it  is  probable  that  one 
of  these  nitrogens  is  attached  to  the  enzyme,  if  not  through  the  metal,  then 
Mjme  other  point;  otherwise  the  superior  inhibiting  power  of  lysine  re- 
mains unexplained. 

It  has  been  shown  that  the  reverse  of  the  arginase  catalyzed  reaction, 

44.  Greenberg,  Ref.  10,  Vol.  I,  Chapt.  25. 

45.  Hellerman  and  Perkins, ./.  Biol  Chem.,  112,  175  (1935). 

46.  Hunter  and  Downs,  /.  Biol.  Chem.,  157,  427  (1945). 


716  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

the  conversion  of  ornithine  to  arginine16a  or  citrulline46b,  may  be  achieved  in 
the  laboratory  by  blocking  the  a-amino  group  through  coordination  with 
copper,  thus  leaving  only  the  co-amino  group  vulnerable  to  attack  by  urea. 
Since  this  reaction  is  reversible,  it  appears  reasonable  to  suppose  that  one 
of  the  functions  of  the  metal  in  arginase  is  to  prevent  the  urea  that  has  been 
removed  from  the  end  of  the  arginine  molecule  from  attaching  itself  to  the 
a-amino  group,  a  process  that  would  result  in  the  formation  of  a  biochem- 
ically unknown  substance. 

Glutathione 

Another  possible  illustration  of  coordination  blocking  may  be  that  con- 
cerned with  the  biosynthesis  of  the  widely  distributed  tripeptide,  gluta- 
thione, 

CH— CHoCHo— CO— NH— CH— CO— NH— CH2— COOH 

/  \  ! 

HOOC  NH2  CH2SH 

Whereas  peptide  bonds  are  generally  formed  between  a-amino  and  a-car- 
boxyl  groups,  the  glutamic  acid  is  in  this  instance  bound  through  the 
7-carboxyl.  It  is  easy  to  visualize  how  such  a  linkage  could  be  achieved  if  it 
is  supposed  that  the  a-carboxyl  group  is  tied  up  along  with  the  a-amino 
group  by  chelation  with  a  metal  ion. 

Stereochemical  Specificity 

Many  of  the  enzymes  that  have  been  discussed  up  to  this  point  are 
specific  for  their  substrates  to  the  extent  that  they  will  act  upon  one  optical 
isomer  and  have  no  effect  at  all  upon  its  antipode47.  Thus  glycyl-L-leucine 
dipeptidase  does  not  attack  glycyl-D-leucine,  and  arginase  splits  L-arginine 
only48.  This  specificity  becomes  plausible  when  it  is  remembered  that  prior 
to  coordination  to  the  substrate  the  metal  is  already  coordinated  to  an  en- 
zyme protein  that  is  optically  active  by  virtue  of  its  being  composed  of 
optically  active  amino  acids.  Further  coordination  with  optically  isomeric 
substrates  would  therefore  result  in  the  formation  of  diastereoisomers.  It 
becomes  evident  from  a  consideration  of  enzyme  specificity  that  only 
one  of  these  diastereoisomers  is  capable  of  existence;  possibly  steric  hin- 
drance between  the  organic  groups  is  responsible  for  the  instability  of  the 
unattainable  diastereoisomer10.  The  influence  exerted  by  coordinated  op- 

Ki;i.  Turba  and  Schuster,  /   phyaiol.  Chem.,  283,  27  (1948). 
liil>.   Kurtz,  ./.  Biol.  Chem.,  122,  477  (1938). 

17.  Bergmann,  Zervas,  Fruton,  Schneider,  and  Schleich,  ,/.  Biol.  Chem.,  109,  325 

1936). 

18.  Reisser,  7. .  Physiol.  Chem.,  49,  210  (1906);  Edlbacher  and  Bonem,  ibid.,  145,  69 

1926). 


COOIWIX AT/OX  COMPOUNDS  IN  NATURAL  riiohUCTS 


717 


tically-active  molecules  upon  entering  optically-active  donors  has  been 
discussed  by  Bailar  ei  ul  See  (Chapter  8.) 

The  Porphyrins 

It  has  been  aoted  in  the  introduction  to  this  chapter  that  the  biochemical 
functions  of  metal  ions  discussed  in  the  preceding  sections  arc  such  thai 
the  complexes  produced  must  be  relatively  labile.  On  the  other  hand,  those 
functions  which  remain  to  be  considered  require  that  the  metal  ion  be  rather 
firmly  held  by  the  donor  molecules.  A  group  of  such  molecules  that  appear 
to  have1  been  especially  constructed  for  this  purpose  are  the  porphyrins. 

These  compounds  are  derivatives  of  the  parent  substance  porphine49, 
which  consists  of  four  pyrrole  nuclei  joined  at  their  a-carbon  atoms  by 
methene  groups.  All  of  the  porphyrin  complexes  to  which  reference  will  be 
made  here  are  derivatives  of  protoporphyrin,  which  has  the  following 
structure50: 


CH=CH2 


CH2COOH 


CH2COOH 


Stability  of  Porphyrin  Complexes 

Porphyrin  molecules  form  complexes  with  metal  ions  by  coordination 
through  the  four  pyrrole  nitrogen  atoms;  since  two  hydrogen  atoms  are 
lost  in  the  process,  the  porphyrin  can  neutralize  a  dipositive  charge  on  the 
metal  ion  in  addition  to  occupying  four  positions  in  its  coordination  sphere. 


19.  Fischer  and  Gleim,  Ann.,  621, 157    L936 
50.  Fischer  and  Orth,  "Die  Chemie  des  Pyrroli 
gesellschaft  M.  B.  H,  L937.  Vol.  II,  p.  396. 


Leipzig,  Alcademische  Verlaga 


718  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

The  porphyrin  complexes  contain  four  six-membered  chelate  rings  (gen- 
erally the  ring  size  of  greatest  stability  when  double  bonds  are  involved) 
which  have  been  fused  together  in  a  manner  such  that  each  nitrogen  atom 
is  part  of  two  of  the  rings.  All  of  the  atoms  in  the  porphine  nucleus  of  the 
porphyrin  molecule  lie  in  the  same  plane51*;  consequently  the  resonance 
stabilization  of  the  organic  molecule  is  very  high,  and,  since  the  coordinated 
metal  ion  must  occupy  a  position  that  is  coplanar  with  the  rest  of  the 
molecule,  it  can  serve  as  a  nucleus  for  enhanced  resonance  stabilization  by 
producing  four  additional  fused  rings.  Another  unique  feature  of  the  por- 
phyrin complexes,  shared  only  with  the  closely  related  phthalocyanine  dyes 
(see  page  73),  is  the  completely  enveloping  cyclization  of  the  organic 
molecule,  a  factor  that  may  also  contribute  to  the  great  stability  of  these 
substances. 

Because  of  this  stability,  the  structure  of  porphyrin-containing  complexes 
is  much  more  completely  known  than  are  the  structures  of  other  naturally 
occurring  coordination  compounds,  since  the  protein  may  be  removed 
without  destruction  of  that  part  of  the  molecule  in  the  immediate  environ- 
ment of  the  metal  ion52.  Another  consequence  of  porphyrin  stability  is  the 
survival  of  the  structure  intact  in  a  variety  of  inanimate  materials  that 
have  their  origins  in  the  prehistoric  decay  of  living  matter53. 

Heme,  Hemin,  and  Hematin 

Because  many  of  the  naturally  occurring  porphyrin  complexes  contain 
iron  as  the  metallic  constituent,  the  iron  complexes  of  the  porphyrins  have 
been  exhaustively  studied.  The  most  common  of  these,  iron(II)  proto- 
porphyrin, or  heme,  presumably  has  two  coordination  positions  above  and 
below  the  plane  of  the  porphyrin  molecule  occupied  by  water  molecules.  The 
magnetic  moment  of  heme  indicates  the  presence  of  four  unpaired  electrons, 
suggesting  ionic  (outer  orbital)  bonding54a;  this  relatively  loose  bonding  in 
heme  is  in  sharp  contrast  to  that  of  the  nickel(II)  porphyrin  complexes55,  for 
which  strictly  covalent  bonding  is  indicated  by  magnetic  measurements.  It 

*  This  conclusion  is  based  on  the  similarity  of  the  structures  of  porphyrins  and 
phthalocyanines. 

51.  Robertson  and  Woodward,  J.  Chem.  Soc  ,  1937,  219;  Robertson  and  Woodward, 

ibid.,  1940,  36. 

52.  Nencki  and  Zaleski,  Z.  Physiol.  Chem.,  30,  384  (1900). 

53.  Treibs,  Angew.  Chem.,  49,  682  (1936). 

54.  Pauling,  and  Coryell,  Proc.  Natl.  Acad.  Sri.  U.  S.,  22,  159  (1936);  Pauling  and 

Coryell,  ibid.,  22,  210  (1936);  Haurowitz  and  Kittel,  Ber.,  66B,  1046  (1933); 
Pauling,  Whitney,  and  Felsing,/.  Am.  Chem.  Soc,  59,  633  (1937). 
56    Haurowitz  and  Klemm,  Bcr.,  68B,  2312  (1935);  Klemm  and  Klemm,  ./.  prakt. 
Chem.,  143,  82  (1935) ; Klemm; Angew.  Chem.,  48,  617  (1935). 


COORD/XMlo.X  COMPiH   \l>s  l\    XATIPAL  PRODUCTS 


719 


i  considerable  interest  that  the  iron  complexes  of  the  porphyrins  are 
among  the  Least  stable  of  the  heavy  metal  porphyrins,  and  have  evidently 
beeD  selected  by  nature  because  their  stability  can  be  enhanced  through 
further  coordination;  such  an  increase  in  stability  cannot  occur  in  other 
porphyrin  complexes,  since  they  have  already  attained  the  maximum  stabil- 
ity of  which  they  are  capable. 

Heme  is  very  sensitive  to  reaction  with  oxygen56,  which  results,  through 
the  formation  of  a  labile  oxygen  complex  intermediate,  in  the  conversion 
to  the  iron  (III)  protoporphyrin,  hemin52;  the  charge  on  this  ion  is  neutral- 
ized by  its  association  with  an  anion56: 


CI 


Treatment  of  hemin  with  base  at  room  temperature  results  in  the  neutral- 
ization of  the  propionic  acid  carboxyl  groups,  and  the  removal  of  a  hydrogen 
ion  from  a  coordinated  water  molecule56-  57: 


OH 


COO 


Titration  of  this  anionic  complex  with  acid  results  in  the  utilization  of  only 
two  equivalents  of  hydrogen  ion,  a  phenomenon  which  has  been  interpreted 
as  indicating  that  one  carboxyl  group  has  regained  its  proton,  and  that  the 
second  hydrogen  ion  has  neutralized  the  hydroxyl  group,  which  has  been 
replaced  from  the  coordination  sphere  by  the  unprotonated  carboxylate 
group  of  a  neighboring  ion,  thus  producing  the  binuclear  complex  a-hema- 

56.  Lemberg  and  Legge,  Reference  la,  p.  166. 

Hamsick,  Z.  Physiol.  Chem.,  182,  117  (1929);  Hamsirk,  ibid.,  190,   199  (1930); 
Morrison  and  Williams,  ./.  Biol.  Chem.,  123,  Ixxxvii  (1938). 


720 


i  II i:\ffSTHY  OF  THE  COORDINATION  COMPOUNDS 


tin56: 


H20 

This  molecule  serves  as  a  simple  model  that  demonstrates  two  types  of 
linkage  commonly  found  in  the  proteinated  natural  porphyrin  complexes. 
The  utilization  of  the  propionic  acid  carboxyl  group  for  chemical  bonding 
is  probably  a  feature  of  hemoglobin  as  well  as  of  peroxidase,  although  in 
these  compounds  the  carboxyl  is  linked  to  a  protein,  rather  than  to  another 
iron  atom.  Carboxylate  coordination  with  iron  may  also  occur  in  peroxidase, 
but  the  carboxyl  group  in  this  instance  is  derived  from  the  protein. 

Hemochromes  and  Hemichromes 

Since  the  iron  in  all  of  its  naturally  occurring  porphyrin  complexes  is 
coordinated  to  a  protein,  complexes  in  which  its  extra  valences  are  occu- 
pied by  simple  monodentate  basic  groups  can  serve  as  models  for  the 
larger  molecules.  For  reasons  already  mentioned,  nickel  porphyrins  are 
incapable  of  further  reaction  with  bases,  but  both  heme  and  hemin  can  fill 
the  coordination  positions  unoccupied  by  the  porphyrin  nitrogens  with 
ammonia,  amines,  cyanide,  etc.,  to  produce  the  so-called  hemochromes  and 
hemichromes,  respectively58.  The  former  are  diamagnetic54a,  and  the  latter 
have  only  one  unpaired  electron54a;  substitution  of  the  water  molecules  of 
heme  and  hemin  by  basic  groups  thus  has  a  profound  effect  upon  the  iron 
to  porphyrin  linkage,  transforming  essentially  ionic  bonds  into  essentially 
covalent  bonds.  Reference  has  already  been  made  to  the  importance  of  this 
transition  in  the  natural  porphyrin  compounds. 

Although  dipyridyl  and  ortho-phenanthroline  are  among  the  strongest 
electron  donors  to  ferrous  ion,  these  molecules  are  incapable  of  hemochrome 
formation,  probably  because  the  donor  atom  is  sterically  hindered  in  its 
approach  to  the  porphyrin  iron  atom59.  Ethylenediamine  does  react,  not  as 
a  chelating  agent,  since  the  replaceable  groups  are  not  in  cis  positions  but 
as  a  monodentate,  as  evidenced  by  the  fact  that  two  molecules  of  the  amine 
coordinate  with  every  iron  atom. 

58.  Lemberg  and  I-egge,  Ref.  la,  p.  174. 
50.  Ibid.,  p.  176. 


COORDINATION  COMPOUNDS  IN  NATURAL  i'h'ODUCTS          721 

The  coordination  of  iron  porphyrin  with  imidazole  is  of  particular  inter- 
est, since  the  linkage  of  iron  to  the  proteins  of  hemoglobin  and  the  cyto- 
chromes has  been  postulated  to  occur  through  an  imidazole  nitrogen  of 
histidine.  Three  molecules  of  imidazole  have  been  found  to  combine  with 
hemin60;  since  the  formation  of  imidazolium  salts  with  the  propionic  acid 
side  chains  can  account  for  a  maximum  of  two  moles,  it  has  been  demon- 
strated that  at  least  one,  and  possibly  two,  imidazole  molecules  may  be 
coordinated  with  the  iron. 

Oxidation  -reduction  Potentials 

Because  hemochrome-hemichrome  systems  are  models  of  the  biochemi- 
cally active  oxidation-reduction  catalysts,  their  oxidation  potentials  are  of 
considerable  interest.  When  the  water  molecules  of  heme  and  hemin  are 
replaced  by  basic  groups,  the  potential  decreases61,  revealing  that  coordina- 
tion with  the  nitrogen  bases  stabilizes  the  iron (II).  The  oxidation  potentials 
of  the  heme-hemin  and  hemochrome-hemichrome  systems  are  pH  de- 
pendent, and  the  slopes  of  potential  vs.  pH  curves  are  independent  of  the 
nature  of  the  coordinated  bases,  approximating  a  value  of  0.059  in  all  com- 
plexes except  those  with  cyanide  and  imidazole115.  This  constancy  has  been 
interpreted  by  Martell  and  Calvin  as  resulting  from  the  displacement  of 
one  of  the  coordinated  hemochrome  bases  by  hydroxide  ion  during  the 
oxidation.  The  constancy  of  the  oxidation  potential  of  the  cyanide  com- 
plexes has  been  attributed  to  the  great  stability  of  these  complexes,  and 
their  consequent  inertness  toward  hydroxide  ion.  The  increased  slope  ob- 
served for  the  imidazole  system  may  be  due  to  the  dissociation  of  a  hydro- 
gen ion  from  the  uncoordinated  nitrogen  in  the  imidazole  molecule. 

Reaction  with  Oxygen,  Cyanide,  Carbon  Monoxide,  and  Hydrogen 
Peroxide 

Carbon  monoxide62,  cyanide62b>  63  and  hydrogen  peroxide64  react  readily 
with  the  simple  iron  porphyrin  complexes  in  reactions  analogous  to  those 
that  occur  in  the  proteinated  biologically  active  materials.  Carbon  monoxide 
reacts  with  heme,  but  not  with  hemin,  whereas  cyanide  ion  coordinates  with 
hemin,  and  not  with  heme.  This  behavior  toward  carbon  monoxide  and 

60.  Hamsick,  Z.  Physiol.  Chem.,  241,  156  (1936). 

61.  Ref.  la,  p.  195. 

62.  Anson  and  Mirsky, ./.  Physiol.  London,  60,  50  (1925) ;  Hill,  Proc.  Roy.  Soc.  London 

100B,  419  (1926);  Hill,  ibid.,  105B,  112  (1930);  Milroy,  ./.  Physiol,  38,  392 
(1909);Pregl,  Z.  Physiol.  Chem.,  44,  173  (1905);  Lemberg  and  Legge,  Ref.  la, 
p.  185. 

63.  Hogness,  Tscheile,  Sidwell,  and  Barron, ./.  Biol.  Cfu  m .,  118,  1  (1937). 

64.  Von  Euler  and  Josephson,  Ann.,  456,  111   (1927);  Haurowitz,  Enzymologia ,  4, 

139  (1937);  Haurowitz,  Brdicka,  and  Kraus,  ibid.,  2,  9  (1937). 


722  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

cyanide  is  applicable  to  the  protein-containing  porphyrin  complexes,  and 
has  been  utilized  to  differentiate  between  natural  porphyrin  complexes  of 
iron  in  the  di-  and  tripositive  states.  Coordination  of  oxygen  with  heme, 
however,  results  in  rapid  oxidation  of  the  iron  to  the  tripositive  state;  one 
of  the  functions  of  the  protein  in  hemoglobin  must  therefore  be  the  stabiliza- 
lion  of  the  iron  (I  I) -oxygen  complex  (see  page  732),  and  the  protein  in 
cytochrome-c  is  apparently  designed  to  prevent  any  kind  of  reaction  with 
molecular  oxygen.  The  functions  of  catalase  and  peroxidase  require  weak 
coordination  of  these  compounds  with  hydrogen  peroxide;  their  protein 
components  are  evidently  responsible  for  weakening  the  rather  strong  bonds 
between  hydrogen  peroxide  and  hemin.  One  of  the  prime  effects,  therefore, 
of  the  presence  of  proteins  in  biologically  active  molecules  is  to  regulate 
the  strength  of  the  bonds  between  the  porphyrin  iron  atom  and  various 
potential  coordinating  donors  to  which  the  iron  becomes  attached  during 
the  course  of  a  catalytic  process. 

Oxidation-Reduction 

Oxidation-reduction  reactions  are  of  fundamental  importance  in  bio- 
chemical processes;  they  are  of  such  wide  occurrence  that  one  of  the  chief 
requirements  of  an  oxidant  is  its  specificity  toward  a  particular  substrate. 
The  role  of  coordination  compounds  becomes  immediately  apparent,  since 
coordination  of  the  same  metal  ion  with  different  donor  molecules  may  re- 
sult in  the  formation  of  complexes  exhibiting  a  wide  variation  in  oxidation 
potentials  (see  Chapter  11).  Another  attribute  of  complexes  which  is  useful 
in  promoting  specificity  is  their  ability  to  attach  themselves  to  the  substrate 

!  through  functional  groups  of  the  donor  molecule. 

The  oxidizing  enzymes  may  be  classified  into  two  categories:  (1)  those 
that  are  directly  responsible  for  the  oxidation  of  a  substrate,  and  (2)  those 
that  participate  in  the  chain  of  transmission  of  the  oxidizing  power  of 
molecular  oxygen  to  the  final  substrate.  Among  the  first  group  are  certain 
enzymes,  the  reduced  form  of  which  can  be  oxidized  by  molecular  oxygen; 
these  will  be  discussed  in  the  following  section. 

Oxidases 

Phenol  Oxidases.  Among  the  enzymes  susceptible  to  oxidation  by 
molecular  oxygen  are  some  that  do  not  contain  a  porphyrin  prosthetic 
group,  but  appear  to  have  the  metal  directly  attached  to  the  protein.  The 
most  thoroughly  investigated  of  these  substances  are  the  phenol  oxidases; 
these  are  capable  of  converting  phenols  or  amines  to  quinones,  which,  ac- 
cording to  Warburg65,  may  in  turn  be  instrumental  in  the  oxidation  of  other 

66.  Warburg,   "Heavy  Metal  Prosthetic  Groups  and   Enzyme  Action,"  Oxford, 
Clarendon  Press,  1949. 


COORDIXATIo.X  COMPOl  NDS  IN   V  iTURAL  PRODI  CTS 


723 


compounds.  The  metallic  component  of  these  enzymes  is  copper88,  and  their 
oxidizing  ability  depends  upon  the  reduction  of  COpper(II)  t<>  cop- 
per(I)wb.  ste. 

Phenol  oxidases  have  been  placed  in  two  groups,  the  monophenol  and 

the  polyphenol  oxidases'17.  The  latter  are  capable  of  the  rapid  oxidation  of 
ortho-diphenolic  compounds,  and  the  slower  oxidation  of  monophenolic 
substances,  to  quinones.  The  oxidation  of  the  monophenols  is,  presumably, 
a  two-step  process,  consisting  of  the1  initial  insertion  of  a  hydroxyl  group  in 
a  position  ortho  to  the  already  existing  one68,  and  a  subsequent  mani- 
festation of  polyphenol ic  oxidase  activity.  Monophenol  oxidases,  whose 
existence  is  a  matter  of  controversy,  are  incapable  of  reacting  with  diphenol 
substances. 

The  behavior  of  the  monophenol  and  diphenol  oxidases  may  be  explained 
by  the  hypothesis  that  a  monophenol  oxidase  contains  one  readily  replace- 
able coordinated  group,  being  firmly  linked  to  the  protein  in  three  positions, 
whereas  diphenol  oxidase  contains  two  labile  donors,  being  securely  at- 
tached to  the  protein  at  only  one  point.  The  failure  of  a  monophenol  oxidase 
to  coordinate  with  diphenols  may  then  be  attributed  to  steric  hindrance, 
and  the  ability  of  a  diphenol  oxidase  to  act  upon  monophenols  or  diphenols 
can  be  explained  on  the  basis  of  the  replacement  of  either  one  or  both  of 
the  labile  groups  in  the  formation  of  the  enzyme-substrate  complex.  Such 
a  scheme  could  have  validity  even  if  the  distinction  between  monophenol 
and  polyphenol  oxidases  is  an  artifact;  it  could  then  explain  the  behavior 
toward  phenolic  substrates  of  the  oxidases  with  various  modifications  of 
their  protein  component. 


PROTEIN 


<:u-OH2+  HQ-<(        y 


PROTEIN 


>-°<Z> 


MONOPHENOL      OXIDASE 


OHP        HO 


PROTEIN. 


\r  + 


DIPHENOL      OXIDASE 


+    HO 


OH2        HO 


PROTEIN       Cu 


PROTEIN 


r^°<-> 


'OH; 


66.  Kubowitz,  Biochem.  Z..  292,  221  (1937);  Kubowitz,  ibid.,  299,  32  (1939);  Keilin 

and  Mann,  Proc.  Roy.  Soc.  London,  B125,  187  (1938). 

67.  Baldwin,  Ref.  28b,  p.  156. 

68.  Dawson  and  Tarpley, Ref.  10,  Vol.  II,  pp.  454-98;  Raper,  Ergeb.  Enzymforsch., 

1,  270  (1932). 


724  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Once  coordination  has  been  achieved,  oxidation  presumably  occurs  through 
electron  transfer  from  the  phenolic  group  to  the  copper69. 

Some  evidence  has  been  accumulated  to  suggest  that  the  difference  be- 
tween monophenol  and  diphenol  oxidases  may  be  artificial,  and  that  di- 
phenolases  may  lose  their  monophenolase  activity  as  a  result  of  structural 
modifications  during  the  purification  procedure6821- 70.  According  to  Dawson 
and  Tarpley68a,  there  are  only  three  well-characterized  phenol  oxidases: 
tyrosinase,  the  enzyme  responsible  for  the  eventual  conversion  of  tyrosine 
to  a  melanine-like  substance  that  accounts  for  plant  and  animal  pigmenta- 
tion, laccase71'72,  a  diphenolase  without  monophenol  oxidase  activity,  and 
ascorbic  acid  oxidase,  a  specific  phenol  oxidase  for  the  conversion  of  its 
substrate  to  dehydroascorbic  acid73. 

Peroxidases  and  Catalases.  In  contrast  to  the  phenol  oxidases,  two 
groups  of  autoxidizable  redox  enzymes,  the  peroxidases  and  catalases,  have 
porphyrin  prosthetic  groups,  and  as  a  result  much  more  is  known  of  the  way 
in  which  iron,  their  metallic  constituent,  is  bound  to  the  substrate  and  to 
the  organic  portion  of  the  molecule. 

Both  types  of  enzymes  are  associated  with  the  degradation  of  hydrogen 
peroxide,  which  arises  as  a  by-product  of  the  oxidation  reactions  catalyzed 
by  other  enzymes  and  must  be  rapidly  transformed  because  of  its  high 
toxicity.  Catalases  are  capable  of  bringing  about  the  decomposition  of 
hydrogen  peroxide  into  water  and  oxygen74  and  of  oxidizing  primary  and 
secondary  alcohols  at  the  expense  of  hydrogen  peroxide75.  Whether  the  first 
of  these  two  processes  is  designed  to  eliminate  hydrogen  peroxide  in  an 
emergency,  after  too  rapid  accumulation,  has  been  a  controversial  issue76. 
Peroxidases  cause  the  oxidation  via  hydrogen  peroxide  of  a  large  number  of 
substances,  e.g.,  aminophenols,  diamines,  diphenols,  and  some  leuco  dyes75b. 

69.  Martell  and  Calvin,  Ref .  lb,  p.  388. 

70.  Mallette  and  Dawson,  Arch.  Biochem.,  23,  29  (1949). 

71.  Bertrand,  Compt.  rend.,  121,  166  (1895). 

72.  Bertrand,  Bull.  Soc.  Chim.  Biol.,  27,  396  (1945);  Bertrand,  Compt.  rend.,  221, 

35  (1945). 

73.  Zilva,  Biochem.  J.,  28,  663  (1934);  Tauber,  Kleiner,  andMishkind,  /.  Biol.  Chem., 

110,211  (1935);Tauber  and  Kleiner,  Proc.  Soc.  Exptl.Biol.  Med.,  32,  577  (1935); 
Srinivasan,  Current  Sci.,  4,  407  (1935);  Ghosh  and  Guba,  /.  Ind.  Chem.  Soc, 
14,  721  (1937);  Johnson  and  Silva,  Biochem.  J.,  31,  438  (1937);  Stotz,  J.  Biol. 
Chem.,  133,  c  (1940);  Lovett-Janison  and  Nelson,  J.  Am.  Chem.  Soc,  62,  1409 
(1940). 

74.  Lemberg  and  Legge,  Ref.  la,  p.  401;  Zeile  and  Hellstroem,  Z.  Physiol.  Chem.., 

192,  171  (1930). 

75.  Keilin  and  Hurtree,  Biochem.  J.,  39,  293  (1945);  Chance,  Ref.  10,  Vol.  II,  p.  448. 

76.  Theorell,  ibid.,  p.  397. 


COORDINATION  COMPOl  NDS  IN  NATURAL  PRODI  CTS         725 

Cataiases  and  peroxidases  are  iron(III)  protoporphyrin  <*< >i  1 1 1 )l<*?s:t*.^T ;1>  •  " 
that  differ  in  the  nature  of  the  protein  component,  the  principal  function 
of  which  appears  to  be  the  regulation  of  the  lability  of  hydrogen  peroxide 
in  the  hydrogen  peroxide  complex;  (a  secondary  effect  of  the  protein  in  the 
case  oi  catalase  is  a  high  degree  of  stabilization  of  iron(III);  unlike  the  iron 
in  peroxidase,  that  in  catalase  cannot  be  reduced  by  the  action  of  dithi- 
onite78).  Hemin  itself  exhibits  some  catalase  activity,  but  the  reaction  is 
very  slow  because  of  the  relative  inertness  of  the  iron-H202  bond64*.  This 
bond  has  been  considerably  weakened  in  peroxidase  to  permit  more  rapid 
reaction,  but  is  most  labile  in  catalase,  which,  according  to  some,  musl 
function  when  too  much  hydrogen  peroxide  has  accumulated.  Another 
apparent  difference  between  the  two  enzyme  types  is  the  existence  of  only 
one  iron  porphyrin  prosthetic  group  in  a  molecule  of  peroxidase79  and  of 
four  such  groups  in  a  molecule  of  catalase80. 

It  has  been  concluded  from  a  study  of  titration  data  that  the  iron  atom 
in  horseradish  peroxidase  is  coordinated  to  a  carboxyl  group  of  the  protein; 
at  the  same  time  one  of  the  propionic  acid  side  chains  of  the  porphyrin  may 
be  tied  to  the  protein  at  another  point81,  possibly  to  a  tyrosine  hydroxyl 
group  to  form  an  ester  type  linkage.  It  has  not  been  definitely  established 
whether  the  sixth  position  of  the  peroxidase,  the  one  to  which  hydrogen 
peroxide  becomes  attached  in  the  catalysis,  is  occupied  by  water82  or  a 
hydroxyl  group81b>  82, 83.  These  features  have  been  incorporated  in  the  fol- 
lowing diagram: 

H.        M 

'"or 


N 


/    X/COOH 


^r  ? 

O-C  — PROTEIN 
ii 
O 

77.  Theorell,  Arkiv.  Kemi.  Mineral.  Geol.,  14B,  No.  20  (1940);  Theorell,  Bergstrioni, 

and  Alleson,  Arkiv.  Kemi.  Mineral.  Geol,  16A,  Xo.  13  (1942);  Stern,  J.  Biol. 
Chem. ,112,  661  (1936). 

78.  Keilin  and  liar-  hem.  J.,  39,  148  (1945). 

79.  Theorell,  Arkiv.  Kemi.  Mineral.  Geol.,  15B,  Xo.  24  (1940). 

BO.  Agner,  ibid;  16 A,  No.  6  (1943);  Theorell,  Ad  .,  7,  265  (1947);  Lemberg 

and  Lef.  La,  p.  41 1. 

81.  Theorell,  Arkiv.  Ken  i.  Mint  al.  Geol.,  16A,  No.  11    1942);  Theorell  and  Paul, 

ibid.,  18A,  No.  12    1944). 

82.  Cham .  Biophys.,  40,  153  (1962). 

gner  and  Theorell,  Arch  .10,321  (1946),  for  catalase. 


720  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Unreacted  peroxidase  contains  five  unpaired  electrons84,  indicating  ionic 
bond  character.  Substitution  of  the  labile  group  with  fluoride  leaves  the 
magnetic  moment  unchanged,  but  coordination  with  cyanide  or  hydrogen 
sulfide  results  in  a  transition  to  the  covalent  type,  as  manifested  by  a  re- 
duction of  the  magnetic  moment  to  that  corresponding  to  one  unpaired 
electron84.  The  nitric  oxide  complex  is  diamagnetic  as  a  result  of  the  pairing 
of  the  unpaired  electrons  of  the  metal  and  donor  molecules  and  reduction 
of  the  iron  (III)  to  iron  (II)  by  the  nitric  oxide.  Carbon  monoxide  produces 
a  diamagnetic,  covalent  complex  with  the  reduced  form  of  the  peroxidase, 
but  it  does  not  inhibit  the  activity  of  the  enzyme,  since  that  depends  upon 
the  availability  of  the  oxidized  form  of  the  molecule. 

Because  of  the  lability  of  the  complexes  of  catalases  and  peroxidases 
with  hydrogen  peroxide  their  investigation  has  proved  to  be  a  more  difficult 
task  than  is  the  study  of  the  complexes  with  the  inhibitors;  Chance  has 
been  able  to  overcome  this  difficulty  with  a  good  deal  of  success  by  applica- 
tion of  a  technique  for  the  study  of  extremely  rapid  reactions85;  he  has 
proposed  the  existence  of  four  types  of  complexes  between  enzyme 
and  peroxides86.  The  most  significant  of  these  are  the  "primary  enzyme- 
substrate  compounds,"  and  the  "secondary  enzyme-substrate  compounds" 
that  are  formed  initially  by  a  change  in  the  structure  of  the  primary  com- 
plexes87. The  spectra  of  the  primary  compounds  suggest  that  the  hydrogen 
peroxide  molecule,  in  addition  to  its  coordination  with  iron,  is  also  some- 
how tied  to  a  methene  bridge  of  the  porphine  ring86.  The  spectra  of  the  sec- 
ondary compounds  resemble  those  of  the  cyanide  and  hydrogen  sulfide 
complexes870;  hence  they  are  probably  simple  coordination  compounds. 
In  peroxidases  the  formation  of  the  primary  and  secondary  compounds  is 
essential  if  the  reaction  with  the  reductant  is  to  occur88,  but  in  catalases 
the  primary  compound  seems  to  be  the  only  catalytically  active  compon- 
ent, and  the  secondary  compound  actually  inhibits  catalase  activity89. 

The  specificity  of  catalases  for  their  substrate  is  considerabty  greater 
than  that  of  the  peroxidases,  probably  because  the  catalase  protein  per- 
mits reaction  only  with  molecules  of  restricted  size  and  shape  (activity 
toward  alkyl  peroxidases  decreases  with  chain  length)  and  the  peroxidase 
prosthetic  group  apparently  lies  exposed,  thus  minimizing  steric  hindrance 
in  the  coordination  with  a  substrate86. 

84.  Theorell,  Arkiv.  Kemi.  Mineral.  Geol.,  16A,  No.  3  (1012). 

86  Chance,  Rev.  Sci.  Instruments,  18,  601  I L947)-. 

86.  Chance,  Kef.  10,  Vol.  II,  p.  440. 

s7  Chance,  ./.  Biol.  Chem.,  179,  L331,  1341  (1040);  Chance,  ./.  Am.  Chem.  Soc,  72, 

L577    1950  ;  Chance,  Arch.  Biochem.,  21,  416  (10-10). 

sx.  Chance,  Arch.  Jiiochem.,  22,  224  (1040). 

89.  Chance,  ./.  Biol.  Chem.,  179,  1341  (1040). 


COORDINATION  COMPOUNDS  l\    S  [TUBAL  PRODUCTS         727 

Dehydrogenases 

Many  redox  enzymes  cannot  read  directly  with  molecular  oxygen,  and 
are  therefore  reoxidized  through  the  cytochrome  system.  Some  of  these 
enzymes  such  as  yeast  lactic  acid  dehydrogenase,  which  catalyzes  the  inter- 
conversion  of  pyruvic  acid  and  lactic  acid90,  maybe  metalloproteins.  The 
hydrogenase  enzymes  can  catalyze  the  reaction  of  molecular  hydrogen 
with  oxygen  to  form  water,  with  carbon  dioxide  to  produce  formic  acid, 
etc.91.  Evidence  for  the  presence  of  a  hematin  prosthetic  group  in  this 
enzyme  consists  of  the  inhibition  by  cyanide  ion  in  the  oxidized,  but  not 
in  the  reduced  form,  inhibition  by  carbon  monoxide93,  but  only  in  the 
dark,  and  the  fact  that  deficiency  of  iron  in  organisms  induces  decreased 
hydrogenase  activity91  •  94. 

It  should  be  pointed  out  that,  of  the  known  dehydrogenases,  those  that 
have  been  shown  to  be  metal  complexes  are  very  much  in  the  minority. 

The  Cytochrome  System 

The  enzymes  that  act  as  the  middlemen  in  the  delivery  of  the  oxidizing 
power  of  molecular  oxygen  to  the  eventual  substrate  belong  to  the  cyto- 
chrome system;  these  are  a  group  of  iron-porphyrin-protein  complexes 
that  differ  from  each  other  in  the  nature  of  the  protein95,  and  possibly  in 
the  attachment  of  the  latter  to  the  prosthetic  group.  The  need  for  the 
cytochrome  system  apparently  arises  from  the  fact  that  autoxidation  of 
most  substrates  would  entail  such  high  oxidation  potentials  that  the  cells 
would  be  damaged  or  destroyed96.  The  existence  of  the  system  thus  sub- 
stitutes a  series  of  redox  reactions  of  low  potential  for  one  such  reaction 
whose  potential  is  too  high.  The  order  in  which  the  various  cytochromes 
take  part  in  the  scheme  is  not  at  all  definite  at  this  time.  It  appears  cer- 
tain that  cytochrome  oxidase  is  oxidized  directly  by  the  oxygen  that  it 
receives  from  oxyhemoglobin.  Cytochrome  oxidase,  in  turn,  may  act 
upon  cytochrome-a,  which  oxidizes  cytochrome-c,  which  in  turn  acts  upon 
•  ytochrome-696. 

<\  tochronie-c.  Of  the  various  components  of  the  cytochrome  system, 
present  structural  knowledge  is  most  adequate  for  cytochrome-c,  because 
that  compound  is  the  only  soluble,  and  therefore  easily  separable,  member 
of  the  group. 

Bach,  Dixon,  and  Zerfas,  Biochem.  ./.,  40,  229  (1946 
91.  Uml  10,  Vol.  II,  Chapt.  54;  Green  and  Strickland,  Biochem.  J.,  28,  898 

L934);  Stephenson  and  Strickland,  ibid.,  26,  712  (1932);27,  1517,  L528    1933 

93.  Boberman  and  EUttenberg, ./.  Biol.  Chern.,  147,  211  (1943). 

94.  Waring  and  Werkman,  Arch.  Biochem.,  1,  425  (1042-3);  4,  75    1944 

95.  Warburg,  Ref.  65,  p 

I.emberg  and  Legge,  Ref    La,  p.  376. 


728  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

The  magnetic  moment  of  ferrocytochrome-c  is  zero97.  That  of  ferricyto- 
chrome-c  is  pH  dependent,  and  five  different  spectrophotometrically 
distinguishable  species  of  the  oxidized  form  of  the  enzyme  have  been  dis- 
covered". Two  of  these  forms  are  found  in  highly  acid  solutions  (pH  =  0.7 
and  1.4)  and  have  five  unpaired  electrons,  but  the  three  species  that  pre- 
dominate at  higher  pH  levels  (starting  at  pH  4.75)  have  only  one  such 
electron".  It  thus  appears  that  the  iron  of  cytochrome-c,  except  in  the 
oxidized  state  in  highly  acid  solution,  is  essentially  covalently  bound,  in 
contrast  to  the  iron  in  peroxidase  and  catalase.  In  line  with  this  indicated 
stability,  cytochrome-c  does  not  react  readily  with  oxygen,  carbon  monox- 
ide, hydrogen  sulfide,  azide,  and  similar  coordinating  agents97;  indeed, 
it  had  been  believed  for  some  time  that  no  such  reaction  occurs.  The  reac- 
tion of  ferrocytochrome-c  with  carbon  monoxide97  and  of  ferricytochrome-c 
with  cyanide100  and  azide101  has  now  been  demonstrated,  but  the  rate  of 
formation  of  the  former,  and  the  stability  of  the  latter,  are  so  low  as  to 
render  any  physiological  importance  of  these  compounds  quite  unlikely". 
The  cytochromes  are  the  only  known  naturally  occurring  iron-porphyrin 
complexes  whose  biochemical  function  may  not  involve  a  change  in  the 
coordination  sphere  of  the  metal  ion. 

Each  molecule  of  cytochrome-c  contains  one  hematin  group102,  which  is 
apparently  bound  to  the  protein  at  four  places.  The  two  coordination 
positions  of  the  iron  that  are  unoccupied  by  the  porphyrin  nitrogens  are 
apparently  attached  to  a  basic  donor  in  the  protein  since  cytochrome-c 
has  a  hemochrome  type  spectrum103;  titration  data  indicate  that  the  donor 
may  be  histidine  imidazole97.  The  other  two  links  between  protein  and 
the  prosthetic  group  involve  the  side  chains  of  the  porphyrin103.  The  par- 
ticular porphyrin  that  can  be  isolated  from  cytochrome-c  resembles  proto- 
porphyrin in  all  aspects  but  one ;  namely,  the  addition  of  two  cysteine 
molecules  across  the  double  bonds  of  the  vinyl  groups104.  These  cysteine 
molecules  are  the  terminal  groups  of  the  protein ;  the  firmness  of  the  attach- 
ment of  protein  to  prosthetic  group  in  this  compound  is  evidenced  by  the 
fact  that  the  iron  may  be  removed  without  disturbing  this  attachment. 
The  structure  of  cytochrome-c  may  then  be  represented  as  follows : 

97.  Theorell  and  Akesson,  /.  Am.  Chem.  Soc,  63,  1804,  1812, 1818,  1820  (1941). 
99.  Paul,  Ref.  10,  Vol.  II,  p.  376. 

100.  Horecker  and  Kornberg,  J.  Biol.  Chcm.,  165,  11  (1946);  Potter,  ibid.,  137,  13 

(1941). 

101.  Horecker  and  Stannard,  ./.  Biol.  Chcm.,  172,  589  (1948). 

102.  Theorell,  Biochem.  Z.,  279,  463  (1935) ;  285,  207  (1936);  Zeile  and  Reuter,  Z.  Phys. 

Chcm.,  221,  101  (1933);  Ref.  97. 

103.  Lemberg  and  Legge,  Ref.  la,  p.  351. 

104.  Hill  and  Keilin,  Proc. Roy. Soc.  London,  107B, 286  (1930);  Theorell,  Biochem.  Z., 

301,  201  (1939);  298,  242  (1938). 


COORDINATION  COMPOUNDS  /A   AM/7  HAL  PRODUCTS 


729 


^COOH 


Although  cytochrome-c  itself  does  not  react  with  molecular  oxygen  it  may 
be  converted  by  the  action  of  pepsin  into  an  autoxidizable  fragment  of 
one-sixth  of  the  total  molecular  weight  of  the  enzyme105. 

Cytochromes  a  and  b.  Not  much  is  known  about  the  structure  of  these 
components  of  the  cytochrome  system.  Both  are  apparently  mixtures  of 
substances,  but  one  of  the  presumed  components  of  cytochrome-a  is  now 
believed  to  be  identical  with  cytochrome  oxidase106. 

Cytochrome  Oxidase.  The  porphyrin  of  cytochrome  oxidase  differs 
from  protoporphyrin  in  the  substitution  of  a  CHO  group  for  the  methyl 
group  in  the  3-position107.  The  properties  of  the  compound  have  been  in- 
vestigated mainly  through  spectrophotometric  measurements  and  in- 
hibition techniques108.  Since  the  oxidase  is  inhibited  by  carbon  monoxide, 
which  prevents  oxidation  of  the  reduced  form108,  109,  and  by  cyanide,109 
sulfide,  and  azide,  which  prevent  reduction  of  the  oxidized  form110,  the  pres- 
ence of  a  labile  coordinate  link  is  indicated,  suggesting  that  oxidation  of 
cytochrome  oxidase  may  take  place  through  the  formation  of  an  unstable 
oxygen  complex  intermediate. 

That  cytochrome  oxidase  is  essential  to  the  oxidation  of  the  cyto- 
chromes111 has  been  demonstrated  by  the  observation  that,  even  though 
complexes  of  cytochrome-c  with  cyanide  and  azide  are  extremely  unstable, 

105.  Tsou,  Nature,  164,  1134  (1949). 

106.  Keilin  and  Hartree,  Proc.  Roy.  Soc.  London,  127B,  167  (1999). 

107.  Paul,  Ref.  10,  Vol.  II,  p.  363. 

108.  Warburg,  Biochem.  Z.,  177,  471   (1920;;  Warburg,  Naturwisa.,  15,   546   (191 

109.  Krebs,  Biochem.  Z.,  193,  347  (1928); 904,  322  (192! 

110.  Keilin,  Proc.  Roy.  Soc.  London,  104B,  206  (1929);  121B,  165  (1936). 

111.  Warburg,  Xaturwiss.,  22,  441  (1934). 


730 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


the  oxidation  of  cytochrome-c  is  inhibited  by  these  ions110.  Actually  it 
must  be  the  oxidase  which  becomes  inhibited,  and  therefore  incapable  of 
the  oxidation  of  cytochrome-c. 

The  Cysteine -Cystine  System 

Many  biological  redox  reactions  are  related  to  the  oxidation  of  the 
sulfhydryl  group  of  cysteine,  and  the  reverse  of  that  reaction,  the  reduc- 
tion of  the  disulfide  link  of  cystine.  These  reactions  may  not  involve  free 
cysteine  or  cystine  molecules;  it  is  more  likely  that  these  substances  func- 
tion as  part  of  a  protein,  their  immediate  environment  in  many  substances 
being  suggested  by  the  tripeptide  glutathione. 

Pure  cysteine,  from  which  heavy  metals  have  been  removed,  is  very 
slowly  oxidized  by  molecular  oxygen113.  The  catalytic  effect  of  metal  ions 
upon  this  oxidation  has  been  investigated  in  comparative  experiments 
with  divalent  iron,  cobalt,  and  nickel114.  The  difference  in  the  behavior  of 
the  three  ions  in  their  reaction  with  cysteine  is  highly  instructive  in  view 
of  the  specificity  of  metal  ions  in  biochemical  reactions.  All  three  metal 
ions  react  with  cysteine  in  the  absence  or  in  the  presence  of  oxygen;  only 
in  the  case  of  nickel  are  the  complexes  produced  under  the  two  conditions 
identical,  indicating  that  the  nickel  complex  is  the  only  one  that  is  not  sus- 
ceptible to  oxidation. 

The  cobalt(II)  complex  does  absorb  oxygen;  quantitative  determina- 
tions of  oxygen  uptake  have  revealed  that  the  amount  of  oxygen  consumed 
depends  upon  the  cobalt  concentration,  if  cysteine  is  in  excess,  and  upon 
the  cysteine  concentration,  if  cobalt  is  in  excess.  In  either  case,  one-half 
mole,  and  no  more,  of  oxygen  is  consumed  per  mole  of  cobalt  or  three  moles 
of  cysteine,  and  no  free  cystine  is  produced.  The  oxidized  molecule  is  ap- 
parently the  1:3  cobalt(III)  cysteine  complex: 


NH2       / 


CH; 


COO" 


113.  Harrison,  Biochem.  J.,  18,  1009  (1924). 

114.  Michaelis  and  Barron,  J.  Biol.  Chem., 

ibid.,  83,  367  (1929);  Michaelis,  ibid. 
Soc,  53,  3851  (1931). 


3,  191  (1929) ;  Michaelis  and  Yamaguchi, 
84,  777  (1929);  Schubert,  /.  Am.  Chem. 


COORDINATION  COMPOUNDS  IN  NATURAL  PRODI  CT8 


731 


The  reaction  with  iron  is  quite  different  from  that  with  cobalt.  The  oxygen 
Uptake  depends  upon  fche  cysteine  concentration,  even  when  the  hitter  is 
present  in  great  excess.  The  introduction  of  oxygen  (air)  into  such  a  solu- 
tion results  in  the  formation  of  a  violet  color  that  fades  upon  standing, 
only  to  he  revived  by  repeated  shaking  with  air,  until  all  of  the  cysteine 
has  been  completely  consumed.  The  violet  complex  is  probably  tris(cys- 
teine)-iron(III),  analogous  to  the  cobalt  complex  pictured  above.  It  is 
apparently  readily  transformed  into  the  1:1:1  iron(II)  cysteine-cystine 
complex : 


COO 


Because  of  the  instability  of  the  three-membered  chelate  ring,  the  cystine 
molecule  is  subsequently  lost,  replaced  by  two  more  cysteines,  and  the 
cyclic  process  is  renewed.  Thus  iron  can  serve  as  a  catalyst  for  the  oxida- 
tion of  cysteine  to  cystine.  Nickel  cannot  take  its  place  because  it  is  too 
difficult  to  oxidize,  and  cobalt  cannot  function  because  of  the  high  stability 
of  the  cobalt  (III)  complex. 

Miehaelis  and  Schubert  postulate  that  the  metal  may  be  bound  to  the 
carboxyl  and  sulfur  groups  of  cysteine  as  it  is  in  the  complexes  of  thiogly- 
colic  acid,  which  they  also  investigated115,  and  which  bear  some  resemblance 
to  the  complexes  of  cysteine.  Martell  and  Calvinlb  have  pointed  out  that 
in  the  light  of  present  knowledge  and  experimental  data  it  is  more  ap- 
propriate to  assume  that  the  amino  groups,  rather  than  the  carboxyl 
groups,  are  coordinated.  Further  support  for  the  latter  theory  may  be 
gained  from  the  observat  ion  that  the  glutathione  sulfhydryl  group  may  be 
oxidized  by  iron  in  a  fashion  that  resembles  the  oxidation  of  cysteine;  a 
similar  violet  color  is  produced  during  the  progress  of  the  oxidation113. 
In  glutathione,  coordination  of  the  carboxyl  group  of  cystine  is  prevented 
through  the  engagement  of  the  latter  in  a  peptide  bond.  It  is  possible  thai 
the  glycine  carboxyl  group  participates  in  the  chelation;  in  any  ease,  the 


115.  Miehaelis  and  Schubert,  ./.  Am.  Chem.  Soc,  62,  4418  (1930);  Schubert,  ./.    im 
Chem.Soc.,U,  1077  <1!>32). 


732  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

possible  structures  (A)  and  (B)  both  involve  nitrogen  coordination : 
NH-CH2-COOH 


o=c; 


CH CH2 


NH        S 
\    / 


Fe 


-CH 


^3 

NH2 
COOH 


COOH 


A  B 

The  cysteine-cystine  system  is  illustrative  of  two  of  the  catalytic  func- 
tions of  metal  ions,  since,  in  addition  to  the  redox  character  of  the  reac- 
tions, they  are  also  concerned  with  bond  formation  and  cleavage,  and  in 
this  sense  may  be  related  to  the  reactions  that  have  been  considered  in  an 
earlier  section. 

Storage  and  Transfer 

A  common  feature  of  the  coordination  compounds  described  up  to  this 
point  is  their  role  as  catalysts  in  chemical  reactions.  The  nature  of  coor- 
dination compounds  would  suggest  another  role — the  storage  and  transfer 
of  either  metal  ions  or  donor  molecules.  Complexes  which  perform  such 
functions,  as  well  as  some  whose  biochemical  function  is  not  yet  understood, 
will  be  considered  in  this  section. 

The  Transportation  of  Oxygen 

Hemoglobin.  Of  all  iron  porphyrin  complexes,  the  hemoglobin  molecule 
is  uniquely  constructed  for  the  purpose  of  oxygen  transport.  Unproteinated 
ferroheme  compounds  form  extremely  unstable  complexes  with  oxygen; 
they  are  easily  transformed  to  the  iron  (III)  complexes.  On  the  other  hand, 
when  the  protein  is  linked  as  in  reduced  cytochrome-c,  the  heme  iron  is 
not  affected  by  oxygen  at  all.  Hemoglobin  represents  an  intermediate  stage; 
it  is  a  molecule  capable  of  complexing  with  oxygen  without  a  resultant 
oxidation  of  the  iron.  The  stability  of  the  iron  to  oxygen  linkage  must  be 
great  enough  to  prevent  decomposition  of  the  oxyhemoglobin  during  its 
circulation  through  the  body,  yet  weak  enough  to  permit  dissociation 
when  contact  with  an  oxidase  has  been  established.  Just  how  the  globin- 


COORDINATION  COMPOUNDS  IN  NATURAL  1'h'ODUCTS         733 

heme  linkage  satisfies  all  of  these  requirements  cannot  be  understood  until 
the  nature  of  the  globin  has  been  further  elucidated. 

The  prosthetic  group  of  hemoglobin  is  iron(II)  protoporphyrin  (heme). 
It  is  believed  that  the  propionic  acid  carboxyl  groups  of  the  porphyrin  are 
tied  to  the  protein  as  in  horseradish  peroxidase116  (page  725).  It  has  been 
established  further  that  the  protein  is  also  linked  to  the  heme  by  coordina- 
tion with  iron,  but  whether  this  occurs  at  one  or  two  points,  and  through 
what  basic  group  of  the  protein,  are  issues  which  have  not  yet  been  settled. 

The  theory  that  iron  is  coordinated  to  globin  through  two  histidine 
imidazole  groups  is  based  upon  studies  of  the  pH  dependent  factor  in  the 
heat  of  oxygenation  of  hemoglobin,  which  corresponds  to  the  heat  of  dis- 
sociation of  histidine117,  and  upon  a  difference  in  the  titration  curves  of 
hemoglobin  and  oxj-hemoglobin,  that  has  been  interpreted  as  reflecting 
the  presence  in  hemoglobin  of  an  imidazole  grouping  whose  acidity  in- 
creases upon  oxygenation  as  a  result  of  removal  from  the  iron  coordination 
sphere118.  According  to  this  view  one  histidine  is  more  tightly  bound  than 
the  other  by  virtue  of  a  more  favorable  spatial  relationship;  upon  oxygena- 
tion the  "proximal"  histidine  remains  coordinated,  while  the  "distal" 
histidine  dissociates. 

There  is,  however,  some  objection  to  the  "imidazole  hypothesis",  based 
on  the  ability  of  the  oxylabile  group  to  react  with  carbon  dioxide  to  pro- 
duce carbamino  compounds,  a  reaction  not  shown  by  imidazole  itself119. 
Moreover,  Haurowitz  has  accumulated  evidence120  in  favor  of  the  theory 
that  globin  is  bound  to  iron  at  only  one  point,  and  that  the  group  displaced 
by  oxygen  is  actually  a  water  molecule.  He  has  shown  that  at  low  water 
vapor  pressures  the  spectrum  of  hemoglobin  is  converted  to  a  hemo- 
chromogen-like  spectrum,  a  phenomenon  that  can  be  reversed  by  raising 
the  humidity,  whereas  the  spectrum  of  oxyhemoglobin  is  independent  of 
the  water  vapor  pressure.  These  facts  lead  to  the  conclusion  that  hemo- 
globin contains  coordinated  water  which  may  be  removed  through  de- 
humidification  of  the  environment,  or  through  displacement  by  oxygen. 

From  the  composition  and  molecular  weight121  of  hemoglobin  it  has 
been  concluded  that  each  molecule  contains  four  heme  groups,  and  it  has 

116.  Granick,  Ckem.  Eng.  News,  51,  668  (1953). 

117.  Wyman,  J.  Biol.  Chem.,  127,  1,  581  (1939). 

118.  Wyman  and  Ingalls,  ibid.,  139,  877  (1941);  Coryell  and  Pauling,  /.  Biol.  Chem., 

132,  769  (1940). 

119.  Roughton,  Harvey  Lectures,  39,  96  (1944);  Lemberg  and  Legge,  Ref.  la,  p.  238. 

120.  Haurowitz,  "Hemoglobin,"  Roughton  and  Kendrew,  p.  53,  Barcroft  Symposium, 

York,  lnterscience  Publishers,  Inc.,   1949;  Haurowitz,  J.  Biol.  C 
193,443  (1951). 

121.  Adair,  Proe.  Hoy.  Soc.  London,  108A,  627  (1924);  Svedberg  and  Nichols,  ./. 

Chem.Soc.,49,  2920(1927]  ;  Bvedberg  and  Fahraeus,  ibid.,  48,  130  (1926). 


UPOCXDS 


been  shown  tint  these  lie  an  the  surface  of  the  giobin  niokeule^.  All  of 
the  known  and  postulated  structural  characteristics  of  the  coordination 
of  hemoglobin  have  been  incorporated  in  the  following  two 


( 


G_ 

O  L" 

N 


/   \    ^:— c— = 

s^ ^\ 

n  =--.   :_ 


^:::- 


Tz.^  -   ......   r 

rested  by  a  magnetic  moment  eotie^mndmg  to  the  presence  of  four  un- 
paired electrons:  consequently  the  molecule  is  susceptible  to  reaction  not 

:zly  -  . ~L         .  :-'-  ~ :--      -_:.   :..■:_:•  -_":--       :j.z.-zzz.z   ±z-zr<^ 

'—'-=2—  -  :— iz.i  L'V.^L:::i-:^:-i:''.^  ::r  ..zz.:z\^~.  mi: ::•-.  _ 
that  the  replacement  of  the  water  molecule  (or  "distal"  imidaxole  group) 
causes  the  iron  to  form  octahedral  coraknt  bonds  with  the  donor  atoms. 
TlrmigMan  (methemoglolni),  which  may  be  produced  by  oxidation  of 
i  whh  a  number  of  andante,  eg.,  potassium  f erricyanide  or 


13.  Son  *m&  Coryril,  /-  A«.  dm.  S«..  O,  136  (1939);  Holdea,  4«sfra&c 
£xy<L  Bitf.  Med.  Set.,  H,  159  (&0);ffi0,  Biceiem.  J.,  IS,  341  (1925). 


: ;e :  II    : 


ba 


Svnthetic  Cbresen-e.ai-rviii£  Chelate- 


Mf   Si^.~. 


• 


73G  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

For  a  discussion  of  the  structural  features  of  these  interesting  materials 
the  reader  is  referred  to  the  treatise  by  Martell  and  Calvinlb  (see  also  Chap- 
ter 1).  It  has  been  concluded  from  a  study  of  the  polarographic  half -wave 
potentials  for  the  reduction  of  oxygen  in  the  presence  of  various  chelating 
agents  related  to  the  oxygen  carriers  that  the  oxygen-carrying  ability  of 
a  molecule  is  related  to  its  ability  to  catalyze  the  reduction  of  oxygen. 
This  property  is  determined  by  the  ability  of  the  metal  to  furnish  electrons 
to  oxygen,  which  may,  in  turn,  be  correlated  with  the  stability  of  the 
complex. 

It  is  noteworthy  that  only  cobalt,  of  all  the  metals  in  the  first  transition 
series,  can  serve  in  these  simple  oxygen-carrying  chelates;  iron(II)  is  ir- 
reversibly oxidized,  and  the  copper(II)  and  nickel(II)  complexes  have 
little  tendency  to  react  with  oxygen  at  all.  The  hemoglobin  molecule  has 
been  so  constructed  that  the  coordination  of  iron (II)  with  oxygen  is  stabi- 
lized ;  there  is  therefore  an  analogy  between  cobalt  in  the  model  compounds 
and  iron  in  hemoglobin.  Because  of  this  stabilizing  ability  of  the  organic 
portion  of  the  hemoglobin  molecule,  cobalt  hemoglobin  does  not  readily 
react  with  oxygen,  and  this  substance  is  therefore  analogous  to  the  copper 
and  nickel  complexes  of  the  models. 


Storage  of  Metal  Ions 

Ferritin.  The  synthesis  of  so  important  and  elaborate  a  molecule  as 
hemoglobin  is  undoubtedly  a  complicated  process,  the  nature  of  which  is 
being  slowly  unravelled.  An  important  advance  in  this  direction  has  been 
the  discovery  of  ferritin,  an  iron  (III)  protein  complex,  whose  sole  function 
appears  to  be  the  storage  of  iron  until  it  is  needed  for  hemoglobin  syn- 
thesis131. 


The  molecule  has  evidently  been  exceedingly  well  constructed  for  the 
efficient  storage  of  iron,  since  from  17  to  23  per  cent  of  its  total  weight 
consists  of  this  metal.  The  iron  may  be  removed  by  treatment  of  ferritin 
with  sodium  thiosulfate  and  by  dialaysis  of  the  iron(II)  as  the  dipyridyl 
complex.  It  is  not  possible  to  reconvert  the  apoferritin  thus  produced  to 
ferritin  by  the  readdition  of  iron  either  in  the  form  of  its  divalent  or  tri- 
valent  salts  or  as  a  colloidal  suspension  of  iron  (III)  hydroxide. 

The  magnetic  moment  of  ferritin,  like  that  of  hematin  and  some  of  the 
methemoglobin  derivatives  (page  734),  corresponds  to  the  presence  of 
three  unpaired  electrons  per  iron  atom.  The  structure  of  ferritin  is  believed 
to  involve  long  chains  or  layers  of  protein  through  peptide  bonds.  Thus 
there  may  be  some  analogy  between  the  structures  of  ferritin  and  the 
chromium  complex  produced  in  the  tanning  of  leather. 

Hemocuprein,  and  the  Requirements  of  Copper  in  Hemoglobin 

131.  Michaelis,  Adv.  Prot.  Chem.,  Ill,  53  (1947). 


COORDINATION  COMPOl  NDS  I\  A  L77  R  II  PRODI  CT8         737 

Synthesis.  Maim  and  Kcilin13-'  have  isolated  from  Mood  cells  a  metallo- 
protein  containing  0.34  per  cent  copper  thai  is  bo  Loosely  held  that  it  Is 
removed  by  treatment  with  trichloroacetic  acid.  The  function  of  "licino- 
cuprein"  is  not  known;  it  is  possible,  however,  that  the  compound  is  con- 
cerned with  the  role  of  copper  in  the  synthesis  of  hemoglobin.  A  Large 
number  oi  experiments  have  proved  that  copper  in  trace  amounts  i- 
sential  for  this  synthesis11-;  for  example,  the  administration  of  iron  does 
not  aid  an  anemic  animal  in  hemoglobin  production  unless  the  iron  is  ac- 
companied by  copper134' 135.  The  latter,  moreover,  is  quite  specific  in  its  ac- 
tion; substitution  of  any  of  a  large  variety  of  other  metal  ions  has  proved 
ineffective13411. 

The  suggestion  that  copper  is  active  in  porphyrin  formation136  and  is 
subsequently  replaced  from  the  porphyrin  complex  by  iron  appears  to  be 
inconsistent  with  the  observation  that  the  addition  of  iron  as  the  porphyrin 
complex  has  no  effect  on  hemoglobin  synthesis  in  the  absence  of  copper137. 
Moreover,  since  copper  forms  more  stable  complexes  with  the  porphyrins 
than  does  iron,  it  is  difficult  to  envisage  such  a  replacement  reaction.  On 
the  other  hand,  it  is  more  plausible  to  assume  that  the  function  of  copper 
is  the  liberation  of  iron  from  ferritin;  perhaps  the  hemocuprein  molecule 
approaches  a  molecule  of  ferritin,  and,  as  a  result  of  the  attraction  of  copper 
for  the  ferritin  protein,  the  latter  becomes  detached  from  iron,  which  is 
then  free  to  enter  into  the  hemoglobin  production  sequence.  It  is  possible 
also  that  copper  is  responsible  for  the  coordination  of  iron  to  globin  at 
the  proper  places  by  blocking  other  positions  on  the  globin,  which  might 
otherwise  become  attached.  Our  understanding  of  the  function  of  hemo- 
cuprein and  the  role  of  copper  in  hemoglobin  synthesis  leaves  much  to  be 
desired. 

<  \anocobalamin.  Another  coordination  compound  that  may  play  a 
part  in  hemoglobin  synthesis  is  the  anti-anemic  cobalt-containing  vita- 
min B12 ,  cyanocobalamin138"143.  Knowledge  of  the  structure  of  the  compound 

_     Mann  and  Keilin,  Nature,  142,  148  (1938). 
L.sephs,  J.  Biol.  Chem.,  96,  559  (1932). 

134.  Elvehjem  and  Hart,  ibid.,  95,  363  (1932);  Keil  and  Nelson,  ibid.,  93,  49  (1931); 

Hart,  Steenback,  Waddell,  and  Elvehjem,  ibid.,  77,  777   (1928);  Elvehjem, 
Physiol.  Rev,  lb,  471  (1935). 

135.  Polonovsky  and  Briakas,  <"„//f/,/.  &  nd  Snr.  Biol.,  129,  379  1 1938). 
<  innningham,  Biochem.  -/.,25,  1267  (1931). 

137.  Kohler,  Elvehjem,  and  Hart, ./.  Biol.  Chem.,  128,  501  (1939). 

138.  Diehl,  Rec.  Chen    I'  13.  9    1952). 

130.  Buchanan,  Johnson,  Mills  and  Todd,/.  Chen  ,  8oe.,  1950.  2846 
Uo.  Schmid,  Abnoether,  and  Karrer,  Helv.  chim.  Acta,  86,  65  (11 

141.  Diehl,  Van  der  Baar,  and  Sealock,  J\  Am.  Chem.  Soc  .  72.  :>:;12  (1950). 

142.  Brink,  Kuehl,  and  Eolker.s,  &     -       112,  354    I960). 

L43.  Brockmann,  Roth,  Broquiat,  Bultquiat,  Smith.  Fahrenbach,  Cosulich,  Parker, 
Stohstad,  and  Jukes,  J.Am.  Chem.  Soc. ,72,  4325  (1950). 


738  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

is  constantly  increasing,  since  this  recently  discovered  vitamin  is  receiving 
a  great  deal  of  attention.* 

Acid  hydrolysis  of  the  vitamin  yields  a  nucleotide  and  two  moles  of 
cthanolamine  in  addition  to  a  large  molecule  to  which  the  cobalt  is  still 
attached139.  Whether  either  the  nucleotide  or  the  ethanolamine  is  coor- 
dinated to  cobalt  has  not  been  ascertained,  although  it  has  been  the  sub- 
ject of  considerable  speculation.  When  the  cobalt-containing  hydrolysis 
product  is  subjected  to  oxidation  by  5  per  cent  aqueous  permanganate, 
eight  organic  acids  are  produced,  among  them  oxalic,  succinic,  and  methyl- 
and  dimethyl-  succinic  acids,  in  addition  to  four  others  whose  structures 
are  undetermined140. 

The  oxidation  state  of  cobalt  in  the  vitamin  is  plus  three,  as  has  been  de- 
duced from  the  fact  that  the  substance  is  diamagnetic141.  An  unusual  feature 
of  the  vitamin,  in  view  of  its  biological  importance,  is  that  one  of  the  co- 
ordination positions  of  the  cobalt  is  occupied  by  a  cyanide  ion142.  The 
cyanide  may  be  replaced  by  hydroxide  through  acid  hydrolysis144- 145,  and 
treating  with  base  yielding  hydroxocobalamin,  another  compound  that  is 
frequently  associated  with  the  vitamin;  it  may  be  reconverted  into  the 
vitamin,  cyanocobalamin,  by  treatment  with  cyanide  ion. 

When  hydroxocobalamin  is  dissolved,  it  is  supposed  that  the  hydrox- 
ide group  leaves  the  coordination  sphere,  and  is  replaced  by  water,  thus 
forming  aquocobalamin  hydroxide145.  This  substance  gives  two  different 
responses  when  it  is  treated  with  various  anions.  Reaction  with  cyanide 
(yielding  the  vitamin),  nitrite,  or  thiocyanate  results  in  the  displace- 
ment of  the  water  molecule  from  the  coordination  sphere.  Chloride 
and  sulfate,  on  the  other  hand,  are  not  capable  of  this  kind  of  substitution, 
and  consequently  they  simply  replace  the  hydroxide  anion,  forming  the 
respective  aquocobalamin  salts146.  The  reaction  of  aquocobalamin  hy- 
droxide with  basic  groups,  e.g.,  ammonia,  amino  acids,  peptides,  etc., 
also  leads  to  the  replacement  of  coordinated  water;  these  substances  have 
been  termed  "cobalichromes,"  in  analogy  with  the  hemichromes145' 147.  It 
has  been  suggested  that  the  biological  action  of  cyanocobalamin  involves 
an  equilibrium  with  cobalichromes,  and  that  the  cyanide  ion  functions  in 
the  inhibition  of  various  enzymes145a  . 

*  Note  added  in  proof:  The  elucidation  of  the  structure  of  vitamin  BJ2  is  an  out- 
standing example  of  the  rapid  progress  made  in  the  coordination  chemistry  of  natural 
products  since  this  chapter  was  written.  See  Nature  176,  325,  328  (1955). 

144.  Veer,  Edelhauser,  Wijmenga,  and  Lens,  Biochem.  Biophys.  Acta,  6,  225  (1950); 

Wijmenga,  Veer,  and  Lens,  ibid.,  6,  229  (1950). 

145.  Cooley,  Ellis,  Petrow,  Beaven,  Holiday,  and  Johnson,  J.  Pharm.  Pharmacol.,  3, 

271  (1951);  Buhs,  Newstead,  and  Trenner,  Science,  113,  625  (1951). 

146.  Ellis  and  Pet  row  ,  ./ .  /'harm.  Pharmacol.,  4,  152  (1952);  Welch  and  Nichol,  Ann. 

Rev.  Biochem.,  21,  646  (1952). 

147.  Petrow,  unpublished  work;  ibid.,  21,  647. 


COORDINATION  COMPOUNDS  IN  NATURAL  PRODUCTS  739 

When  cyanocobalamin  us  treated  with  an  excess  of  cyanide  ion,  one 
other  coordinated  group  is  replaced,  yielding  thedicyano  complex146*;  this 

reaction  reveals  that  the  vitamin  contains  only  one  weak  coordinate 
covalent  bond.  X-ray  studies  have  indicated  that  the  four  strongly  co- 
ordinated groups,  irreplaceable  by  cyanide,  are  coplanar;  as  a  result  it  has 
been  proposed  that  cohalainin  may  be  a  porphyrin  complex.  A  recent 
study  has  shown  that  hydrogenation  of  vitamin  Bu  in  the  presence  of 
PtOo  results  in  the  liberation  of  five  to  six  moles  of  ammonia140;  the  sub- 
sequent discovery  that  a  large  number  of  cobalt  ammines  lose  their  am- 
monia when  subjected  to  the  same  treatment  led  to  the  suggestion  that 
ammonia  may  also  be  coordinated  to  cobalt  in  the  vitamin.  It  appears, 
however,  that  such  coordination  is  unlikely,  in  view  of  the  inertness  of  the 
coordinated  groups  in  question  toward  reaction  with  cyanide  ion. 

Polarographic  reduction  of  the  vitamin  has  been  interpreted  as  indicating 
a  two-electron  reduction  to  a  cobalt(I)  complex138.  Reduction  via  platinum 
catalyzed  hydrogenation138  leads  to  a  complex  of  cobalt (II)  that  can  be 
reoxidized  to  cobalt(III)  with  ferricyanide,  or  by  treatment  with  excess 
cyanide  ion,  which  results  in  the  formation  of  the  dicyano  complex.  The 
polarographic  wave  of  the  cobalt (II)  complex  indicates  two  one-electron 
reductions,  and  the  ultimate  conversion  to  metallic  cobalt. 

Calcium  Proteinates.  It  has  been  estimated  that  half  of  the  calcium 
present  in  blood  plasma  is  in  the  form  of  ionic  calcium,  and  that  the  other 
half  is  coordinated  to  a  protein.  It  has  been  proposed  that  the  function  of 
the  calcium-protein  complex  is  the  regulation  of  the  ionic  calcium  content148. 

Transmission  of  Energy — Chlorophyll 

Many  of  the  coordination  compounds  that  have  been  discussed  through- 
out this  chapter  are  important  both  in  plant  and  animal  metabolism.  The 
best  known  and  most  unique  complex  of  plant  materials  is  the  chlorophyll 
molecule,  whose  function  is  the  capture  of  photons  of  light  and  their  trans- 
mission to  a  system  which  may  convert  them  into  the  energy  required  for 
a  chemical  reaction. 

Calvin149  has  pointed  out  that  a  possible  specific  point  to  which  the  light 
energy  may  be  transferred  by  chlorophyll  is  the  disulfide  link  in  6 ,8-thioci  ic 
acid. 

(II 

/        \ 
CH2  (II  -(CH2)4— COol I 

\        / 

S    9 

a  compound  capable  of  promoting  the  oxidative  decarboxylation  of  pyruvic 

148.  Greenberg,  Adv.  Prot.  Chem.,  I,  147  (I'M  I 
140.  Calvin,  Ind.  Eng.  News,  31,  1735  (1953). 


740 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


acid  into  acetyl,  which  may  then  be  fed  into  the  tricarboxylic  acid  cycle 
of  the  "dark  reaction"  of  photosynthesis.  Through  the  agency  of  chloro- 
phyll, however,  the  energy  for  the  dissociation  of  the  disulfide  bond  may 
be  delivered  to  this  molecule  in  the  presence  of  light.  Since  the  "dark 
reaction"  depends  upon  the  existence  of  the  disulfide,  the  "dark  reaction" 
stops;  at  the  same  time  the  free  radical  sulfur  atoms,  produced  as  a  result 
of  the  cleavage,  become  active  in  the  reducing  portion  of  the  photosyn- 
thetic  cycle,  the  so-called  "light  reaction,"  whose  ultimate  goal  is  the  fixa- 
tion of  carbon  dioxide,  and  which  is  accompanied  by  the  elimination  of 
molecular  oxygen. 

The  structure  of  the  chlorophyll  molecule  as  it  occurs  in  the  natural 
state  is  not  known,  since  the  protein  component  is  dissociated  from  the 
prosthetic  group  during  the  extraction  of  chlorophyll.  The  prosthetic 
group  itself  exists  in  various  modifications,  all  of  which  are  complexes  of 
magnesium  with  porphyrins.  The  predominant  chlorophyll  type  in  green 
plants  is  chlorophyll-a,  which  has  the  structure150; 


.CH2 


2n5 


CH,-  CH 


?    7C.H 

^20^39    ^  CHg    C^Hg 


COOCH3 


The  other  chlorophylls  have  prosthetic  groups  that  differ  in  only  a  few 
respects;  chlorophyll-??,  for  example,  has  a  formyl  group  substituted  for 
the  3-methyl,  and  bacteriochlorophyll  has  an  acetyl  in  place  of  the  vinyl 
group  at  position  2,  while  the  3  and  4  pyrrole  carbon  atoms  have  been 
reduced151. 


150.  Fischer,  Naturwiss.,  28,  401  (1940). 

151.  Rabinowitsch,  "Photosynthesis,"  I,  Chapt.  16,  New  York,  Interscience  Pub- 

lishers, Inc.,  1945;  Loomis,  Ref.  10,  Vol.  II,  p.  1059. 


COORDINATION  COMPOUNDS  IN  X  ITURAL  PRODUCTS  711 

The  ability  of  the  chlorophyll  molecule  to  act  as  an  agent  for  the  trans- 
mission of  light  energy  is  due  to  its  capacity  to  absorb  light  and  to   be 

raised  to  an  excited  energy  state.  The  factors  that  influence  this  excitation 
are  also  the  factors  that  determine  the  absorption  spectrum152;  analyses  of 
the  spectra  of  chlorophyll  and  related  compounds  have  shown  that  the 
reduction  of  one  of  the  pyrrole  rings153  and  the  introduction  of  magnesium154 
are  the  two  most  important  structural  modifications  of  protoporphyrin 
that  affect  the  absorption  spectrum  and  give  the  characteristic  green 
color1'5. 

The  chlorophyll  molecule  in  the  excited  state  may  regain  its  ground 
state  condition  by  a  variety  of  paths156;  among  these  are  luminescence, 
and  the  transfer  of  energy  to  a  chemical  reaction  system.  Models  have 
been  devised,  in  which  chlorophyll  has  been  permitted  to  initiate  reac- 
tions ether  than  photosynthesis157;  of  more  importance  from  the  point  of 
view  of  coordination  chemistry,  it  has  been  demonstrated  that  the  mag- 
nesium complex  of  phthalocyanine,  in  hot  hydrocarbon  solvents,  exhibits 
both  the  phenomenon  of  luminescence158  and  the  ability  to  stimulate 
chemical  reactions,  such  as  the  conversion  of  tetralin  hydroperoxide  to 
a-tetralone159.  The  substitution  of  zinc  for  magnesium  yields  a  compound 
that  luminesces,  but  not  nearly  to  the  extent  of  the  magnesium  complex; 
the  iron,  copper,  and  nickel  complexes  do  not  luminesce  at  all159, 160.  The 
reason  for  the  metal  specificity  in  the  production  of  luminescent  porphyrin 
and  phthalocyanine  complexes  cannot  be  clearly  understood  until  the 
phenomenon  of  luminescence  itself  has  been  more  thoroughly  elucidated. 
Hill160  has  observed  that  the  magnesium  and  zinc  complexes,  which  ex- 
hibit this  property,  possess  an  inert  gas  configuration,  wThereas  the  iron, 
copper,  and  nickel  complexes  do  not. 

It  is  possible  to  make  some  further  observations  of  differences  in  the 
structures  of  luminescing  and  nonluminescing  complexes,  which  may  or 
may  not  prove  helpful  in  the  correlation  of  this  property  with  structure. 
Most  of  the  complexes  of  iron,  copper,  and  nickel,  whose  structures  have 
been  determined,  are  octahedral  or  square  planar;  in  either  case  four  of 
the  bonds  connecting  the  metal  to  the  coordination  donors  are  coplanar. 
In  general,  the  influence  of  the  metal  ion  is  all-important  in  the  determina- 

152.  Rabinowitsch,  "Photosynthesis,"  Vol.  II,  p.  619. 

153.  Stern  and  Wenderlein,  Z.  Physik.  Chem.,  174 A,  81  (1935). 

154.  Stern  and  Wenderlein,  ibid.,  176A,  81  (1936). 

155.  Rabinowitch,  Ref.  151,  Vol.  II,  p.  619. 
166.  Rabinowitch,  ibid.,  Vol.  II,  p.  796. 

157.  Warburg  and  Luettgens,  Biokhimija,  11,  303  (1946). 

158.  Helberger,  Naiurwiss.,  26,  316  (1938). 
L59.  Helberger  and  Hever,  Ber.,  72,  11  (1939). 
160.  Hill,  Ado.,  Enzym.,  12,  1,  (1951). 


742  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

tion  of  the  geometrical  configuration  of  the  complex.  The  porphyrin  and 
phthalocyanine  molecules,  however,  have  the  unusual  property  of  forcing 
the  planar  configuration  upon  the  metal  ion,  if  coordination  is  to  take 
place,  since  the  donor  atoms  are  held  in  the  same  plane  by  the  rigid  struc- 
ture of  these  molecules.  Magnesium  and  zinc  generally  form  only  tetra- 
hedral  complexes;  magnesium,  in  particular  has  no  d-orbitals  available 
for  strong  planar  bond  formation.  Therefore  the  planar  bonds  in  these 
complexes  must  be  strained,  and  the  electrons  that  make  up  these  bonds 
may  be  partly  responsible  for  the  ability  to  absorb  and  to  reemit  energy. 
It  is  significant  that  the  magnesium  phthalocyanine  complex  in  the  solid 
state  is  combined  with  two  molecules  of  water  that  are  not  thermolabile161, 
thus  defying  the  usual  coordination  number  of  four  for  magnesium.  Chloro- 
phyll itself  is  very  hygroscopic,  and  the  presence  of  half  a  mole  of  water 
per  mole  of  chlorophyll  has  been  noted162.  Evidently  enough  electron 
density  resides  outside  the  plane  of  the  molecule  to  make  such  bonding 
possible;  perhaps  the  excited  and  unexcited  states  of  chlorophyll  are 
differentiated  by  the  presence  or  absence  of  coordinated  molecules  of 
water. 

161.  Linstead  and  Lowe,  J.  Chem.  Soc,  1934,  1022. 

162.  Rabinowitch,  Ref.  151,  Vol.  I,  p.  450. 


A/..  Dyes  and  Pigments 

Roy  D.  Johnson 

American  Embassy,  Melbourne,  Australia 

and 

Niels  C.  Nielsen 

University  of  Missouri,  Columbia,  Missouri 

The  importance  of  coordination  in  dyeing  has  been  systematically  in- 
vestigated only  during  the  past  few  decades.  Although  Werner1  called  atten- 
tion to  it  in  1908,  Morgan  and  his  co-workers  must  be  credited  with  the 
first  complete  studies  in  the  field. 

Purely  inorganic  coordination  compounds  comprise  only  a  small  fraction 
of  the  pigments  and  dyes  being  used.  Most  dyestuffs  are  synthetic  organic 
compounds;  and,  of  these,  the  large  class  of  metal-dye  compounds  called 
"dye  lakes"  are  of  greatest  interest  to  the  coordination  chemist.*  The  lakes 
are  of  two  types:  coordination  compounds  and  metal  salts  of  dyes.  Many 
commercial  dyes  contain  both  types  of  lakes. 

Although  the  term  "mordant  dyeing"  has  been  applied  to  any  process 
which  involves  the  application  of  some  compound  in  addition  to  the  or- 
ganic dyestufY,  there  is  now  a  tendency  to  consider  mordant  dyes  as  those 
which  contain  groups  capable  of  acting  as  electron-pair  donors  in  the 
formation  of  coordinate  covalent  bonds.  Work  which  is  now  in  progress  on 
the  role  of  metal  ions  in  dye-fiber  interactions  makes  it  appear  certain  that 
coordination  phenomena  are  involved  in  that  aspect  of  dyeing,  also. 

Mineral  Colors  and  Inorganic  Complexes  as  Mordants 

Many  coordination  compounds  are  highly  colored,  but  few  of  them  have 
found  use  as  coloring  agents.  One  inorganic  pigment  which  is  used  ex- 
tensively, except  in  the  United  States,  is  mineral  khaki,  which  is  formed 

1.  Werner,  Ber.,  41,  1062  (1908). 

*  A  review  of  the  literature  on  color  lakes  containing  an  extensive  bibliography 
has  been  presented  by  W.  B.  Blumenthal  in  -1///.  DyettuffReptr.,  35,  520  1 1946).  (  tther 
reviews  may  be  found  in  liefs.  18,  45,  48b. 

743 


I 


744  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

by  the  precipitation  of  mixed  iron  and  chromium  hydroxides  on  cotton 
fabrics.  The  cloth  is  impregnated  with  the  metal  salts,  treated  with  an 
alkaline  solution,  and  aged.  Polynuclear  complexes,  related  to  those  used 
in  chrome  tanning,  are  formed  by  oxolation  and  olation  (Chapter  13).  Cane 
sugar,  glucose,  glycerol,  and  other  nonelectrolytes  containing  OH  groups 
are  added  to  prevent  precipitation  of  the  pigment  by  forming  complexes 
with  the  metal  ions2- 3.  The  colloidal  behavior  of  these  solutions  also  indi- 
cates complex  formation. 

Complex  iron  cyanides  such  as  Prussian  Blue  have  been  used  in  the  dyeing 
of  textiles.  Although  early  investigations  of  the  chemical  nature  of  these 
complexes  produced  conflicting  evidence,  x-ray  analysis4  shows  that  Prus- 
sian Blue  has  a  cubic  lattice  with  Fe(II)  and  Fe(III)  ions  placed  alter- 
nately at  the  corners  of  the  cube  (p.  90).  The  cyanide  groups  are  situated 
along  the  edges  of  the  cube  and  serve  to  join  neighboring  metal  ions.  Alkali 
metal  ions  appear  at  the  centers  of  alternate  cubes.  Numerous  studies  of 
these  compounds  are  indicative  of  the  variations  in  composition5.  Salts  of 
the  [Fe(CN)6]4_  and  [Fe(CN)6]3_  complex  ions  may  be  formed  with  many 
metals  to  produce  colored  materials  whose  insolubility  suggests  their  use- 
fulness as  pigments.  The  familiar  Iron  Blues  are  well  known  examples  of 
these  compounds6.  A  newer  pigment,  Inorganic  Maroon,  has  the  approxi- 
mate composition  K2Cu[Fe(CN)6]7.  The  high  tinctorial  power  of  this  com- 
pound suggests  further  investigation  of  the  heavy  metal  salts  of  the  com- 
plex iron  cyanides  which  may  be  applicable  in  the  dyeing  of  the  newer 
synthetic  fibers  (see  page  766).  Heavy  metal  cyanides  also  have  been  em- 
ployed for  the  production  of  colored  gold  plating8. 

(The  heavy  metal  ferro-  and  ferricyanides  can  be  characterized  as  poly- 
nuclear coordination  compounds.  This  can  be  explained  by  the  tendency 
of  the  cyanide  group  to  complex  with  most  of  the  heavy  metal  ions  and 
to  its  unparalleled  ability  to  behave  as  a  bridging  group.  Hydroxide  groups 
behave  in  the  same  manner,  but  the  number  of  metal  ions  which  form 
stable  OH  bridges  is  very  much  smaller9.  Often  the  OH  group  losses  pro- 

2.  Daruwalla,  and  Nabar,  J.  Soc.  Dyers  Colourists,  68,  168  (1952);  Bhende  and 

Ramachandran,  /.  Sci.  Ind.  Research  {India),  7B,  176  (1948) ;  8B(1),  10  (1949). 

3.  Daruwalla  and  Nabar,  Kolloid.  Z.,  127,  33  (1952). 

4.  Keggin,  Nature,  137,  577  (1936). 

5.  Schaeppi  and  Treadwell,  Helv.  Chim.  Acta,  31,  577  (1948);  Saxena  and  Bhat- 

tacharya,  J.  Indian  Chem.  Soc,  28,  703  (1951);  Bhattacharya  and  Sexton, 
J.  Indian  Chem.  Soc,  29,  263  (1952);  Bhattachar}ra  and  Saxena,  J.  Indian 
Chem.  Soc,  29,  284,  529,  535,  632  (1952);  Bhattacharya  and  Saxena,  J.  Indian 
Chem.  Soc,  28,  141,  221,  (1951). 

6.  American  Cyanamid  Co.,  Nitrogen  Chemicals  Digest,  Volume  VII,  "The  Chem- 

istry of  the  Ferrocyanides,"  New  York,  American  Cyanamid  Company,  1953. 

7.  Gessler  and  Goepfert,  U.  S.  Patent  2564756  (1951);  cf.  Chem.  Abs.,  45,  10613  (1951). 

8.  Thews,  Metal  Finishing,  49  (9),  80  (1951). 

9.  Scott  and  Audrieth,  J.  Chem.  Ed.,  31,  168  (1954). 


DYES  AM)  ricuk'.xrs  746 

tons,  leaving  oxide  ion  linkages  between  the  metal  ions.  Certain  well 
known  inorganic  pigments  may  bo  coordination  compounds,  for  simple 

ratios  of  hydrated  oxides  to  normal  metal  salts  prevail  in  practically  all 
basic  salts  such  as  white  lead  and  malachite10.  This  hypothesis  has  been 
verified  in  some  cases11,  but  other  explanations  have  also  been  given  to 
account  for  the  formation  of  complex  basic  salts1-'-  l8.  These  inorganic  poly- 
mers illustrate  a  modification  rather  than  a  contradiction  of  Werner's 
hypothesis. 

Among  the  inorganic  complexes  used  as  mordants  are  the  familiar  phos- 
photungstic  and  phosphomolybdic  acids  (see  Chapter  14).  The  complexity 
of  these  materials  has  made  it  difficult  to  evaluate  their  exact  behavior  in 
mordanting  operations.  Several  formulas  for  the  mordanted  products  have 
been  suggested14.  The  addition  of  the  acid  to  the  dye  produces  both  physical 
and  chemical  changes,  the  latter  probably  involving  coordination  of  several 
dye  molecules  (R)  to  the  complex  acid  to  give  structures  of  the  type: 

R  R 

\  / 

R — Complex  Acid — R 

/  \ 

R  R 

Some  basic  dyes  are  susceptible  to  mordanting  with  potassium  ferro- 
cyanide  and  sodium  sulfite,  if  copper  sulfate  is  first  added  to  the  dye  solu- 
tion. The  use  of  the  tannin-tartar  emetic  mordant  system  is  well  known. 
After  initial  interaction  between  tannic  acid  and  the  basic  dye  molecule, 
the  antimony  salt  combines  with  the  tannic  acid  portion  of  the  molecule 
or,  more  specifically,  with  the  or^/io-hydroxy  groups  present  in  the  digallic 
acid  constituent  of  the  tannic  acid15. 

A  recent  patent  proposes  the  use  of  metal  carbonyls  of  the  iron  group 
for  mordanting  acetate  rayons.  The  process  is  suitable  for  a  large  number 
of  lake-forming  dyes  which  contain  nitro  groups16. 

Metal  Complexes  of  Organic  Dyestuffs 

Any  organic  compound  containing  intramolecular  hydrogen  bonds  will, 
in  general,  react  with  metal  ions  to  form  coordinate  covalent  bonds.  Co- 

10.  Werner,  Ber.,  40,  4441  (1907). 

11.  Weinland,  Stroh,  and  Paul,  Ber.,  55,  2706  (1922). 

12.  Feitknecht,  Helv.  C him.  Acta.,  13,  22  (1930);  16,  427,  1302  (1933);  18,  28,  40  (1935); 

19,  448,  467,  831  (1936). 

13.  Thomas,  "Colloid  Chemistry,"  New  York,  McGraw-Hill  Book  Co.,  1934. 

14.  Pratt,  "The  Chemistry  and  Physics  of  Organic  Pigments,"  NTe*  York,  John  Wiley 

&  Sons,  Inc.,  1947. 

15.  Ref.  14,  p.  178. 

16.  Grimmel,  British  Patent  631,765. 


746 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


ordination  can  occur  with  any  class  of  dyes  which  has  derivatives  containing 
the  necessary  donor  groups  in  the  proper  positions.  The  most  characteristic 
groupings  found  in  commercial  dyes  are  — OH,  — COOH,'=0,  =NOH, 
and  — NH2  in  ortho  or  peri  positions  with  respect  to  each  other  or,  in  the 
case  of  the  azo  or  azomethine  dyes,  in  the  ortho  positions  with  respect  to  the 
— N=N—  or  — N=C—  linkages. 

—NO,  —OH  Substituted  Dyes 

Naphthol  Green  B  (structure  I)  was  the  first   commercially  available 


Na03S 


soluble  acid  dye  containing  a  coordinated  metal  ion17.  The  — NO,  — OH 
groups  characteristic  of  this  dyestuff  occur  in  many  metallized  dyes. 

The  o-nitrosophenols  are  polygenetic  dyes  with  colors  ranging  from 
green  (with  Fe)  to  brown  (with  Cr)  and  yellow  (with  Zn)18.  The  similarity 
between  the  zinc  and  barium  compounds  suggests  that  salt  formation,  rather 
than  coordination,  may  occur.  Pigment  Green  B,  the  bisulfite  compound  of 
l-nitroso-2-naphthol  complexed  with  iron,  is  suitable  for  filling  rubber19. 
Various  substituents  have  led  to  numerous  other  dyes  in  the  Pigment  Green 
series. 

The  coordination  phenomena  occurring  with  the  nitrosophenols  have 
been  investigated20.  When  Gambine  Y  (1,2-naphthoquinone-l-oxime)  was 
allowed  to  react  with  [Co(NH3)6]Cl3  at  room  temperature,  a  simple  salt  was 
formed.  Upon  warming  the  salt,  six  molecules  of  ammonia  were  evolved 
and  the  chelate  compound  (structure  II)  was  formed. 


17.  Ilot'mann,  Her.,  24,  3741  (1891). 

18.  Venkataraman,  "The   Chemistry  of  Synthetic  Dyes,"  New  York,   Academic 

Press,  1952. 

19.  E.  I.  duPont  de  Nemours  and  Co.,  U.  S.  Patent  2092750  (1937). 

20.  Morgan  and  Main  Smith,  J.  Chem.  Soc,  119,  704  (1921). 


DYES  AND  PIGMENTS  7  Vt 

Morgan  and  Main  Smith  reported  that  air  oxidation  of  a  mixture  of 
7-hydroxy-l  ,2-naphthoquinone-l-oxime  and  a  cobalt  salt  gave  the  com- 
pound shown  in  Btructure(III),  while  oxidation  by  hydrogen  peroxide  in 
the  presence  of  ammonia  gave  a  more  complex  sail  (structure  IV.).  Accord- 
ing to  them,  the  [CoCNHj)*]*1"  ion  neutralized  the  three  charges  on  tin- 
complex  with  the  sixth  coordination  position  of  the  pentammine  l)ein<»; 
filled  by  one  of  the  phenolic  oxygens.  This  is  not   clearly  shown  by  their 


[Co  (NHj 


formulation  (IV).  The  formation  of  the  three  chelate  rings  widely  separates 
the  three  hydroxyl  groups  in  position  7  so  that  not  more  than  one  of  them 
could  possibly  satisfy  a  secondary  valence  of  a  given  cobalt.  Analysis 
showed  that  the  compound  contained  a  mole  of  water,  and  Lamb  and  Lar- 
son-1 have  shown  that  the  [Co(XH3)5H20]3+  ion  is  more  stable  than  the 
[Co(XH3)6]3+  ion.  This  suggests  that  the  lake  is  probably  a  simple  salt  of  the 
former.  Under  more  stringent  conditions,  the  dye  might  replace  the  water 
molecule  as  in  the  analogous  reaction: 

[Co(NH3)5H20]Cl3->  [Co(NH3)5Cl]Cl2  +  H20 

In  a  study  of  the  cobaltammine  and  iron  lakes  of  dinitrosoresorcinol  the 
cobaltammine  lakes  were  shown  to  be  monochelate.  Evidently,  the  chelate 
ring  is  formed  with  the  two  intermediate  functional  groups,  leaving  the  salt 
forming  function  to  the  terminal  functional  groups.  Similar  results  were 
obtained  with  the  green  iron(III)  lakes22. 

o-Nitrosophenol  combines  quantitatively  with  copper(II),  mercury(I), 
nickel(II),  palladium(II)  and  cobalt(III)23,  while  2-nitroso-l-naphthol  and 
the  related  Nitroso-R  salt  have  been  suggested  as  analytical  reagents  for 
cobalt24  and  for  the  colorimetric  and  photometric  determination  of  iron-'. 

21.  Lamb  and  Larson,  •/.  Am.  Chem.  Snc,  42,  2024  (1920). 
.'_>    Morgan  and  Moss,  ./.  Chem.  Soc.,  121,  2857  (1922). 

23.  Cronheim.  /.  Org.  Chem.,  12,  1  (1947). 

24.  Jung,  Cardini,  and  Fuksman,  Anales  Assoc,   quim.   Argentict,   31,    122    (1943  ; 

Haywood  and  Wood. ./.  8oc.  Chem.  Ind,,  62,  37    L943  ;  Willard  and  Kaufmann, 
Anal.  Chem.,  19,  505  (1947  . 

25.  Sideris,  Young,  and  Chun,  Ind.  Eng.  Chem.,  Anal.  Ed.,  16,  276  (1944). 


748  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

The  a-oximinoketones  form  metal  complexes  of  the  type 


t=o. 


Fe  (where  n  =  2  or  3) 


N  —  O 


These  have  been  patented  for  use  on  photo  images26.  Nilssen27  has  reported 
that  iron  forms  complexes  with  the  compound 


OCPL 


O     NOH 

II       II 
-N— C— C— CH, 
H 


The  stoichiometry  and  structure  of  the  resulting  complex  have  not  been 
investigated,  but  it  seems  possible  that  the  oxime  group  is  not  involved 
in  the  coordination28. 

In  the  case  of  the  1-nitroso  derivatives  of  2-hydroxy-3-naphthoic  acid 
arylamides,  two  ferric  compounds,  formulated  as  structures  (V)  and  (VI), 
have  been  prepared29. 


ZZ  21 

The  formation  of  compound  (V)  requires  "enolization"  in  the  arylamide 
group.  Evidence  for  this  comes  from  the  preparation  of  the  iron  lake  of  the 
N-benzyl  derivative  in  which  "enolization"  cannot  occur,  and  only  com- 
pound (VI)  is  formed30. 

The  commercial  use  of  the  iron  complexes  of  the  o-nitrosophenols,  to  the 

26.  Sargent,  U.  S.  Patents  2533181  and  2533182. 

27.  Nilssen,  Soc.  Dyers  and  Colourists,  Symposium  on  Fibrous  Proteins,  1946,  142. 

28.  Ref.  18,  p.  404. 

29.  Unpublished.  See  Ref.  18,  p.  404. 

30.  Forster,  Kudva,  and  Venkataraman,  J.  Indian  Chem.  Soc,  Ind.  and  News  Ed.,  6, 

119  (1943). 


DYES  AND  PIGMENTS 


74!) 


exclusion  of  other  well  known  metal  complexes,  is  indicative  of  the  stability 

of  these  materials. 

Ortho-Dihydroxy  Substituted  Dyes 

Numerous  dyes  of  all  classes  contain  the  or^o-dihydroxy  group  or  the 

related  quinoid  structure  (=0,  ■ — OH);  the  most  important  of  these  are  the 
alizarin  dyes.  An  understanding  of  the  coordination  phenomena  involved 
has  resulted  from  investigations  of  simpler  ring  systems  and  of  derivatives 
of  anthracene.  Colorless  2,4,5-trihydroxytoluene  will  complex  with  cop- 
per(II),  iron(II)  and  cobalt(II)  to  give  wool  dyes  ranging  from  medium 
brown  to  black  in  color31.  The  compounds  are  formulated  as 


The  oxidation  of  the  organic  molecule  is  analogous  to  that  observed  in 
the  complexes  of  Diamond  Black  PV(VII)32. 


S03Na 


VII 


When  treated  with  chromic  acid,  this  type  of  dye  oxidizes  to  a  quinoid  form 
with  which  the  chromium,  in  its  reduced  state,  can  coordinate.  The  evi- 
dence for  this  mechanism  is  neither  extensive  nor  accurate  enough  to 
warrant  assignment  of  specific  structures  to  the  resulting  compound-. 
most  of  which  are  impure. 

Alizarin  is  a  polygenetic  dye  with  colors  ranging  from  rose-red  with 
aluminum  salts  to  violet-black  with  iron  compounds.  Turkey-Red  lake  is 
the  most  important  commercial  dye  of  this  series.  The  lakes  of  alizarin  are 

31.  Burton  and  Stoves,  •/.  Soc.  Dyers  Colon  lists.  66,  17  1    I960). 
f2.  Morgan  and  Main  Smith,  ./.  Chem.  Soc  .  125,  1731  (1924 


750 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


often  regarded  as  adsorption  complexes33,  but  a  pure  compound  has  been 
isolated  and  assigned  structure  (VIII)34. 


O 


UvM — ° Ca — ° TT^ 


O    O 

\/ 

HP0  — Al O 

A 

O    O 

o 


r     v 

Ca  —  O Al HP0 

A 


hLO 

r 

Ca  — 


O  O 


^Co 


~VTTT 


Alizarin  forms  a  cobalt  (III)  complex  containing  two  cobalt  atoms  for 
each  five  ammonia  molecules35.  This  was  first  reported  to  have  structure 
(IX),  but  is  probably  the  salt  shown  in  structure  (X). 


[Co(NH3)5] 


IT 


[co(nh3)5h2o] 


An  interesting  complex  analogous  to  Turkey  Red  contains  both  di-  and 


33.  Bancroft, ./.  Phys.  Chcm.,  36,  3137  (1932) ;  Reference  14,  p.  110. 

34.  Fierz-David  and  Rutishauser,  Helv.  Chim.  Acta,  23,  1298  (1940). 

35.  Morgan  and  Main  Smith,  /.  Chem.  Soc,  121,  160  (1922). 


D)  ES  AND  PIOMEh  TS 


751 


trivalent  cobalt  (XI) 


-12 


Purpurin  gives  a  mixture'  of  two  cobalt  lakes  in  approximately  equal  pro- 
portions, while,  with  alizarin  cyanine,  cobalt  is  reported  to  form  a  lake 
containing  two  chelate  rings  (XII). 


NNH3)3 


[Co(NH3)3 


211 

A  similar  structure  results  when  an  amine  group  is  substituted  in  the 
3-position;  however,  2-nitroalizarin  reacts  with  cobalt  to  form  only  a  single 
chelate  ring.  Many  complexes  of  alizarin  are  salts  rather  than  coordination 
compounds36- 37. 

Complexes  of   1-hydroxyanthraquinone   with  several  transition  metal 
ions  have  been  investigated38  and  formulated  as 


on  the  basis  of  analytical  and  spectral  data.  Beryllium  forme  similar  com- 
pounds with  naphthazarin  and  alkannin.  It  also  forms  a  polymer  with  ;i 
metal-ligand  ratio  of  1 :  l39. 

M    Dorta-Schaeppi,  Hurzeler,  and  Tread  well,  Helv.  Chim.  Acta,  34,  797  (1961). 

Liebhafsky  and  Winslow,  ./.  Am.  Chem.  Soc.,  60,   1776     L938  .  69.   1130     I" 
Flagg,  Liebhafsky,  and  Winslow,/.  -1///.  Chem.  80c.,  71,  363d    L94fl 
38.  Geyer  and  Smith,/.  Am.  Chem  80c  .64,  1649    L942). 

aderwood,  Toribara  and  Neuman,  ./.  Am.  Chm*.  Sue,  72,  .v>!»7  MdoOj. 


752 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Many  compounds  related  to  alizarin  are  of  commercial  importance  as 
dyes,  and  most  of  them  are  applied  in  conjunction  with  metal  salts.  Typical 
examples  are  anthragallol,  Alizarin  Cyanine  NS,  Anthracene  Blue  WR, 
Bordeaux  B,  and  Alizarin  Red  S.  More  complex  derivatives  such  as  Alizarin 
Irisol  R  (XIII)  are  also  useful  for  the  preparation  of  barium  and  aluminum 
lake  pigments. 

O      OH 


-CH: 


O 


N         | 
H        S03Na 
XIII 


The  presence  of  or^/io-dihydroxy  groups  in  other  classes  of  dyes  plays  an 
important  role  in  mordanting  operations  with  the  indication  that  complex 
formation  occurs  during  the  application  of  the  dyestuffs.  Gallocyanine 
(XIV),  a  member  of  the  oxazine  class  of  dyes,  is  applied  on  a  chrome  mor- 
dant. Among  the  xanthenes,  Gallein  (XV)  and  Coerulein  (XVI)  are  ap- 
plied on  chromed  wool. 

HO         OH 


HO 


HO- 


N(CH 


3/2 


HO 


fr00 


OONa 


TTV 


32: 


xvl 

The  thiazine  class  of  dyes  is  represented  by  Brilliant  Alizarin  Blue  3R 
(XVII),  which  yields  blue  chromium  lakes. 

SOT 


(CH,)2N^^S--Y-0H 
+      OH 


XVII 


DYES  WD  riGMEXTS 


753 


In  dyeing,  the  variations  in  color  or  shade  resulting  from  changes  in  the 
metal  ions  present  in  the  bath  or  on  the  fiber  suggest  the  formation  of 
coordination  compounds  rather  than  salts.  The  presence  of  the  ortho- 
dihydroxy  group  characterizes  all  members  of  each  class  which  are  useful  in 
mordanting  operations.  It  is  reasonable  to  assume  that  stable  coordination 
compounds  could  he  prepared  and  characterized  in  order  to  clarify  the  role 
of  complex  formation  in  the  dyeing  process. 

— COOH,  —OH  Substituted  Dyes 

Azosalicylic  acids  constitute  the  largest  class  of  com  ercial  dyes  which 
are  characterized  by  the  presence  of  — COOH  and  OH  groups  on  ad- 
jacent carbons  and  are  suitable  for  the  dyeing  of  fabrics  by  the  chrome 
process.  The  simpler  dyes  include  the  Alizarin  Yellows,  Ergansoga  Brown 
3R,  Diamond  Flavine  G,  and  Eriochrome  Flavine  A.  All  are  formed  by 
coupling  diazonium  salts  with  salicylic  acid. 

The  constitution  of  some  of  these  complexes  has  been  determined40. 
Alizarin  Yellow  2G  reacts  with  chromium  compounds  to  form  the  complex 
ion 


OgN 


XJX-  N=N  -Cj£  c  -  O 


which  has  been  isolated  as  the  chromium  (III)  salt.  Other  compounds  having 
different  Crrdye  ratios  have  also  been  prepared41.  One  of  these  has  been 
assigned  the  structure 


H20 
■N  =  N-/~)—  O— ;  Cr— 0^(~V-N=N-R 

6     h*°   a 


The  two  coordinated  water  molecules  may  be  replaced  by  ammonia. 

Drew  and  Fairbairn42  prepared  chromium  complexes  of  azosalicylic  acids 
containing  both  two  and  three  salicylic  acid  groups  per  chromium  ion. 
More  recently,  coordination  compounds  were  prepared  from  tetramminc 
COpper(II)     sulfate    and    aquopentammineeobalt(III)    chloride    and    the 

10.  Morgan  and  Main  Smith,  J.Chem.  Nor.,  121,  2866  (1922) ;  J   8ot    Dyt      ColourUts, 

41,  223  (1925). 

41.  Brass  and  Wirtnitzer,  Alii  X  congr.  intern,  ekim.,  3,  46  (1939). 

42.  Drew  and  Fairbairn,  J.  Chem.  Soc,  1939,  823. 


754  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

azosalicylic  acid  dye,  Mordant  Yellow  O43.  Two  ammonias  in  the  copper 
complex  were  replaced  by  one  dye  molecule,  while  all  of  the  ligands  in  the 
simple  cobalt  complexes  were  replaced  to  yield  a  complex  ion  [Co(dye)3]6-. 

Many  triphenylmethane  derivatives  contain  salicylic  acid  residues,  and 
lake  formation  has  been  indicated  by  several  workers44.  Xo  evidence  is 
available  regarding  the  structure  of  these  compounds45.  A  group  of  dyes 
known  as  the  Chromoxanes  is  especially  useful  for  application  with  chrome 
mordants.  By  heating  the  chromium  ammonium  salt  of  salicylic  acid  with 
the  dye  Eriochrome  Azurol  B  (XVIII),  a  compound  is  formed  which  will 
dye  blue  on  both  protein  and  animal  fibers46. 

In  the  xanthene  class,  compounds  such  as  Chromogen  Red  B  (XIX)  are 
useful  for  chrome  printing  on  cotton. 


O 
COOH 


(^VCOOH 

•xvnr  3EC 


Azine  dyes  can  also  be  adapted  for  chrome  printing  on  cotton  by  substitu- 
tion of  a  salicylic  acid  group  on  a  ring  nitrogen. 

Because  of  the  complexity  of  the  metal  derivatives  of  the  ortho-hy- 
droxy-carboxy  triphenylmethanes  and  azosalicylates,  it  is  difficult  to 
isolate  them  in  pure  enough  form  to  allow  study  of  their  structures.  Further 
work  is  needed.  Some  of  these  compounds  may  well  be  simple  salts,  but 
others,  having  either  the  — OH  or  — COOH  group  adjacent  to  the  azo 
bond,  afford  the  possibility  of  coordination  with  the  azo  group. 

Ortho-Substituted  Azo  Dyes 

Most  commercially  important  azo  dyes  are  characterized  by  the  follow- 

43.  Ref.  18,  p.  567. 

U    Middleton,  J.  Am.  Chem.  Soc,  48,  2125  (1926),  Hammett  and  Sottery,  J.  Am. 

Chem.  Soc.,  47,  142  (1925);  Corey  and  Rogers,  J.  Am.  Chem.  Soc,  49,  216,  (1927 
L£    Wttenberger,  Melliand  Textilber.,  32,  454  (1951).  See  ret.  s:>  and  88. 
L6    Ref.  18,  p.  731 


DYES  AM)  PIGMENTS  755 

ing  substituents47-48: 

v  Y 

X  Y 


-OH 

-OH 

-OH 

— COOH 

-OH 

— NH2 

-OH 

— H 

— NH2 

— H 

The  aromatic  nuclei  containing  the  or&o-substituents  may- be  either  ben- 
zene, naphthalene,  or  pyrazalone  rings.  The  latter  two  are  encountered 
most  frequently  in  the  patent  literature.  The  mordanting  metals  commonly 
used  are  chromium  for  wool  dyes  and  copper  for  cotton  dyes,  but  com- 
pounds of  manganese,  iron,  cobalt,  nickel,  vanadium,  tungsten,  molyb- 
denum, tellurium,  zirconium,  and  titanium  have  also  been  patented. 
Boyle49  has  reviewed  the  patent  literature  on  soluble  chromium  dyes  up  to 
1939.  A  more  recent  compilation  of  commercially  available  metal-complexes 
of  azo  dyes  includes  the  Benzo  Fast  Copper,  the  Chlorantine  Fast,  the 
Palatine  Fast,  and  the  Coprantine  dyes50. 

The  Palatine  Fast  and  Xeolan  colors  have  one  metal  atom  per  dye  mole- 
cule. Palatine  Fast  Blue  CGN  (XX)  may  be  formulated  as51 


SOaH 


,0- 


H03S— <V- 


■£%p 


2Z 


These  two  classes  of  dyes  include  fifty  individual  compounds  ranging  in 
shades  from  yellow  to  black52.  Most  of  the  colors  are  chromium  complexes, 
although  copper  was  once  employed  in  preparing  several  members  of  the 
group. 

Xeolan  Red  B  is  the  chromium  complex  of  Eriochrome  Red  B  (XXI) 
while  the  complex  formed  by  chromium  and  Eriochrome  Blue  Black  II 

17.  Knight,  ./.  Soc.  Dyers  Colourists,  66,  34  (1950). 

18.  Mackenzie.  Millson,  and  West,  Ind.  Eng.  Chem.,  44,  1017  (1952) ;  Pfitzner,  Angew. 

Chem.,  62,  242  (1950). 

49.  Boyle.  Am.  Dyestuff Reptr.,38,  741  (1939). 

50.  Specklin,  Teintez,  16,  451  (1950). 

51.  Valko,  Oesierr.  dun,.  Ztg.,  40,  405  (1937). 

52.  Ref.  18,  pp.  534-9. 


756  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

(XXII)  is  sold  as  Neolan  Blue  B.  Some  Palatine  Fast  colors  are  also  being 
marketed  for  leather  dyeing  under  the  name  Erganil  dyes. 

H3C-C-C-N=N-(      VSOjNa 

II     I               >— ( 
N   CO  \ / 

N 


6 


x  XI 


XZJT 


Knowledge  of  the  constitution  and  structures  of  the  metal  complexes  of 
azo  dyes  is  more  extensive  than  for  any  other  class  of  coloring  agents.  As 
early  as  1900,  an  alcohol-soluble  copper  compound  of  o-hydroxyazobenzene 
which  contained  two  azo  dye  molecules  for  each  copper  atom  was  re- 
ported53. Werner54  included  this  compound  in  his  newly  developed  theoty; 
however,  the  exact  formulation  of  the  azo  dye  lakes  was  not  attempted  until 
a  much  later  date  when  Morgan  and  his  students  initiated  a  systematic 
investigation55.  Eriochrome  Red  B  (XXI)  and  Palatine  Chrome  Black  6B 
each  contain  two  hydroxyl  groups  in  positions  ortho  to  the  azo  bond. 
With  Eriochrome  Red  B,  three  different  compounds  were  isolated;  these 
had  dye: metal  ratios  of  3:1,  3:2,  and  1:1.  Palatine  Chrome  Black  6B, 
HO-CioH6-N2-CioH5(OH)-S03H,  formed  two  lakes  having  dye: metal 
ratios  of  3:1  and  1:1.  Because  of  the  presence  of  the  sulfonic  acid  groups, 
the  ratios  are  not  representative  of  the  number  of  metal  ions  coordinated 
with  a  single  azo  group.  In  the  above  dyes,  there  are  three  azo  groups  for 
each  coordinated  metal  ion.  The  same  ratio  was  obtained  for  the  cobalt 
complex  of  an  o-amino,  o'-hydroxyazo  dye,  Metachrome  Brown  B.  These 
results  led  Morgan  to  conclude  that  only  one  hydroxy  group  was  included 
in  the  coordination  sphere  of  the  metal  ion.  The  error  in  his  interpretation 
resulted  from  the  presence  of  the  sulfonic  acid  groups  which  also  interacted 
with  the  metal  ammine  complexes  used  in  the  preparations. 

Drew  and  his  co-workers  may  be  credited  with  clarifying  the  structures 
of  the  azo  dye  complexes.  Copper  lakes  of  2-hydroxy-5-methylazobenzene, 
o-hydroxyazobenzene,  2-hydroxy-5 ,5'-dimethylazobenzene,  benezeneazo- 
i8-naphthol,  and  ra-tolylazo-/3-naphthol,  showed,  on  analysis,  a  dye: copper 
ratio  of  2:156.  All  of  the  compounds  were  anhydrous  and  did  not  add  or- 
ganic amines,  so  the  two  molecules  of  dye  in  each  compound  must  have 
formed  four  bonds  with  the  copper  ion,  thus  satisfying  its  normal  coordina- 

53.  Bamberger,  Ber.,  33,  1951  (1900). 
:.l.  Werner,  Ber.,  41,  2383  (1908). 

55.  Morgan  and  Main  Smith,  J.  Chem.  Soc.,  125,  1731  (1924). 

56.  Drew  and  Landquist,  ./.  Chem.  Soc.,  1938,  292. 


DYES  AND  PIGMENTS  7:»7 

tion  number.  The  general  structure  of  these  lakes  may  be  represented  as 


2 

Analogous  results  were  obtained  with  dyes  having  a  single  ortho-c&rboxy 
group,  except  for  a  marked  decrease  in  the  stability  of  the  complexes.  The 
dye: metal  ratio  was  the  expected  2:1,  but  dihydrates  also  formed,  and  the 
water  could  be  replaced  by  pyridine  or  aniline.  Since  or//?o- carboxy  and 
orMo-hydroxy  complexes  should  be  identical  with  respect  to  coordinative 
saturation,  it  is  difficult  to  understand  the  ability  of  the  former  to  add 
additional  donor  molecules. 

The  copper  lake  of  2,2'-dicarboxyazobenzene  (dye: metal  =1:1)  formed 
a  stable  monohydrate,  thus  satisfying  the  coordination  number  of  four  for 
the  copper  ion.  The  copper  derivatives  of  o-carboxybenzeneazo-p-cresol 
and  o-carboxybenzeneazo-jS-naphthol  also  gave  a  ratio  of  1:1  and  added 
one  molecule  of  either  pyridine  or  aniline.  The  o,o'-dihydroxyazo  and  azo- 
methine  dyes  formed  copper  complexes  containing  one  metal  ion  per  dye 
molecule  and  capable  of  giving  monopyridine  and  monoquinoline  deriva- 
tives. Pfeiffer's57  work  supports  that  of  Drew. 

Investigations  of  the  chromium,  iron,  nickel,  and  zinc  compounds  of 
mono-  and  di-or/Zio-substituted  azo  dyes  were  also  made42.  By  treating 
o-hydroxybenzeneazo-jS-naphthol  with  chromium(III)  chloride,  a  salt-like 
material,  Cr(dye)Cl,  containing  water,  was  formed.  It  could  be  converted 
to  a  compound  containing  non-ionic  chlorine  by  heating.  A  dipyridine 
derivative  was  also  prepared  in  which  chromium  has  its  preferred  coordina- 
tion number  of  six. 

The  chromium  lakes  of  2-hydroxy-5-nitrobenzeneazo-/3-naphthol  and 
'J-hydro\y-5-sulfobenzeneazo-/3-naphthol  gave  the  same  dye: metal  ratios 
as  those  which  had  only  ortho  substituents.  The  only  differences  noted  were 
in  the  solubility  of  the  complexes  and  the  high  water  content  of  the  solid 
material.  A  single  hydroxy  group  in  the  ortho  position  was  not  capable  of 
holding  a  chromium(III)  ion  in  stable  union  with  the  dye.  All  of  the  dihy- 
droxy  dyes  gave  the  expected  1:1  complexes  with  nickel' II  >.  zinc (II),  and 
iron(III).  In  addition,  the  iron(III)  lakes  gave  other  dye:metal  ratios  sim- 
ilar to  those  given  by  the  chromiumdll)  compounds.  The  nickel(II)  and 
zinc(II)  complexes,  like  those  of  copper (II),  formed  monopyridine  deriva- 

57.  Pfeiffer,  Hesse,  Pfitzner,  SchoU  and  Thielert,  •/ .  prakt.  Chem.,  149,  217    i 


758  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

tives,  thus  demonstrating  a  coordination  number  of  four,  the  azo  group 
taking  part  in  the  formation  of  one  coordinate  covalent  bond. 

With  o-carboxy,  o '-hydroxy  dyes,  nickel(II),  chromium(III),  and  iron 
(III)  compounds  containing  one  mole  of  dye  per  metal  ion  were  isolated. 
Copper  and  zinc  ions  combined  with  this  structure  to  give  salts,  one  of 
which  Drew  formulated  as  Cu[Cu(dye)2NH3]-6H20.  Analogous  aluminum 
lakes  were  also  prepared58  but  in  the  case  of  chromium,  definite  compounds 
of  monohydroxy  dyes  were  not  obtained.  The  lake  from  o-hydroxybenzene- 
azo-/3-naphthol,  formulated  as  [Al(dye)]Cl-5H20,  was  not  stable  to 
treatment  with  ammonium  hydroxide  or  potassium  chromate.  With  2'-hy- 
droxy^'-sulfobenzene^-azo-l-phenyl-S-methyl-l-pyrazol-S-one,  a  com- 
pound having  the  composition  Al(dye)-6H20  was  isolated. 

o-Hydroxybenzeneazo-/3-naphthol  gives  hydrated  V(dye)2  which  is 
readily  converted  to  VO(dye).  The  latter  adds  one  mole  of  pyridine,  and, 
like  the  other  vanadyl  complexes  which  were  prepared,  it  is  similar  to  the 
complexes  of  chromium(III)59. 

Beech  and  Drew60  investigated  the  effect  of  sulfonic  acid  groups  on  the 
coordinating  tendencies  of  the  o ,  o'-dihydroxyazo  dyes.  By  permitting 
copper(II)  chloride  to  react  with  2'-hydroxy-5'-sulfobenzeneazo-/3-naphthol, 
an  unusual  compound  was  formed: 


4H20  •  03S 


1  R/N>.^ 


^O  —  Cu+  S03-4H20 

2H20 

A  similar  dye,  containing  an  additional  sulfonic  acid  group  on  the  naphtha- 
lene ring,  may  be  metallized  with  copper(II)  chloride  to  give  a  compound 
which  has  been  assigned  a  structure  having  two  copper(II)  ions  coordinated 
to  a  single  azo  group. 

These  results  suggested  that  the  sulfonic  acid  groups  present  on  the  dye 
nucleus  serve  to  neutralize  part  of  the  charge  on  the  metal  ion.  The  latter, 
therefore,  does  not  require  both  hydroxyl  groups  for  neutralization,  and  it 
is  possible  for  two  metal  ions  to  be  attracted  to  the  vicinity  of  a  single  azo 

58.  Beech  and  Drew,  J.  Chem.  Soc,  1940,  603. 

59.  Drew   and  Dutton,  ./.  Chem.  Soc,  1940,  1064. 

60.  Beech  and  Drew,  ./.  Chem.  Soc.,  1940,  608. 


DYES  AND  PIGMENTS 


759 


Table  22.1.  Metal  Complexes  of  Azo  and  Azomethine  Dyes 


Dye 

Composition  of  Lake 

Configuration 

Benzeneazo-0-naphthol 

Co  (dye)  3 
Ni(dye), 

planar 

5-ChIoro-2-hy(lroxyljenzeneazo-/S- 

Co  (dye)-. 

betrahedral 

aaphthylamine 

\i.dye)OH 
Ni(dye)OH-II  0 

2  '-Hydroxy  henzal-2-hy  droxy-o-chloro- 

Co(dye)-2H,0 

letrahedral 

aniline 

\|m1v<>J-H20 

2'-(';irho\ybenzene-l  azo-1 -phenyl -3- 

Co(dye)-H,0 

tetrahedral 

methvlpvrazole-5-one 

Xi(dye)-H20 

group,  each  forming  a  coordinate  covalent  bond  with  one  of  the  nitrogen 
atoms.  Subsequent  evidence  fails  to  support  this  conclusion. 

The  chromium  complex  of  2'-hydroxy-3'-sulfo-5'-methylbenzene-4-azo- 
l-phenyl-3-methyl-l-pyrazol-o-one  and  related  d3res,  when  prepared  with 
disalicylato  chromic  acid  or  its  ammonium  salt,  contain  a  salicylaldehyde 
residue  which  completes  the  coordination  sphere  of  the  chromium  ion61. 
Similarly,  nickel  and  copper  complexes  of  formazyl  compounds  of  the  type 
shown  below  (XXIII)  add  a  mole  of  ammonia,  ethanolamine,  or  pyridine62. 


O-Cu 

i 


V^> 


0=C      \  N 


3xnr 

In  recent  years,  several  workers  have  made  use  of  magnetic  measure- 
ments and  complete  analyses  to  establish  the  composition  and  structure  of 
a  -cries  of  dyes  representing  a  variety  of  substituents.  Some  of  the  results 
are  summarized  in  Table  22. I63.  In  addition,  the  replacement  of  coordinated 
groups  from  cobalt  complexes  by  dye  molecules  was  examined64.  Table  22.2 
lists  some  of  the  compounds  obtained  in  this  investigation.  The  studies  also 
included  dyes  in  which  the  "ortho"  substituent  is  a  nitrogen  atom  in  a 
heterocyclic  ring65.  Simple  salts  were  used  in  most  cases,  so  the  coordination 
positions  remaining  unfilled  after  the  formation  of  the  metal-dye  complex 
contain  water  molecules  as  indicated  in  Table  22.3. 

Except  for  the  work  with  dyes  containing  sulfonic  acid  groups,  and  the 
behavior  of  organometallic  compounds  with  respect  to  the  azo  bond,  all 

61.  Shetty,  Helv.  Chim.  Acta,  35,  716    1962 

62.  Wisinger  and  Biro.  Helv.  chin,.  Acta,  32,  901    L940). 

63.  Caliis,  Nielsen,  and  Bailar,  ./.  .1//'.  Chem.  Sue.  74,  3461   (1! 

64.  Bailar  and  Caliis,  ./.  .1///.  Chem.  Sac.,  74,  6018  (1952). 

65.  Liu.  thesis.  University  of  Illinois,  1961. 


760  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Table  22.2.  Metal  Complexes  of  Azo  and  Azomethine  Dyes 

Dye  Metal  Salt  Composition  of  Lake 

Benzeneazo-j8-naphthol  [Co(NH3)6]Cl3  Co  (dye)  3 

2'-hydroxybenzal-2-hy-  Co(OAc)2-4H20  Co(dye)  and 

droxy-5-ehloroaniline  Co(dye)3 

[Co(NH3)6]Cl3  Co2(dye)3(NH3)3 

[Co(dien)2]Cl3  [Co  dien  dye]  CI 

[Co(NH3)5SCN]Cl2  Co2(dye)3(NH3)3 

Na3Co  (N02)  6  Co2  (dye)  3  (NH3)  3 

[Cr(NH3)6](N03)3  Cr2(dye)3(NH3)2H20 

Zn(OAc)2-2H20  Zn(dye) 

ZnCl2  Zn(dye) 

evidence  indicates  that  the  azo  group  occupies  only  one  of  the  coordination 
positions  available  in  the  sphere  of  a  metal  ion.  Consideration  of  this  fact 
is  important  in  the  choice  of  other  coordinating  agents  which  might  be 
added  to  dye  baths,  or  in  evaluating  interactions  between  metallized  dyes 
and  fibers. 


Miscellaneous  Dyes 

Phthalocyanines.  The  phthalocyanines  constitute  an  important  series 
of  fast  blue  to  green  pigments66.  Although  earlier  workers  had  apparently 
prepared  a  copper  phthalocyanine,  it  was  the  excellent  work  of  Linstead 
and  his  students67  which  resulted  in  a  complete  picture  of  the  structure 
and  properties  of  this  new  chromophore.  The  work  has  since  been  confirmed 
by  the  x-ray  studies  of  Robertson  and  others68. 

The  structure  of  the  phthalocyanines  was  found  to  be  similar  to  that  of 
porphin,  the  fundamental  nucleus  of  chlorophyll  (page  74)  and  hemin 
(page  74).  The  phthalocyanine  nucleus  may  be  derived  by  replacing  the 
methine  groups  by  nitrogen  atoms.  The  products  are  known  as  azaporphins. 
All  attempts  to  prepare  the  simple  azaporphins  appear  to  have  failed. 

The  phthalocyanines  have  a  coplanar  structure  and  are  capable  of  occu- 
pying four  coordination  positions  and  neutralizing  two  charges  of  a  metal 
ion.  The  stability  of  complexes  of  the  chromophore  has  been  demonstrated 
by  preparing  derivatives  of  more  than  twenty  elements.  These  include 
representatives  of  each  group  of  the  periodic  table.  Divalent  metals  dis- 

66.  For  reviews,  see:  Dahlen,  Ind.  Eng.  Chem.,  31,  839  (1939) ;  Haddock,  J.  Soc.  Dyers 

Colourists,  61,   68   (1945);  Haddock  and  Linstead,    "Thorpe's  Dictionary  of 
Applied  Chemistry,"  p.  617,  4th  ed.,  Vol.  IX,  London,  Longman's. 

67.  Linstead  et  al.,  J.  Chem.  Soc,  1934,  1016,  1017,  1022,  1027,  1031,  1033;  1936,  1719,* 

1725,  1737,  1739,  1744;  1937,  911,  922,  929,  933;  1938,  1157;  1939,  1809,  1820;  1940, 
1070,  1076,  1079;  Brit.  Pat.  389,842  (1933) ;  390,148  (1933) ;  410,814  (1934) ;  441,332 
(1936);  Dent,  J.  Chew.  Nor.,  1938,  1. 
88.  Robertson,  ./.  Chem.  Soc,  1935,  615;  1936,  1195,  1736;  1937,  219;  1940,  36;  Ender- 
mann,  Z.  physik.  Chem.,  190,  129  (1942). 


DYES  .l.\7>  PIGMENTS 


761 


Table  22.3.  Metal  Complexes  oi   Axo  Dyes 


1  >vo 

-Pyridylaio-0-naphthol 


a-Pyridylazoresorcinol 


(o-Carboxyaiobenzene)-o'-chloroace- 
toacetanilide 


-Carboxybenzene-4-azo-l -phenyl -3- 
methylpyrazol-5-one 


Prepared  from  [Co(NH3)6]Cl; 


Composition  of  Lakes 

[Cu  dye  H,OJNOi 

[Cu  dye]NO, 

[Ni  dye  H,0]N03 

[Ni  dye]N03 

ICo(dye)2]Cl* 

[Cr(dye),lNO 

[Cu  dye  H20] 

[Cu  dye] 

H2[Ni(dye)2]-H20 

Co[Co(dye)>]J-3H20 

[Cr(dye)2] 

[Cr  dye  (H20)3]-3H20 

[Cu  dye  H20] 

[Cu  dve] 

H2[Ni(dye)2] 

H[Cr(dye)2l 

[Cu  dye] 

[Ni  dve  (H20)2] 

[Co  dye  H20] 

[Co  dye  (H20)3]* 

[Co(dye)2]-2H,0* 


place  the  two  hydrogen  atoms  to  form  a  nonionic  complex.  Trivalent  ions 
form  compounds  of  the  type  (Phthalocyanine  MX),  while  tetravalent  ions 
give  (Phthalocyanine  MX2)  compounds.  The  metal  phthalocyanine  may 
be  used  directly  or,  in  some  cases,  the  metal  may  be  removed  by  treatment 
with  acid. 

Although  a  great  many  phthalocyanines  have  been  synthesized,  the 
copper  derivative  is  the  most  important  and  is  sold  commercially  in  the 
Monastral  Fast  Blue,  Heliogen  Blue,  and  Vulcan  Blue  series.  These  arc 
valuable  because  of  their  brilliant  shades,  high  tinctorial  strength,  insolu- 
bility in  water,  and  stability.  In  the  usual  organic  solvents,  they  vary  from 
total  insolubility  to  very  slight  solubility.  They  are  soluble  in  most  strong 
acid-  but  reprecipitate  upon  dilution.  The  pigments  are  relatively  stable  to 
heat,  light,  and  chemical  reagents.  The  pigment  properties  have  Keen  suc- 
-lully  modified  by  halogenation  and  sulfonation.  The  soluble  sulfonated 
phthalocyanine-  thus  produced  are  somewhat  less  stable  than  the  insoluble 
pigments.  Helberger*  has  shown  that  some  metal  phthalocyanines  exhibit 
brilliant  chemiluniinescence  when  oxidized  under  certain  conditions.  The 
phthalocyanine-  have  numerous  applications  wherever  coloring  materials 
are  used. 


f>".  Helberger,  NaturwUsenschaften,  26,  316  (1938). 


762 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Other  Nitrogen -donor  Dyes.  Patents  have  been  issued  on  dyes  from 

2 , 4-diarylpyrroles  such  as  2 , 2' ,  4 , 4/-tetraphenylazadipyrromethine 

— — Ph  Ph- 


-N= 


■\N^\ 
Ph        H  Ph 

This  compound  forms  metal  complexes  similar  to  those  of  the  phthalo- 
cyanines70. 

Kunz  prepared  the  copper  and  iron  compounds  of  indigo71.  The  structure 
of  the  copper  compound  has  been  given  as72 


a: 


V 
Cur-  6 

2 


Drew  and  Kelly73  obtained  highly  colored  metallic  compounds  of  dithio-/3- 
isoindigo. 

The  primary  application  of  these  results  has  been  in  the  solubilization  of 
indigo  and  other  vat  dyes  through  complex  formation.  In  the  reaction,  the 
active  groups  are  the  carbonyl  functions74. 

Sulfur  Containing  Dyes.  These  dyes  are  probably  the  least  under- 
stood from  the  point  of  view  of  the  structure  of  the  organic  compounds 
present  in  the  commercial  products;  however,  the  extensive  use  of  metal 
salts  in  the  preparation  of  these  materials  suggests  that  coordination  phe- 
nomena are  involved75.  Thionyl  Purple  2B  forms  bordeaux  red  lakes  when 
copper,  cobalt,  or  nickel  salts  are  added76.  Structures  have  been  proposed 
for  several  sulfur  dyes  including  Pyrogene  Green77 


H03S 


SO3H 


s-o 


CuS 


-■x 


70.  Rogers,  J.  Chem.  Soc,  1943,  590,  596,  598;  British  Patents  562,754-61  (1950)  and 

others. 

71.  Kunz,  Ber.,  55,  3688  (1922). 

72.  Kuhn  and  Machemer,  Ber.,  61,  118  (1928). 

73.  Drew  and  Kelly,  J.  Chem.  Soc,  1941,  625,  630,  637. 

74.  Ref.  18,  pp.  1047-48. 

75.  Ref.  18,  pp.  1063-4,  1071  ff. 

76.  Vlies,  J.  Soc.  Dyers  Colourists,  29,  316  (1913). 

77.  Fierz-David  et  al.,  Helv.  Chim.  Acta,  15,  287  (1932);  16,  585  (1933);  J.  Soc.  Dyers 

Colourists,  51,  50  (1935);  Naturwissenschaften,  20,  945  (1932). 


DYES  AND  PIGMENTS  763 

Copper,  nickel,  and  cobalt  lakes  of  two  0-mercaptoazo  compounds  contain- 
ing the  grouping 

SH  HO 

show  a  dye: metal  ratio  of  2:1TO.  The  sulfur-containing  dyes  offer  a  fertile 

field  of  research  for  the  coordination  chemist. 

The  Dye-Metal-Fibee  Interactions* 

In  practice,  the  application  of  a  dye  involves  both  physical  and  chemical 
changes.  The  physical  phenomena  involved  appear  to  be  independent  of 
the  type  of  fiber,  while  chemical  changes  are  related  to  the  structure  of  the 
material  being  dyed.  Textile  fibers  may  be  divided  into  four  classes  on  the 
basis  of  their  chemical  structure:  cellulose  and  rayons;  proteins,  which 
include  wool  and  silk;  synthetic  polyamides  which  are  chemically  related 
to  the  proteins;  and  miscellaneous  polymers. 

Cotton,  which  is  nearly  pure  cellulose,  may  be  dyed  by  colors  having  the 
chromophore  in  the  anion.  The  principal  attraction  involves  hydrogen  bond- 
ing with  the  possibility  of  some  electrostatic  forces  if  the  hydroxyl  groups 
of  the  cellulose  have  some  acidic  character.  The  direct  cotton  dyes  are  often 
o-hydroxy-  or  o-aminoazo  dyes  in  which  chelation  assists  in  the  formation 
of  hydrogen  bonds  between  the  dye  molecules  and  the  cellulose  chain: 

H  — O  —  Cellulose 

6-H 


This  bonding  implies  that  chelation  of  the  proton  with  the  azo  group  in- 
creases the  accessibility  of  the  electron  pair  involved  in  the  formation  of 
the  hydrogen  bond  with  the  cellulose.  The  chelation  of  a  metal  ion  would 
probably  result  in  the  formation  of  a  more  stable  chelate  ring  but  would 
also  introduce  the  probability  of  delocalizing  the  electron  pair  as  well  as 
converting  the  dye  to  a  cation.  Evidence  suggests  that  the  presence  of  a 
metal  ion  results  in  the  formation  of  a  chemical  bond  between  it  and  the 
cellulose  group-.  Systems  containing  [Cu(NH3)4]'f^  show  a  decrease  in  pll 
upon  addition  of  polyhydroxy  compounds  such  as  cellulose  or  sucrose79.  The 

78.  Burawoy  and  Turner,  ./.  Chem.  Soc.t  1952,  1286. 

*  See  Ref.  18,  Chap.  VI,  XI. I,  p.  567;  Race,  Rowe,  and  Speakman,  •/.  8oe.  Dyera 
Colourists,  62,  372  (1946);  Giles,  ./.  Soe.  Dyera  CoUmrists,  60,  303  (1944);  Justin 

Mueller.  Teintex,  15,  r>7  (1950). 

Vrkhipov  and  Kharitonova, ./.  Appl.  Chem.  U.S.S.R.  24,  733  (1961   ;  •/.  Sot  D 

Colourists,  67,  471  (1951). 


764  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

following  reaction  has  been  suggested: 

I  H— C— O 


H— C— OH 

+  [Cu(NH3)4](OH)2 
H—  C— OH 


H— C— O 


Cu(NH3)4  +  2H20 
/ 


Rayons,  which  are  derivatives  of  cellulose,  may  be  classified  into  two 
groups:  nitro  rayon,  cuprammonium,  viscose;  and  cellulose  acetate.  The 
first  group  may  be  dyed  in  the  same  manner  as  cotton.  Cellulose  acetate, 
however,  is  dyed  by  materials  which  dissolve  in  the  fiber.  Most  cellulose 
acetate  dyes  are  sparingly  soluble  in  water  and  are  handled  as  dispersions. 

Wool  and  silk  have  similar  dyeing  properties  since  both  consist  of  pro- 
tein chains.  Wool  contains  sulfur  in  the  form  of  cystine  and  as  disulfide 
linkages  between  the  keratin  residues.  The  latter  may  also  be  joined  by 
salt  groups.  Wool  is,  therefore,  capable  of  reacting  with  both  anionic  and 
cationic  dyes. 

In  the  dyeing  of  wool,  as  in  the  case  of  cotton,  hydrogen  bonding  seems 
to  be  involved.  Much  evidence  has  also  been  found  for  direct  chemical 
combination  between  metal  ions  and  protein  fibers.  Dichromate  ions  are 
absorbed  and  are  reduced  to  chromium  (III)  ions  on  heating.  The  combina- 
tion of  chromed  wool  with  a  dye  may  involve  chemical  bonding,  but  many 
chrome  dyes  have  no  salt  or  chelating  groups,  and  the  interaction  probably 
involves  adsorption.  Where  lake  formation  with  a  dye  is  possible,  it  is  nec- 
essary to  have  the  chromium  present  as  the  chromium  (III)  ion80.  A  syste- 
matic investigation  of  the  interaction  of  chromium  complexes  with  collagen, 
collagen  with  the  amino  groups  blocked,  silk  fibroin,  and  polycaprolactam 
led  to  the  conclusion  that  cationic  chromium  reacts  with  carboxy  groups 
wiiile  chromium  anions  react  with  amino  groups  in  protein  fibers81.  Others 
have  questioned  these  results82,  but  Shuttleworth83  appears  to  have  re- 
solved the  conflicting  data  by  examining  the  adsorption  of  eighteen  chro- 
mium complexes  on  amino,  sulfonic  acid,  and  carboxylic  resins.  The  chief 
mechanism  is  coordination  of  the  complexes  with  carboxy  groups;  it  can  be 
related  to  the  dissociation  constants  of  the  ligands. 

Wool  absorbs  nickel  ions  from  solutions  of  [Ni(NH3)4](OH)2  with  no 
increase  in  the  nitrogen  content  of  the  wool84.  The  coordinated  ammonia 

80.  Gaunt,  J.  Soc.  Dyers  Colourists,  65,  429  (1949). 

81.  Strakhov,  J.  Appl.  Chem.  U.S.S.R.,  24,  142  (1951);  ./.  Soc.  Dyers  Colourists,  67, 

292  (1951). 

82.  Gustavson,  J.  Soc.  Leather  Trades  Chem.,  36,  182  (1952). 

83.  Shuttleworth,  ./.  Amer.  Leather  Chemist's  Assoc,  47,  387  (1952). 

84.  Bell  and  Whewell,  ./.  Soc.  Dyers  Colourists,  68,  299  (1952). 


DYES  AND  PIGMENTS 


765 


molecules  may   be  replaced   by  the  amine  groups  of  the   wool;  however, 

mollification  of  the  amine  groups  does  not  decrease  the  amount  of  nickel 
ion  absorbed  although  it  dot's  decrease  the  rate  of  the  process.  Similar  modi- 
fications of  the  disulfide  and  earboxy  groups  have  little  effect  on  the  ad- 
sorption of  nickel  ion,  and  it  appears  that  main  chain  >(()  and  >NH 
groups  are  involved. 

Another  investigation  of  the  interaction  between  metal  ions  and  wool 
indicates  that  bonding  is  dependent  on  the  nature  of  the  metal  ion  in- 
volved*. Wool  was  treated  with  salts  of  lead,  cadmium,  zinc,  copper,  iron, 
bismuth,  and  mercury.  Upon  treatment  with  ions  of  the  first  four  metals, 
the  cystine  content  of  the  wool  decreased  and  the  nitrogen  content  of  the 
hath  increased.  X-ray  studies  suggested  that  the  metal  ions,  except  perhaps 
copper,  were  present  in  the  wool  as  metal  sulfides.  In  all  cases,  the  metal 
content  of  the  wool  was  in  excess  of  the  noncystine  sulfur  present,  and  some 
of  the  metal  must  have  been  bound  by  functional  groups  of  the  keratin. 

The  dyeing  of  synthetic  fibers  has  presented  many  problems  which  vary 
with  the  chemical  nature  of  the  materials86.  A  survey  has  been  made  of  the 
dyeing  methods  suitable  for  three  typical  products:  "Nylon,"  "Orion" 
acrylic  fiber,  and  "Dacron"  polyester  fiber87.  Of  the  three,  "Nylon"  com- 
pares favorably  with  wool  in  ease  of  dyeing. 

A  -tries  of  metal  complexes  of  azo  dyes,  known  as  the  Perlon  Fast  colors, 
has  been  developed  for  the  dyeing  of  Perlon,  a  nylon-type  fiber88.  Examples 
are  Perlon  Fast  Yellow  G  (XXIV)  and  Perlon  Fast  Red  3BS  (XXV). 


y 


N 


// 


N 


c  =  c 


/ 


CH3-C         N— C6H5 
N 


xxrv 


Nd 


Schoberl,  MelliandTextilber.,**,  1(1962  \J.Soc.Dyi      Colov      •     68. -'_'«,    19* 

86.  Baumann,  Am.  D  :■  ptr.,  41,  P.  153  (1952). 

B7.  Turnbull,  Am.  Dyestuff  Rept      41,  P.  7.5,  P.  82     L962  . 

88.  Anacker,  MeUiand  TextiWer.,  30.  256  (1949);  Knight, ./.  Soc.  Dyers  Colourists,  66, 

169  (1950). 


766 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


NdL 


X22 

Nylon  may  be  chromed  prior  to  the  addition  of  the  dye;  whereas  wool 
reduces  the  dichromate  to  chromium  (III)  on  heating,  this  reduction  does 
not  occur  on  nylon  fibers  without  the  addition  of  a  reducing  agent.  The  re- 
duction is  catalyzed  by  the  presence  of  a  dye  which  forms  a  complex  with  the 
reduced  chromium  ion.  Once  the  chromium  ion  has  been  fixed  on  the  nylon, 
chelation  with  a  lake-forming  dye  follows.  If  nylon  is  treated  directly  with 
CrF3  or  Cr2(S04)3 ,  there  is  a  strong  tendency  for  the  metal  ions  to  migrate 
into  the  dye  solution  and  form  insoluble  complexes. 

Undoubtedly,  the  fixation  of  chromium  on  a  fiber  is  more  than  a  simple 
interaction  between  chromium(III)  ions  and  donor  groups.  The  necessity  for 
starting  with  an  oxyanion  suggests  the  occurrence  of  an  olation-type  reac- 
tion with  chains  of  — Cr — 0 — Cr — O —  groups  being  bonded  to  evenly 
spaced  groups  on  the  material  being  dyed.  This  would  result  in  the  proper 
distribution  and  bonding  of  chromium  atoms  prior  to  their  reduction  to  a 
lower  oxidation  state. 

The  principle  of  impregnating  a  synthetic  fiber  with  copper  ion  prior  to 
application  of  a  dye  has  proved  very  useful  in  the  dyeing  of  acrylonitrile 
fibers  such  as  "Orion,"  "Dynel,"  and  "Acrilan."  The  copper(I)  ions  form 
coordinate  covalent  bonds  with  the  nitrile  groups,  and,  upon  addition  of 
the  dye,  probably  form  copper-dye  linkages.  This  suggests  that  the  copper 
ions  must  be  spaced  at  intervals  in  order  to  permit  discrete  bonding  with 
the  larger  dye  molecules.  In  connection  with  this  point,  it  may  be  noted 
that  "Dynel,"  which  contains  only  40  per  cent  acrylonitrile,  is  dyed  more 
effectively  by  this  process  than  is  the  100  per  cent  acrylonitrile  polymer, 
"Orion"89. 

Although  copper(I)  salts  may  be  added  directly,  it  is  preferable  to  use  a 
i-opper(II)  salt  and  reduce  it  with  hydroxylamine  hydrogen  sulfate.  The 
use  of  the  hydrochloride  tends  to  retard  the  process.  This  may  be  due  to 
the  formation  of  chloride  compounds  with  the  copper(I)  ion.  The  copper 
may  also  be  applied  in  the  form  of  a  salt  of  an  acid  or  a  direct  dye  having 

89.  Douglas,  ./.  Soc.  Dyers  Colourists,  67,  133  (1951);  Hatfield  and  Sharing,  J.  Soc. 
Dyers  Colourists,  64,  381  (1948). 


DYES  AND  PIGMENTS  767 

one,  bill  not  more  than  one,  sulfonic  or  carboxy  group  in  the  molecule'"1. 
From  this  brief  discussion  of  the  dye-metal-fiber  interactions,  it  appears 
certain  that  much  work  remains  to  he  done  to  insure  a  more  complete  under- 
standing o\  the  chemical  reactions  which  are  taking  place.  'The  information 
concerning  dye-metal  interactions,  while  far  from  complete,  is  sufficiently 
advanced  to  enable  reasonable  predictions  of  the  behavior  of  metal  ions 
with  numerous  classes  of  dyes.  A  more  concentrated  effort  in  the  direction 
of  metal-fiber  bonding  seems  indicated. 

90.  Blaker  and  Laucius,  .1///.  Dyestuff  Reptr.  t  41,  I'.  39  I L952);  Fronmuller,  .1///    l>n<  - 
stuf  Reptr.,  41,  1'.  578  ^  L962);  Szlosberg, Am.  Dyestuff  Reptr., 41,  P.  510  (1952); 

Field  and  Fremon,  Text.  Research  ./..  21,  531  (1951);  Field,  Am.  Dyestuff  Reptr., 
41,  P.  475  (1952). 


AO.   Water  Softening  Through  Complex 

Formation 


Roy  D.  Johnson 

American  Embassy,  Melbourne,  Australia 

and 

Clayton  F.  Callis 

Monsanto  Chemical  Co.,  Dayton,  Ohio 

Water  softening  may  be  defined  as  the  process  of  effectively  removing 
ions,  such  as  calcium  and  magnesium,  which  cause  the  precipitation  of 
soaps.  It  is  evident  that  water  softening,  thus  defined,  is  somewhat  simpler 
than  water  conditioning  in  boiler  systems1  where  heating  and  evaporation 
complicate  the  precipitation  problem.  The  general  methods  used  for  water 
softening  are  distillation,  precipitation,  ion  exchange,  and  the  effective 
removal  of  ions  from  solution  by  the  formation  of  soluble  complexes.  This 
discussion  will  be  confined  to  softening  of  water  through  complex  formation. 

This  phenomenon  of  utying-up"  alkaline  earth  ions  in  soluble  complex 
ions,  and  thus  preventing  the  formation  of  precipitates,  is  generally  termed 
"sequestration"2.  The  tests  commonly  used  for  determining  the  sequester- 
ing ability  of  a  "sequestering  agent"  depend  upon  the  prevention  or  diminu- 
tion of  precipitation  as  measured  by  nephelometry  or  by  the  formation  of 
soap  foams3, 4.  The  weight  of  sequestering  agent  per  unit  quantity  of  multi- 
valent positive  ion  needed  to  prevent  the  precipitation  of  alkaline  earth 
salts  under  operating  conditions  is  known  as  the  sequestration  value.  Ma- 

1.  Schwartz  and  Munter,  Ind.  Eng.  Che?n.,  34,  32  (1942). 

2.  Hall,  U.  S.  Patent  1,956,515  (1934) ;  Reissue  19,  719  (1935). 

3.  Van  Wazer,  ''Encyclopedia  of  Chemical  Technology,"  Vol.  XI,  pp.  403-41.  New- 

York,  Interscience  Publishers,  Inc.,  1953. 
1.  For  example,  Andress  and  Wiist,  Z.  anorg.  allgem.  Chem.,  237,  113  (1938);  241,  196 
(1939) ;  Rudy,  Schloesser  and  Watzel,  Angew.  Chem.,  53,  525-31  (1940) ;  Hafford, 
Leonard,  and  Cummins,  Ind.  Eng.  Chem.,  Anal.  Ed.,  18,  411-15  (1946);  Miles 
and  Ross,  ./ .  Amer.  Oil  Chem.  Soc,  24,  23  (1947);  Davies  and  Monk,  J.  Chem. 
Sac.,  1949,  413-22.  Also,  private  communication  from  R.  K.  Skaar,  Food  Ma- 
chinery and  Chemical  Corporation. 

768 


WATER  SOFTENING  THRO  I  GH  COMPLEX  FORMATIOA  769 

terials  useful  as  sequestering  agents  include  the  chain  or  polyphosphates 
ami  certain  polyamino  acids.  The  phytates  have  also  been  suggested. 

THE   ClIAIX    OB    POLYPHOSPHATES   5_,° 

The  phosphates  most  commonly  used  as  sequestering  agents  for  water 
Boftening  are  the  sodium  salts  of  the  chain  phosphates,  i.e..  sodium  acid 
pyrophosphate,  tetrasodium  pyrophosphate,  sodium  tripolyphosphate,  and 
the  sodium  salts  of  the  low  and  high  molecular  weight  glassy  phosphates11. 

Polyphosphates,  One  of  Three  Groups  of  Condensed  Phosphates 

On  the  basis  of  the  present  evidence  (re  viewed  in  reis,  3, 8,  and  10)  includ- 
ing x-ray  studies  of  crystalline  phosphates  and  physical-chemical  studies  of 
solutions  of  the  phosphates,  it  is  believed  that  the  so-called  "condensed 
phosphates"  are  built-up  by  sharing  oxygen  atoms  between  structural  units, 

each  unit  consisting  of  a  tetrahedral  grouping  of  four  oxygen  atoms  around 
a  central  phosphorus  atom.  It  has  been  shown3  that  the  condensed  phos- 
phates can  be  conveniently  divided  into  three  groups:  the  chain,  the  ring, 
and  the  branched  phosphates,  depending  on  the  number  of  shared  oxygens 
per  tetrahedron. 

The  chain  phosphates  are  generally  called  polyphosphates  and  consist  <>i 
unhranched  P-O-P  chains.  The  ring  phosphates  consist  of  simple  rings  of 
interconnected  phosphorus  and  oxygen  atoms,  and  are  included  in  the  class 
of  metaphosphates.  At  present  only  the  six-  and  eight -membered  rings  are 
known  (trimeta-  and  tetrametaphosphate).  The  branched  phosphates,  often 
referred  to  as  ultraphosphates,  include  structures  in  which  one  or  more 
P04  groups  share  oxygen  atoms  with  three  neighboring  groups.  These 
branched  phosphates,  on  dissolution  in  water,  are  rapidly  converted  into 
groups  in  which  no,  one,  or  two  oxygens  are  shared12.  This  means  that  only 

5.  Graham,  IJroc.  Royal.  Sue.,  123,  253  (1833). 

6.  Partridge.  Hicks  and  Smith,  J .  Am.  ('hem.  Soc,  63,  454  (11141  I;  Morey  and  Inger 

Bon,  Am.  ./.  Set.,  242,  1     1944  . 

7.  Quimby,  Chem.  Revs.,  40,  141  (1947);  "Thorpe's  Dictionary  of  Applied  Chem 

istry/'  4th  ed.,  Vol.  '.'.  p.  508,  New  York,  Longmans,  Green  and  Co.,  1949; 
Toplej  ,Qua  I    /,'•   t  ,3,345    194 

8.  Van  Wazer,  et  oi.,  J  Soc.,  72,  639, 644, 647, 906    1950  ;  75, 1563    1" 

!  limb;    •/    PI  58.  603    L954 

t His .  Van  Wazer  and  Aryan,  <'h>m.  Revs.,  54,  777    1964 

11.  ••Sodium  Phosphates  for  Industry,"  Catalog  of  the  Monsanto  Chemical  Com 

pany.  Lnorganic  Chemicals  Division;  "Victor  Chemicals/1  Catalog  of  Victor 
Chemical  Works;  "BlocksoD  Chemicals/'  Catalog  of  the  Blockson  Chemical 
Company;  "Westvaco  Chemicals/1  Catalog  of  Westvaco  Chemical  Division, 
Food  Machinery  and  Chemical  Corporation. 

12.  Pfanstiel  and  Her.  ./.  A  -         74,      W  64     1952  ;  Straus*    Smith  and 

Winem  N      135  10    196 


770  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

orthophosphates,  simple  rings,  or  unbranched  chains  are  present  a  short 
while  after  dissolution,  and  of  these  only  the  unbranched  chains  or  poly- 
phosphates are  effective  in  alkaline  earth  ion  sequestration. 

The  chain  phosphates  constitute  a  homologous  series  of  polymeric  com- 
pounds represented  by  the  formula  M(n+2)P«0(3n+i)(l  <  M20/P205  <  3), 
in  which  M  represents  an  equivalent  of  metal,  and  n  is  the  number  of  phos- 
phorus atoms  in  the  chain.  Thus,  the  monomer  is  the  orthophosphate  (not 
one  of  the  phosphates  which  softens  by  sequestration),  the  dimer  is  the 
pyrophosphate,  and  the  trimer  is  the  triphosphate  or  tripoly phosphate.  In 
the  sodium  system,  higher  crystalline  polymers  are  not  known,  and  Par- 
tridge, Hicks,  and  Smith6  have  shown  from  an  equilibrium  phase  diagram 
that  triphosphate  is  the  only  crystalline  compound  between  the  pyro-  and 
metaphosphate  compositions.  However,  thermal  evidence  for  the  forma- 
tion of  a  crystalline  lead  tetraphosphate  has  recently  been  published13,  and 
all  possible  chain  lengths  up  to  several  hundred  are  present  in  solutions  of 
the  glassy  phosphates80. 

The  sodium  phosphate  glasses,  introduced  as  water  softeners  by  Hall2 
in  1932,  were  the  first  phosphates  used  in  this  application.  They  are  pre- 
pared by  quenching  sodium  oxide -phosphoric  oxide  melts  in  the  composi- 
tion range,  1  <  Na20/P205  <  1.34.  An  infinite  number  of  products  may  be 
produced  within  this  range.  It  has  been  shown  from  solubility  fractionation 
and  end-group  titration  studies80  that  in  aqueous  solution  these  glasses 
exhibit  a  size  distribution  of  linear  molecule-ions,  the  average  of  which  is  a 
first-order  function  of  the  Na20/P206  mole  ratio,  i.e.,  theoretically, 

Na2Q  +  H2Q  =  n  +  2 
P205  n 

where  Na20  ^>>  H20  and  n  is  the  number-average  number  of  phosphorus 
atoms  in  the  chain.  As  n  approaches  infinity,  the  general  formula  of  the 
chain  phosphates  approaches  that  of  the  metaphosphate  composition, 
MnPn03n  .  This  metaphosphate  composition  is  the  limiting  composition  for 
both  the  chain  and  branched  regions,  as  well  as  being  the  empirical  com- 
position for  the  ring  compounds.  Actually,  high-molecular  weight  chain 
compounds  with  empirical  compositions  analytically  indistinguishable  from 
that  of  the  ring  compounds  are  known,  and  the  thermal  interrelationships 
of  a  number  of  crystalline  varieties  of  this  metaphosphate  composition  have 
been  studied7.  These  crystalline  and  glassy  chain  phosphates,  with  com- 
positions near  that  of  the  metaphosphate,  are  not  used  in  commercial  wTater 
softening  primarily  because  of  undesirable  physical  properties,  such  as  slow 
rate  of  dissolution.  The  Na20/P206  mole  ratios  generally  chosen  for  the 
commercial  glasses  are  1.11  and  about  1.33  for  the  high-  and  low-molecular 

13.  Osterheld  and  Langguth,  J.  Phys.  Chew.,  59,  76  (1955). 


WATER  80FTENINQ  THROUGH  COMPLEX  FORMATION  771 

Table  23.1.   Relative  Sewi  k^tkhi.nc   Ability   ok  SevbraL    POLYPHOSPHATES.    \r 

Room   TXMPKBATUBl 


Grams  of  Ca!l  ncr      Cr.uib  of  M-*1  per   Or.tnis  of  Iron'    Pel 

100  Grams  ol  100  Grams  ol  lOOGrami 

Polyphosphate  Phosphate  Phosphate  Phosphate 

Sodium  triphosphate  13.4  6.4  0.184 

■odium  phosphate  glass  with  18.5  0.092 

NasO/PjOi  =  ca.  1.3 
Bodium  phosphate  glass  with  19  :,  2.9  0.031 

NasO/PsOi  =  1.1 
Fetrasodium  pyrophosphate  4.7  8.3  0.273 

\t  optimum  pH  of  10  to  11.  See  reference  4<l  for  details. 

b  pH  adjusted  to  10,  soap  present  *. 

1  Ferric  sulfate  solution  mixed  with  phosphate  in  sodium  sesquicarbonate  solution 
followed  l>y  addition  of  hydrogen  peroxide48. 

weight  glasses,  respectively11.  The  average  number  of  phosphorus  atoms  in 
the  chains  can  be  estimated  from  equation  (1).  Glasses  with  a  1.11  ratio 
have  an  average  chain  length  of  about  14,  and  those  with  the  higher  ratio 
have  an  average  chain  length  of  approximately  6.  Some  products  of  inter- 
mediate composition  are  also  marketed. 

The  Sequestering  Action  of  the  Polyphosphates 

The  addition  of  a  polyphosphate  to  water  containing  calcium  or  mag- 
nesium ions  leads  to  precipitation  of  calcium  or  magnesium  phosphate. 
This  precipitation  continues  until  an  excess  of  the  phosphate  has  been 
added.  Then  the  precipitate  is  peptized,  dispersed,  and  redissolved  in  a 
sequestering  action.  The  sequestering  ability  of  the  phosphates  is  dependent 
upon  many  factors,  the  principal  ones  of  which  are  discussed  below. 

Factors   Affecting  the   Sequestering  Ability   of  Polyphosphates 

Nature   of  the   Polyphosphate    (or  Precipitating   Anion)    and   the 
Metal  Ion 

Measurements  of  the  sequestering  ability  of  the  polyphosphates  give 
widely  different  results  depending  upon  the  anion  used  (sometimes  a  pre- 
cipitating anion  other  than  phosphate  is  added),  the  metallic  ion  and  the 
pH.  One  common  test  consists  of  measuring  the  amount  of  a  soluble  Bait 
of  the  metal  in  question  which  can  be  added  to  a  solution  of  the  phosphate 
before  precipitation  occurs.  Table  23.1  lists  values  for  several  polyphos- 
phates1111 ■  14.  By  this  test,  the  glassy  phosphates  are  better  sequestrants  for 
soluble  calcium  Baits  than  are  tri-  or  pyrophosphates.  However,  with  mag- 
nesium ion  and  soap  present  (Table  23.1),  the  tetrasodium  pyrophosphate 
and  sodium  triphosphate  show  up  as  better  sequestering  agent-. 

14.  ''Technical  Bulletin  number  ^O.s.  Sodium  Tri  polyphosphate,"  Neil    York,  % 

vaco  Chemical  Division,  Food  Machinery  and  Chemical  Corporation. 


772  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

Table  23.2.  Natural  pH  and  Free  Alkalinity  of  the  Polyphosphates118 

Polyphosphate  Natural  pH  of  1%  Soln.         %  Free  Alkalinity  as  Na:>0 

Tetrasodium  pyrophosphate 
Sodium  acid  pyrophosphate 


0.25 

23.3 

4.2 

Equal  to  tetrasodium  py- 

rophosphate   in    buffering 

ability 

9.9 

16.7 

7.9 

8.5 

Sodium  triphosphate 

Sodium  phosphate  glass  with  Na20/ 

P205  =  ca.  1.33 
Sodium  phosphate  glass  with  Na20/  6.9  2.7 

P205  =  ca.  1.11 

In  the  presence  of  anions  such  as  fluoride  and  oxalate,  which  form  highly 
insoluble  precipitates  with  calcium,  the  sequestering  powers  of  the  poly- 
phosphates are  more  nearly  equal,  and,  in  fact,  the  differences  in  sequester- 
ing abilities  are  negligiblella.  It  is  obvious  that  an  indiscriminate  comparison 
of  these  sequestering  values  will  lead  to  confusing  conclusions. 

pH  of  Solutions.  The  phosphates  differ  greatly  in  their  natural  al- 
kalinity and  in  their  ability  to  control  the  pH  of  a  solution  by  buffering 
action.  The  natural  pH  of  one  per  cent  solutions  and  the  free  alkalinity  of 
the  sequestering  polyphosphates  are  given  in  Table  23.2.  The  sodium 
phosphate  glasses  are  not  good  buffering  agents,  as  shown  by  their  low  free 
alkalinity;  however,  if  the  pH  buffering  requirements  are  neglected,  the 
glasses  sequester  as  well  as  the  crystalline  phosphates  under  most  condi- 
tions, and  better  under  some  conditions,  as  shown  by  the  data  of  Table 
23.1. 

The  pH  of  the  solution  has  an  important  effect  on  the  stability  of  the 
phosphates.  The  condensed  phosphates  react  with  water  to  form  less  con- 
densed phosphates  and  ultimately  orthophosphates  through  rupture  of 
P-O-P  linkages.  The  hydrolytic  degradation  of  pyro-  and  triphosphate  has 
been  carefully  studied  by  Van  Wazer,  Griffith,  and  McCullough15.  The  hy- 
drolyses  follow  the  first-order  rate  law  and  are  catalyzed  by  acid  and  not 
by  base.  The  degradation  of  the  polyphosphates  is  extremely  slow  at  neu- 
tral or  alkaline  pH  and  room  temperature,  but  is  accelerated  by  a  number 
of  factors,  the  more  important  of  which  are  increasing  temperature  and 
decreasing  pH.  The  presence  of  cations  (other  than  tetramethyl  am- 
monium), colloidally  precipitated  metal  oxides,  and  the  enzymes  known  as 
phosphatases  also  accelerate  the  breakdown. 

Comparisons  of  the  rates  of  reversion  of  the  polyphosphates  to  ortho- 
phosphates  in  dilute  solutions110  without  pH  control  or  adjustment  have 
shown  tetrasodium  pyrophosphate  to  be  the  most  stable,  followed  in  order 
by  sodium  triphosphate,  the  sodium  phosphate  glasses,  and  sodium  acid 
pyrophosphate.  The  reversion  in  one  hour  at  100°C,  as  measured  by  the 


15.  Van  Wazer,  Griffith  and  McCullough,  J.  Am.  Chem.  Soc,  77,  287  (1955). 


WATER  SOFTENING  THROUGH  COMPLEX   FORMATIOA  773 

build-up  of  orthophosphate,  varies  from  less  than  1  per  oenl  for  tel rasodium 
pyrophosphate  and  s  per  cent  for  sodium  t  riphosphate,  to  about  55  per  ••cut 
for  sodium  acid  pyrophosphate.  Hie  products  of  the  degradation  may  <>r 
may  not  possess  complexing  ability.  Sodium  triphosphate  gives  one  mole  of 
pyro-  and  one  mole  of  orthophosphate,  the  former  having  sequestering 
ability.  Both  Hell"1  and  Thilo17  report  trimetaphosphate  aa  one  of  the 
products  of  the  hydrolysis  of  the  long  chain  phosphates. 

Tin:  Nature  <>f  the  Sequestering   Reaction  and  the  Stability  of 
the  Complex  Eons  Formed 

In  the  sequestering  tests  described  above,  the  amount  of  an  ion  needed 

to  form  a  barely  discernible  precipitate  depends  upon  the  solubility  of  the 
precipitate  and  the  formation  of  a  soluble  complex  ion.  By  neglecting  the 
dispersing  action  and  colloid  stabilization  of  phosphates,  we  can  represent 

this  action  as  follows: 

Ca++  +  polyphosphate  molecule-ion  ^  Ca-polyphosphate  complex  (2) 

( ' .-i ' '  +  precipitating  anion  ;=±  Ca  precipitate,  (3) 

in  which  the  precipitating  anion  can  be  the  phosphate  or  some  other  anion. 
Precipitation  of  calcium  will  occur  if  the  equilibrium  concentration  of  cal- 
cium is  great  enough  to  exceed  the  solubility  product  of  the  precipitate. 
Thus,  the  differences  noted  with  different  anions  can  be  correlated  with 
the  respective  solubility  products.  A  number  of  studies  of  the  complex 
ions  of  polyphosphates  have  been  reported18,  but  most  of  them  fail  to  de- 
scribe accurately  the  complex  ions  by  chemical  formulas  and  true  equilib- 
rium constants  primarily  because  (a)  the  theoretical  treatment  for  chain 
molecule-ions  has  not  been  thoroughly  developed,  (b)  electrochemical 
measurements  are  often  complicated  by  irreversibility  of  the  reactions,  (c) 
the  available  range  of  concentrations  is  restricted  because  of  precipitate 
formation,  and  (d)  it  is  difficult  to  obtain  single  species  of  the  chain  phos- 
phates with  a  degree  of  polymerization  greater  than  three.  In  addition  to 
the  precipitation  tests  discussed  earlier,  a  number  of  other  techniques,  in- 
cluding pll  titration,  membrane  potentials,  conductivity,  transference 
number  measurements,  polarography,  ion-exchange  equilibrium  and  colori- 
metric  studies  have  been  applied  to  these  systems. 

16    Bell,  Ind   Eng.  Chem.,  39,  137    1947 

17.  Thilo.  Chem.  Technik.,  4,  345  .">1     1962 

18.  Van  Wazer  and  Campanella, /.  Am.  Chi  72,  655    1950  ; Rogers  and  Reyn 

old>.  ./.  An     Ch       8        71,  2061     1945  ;  Rosenheim,  Frommer,  Glaser  and 
Sandler,  '/.   anorg.  Chem.,  153,  126    1926  ;  Baasett,  Bedwell  ;,i„i  Hutchinson, 
./    Chen    Sd     1986,  1412;  Kolthoff  and  Watten    Ind    /-.  -      Chem.,  Anal 
15,  8  (1943);  Laitinen  and  Onstott,  ./.  A  71,  Bob 

telsky  and  Kertes,  ./.  Appl.  Chem.  4,  119    1954). 


774 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Gray  and  Lemmerman  (as  reported  by  Quimby9)  carried  out  a  con- 
ductometric  study,  using  Job's  method  of  continuous  variation,  on  the 
calcium  triphosphate  system  at  concentrations  low  enough  to  prevent 
precipitation  at  any  ratio  of  calcium  to  triphosphate.  Their  results  are 
consistent  with  the  existence  of  a  soluble  1:1  calcium  triphosphate  com- 
plex. The  boundary  between  homogeneous  and  heterogeneous  regions  at 
()0°C  was  determined  turbidimetrically,  after  attainment  of  steady  state. 
(Figure  23.1).  The  homogeneous  region  comes  close  to  the  calcium  axis 
and  is  usually  not  detected  upon  adding  sodium  triphosphate  to  calcium 
solutions.  On  branch  DE  of  the  curve,  more  than  one  mole  of  triphosphate 
per  mole  of  calcium  is  required  to  prevent  precipitation.  The  shift  of  the 
curve  to  the  right  as  sodium  salts  are  added  suggests  that  the  precipitates 
contain  sodium,  but  the  equilibrium  solid  phases  have  not  been  com- 
pletely characterized.  From  measurements  of  the  clarification  of  calcium 
oxalate  suspensions,  Gray  and  Lemmerman9  have  estimated  the  dissocia- 
tion constant  to  be  3.1  X  10-7  at  30°C,  assuming  that  the  1:1  calcium  tri- 
phosphate complex  is  the  only  one  involved. 

Rogers  and  Reynolds181*  report  that  pyrophosphate  forms  complexes  of 
the  type  MII(P207)=  with  divalent  ions  such  as  magnesium,  and   com- 


0.2         0.4  0  2  4  6  8  10 

TRIPOSPHATE     ION  CONCENTRATION,  MILLIMOLES  /LITER 


Fig.  23.1.  Homogeneous  and  heterogeneous  regions  at  60°C  for  the  CaCh- 
NajPiOio-HjO  system.  Solid  curve  obtained  turbidimetrically.  Dashed  curve  FG 
gives  sal  united  solutions  obtained  from  compositions  between  dotted  line  and  curve 
DE.  (Reproduced  from  J.  Phys.  Chem.,  58,  613  (1954)) 


WATER  SOFTENING  THROUGH  COMPLEX  FORMATIOh  77;') 

Table  23.3.  Dissociation  Constants  for  Several  Condensed  Phosphates  i  bom 

Condug  1 1\  i  it  Data19 


iation  Constant  for 

the 

Polyphosphate 

Add 

Na  Salt 

-.ill 

Trimetaphosphate 

-   ;  x  in 

68  X  mi 

0.33  X  10  ■ 

Tetrametaphosphate 

1.8  X  10 

!)  X   10-3 

K,  2.2  X  10  3 
K    1.3  X  I')   ■ 

Pyrophosphate 

K,  2.7  X  10  7 

K,  2.4  X  10~10 

4.5  X  10"3 

Triphosphate 

3.0  X  10~3 

Table  23.4.  Apparent 

Dissociation  Constants  of  Calcium  Complexes20 

(Ionic  strength  0.15,  pH  7.4,  temp.  37°C) 

Phosphate  />KC(=  —  log  Kc) 

Triphosphate  4.14 

Pyrophosphate  3.47 

Tetrametaphosphate  3.06 

Trimetaphosphate  2.32 

plexes  of  the  types  Min(P207;r  and  Min(P207)2~5  with  such  metals  as 
iron  and  aluminum.  Considerable  data  on  the  relative  stability  of  polyphos- 
phate complexes  were  obtained  by  Monk19  from  solubility  and  conductivity 
measurements  in  solutions  of  low  ionic  strength.  Some  of  the  data  are  re- 
produced in  Table  23.3.  Gosselin  and  Coghlan20  measured  the  apparent 
dissociation  constants  of  a  number  of  calcium  phosphate  complex  ions, 
utilizing  the  equilibrium  technique  of  ion  exchange21.  Linear  variation  of  the 
distribution  coefficients  with  the  molar  concentration  of  the  phosphate  was 
cited  as  evidence  that  the  complex  ions  formed  were  of  the  1 : 1  type.  The 
values  reported  (Table  23.4)  are  not  true  dissociation  constants  because  the 
identity  and  concentration  of  the  phosphate,  which  enters  into  the  calcula- 
tions, cannot  be  inferred  from  the  available  information,  so  the  values  in 
the  table  are  smaller  than  the  true  pKc's,  and  smaller  than  the  values  re- 
ported by  Monk  from  conductivity  data. 

As  would  be  expected  from  modern  electrochemical  theory22,  both  the 
ring  and  the  chain  phosphates  undergo  association  with  cations  at  relatively 
low  concentration.  In  spite  of  the  relatively  high  negative  charge  on  the 
ring  compounds,  the  ring  phosphates  form  less  stable  complexes  than  do 
the  chain  phosphates.  This  difference  is  in  accord  with  the  known  fact 
that  ring  phosphates  are  not  effective  in  the  prevention  of  precipitation  in 
water  softening.  Van  Wazer  and  Campanella18a  suggested  that  the  chain 

L9.  Monk, etal.,  J.  Chem. Soc.,l»&,  123  27,  127  29,2693  96;  1950,  3475  78;  1962, 1314 
17,  1317-20. 

20.  Gosselin  and  Coghlan,  Archil .  Biochem.  find  Biophys.,  45,  301     Lfl 

21.  Schubert,  Russell,  and  Myers,  •/.  Biol.  Chem.,  185.  :;^7    I960 
.  nose,  Chen     /:■    ..  17,  27    1" 


776  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

phosphate  complexes  are  more  stable  because  the  chain  compounds  can 
form  chelate  rings  with  the  metal  atom,  as  in  (I)  or  (II),  whereas  the  ring 
compounds  cannot  do  so  because  of  mechanical  constraint.  Polarographic  and 
pH  studies18a  indicate  that  to  a  first  order  approximation  the  complexing 
ability  of  a  chain  phosphate  is  proportional  to  the  total  number  of  phos- 
phorus atoms  in  the  polyphosphate,  regardless  of  chain  length.  It  is  also 

o-o  -oo 

I  II  I  II 
— 0— P— O— P— 0— etc. 

II  I 

o         o 

\  / 

M 


etc. 

— 0—  P— 0— P—  0— etc 

II                1 

0             0 

\  / 

M 

/    \ 

0             0 

II      II 

etc.- 

— 0— P— 0— P  — 0— etc 

1                1 

1               1 

o-        o- 

(II) 

(I) 

postulated  that  the  formation  of  polydentate  structures  is  inhibited  by  the 
presence  of  negative  charges  on  the  individual  P04  groups  which  tend  to 
prevent  coiling,  and  cross-linking  of  chains  through  the  metal  atom  is  sup- 
ported by  changes  in  polarographic  diffusion  currents.  Estimates  of  molecu- 
lar weights  range  from  103  to  105  for  complex  ions  formed  from  barium  and 
a  glass  with  an  average  chain  length  of  five  and  of  103  to  107  for  complexes 
from  a  long-chain  glass  with  approximately  the  metaphosphate  composition. 
From  this  wTork,  it  is  also  shown  that  the  barium  ion  is  associated  with  four 
phosphorus  atoms  and  the  sodium  with  two  phosphorus  atoms. 

The  pH  titration  studies  of  Van  Wazer  and  Campanella18a  also  indicate 
that  cations  can  be  divided  into  three  groups  based  on  their  ability  to  form 
complexes  with  the  polyphosphates:  (1)  quaternary  ammonium  ions,  which 
form  no  complexes;  (2)  alkali  and  similar  cations,  which  form  weak  com- 
plexes; and  (3)  the  other  metal  ions  which  form  strong  complexes.  Esti- 
mates of  the  dissociation  constants  were  made,  but  the  assignment  of  def- 
inite structures  and  the  establishment  of  the  relative  covalent  and  ionic 
contributions  to  the  stability  of  the  complexes  is  uncertain  on  the  basis  of 
the  available  evidence. 

Threshold  Treatment 

A  complementary  phenomenon  to  sequestration  is  used  in  "threshold" 
water  conditioning.  Here,  very  low  concentrations  of  condensed  sodium 
phosphates  act  as  deterrents  to  the  crystallization  of  calcium  carbonate. 


WATER  80FTBNINQ  THROUGH  COMPLEX  FORMATION  111 

The  triphosphate.  Xa:)lM  >Ul ,  an<l  the.  phosphate  glasses  may  be  used  BUC- 

sfully  in  concentrations  of  1  to  5  parts  per  million23.  The  "threshold"  ifi 
the  point  at  which  sufficient  sodium  phosphate  has  been  added  to  prevent 
crystallization.  The  concentrations  required  are  considerably  below  the 
amounts  required  to  completely  complex  the  calcium.  Presumably,  the 
phenomena  are  due  to  the  adsorption  of  the  complex  phosphate  on  the 
Bubmicroscopic  nuclei233 • 23c- 23f  in  such  a  manner  as  to  prevent  crystal 
growth  and  precipitation.  Microscopic  studies  indicate  that  the  sodium 
phosphates  cause  distortion  of  the  calcite  crystals,  the  amount  of  distortion 
increasing  as  the  amount  of  phosphate  is  increased.  In  addition  to  pre- 
venting precipitation,  solutions  of  threshold  concentrations  slowly  remove 
old  calcium  carbonate  scale  if  the}-  are  circulated  through  a  given  system 
for  a  period  of  several  months.  Crystalline  sodium  trimetaphosphate  has 
little  or  no  inhibiting  effect  except  in  the  presence  of  alkalies,  which  pre- 
sumably convert  it  to  the  triphosphate. 

Poly  amino  Acids24 

The  use  of  synthetic  polyamino  acids  as  sequestering  agents  is  relatively 
recent.  The  most  important  of  these  substances  are  triglycine  (III),  and 
ethylenediaminetetraacetic  acid  (IV).  Ender  named  these  compounds 
Trilon  A  and  Trilon  B,  respectively25. 

CH2C02H  H02CCH2  CH2C02H 

/  \  / 

X— CHoC02H  NCH*CH«N 

\  /  "\ 

CH2C02H  H02CCHo  CH2C02H 

(III)  (IV) 

Trilon  B  is  one  of  the  most  powerful  coordinating  agents  known,  and 
its  disodium  salt  is  widely  used  under  the  trade  names  "Versene,"  "Seques- 
trene,"  and  "Xullapon."  It  is  significant  that  it  forms  stable  complexes  with 
calcium  and  magnesium — elements  which  do  not  react  strongly  with  most 

23.  Buchrer  and  Reitemeier,  J.  Phys.  Chem.,  44,  552  (1940);  Fink  and  Richardson, 

U.  S.  Patent  2358222  (1940);  Hatch  and  Rice,  Ind.  Eng.  Chem.,  31,  51  (1939); 
Reitemeier  and  Buchrer,  J.  Phys.  Chem.,  44,  535  (1940);  Rice  and  Partridge, 
Ind.  Eng.  Chem.,  31,  58  (1939) ;  Raistrick,  Disc.  Faraday  Soc. ,  1949,  234. 

24.  Martell  and  Bersworth,  Proc.  Sci.  Sect.  Toilet  Goods  Assoc,  No.  10,  Dec.  1948; 

Martell,  Plumb,  and  Bersworth,  Technical  Bulletin  Bersworth  Chemical  Co., 
Framingham,  Mass.;  "Sequestrene,"  Technical  Bulletin,  Alrose  Chemical  Com- 
pany, Providence,  Rhode  Island;  "The  Versenes,"  Technical  Bulletin  #2,  4th 
Ed.,  Bersworth  Chemical  Company,  Framingham,  Mass.,  1952. 

25.  Ender,  Fette  und  Seifen,  45,  144  (1938) ;  Ley,  Ber.,  42,  354  (1909). 


778 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


complexing  agents.  A  considerable  literature  has  grown  up  about  its  use 
in  water  softening,  both  in  the  technical  journals,  and  in  patents250. 

Ethylenediaminetetraacetic  Acid 

Schwarzenbach  and  Ackermann26g"26i  have  measured  the  dissociation 
constants  of  ethylenediaminetetraacetic  acid,  and  have  shown  that  two 
hydrogens  are  held  in  the  form  of  zwitter  ions: 

'OOCCH,  CH2COO' 

\H  H/ 

N— CH2CH2— N 

/  \ 

OOCCH2  CH2COO. 

(V) 
Three  structures  have  been  postulated  for  calcium  salts  of  Trilon  B,  (VI), 
(VII),  and  (VIII).  Structure  (VI)  is  of  the  type  usually  associated  with 
divalent  ions.  Structure  (VII)  was  proposed  by  Pfeiffer26a-26c. 


25a.  For  example,  I.  G.  Farbenindustrie  A.  G.,  French  811938  (1937) ;  German  718981 
(1942);  Munz,  U.  S.  Patent  2240957  (1941);  Bersworth,  U.  S.  Patent  2396938 
(1946). 

26.  Pfeiffer  and  co-workers,  Ber.,  75B,  1  (1942);  76B,  847  (1943);  Z.  anorg.  allgem 
Chem.,  258,  247  (1949) ;  Brintzinger  and  co-workers,  Z.  anorg.  allgem.  Che?n.,  249, 
113  (1942);  251,  285  (1943);  256,  65  (1948);  Schwarzenbach  and  Ackermann, 
Helv.  Chim.  Acta,  30,  1798  (1947);  31,  459,  1029  (1948);  32,  839  (1949). 


WATER  SOFTENING  THROUGH  COMPLEX  FORMATION 


779 


CH  —  N 


CH; 


,CH*COO 


riizcoo 


VTTT 

Formula  (\'III)  was  suggested  because  calcium  .shows  little  tendency  to 

rm  complexes  with  amines  i  11  a([ueoussohitioi).  Mart  ell  and  his  associates84*  »b 
ve  shown,  however,  that  the  addition  of  calcium  chloride  to  a  solution 
the  disodium  salt  (V),  results  in  a  marked  drop  in  the  pH  of  the  solution, 
the  nitrogens  were  not  involved  in  complex  formation,  there  should  lie 
i  change  in  the  pH  of  the  solution.  Further,  (VII)  is  favored  over  (VI) 
the  basis  of  titration  of  one  mole  of  the  amino  acid  in  the  presence  of 


- 

A-  ACID 

- 

B=   1   MOLE   Ca(0Ac)2  /MOLE  ACID 

A 

- 

C  =  1  MOLE  CaCI2//MOLE    ACID 

H 

- 

A, 

- 

B^-^^y 

"B 

_C_ 

/& 

fc 

■ 

1 L                         i                           1                          i 

i           i           i 

2  3 

EQUIVALENTS     OF    BASE 


Pig.  23.2.  The  effect  of  calcium  suits  on  the  Demoralization  curve  of  ethylenedi 

linetetraacetic  arid. 


u 
z 
< 

b 

D 
Q 

Z 

8 

h 
Z 

LU 

-I 

I 

o 


780 
400 

350 

300 

250 


200- 


150 


100 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


50 


- 

NaOH         / 

ADDED        / 

\\A 

/ 

B 

Ca(0H)2                / 

ADDED        ^^/ 

i I               i 

i             —  mm 

12  3  4 

EQUIVALENTS     OF    BASE 
Fig.  23.3.  Conductometric  titration  of  ethylenediaminetetraacetic  acid  with  so- 
dium hydroxide  and  calcium  hydroxide. 


one  mole  of  calcium  salt24b  (Fig.  23.2).  The  considerable  change  in  pH  values 
in  the  presence  of  acetate  ion  supports  the  hypothesis  that  all  of  the  car- 
boxyl  groups  in  ethylenediaminetetraacetic  acid  tend  to  coordinate.  If  two 
of  the  carbox}rl  groups  were  free  to  act  as  proton  acceptors,  the  presence  of 
the  acetate  ion  should  make  little  or  no  difference  in  the  titration  curve. 
Structure  (VI)  should  be  optically  active,  but  Pfeiffer  was  unable  to  resolve 
either  the  strychnine  or  brucine  salts  of  the  calcium  complex.  However,  the 
analogous  cobalt  (III)  complex  has  been  resolved27  and  the  hexadentate 
nature  of  the  ethylenediaminetetraacetato  group  in  the  cobalt  complex  was 
demonstrated  by  means  of  the  infrared  spectrum27.  Isolation  of  anhydrous266 
sodium  ethylenediaminetetraacetatocobaltate(III)  lends  support  to  struc- 
ture (VII). 

Additional  data  on  the  calcium  complex  are  shown  by  studies  of  equiva- 
lent conductance24*1  (Fig.  23.3).  When  the  acid  is  titrated  with  calcium 
hydroxide,  the  equivalent  conductance  decreases  until  nearly  two  equiva 

'27.    Busch  and  Bailar,  ./.  .1///.  Chem.  Soc,  75,  4574  (1953). 


WATER  SOFTENING  THROUGH  COMPLEX  FORMATION  781 

Table  23.5.  Formation  Constants  roa  Alkaline  Eartb   Elembni   Compli 

WITH    I :.  IIIVI.KN'KDI  \mi.\k  i  i:  i  u  \  ICETIC    ACID 
Divalent  ion  log  Kki  log  Kk: 

Mg  2  28  8.69 

Ca  3.51  10.69 

Sr  2.30  S.63 

Ba  2.07  7.76 

Lents  of  base  have  been  added,  and  remain.-  constant  until  four  equivalents 
have  been  added.  Presumably,  the  decrease  represents  the  removal  of  the 
two  strongly  acidic  hydrogens,  and  the  flat  portion  of  the  curve  denotes  the 

removal  of  calcium  ions  and  the  neutralization  of  the  third  and  fourth 
hydrogens  of  the  arid.  Addition  of  excess  calcium  hydroxide  increases  the 
equivalent  conductance. 
Schwarzenbach  and  Ackermann2<*'2,i  studied   the   relative  complexing 

tendencies  of  ethylenediaminetetraacetic  acid  and  homologous  compounds 
with  three,  four,  and  five  carbon  atoms  between  the  nitrogen  atoms.  The 
trimethylenediamine  (C3)  compound  showed  strong  complex  formation, 
but  not  as  strong  as  the  ethylenediamine  compound.  The  higher  homologs 
were  much  less  effective.  Consequently,  they  concluded  that  the  fused  ring 
system  was  not  obtained  with  the  molecules  containing  four  or  five  carbon 
chains,  and  that  in  the  formation  of  complexes  of  them,  the  aminodicar- 
boxylic  groups  act  independently.  Qualitatively,  the  sequestering  action  of 
the  tetrasodium  salt  of  ethylenediaminetetraacetic  acid  is  strong  enough 
to  dissolve  precipitates  such  as  Ca3(P04)2 ,  CaC204 ,  MgCO> ,  BaSO*  ,  and 
alkaline  earth  salts  of  soaps'24'1.  Schwarzenbach  and  Ackermann26g-26i  have 
obtained  equilibrium  constants  for  the  formation  of  a  number  of  complexes 
(Table  23.5).  Kki  is  the  equilibrium  constant  for  the  reaction 

M+++  HY^MHY"  (4) 

and  Kk2  is  the  constant  for  the  reaction 

M++  +  Y*"  ^  MY-  (5) 

The  values  wen-  obtained  by  titrating  ethylenediaminetel  raacetic  acid  with 
potassium  hydroxide  in  the  presence  of  the  various  metal  ion-.  Similar 
titrations  with  sodium  and  lithium  hydroxides  indicated  slight  complex 
formation  with  these  metals.  The  investigators  assumed  no  complex  forma- 
tion with  potassium,  apparently  without  investigating  the  behavior  of 
rubidium  and  cesium.  It  i-  difficult  to  use  the  data  in  a  quantitative  Bense 
Bince  the  equilibria  are  very  sensitive  to  the  addition  of  ions  to  the  solution. 
It  is  evident  that  the  complexes  are  more  stable  in  alkaline  solution.  Pur 
ther,  when  arid  salts  are  used,  complex  formation  is  accompanied  bya  drop 
in  the  pH  of  the  solution. 
The  great  stability  of  the  calcium  and  magnesium  compounds  of  ethylene- 


782  CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 

diaininetetraacetic  acid  is  the  basis  for  an  excellent  method  of  determining 
hardness  in  water26*  • 28.  The  water  is  titrated  with  a  standard  solution  of 
disodium  ethylenediaminetetraacetate,  using,  as  the  indicator,  the  wine-red 
magnesium  complex  of  the  dye  Eriochrome  Black  T.  The  calcium  ion  is 
first  tied  up  by  the  complexing  agent,  and  then  the  "free"  magnesium  ion. 
The  next  drop  of  the  ethylenediaminetetraacetate  solution  destroys  the 
magnesium-dye  complex,  and  the  color  of  the  solution  becomes  a  clear 
blue.  Alternatively,  the  end  point  can  be  determined  by  pH  indicators  or 
by  potentiometric  methods29. 

Triglycine 

The  complexing  action  of  triglycine  is  similar  to  that  of  ethylenediamine- 
tetraacetic  acid,  two  moles  of  triglycine  being  required  per  mole  of  calcium 
•on.  By  analogy,  we  would  expect  the  complex  structure  (IX). 

CH2COO  OOCCH2 

I  \     /  I 

OOCCH2— N Ca N— CH2COO 

I  /   V      I 

CH2COO  OOCCH2 

(IX) 

Extent  of  the  Sequestering  Ability  of  the  Polyamino  Acids 

The  ability  of  the  polyamino  acids  to  form  complexes  with  metals  varies 
widely.  Complexes  similar  to  those  of  calcium  have  been  obtained  with 
magnesium,  strontium,  barium,  copper(II),  mercury (II),  cadmium,  zinc, 
and  nickel.  Of  the  tripositive  ions,  bismuth,  cobalt,  and  chromium  give 
stable  complexes,  while  iron  forms  relatively  weak  ones.  Lead,  lanthanum, 
neodymium,  thorium  and  uranium (IV)  have  little  tendency  for  complex 
formation  with  these  compounds. 

The  polyamino  acids  are  strong  sequestering  agents  above  a  pH  of  5, 
and  the  higher  the  pH  the  stronger  their  sequestering  power.  The  poly- 
amino acids  may  be  used  independently  as  water  softeners  and,  in  addition, 
may  be  incorporated  in  liquid  or  solid  soaps  to  give  them  a  detergent-like 
action  in  hard  water. 

Phytates 
Phytic  acid  is  the  hexaphosphate  ester  of  the  inactive  form  of  inositol30. 


28.  Bredermano  and  Schwarzenbach,  Chimia,  (Switz.),  2,  56  (1948);  Diehl,  Goetz 

and  Bach,  ./.  .1///.  Waterworks  Assoc,  42,  40  (1950);  Goetz,  Loomis  and  Diehl, 
Anal.  Ckem.,  22,  796  (1950). 

29.  Halm,  Anal.  Chim.  Aria,  4,  583  (1950). 

30.  Suzuki,  Yoshemura,  and  Takaishi,  Bull.  Tokyo  Imper.  Univ.,  College  of  Agric.,  7, 


WATER  SOFTENING  THROUGH  COMPLEX  FORMATIOh  783 

H^OjPO  OPOjH* 

^  OPOjH, 
'OPOjHs, 

>0 


OP03H2 

The  phytate  ion  is  known  bo  form  metal  complexes,  bul  few  of  its  deriva- 
tives have  been  studied,  and  apparently,  it  has  not  been  used  commer- 
cially. Aryan"  studied  the  behavior  of  calcium  ion  in  the  presence  of  phytate 
ion.  He  found  that  immediate  precipitation  resulted  if  the  calcium isodium 
phytate  ratio  exceeded  1:1.  Even  at  lower  ratios,  a  substance  of  the  com- 
position Ca^XasCeHeC^Pe-BH^O  slowly  precipitated  after  36  hours.  Addi- 
tion of  sodium  carbonate  or  sodium  oxalate  to  the  solutions  did  not  give 
immediate  precipitation,  although  it  did  so  at  the  same  pH  in  the  absence 
of  phytate. 

The  possibility  of  complex  formation  indicated  by  this  chemical  evidence 
was  not  supported  by  Aryan's  spectrophotometry  studies  in  the  ultraviolet 
region,  and  it  is  possible  that  the  solubility  of  calcium  in  concentrations  less 
than  or  equal  to  phytate  concentration  is  due,  wholly  or  in  part,  to  crystal 
distortion  of  the  type  described  under  threshold  treatment  for  water  condi- 
tioning. 

405  (1907);  Newberg,  Biochem.  Z.,  9,  557  (1908);   Anderson,  thesis,  Cornell 
University,  1919;  Starkenstein,  Biochem.  Z.,  30,  56  (1910);  Vorbrodt,  Bull. 
intern,  acad.  sci.  Cracovie,  ser.  A,  414  (1910). 
31.  Arvan,  thesis,  University  of  Illinois,  (1949). 


Index 


A.  see  Ammonia 

Abbreviation  for  names  of  donor  mole- 
cules, 90 
Absolute  asymmetric  synthesis,  350,  351 
Absorption  bands,  565 

relation  to  coordination  groups,  566 
relation  to  geometric  isomerism,  294- 
297 
Absorption   of   light,   see   also   Infrared 
and  Ultraviolet 
sources  of,  567 
theories  of,  565 
Absorption  spectra,  564-580 
color  related  to,  564 
structure  related  to,  364 
ac,  see  Acetate  ion 
acac,  see  Acetylacetonate  ion 
Acceptor,  1 

Acet amide,  elect rodeposition  from  solu- 
tions in,  670 
Acetate  as  bridging  group,  33 
acetate  ion,  96 

Acetato  group,  bridging,  33,  34,  462,  463 
Acetatopentamminecobalt(III)  ion,  33 
Acetic  acid,  electrodeposition  from  solu- 
tions in,  670 
Acetoacetic    ester,    coordination    com- 
pounds of,  41 
Acetone,    electrodeposition    from    solu- 
tions in,  670 
Acetonitrile,    stabilization  of  copper(I) 

by,  75 
Acetylacetonate  ion,  41,  96 
Acetylacetone  as  solvent  for  extraction 

of  complexes,  45 
Acetylacetone  complexes,  structures  of, 

42-44 
Acetylacetone   metal   complexes,   infra- 
red study  of,  577 
Acetylene 
complex  with  aluminum,  497 
complex  with  copper(I)  chloride,  495 
derivatives,    platinum    complexes    of, 
492 


Acetylene  complexes,  Btudied  by  Raman 
spectra,  597 

Acid-base  phenomena  in  coordination 
compounds,  121,  416-447 

Acid-base  strengths 
relation  to  ionic  potential,  423 
relative  to  different  bases,  432 
steric  effects,  433 

Acidity,  from  conversion  of  aquo  to  hy- 
droxo  group,  418,  424-431,  451 

Acidopentammine  cobalt  (III)  complexes 
containing  chlorate,  bromate,  io- 
date,  and  perchlorate,  29 

Acridines,  complexes  of,  238 

Acrilan,  dyeing  of,  766 

Acrylonitrile  fibers,  dyeing  of,  766 

Activation  energy  in  electrode  processes, 
635 

Active  racemates,  583  - 

Actomyosin,  activation  by  magnesium, 
710 

Acyloin  oxime  group  in  analysis,  679 

Addition  agents  in  electrodeposition,  642 

Adenosine  triphosphate,  magnesium 
complex  of,  709,  710 

Adsorption 
of  complexes  on  ion  exchange  columns, 

622 
of  complexing  ions  on  electrode  sur- 
faces, 641 
of  ions  on  electrodes,  633 

Aggregation  in  basic  aluminum  salt  solu- 
tions, 457 

Aging  of  olated  solutions,  456 

Aging  of  precipitated  hydrous  metal 
oxides,  470 

4>  ala,  see  Phenylalanine  anion 

alan,  see  alanine  ion 

Alanine  ion,  96 

Alcohols,  coordinating  ability  of,  23, 
123,  129 

Aliphatic  amines,  coordinating  ability 
of,  62,  180 


785 


786 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Alizarin,  749 

cobalt  complex  of,  750 
Alizarin,  lakes  of,  749 
Alizarin  cyanine,  cobalt  complex  of,  751 
Alizarin  Red  S,  752 
Alizarin  Yellow   G,   chromium    complex 

of,  753 
Alizarin  Yellows,  753 
Alkali  complexes,  177 

stabilities  of,  176 
Alkali  metal  ions,  complexes  of,  2 
Alkali  metal  reduction  hypothesis,  626 
Alkaline  earth  complexes,  177 
Alkaline  earth  complexes 

of     ethylenediaminetetraacetic     acid, 
778-782 
formation  constants  of,  281 

of  phosphates,  773-777 

of  phytates,  783 

stabilities  of,  176,  181,  281 
Alkaline  earth  metals,  ammoniates  of,  150 
Alkyl  gold  cyanides,  structure  of,  89 
N,  N'-Alkylsubstituted         ethylenedia- 
mines,  stabilities  of  complexes,  236, 
237 
Alkyl  substitution,  effect  on  coordinat- 
ing ability  of  ligand,  78,  123 
Alloys,  electrodeposition  of,  666-669 
Allyl  alcohol 

coordination  compounds  of,  488 

platinum  complexes  of,  488 
Allylamine 

as  a  bidentate  group,  491 

complexes  of,  490 

platinum  complexes  of,  490 
Aluminum 

acetylacetonate,  222 

borohydride,  x-ray  structure  of,  608 

bromo  complexes  of,  6 

chloro  complexes  of,  6 

dye  complexes  of,  749,  750,  752,  758 

electrodeposition  of,  670 

oxolated  complexes  of,  457 

resolution    of   hexacovalent    complex, 
321 

stereochemistry  of  tetracovalent  com- 
plexes, 374 
Aluminum  chloride 

dimer,  608 

structure  of,  18 

uses  in  organic  reactions,  499 


Aluminum  compounds 
crystalloidal  and  colloidal,  466 
in  tanning,  471 
of  unsaturated  hydrocarbons  and  de- 
rivatives, 497 
Aluminum  halides 
dimeric  nature  of,  608 
structure  of,  365 
Aluminum  hydroxide 

deolation  in  peptization,  464 
effect  of  anions  on  precipitation,  471 
olation  in,  464 
Aluminum-iron  alloy,   electrodeposition 

of,  670 
Aluminum  oxide  hydrosols,  464 

effect  of  neutral  salts  on  pH,  465 

Aluminum  oxide,  hydrous,   decrease  in 

chemical  reactivity 

on  aging,  470 

on  heating,  470 

Aluminum  oxide,  sols,  factors  affecting 

pH,  464,  465 
Aluminum  oxychloride  sols,  effect  of  ag- 
ing on  pH,  465 
Aluminum  oxyiodide  sols,  catalytic  ac- 
tivity in  decomposition  of  hydrogen 
peroxide,  471 
Aluminum  salts,  basic,  455 

degree  of  aggregation,  457 
Aluminum  tanning,  471 
amac,  see  amino  acid  anion 
Amide  group,  coordination  and  bridging 

by,  62 
Amines,     anhydrous,    electrodeposition 

from  solutions  in,  670 
Amines,  complexes  with  nickel  cyanide, 

137 
Amines,  relative  coordinating  ability  of 
primar3r,    secondary    and    tertiary, 
123,  128 
Amino  acids 
from  natural  products,  712-716;  730- 

735 
use  in  water  softening,  777-782 
uses  in  analytical  chemistry,  680 
a-Amino  acids 

chromium(III)     complexes,    stability 

and  hydrolysis  of,  37 
cobalt  (III)  complexes,  37 
platinum (II)  complexes,  38 


INDEX 


7S7 


polymeric  cobalt  1 1 1    complexes 
stability  of  copper  complex* 
0-Amino  acids,  chelation  by,  39 

j .  i  and  t  Amino  acids,  failure  to  chelate, 

2-Aminoethanethiol,  gold  derivative  of, 

52 
Amino  group  coordination   in  enzyme 

Bubstrate  complex.  704 
L-Amino-4-hydroxyanthroquinone  in  de- 
termination  of  beryllium  and   tho- 
rium, 695 
Aminopeptidases,    metal    activation    of, 

705 
o- Amino thiophenol,   nickel   complex  of, 

56 
Ammines  and  hydrates,  color-  of,  60 
Am  mines 
as  acids,  426 

dissociation  constants  of,  428 
effect  of  anion  on  stability  of,  61 
formation  from  action  of  ammonia,  60 
early  theories  of,  100 
explosive  character  of,  61 
loss  of  hydrogen  ion  from,  59,  426 
of  alkali  and  alkaline  earth  fluorides — 

non-existence  of,  142 
of  fluoride  salts.  142 
preparation  and  relative  stability   in 
water,  61 
Ammonate,  100,  110 
Ammonia 
dipole  moment  and  polarizability  of, 

124 
electrodeposition  from  solutions  in,  669 
physical  properties  of,  127 
resemblance  to  water.  418 
solvent  properties  of,  59 
bilization  of  valence  by,  60 
Ammonium    chloropalladate(I]  .    -tinc- 
ture of. 
Ammonium  chlorozincate,  si  ructure  of,  1 
Ammonium     dithiocarbazide,      reaction 

with  platinum  II),  54 
Ammonium  theory  of  ammonate-,  lol 
Amphoteric  metal  ion-,  titration  of.    137 
Amperometric  titration-.  501 
Amphoterism,  134  145 
dialysis  studies  of.  hi 
hydroxy -complex  theory  of ,  138  Hi 


in  Don-protonic  bj  9 terns,  1  )<• 
0x3  acid  theorj  of.  i:;t;  138 

peptization  theory  of, 

relation  to  solvent .   12(1 

Amylene 
plat  ilium  complexes  of,  iss 

zinc  chloride  complex  of.   197 
Analytical  chemistry,  coordination  com 

pound-  in.  672 
Anderson's    formulation    of    poly-acids, 

483 
Anhydro  acid.  1 17 
Anhydro  base,  417 
Aniline,  coordination  with  platinum   II    . 

85 
Anionic  complexes 
acid-base  properties  of,  431 
eathodic  reduction  of,  629 
formation  by  chelation  by  dicarboxylic 

anions,  461 
nomenclature  of,  94 
Anion  penetration,  448,  458 
by  molybdate  ion,  458 
effect  of  chelation  on,  460 
effect  of  concentration  of  react  ant-  on, 

459 
effect  of  coordinating  power  of  anion 

on,  459 
formation  of  chelate  rings  by,  461 
in  deolation,  469 
in  dissolution  of  hydrous  metal  oxide-. 

468 
in  formation  of  hydrous  metal  oxides, 

463 
in  metal  oxide  hydrosols,  165,  466 
in     precipitation     of    hydrous     metal 

oxides,  470 
order  of  effectiveness  of  various  ions, 
459,  461,  465,  466,  (69 
Anions 
effect  of  corrdinating  tendency  on  re- 
moval     from     precipitated     metal 
oxide-.   17'i 

organic 
a-  bridging  groups,  163 
effect   of  isomerism  on  penetrating 

power  of.    (61 

penetration  bj  .  !»'>" 

relative  penetrating  ability  of.   161 

reduction  at  cathode,  I 

relal ive  coordinating  power  of 


788 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Anions — Cont. 

role  in  stability  of  solid  complexes,  139 
Anthracene  Blue  WR,  752 
Anthragallol,  752 
Anthranilic  acid  as  a  complexing  agent, 

681 
Antibonding  orbitals,  199  et  seq. 
Antimony 

electrode  polarization  of,  638 

electrodeposition  of,  631,  648 

explosive,  631 
electrodeposition  of,  648 

separation  from  other  metals,  666 
Antimony  (III) 

ability  to   act   as   donor  or  acceptor 
atom,  7 

chloride,  reaction  with  nickel  and  iron 
carbonyls,  86 

halo  complexes  of,  8 

possible  polybromides  of,  8 

stereochemistry  of  CsSb2F7  ,  8,  375 
Antimony  (IV),  existence  of,  573 
Antimony  (V) 

halides,  configuration  of,  388 

halo  complexes  of,  8 

reduction  of,  404 
Aquo  acid,  417 
Aquo  base,  417 
Aquo  bridge,  46,  391 
Aquochloroammines,  isomers  of,  297 
Aquo  complexes  as  acids,  425 

dissociation  constants  of,  428 

effect  of  pH,  574 
Aquo  group 

conversion  to  hydroxo  group,  22,  418, 
425-428,  451,  465 

in  olation,  22,  449,  451 

in  precipitated  h}rdrous  metal  oxides, 
470 
Aquotization  in  formation  of  beiyllate 

hydrosols,  469 
Arginase,  metal  activation  of,  715 
Arginine,  metal  complex  of,  715 
Aromatic  diamines,  complexes  of,  67 
Arsenic,  electrodeposition  of,  648 
Arsenic,  halo  complexes  of,  8 
Arsenic  (III),  ability  to  act  as  donor  or 

aceptor  atom,  7 
Arsenic (V),  halides,  configuration  of,  388 
Arsine,  physical  properties  of,  127 


Arsine  complexes 

double  bonding  in,  81 

of  copper  (I)  and  gold  (I),  79 

of  platinum  and  palladium,  81 
Arsines,  organic 

complexes    containing    two    different 
metals,  83 

donor  properties  of,  78 

reaction  with  copper  (I),  407 
Ascorbic  acid  oxidase,  724 
Asymmetric  induction,  352,  581 
Asymmetric  synthesis,  316,  350,  351,  583 
Asymmetry,  molecular  and  crystalline, 

580 
Atomic  number,  relation  to  spectra,  567 
Atomic  orbitals,  shapes  of,  163,  200,  201 
Atomic  orbital  theory,  164,  198 

and  stability  of  complexes,  174 
Atomic  volume  and  coordinating  ability, 

120 
Auxiliary  valence,  109 
Azobenzene    coordination    in    terms    of 

molecular  orbitals,  207 
Azaporphins,  760 
Azide  complexes 

of  cobalt  (III),  76 

of  copper (II),  76 
Azide  ion 

reaction    with   porphyrin    complexes, 
728,  729 

structure  of,  580 
Azine  dyes,  754 
Azo  dyes 

as  indicators  for  metal  ions,  684 

in  polarographic  analysis  of  aluminum, 
697 

metal  complexes  of,  499,  755,  761 

o-substituted,  754 
Azo  group,  donor  properties  of,  74,  207, 

754-760 
Azomethine  dyes,  metal  complexes  of, 

499,  759,  760 
Azosalicylic  acid 

dyes  from,  753 

metal  complexes  of,  753 

Base  strength  of  ligand  and  coordinating 

ability,  141,  180 
Basic  beryllium  acetate,  and  homologs, 

34 


INDEX 


V> 


Basic  salts 

ol  bridges  in.  22,  4 15 

structures  based  on  coordination  the 

ory,  44 5-447 
x-ray  Btudies  of,  447 
Basic  zinc  acetates,  and  homologs,  ;^i 
Basic  zirconium  acetates,  and  homologs, 

34 
Bathochromic  effect,  565 
in  picrates,  554 
in  quinhy  drones,  550 
Benzac.  sec  bensoylacetate  ion 
Benseneaso-6-naphthol,  copper  lake  of, 

766 
Benzene,    electrodeposition    from    solu- 
tions in,  670 
Benzidine,  96 

complexes  of,  67,  254 
Benzo  Fast  copper  dyes,  755 
Benzoin  in  determination  of  zinc,  695 
-Benzoinoxime    for    determination    of 
copper,  679 
Benzoylacetate  ion,  96 
Benzoylacetone,       coordination       com- 
pounds of,  41 
Benzoylcamphoraluminum(III),    muta- 

rotation  of,  349 
Benzylamine,  96 

coordinating  ability-  of,  180 
Benzylmethylglyoxime,  isomers  of  nickel 

complex  of,  677 
Berlin  green,  structure  of,  90 
Beryllium  acetylacetonate,  42,  222 
Beryllate  hydrosols,  469 
Beryllium 
acetylacetonate,  2 
dye  complexes  of,  751 
electrodeposition  of,  669,  670 
planar  phthalocyanine  complex  of,  243, 

362 
polymeric  complexes  of,  42 
stereochemistry  of  tetracovalent  com- 
plexes of,  372 
Beryllium  oxide,  hydrosols 
cationic,  comparison  of  with  anionic 

sols,  469 
precipitation  of,  468 
Beryllium    oxide,    solubility    in    beryl- 
lium sulfate  solution,  28 
Beryllium    oxychloride    sols,    efiV 
anions  on  conductivity  of,  466 


Beryllium  salts,  oxolation  of,  460 
Beryllium  salt  solutions,  effect  of  aging 

on  pB  of,  160 
Berselius' conjugate  theory,  LOO 
Biacetyl,  determination  of,  677 
Biacetyldioxime    in    determination    of 

nickel,  674 

Bidentate  group,  220 

bridging  by,  234,  463 
BigH,  see  Biguanide 
<t>  BigH,  see  Phenylbiguanide 

Biguanide,  complexes  of,  70,  96 
Biguanides,  substituted,  70 
Bile  .icids,  559 

2,2'-Biphenol  complexes,  255 
Biplumbite  ion,  formula  of,  587 
Biological  importance  of  chelates,  221, 

698-742 
Birefringence,  relation  to  structure,  364 
Bis-benzonitrile  palladium(II)  chloride, 

493 
Bis(cyclopentadienyl)  iron  (II),  494 

structure  of,  507 
Bis  (ethylenediaminedesalicylaldehyde)  - 

M-aquo  cobalt  (III),  391 
Bis-ethylenediamine  disilver(I)  ion,  234 
Bis(isobutylenediamine)palladium(II) 

and  platinum (II)  ions,  reported  re- 
solution of,  369 
Bis(methylbenzylglyoxime)nickeKII) 
diamagnetism  of,  211 
failure  to  exchange  with  radioactive 

Xi(II)  ion,  211 
geometric  isomerism  of,  211,  677 
Bis-a-methyl-/3-indylmethene        copper- 

(II),  257 
Bismuth 

electrode  polarization  of,  638 
electrodeposition  of,  649 
halo  complexes  of,  8 
separation  from  copper,  666 
Bismuth  ions,  complex,  aggregation  from 

olation,  453 
Bismuth   thiosulfate,   double   salt    with 

potassium  thiosulfate,  59 
trans  Bis-oxalato dipyridine iridate  III  . 

281 
Bis-pentadiene    dichloro    platinum  II  . 

240 
Bis  salicylaldehyde  -,.-,'  diaminodipro- 

pylamine  coball  III),  391 


790 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Bis  (salicylal)ethylenediamine        cobalt- 

(II),    absorption    of    oxygen,    nitric 

oxide  and  nitrogen  dioxide  by,  45, 

46 
l,8-Bis(salicylideneamino)-3,6-dithiaoc- 

tanecobalt(III),  287,  320 
Bis-sulfamido-diaquo  rhodiate(III)  ion, 

resolution  of,  323 
Bis-thiourea  copper(I)  ion,  383 
Bis-(a,/3,y-triaminopropane)cobalt(III) 

ion,  288,  318 
Bjerrum,  method  of,  572,  592 
Black  nickel,  electrodeposition  of,  656 
Blocking  of  functional  groups  by  metal 

ions,  714 
Blomstrand 
chain  theory,  102 
formulation  of  poly-acids,  473 
Blueprints,  544 
bn,  see  2,3-butanediamine 
Bodecker  reaction,  539 
Bohr  magneton,  600 
Bond  classification,  207 
disagreement    between    different    cri- 
teria, 212 
Bond  cleavage,  role  of  metal  ions  in,  702 
Bond  formation,  role  of  metal  ions  in, 

702 
Bond  lengths  as  a  criterion  of  double 

bond  character,  205 
Bond,  metal-metal,  160,  525,  534,  536 
Bond  orbitals,  significance  in  complex 

formation,  414 
Bond  stability  and  rate  of  substitution 

reactions,  213 
Bond  strengths,  relative,  170 
Bond  type 
in  halide  complexes  of  metals  of  first 

transition  series,  11 
determination  by  infrared,  576 
Bonding,  relation  to  second  band,  566 
Bonding  orbitals,  199-201 
Bordeaux  B,  752 
Borine-phosphorus   trifluoride   complex, 

194,  206 
Borine-trimethylphosphine  complex,  194 
Boron  complexes 
with  acetylacetone,  43 
with   carbon    monoxide,    stability    of, 

194 


Boron  trichloride-halide  bonding,  598 

table  of,  599 
Bragg,  method  of,  606 
Brass,  electrodeposition  of,  666,  667 
Bridged  complexes,  coordination  number 

four  in,  365 
Bridged    complexes    of   palladium   with 

phosphines  (halogen  bridges),  81 
Bridged  halo  complexes,  interaction  ab- 
sorption in,  19 
Bridges,  mixed,  451,  462 
Bridging  group 

bidentate,  463 

halo  as,  7,  8,  17,  81,  365 

hydroxo  as,  22,  23,  323,  448 

nitro  as,  451 

oxo  as,  448 

nomenclature  of,  94 

peroxo  as,  26,  47,  451 
Bridging,  maximum  number  of  groups, 

with  octahedral  atoms,  450 
Bright  electrodeposits,  640 
Brighteners  in  electrodeposition,  642 
Brightness  of  electrodeposits  and  irre- 
versibility of  deposition,  644 
Brilliant  Alizarin  Blue  3R,  752 
Bromate  ion 

coordination  compounds  of,  29,  272 

structure  of,  580 
Bromide  ion,  donor  properties  of,  4-20 
Bromine (V)  fluoride,  388 
Bromocadmium  complexes,  405 
Bromo  chloro  ethylene    ammine    plati- 
num (II),  490 
Brompentamminecobalt(III)  sulfate,  267 
a-Bromopropionic    acid,    resolution    by 

complex  formation,  33 
Bromopurpureo  salts,  98 
Buffers  for  metal  ions,  221 
Butadiene 

absorption  by  CuCl,  495 

complex  with  copper  (I)  chloride,  495 

complexes  with  platinum,  489 
Butadiene  tricarbonyl  iron(0),  493 
2,3-Butanediamine,  96 
Butene 

platinum  complex  of,  501 

silver  complexes  of,  496 
Butylene,  palladium  complexes  of,  493 
2,3-Butylenediamine,  chelates  of,  228 


INDEX 


7!)  I 


bzd,  see  Benzidine 
bil,  Bee  Bensylamine 

Cacodyl  oxide,  coordination  by,  84 
Cadmium 
amphoteriam  of,  1 i- 
eleotrodepoeitioD  of,  6 19 
from  ammines,  638 
from  cyano  complexes,  638 
from  thiosulfate  complexes,  630 
silver  alloys,  667,  669 
Cadmium  complexes 
Cyanide,  refractometric  Btudy  of,  583 
dissociation  constants  of,  413 
formulas  of,  592 
halo,  5 

polarographic  behavior  of,  405 
stereochemistry  of  tetracovalent  com- 
plexes, 372 
structure  of  Cd(NH3)2Cl2  ,  367 
Calcium  carbonate  scale,  removal  of,  777 
Calcium      ethylenediaminetetraacetate, 

proposed  structures  for,  778 
Calcium  phosphate  complexes,  dissocia- 
tion constants  of,  775 
Calcium  proteinates,  739 
Calcium  triphosphate  complex,  dissocia- 
tion constant  of,  774 
Camphorene,  compound  with  palladium- 

(II)  chloride,  493 
Carbonate  exchange  by  carbonato  am- 
mines of  cobalt  (III),  32 
Carbonato  group 
bridging  by,  463 
chelation  by,  32 
Carbonatopentammine  cobalt  (III)  chlo- 
ride  1-hydrate,  nature  of  the  car- 
bonato group,  32 
Carbonatotetrammine    cobalt  (III)     ion 
bidentate  coordination   by  carbonate 

in,  32 
preparation  of,  17 
use  in  synthesis,  278 
( Sarbon  coordination  and  stabilization  of 

low  oxidation  states,  91 
( larbon  coordinators,  3 
Carbon,  donor  properties  of, 
Carbonic  anhydrase,  sine  in,  708 
Carbon  monoxide 
reaction   with   metalfl   and   Baits,   509- 
518,  540,  542-544 


reaction  with  osinimn  let  roxide,  513 

reaction  with  peroxidase,  726 
reaction    with    porphyrin    complexes, 

721,  726,  728,  729,  735 
trans  influence  of,  148 
( Sarbonyl,  Bee  Metal  carbonyl 

Carbony]  group,  donor  properties  of,    11 

( 'arbonyl  metals,  641 

Carbonyl  phosphine  complexes  of  nickel 

(0),  cobalt (0)  and  iron(0),  84 
( Sarbonyls 
halo,  160 

multiple  bonding  in,  192 
thio,  160 
o-Carboxybenzeneazo-p-cresol,      copper 

lake  of,  757 
o-Carboxybenzeneazo-/9-napht  hoi ,     cop- 
per complex  of,  757 
Carboxylase,  706 
Carboxylate  ion,  as  bridging  group,  34, 

462,  463 
Carboxypeptidases,  metal  activation  of, 

705 
Carrier  agents  in  nickel  electrodeposi- 

tion,  643 
Carrier,  radioactive,  612 
Catalase,  724 

Catalysts,  metal  carbonyls  as,  542 
Catechol,  complexes  of,  25 
Cathodic  reduction 
of  complex  ions,  402-406,  628,  632 
of  negative  ions,  629 
Cation  charge  and  energy  of  coordina- 
tion, 126 
Cation  deformation,  role  in  coordination. 

125 
( "at ionic  complexes 
as  acids,  425 
as  bases,  425 

formed  from  complex  anions,  630 
table  of  acid  strengths  of,  427 
Cerium(III)-Cerium(IV)  couple,  100 
Cerium(IV)  complexes  with  nitrate  and 
perchlorate,  29,  I'd 

( 'eriuinf  III  I  nitrate-,  oxidation  of,   MX) 

( leriumi  IV     perchlorate,  hydrolj  -i-  of, 

Mil 

(  'erium    III     sulfate,  oxidation  of.  400 

Cesium  chloroaurate  1.    IN  .   structure 

of. 


'92 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Chain  length,  determination  by  infrared, 

576 
Chain  phosphates,  769 

complexing  ability  of,  776 
Chain  theory  of  metal  ammines,  102 
Charge  on  complex  beryllium  cations, 

effect  of  anion  penetration,  466 
Charge    reversal   on   micelles    in   metal 

oxide  hydrosols,  466,  467 
Charge-size  ratio,  importance  in  coordi- 
nation, 120 
Chelate    complex   formation    constants, 

178-183,  237,  241,  246 
Chelate-containing  cations  in  poly-acids, 

486 
Chelate,  definition  of,  220 
Chelate  effect,  221 
and  chain  length,  250 
definition  of,  223 
for  polydentate  ligands,  251 
in  terms  of  statistical  model,  250 
relation  to  metal,  251 
thermodynamics  of,  251 
Chelate  rings,  220-252 
formation  by  anion  penetration,  461 
formation  of,  steric  factors  in,  225 
in  complexes  in  chrome  tanning  solu- 
tions, 461 
sizes  of,  225 
Chelate  stability,  factors  in,  224 
Chelate  structures,  stability  of,  221 
Chelates,     synthetic,     oxygen-carrying, 

735 
Chelation  as  a  factor  in  anion  penetra- 
tion, 460 
Chelation,  220-252 
bidentate  groups  occupy  cis-positions, 

277 
by  anions  of  dicarboxylic  acids,  461 
by  anions  of  a-hydroxy  acids,  467 
effect  on  stability  of  complexes,  40,  413 
entropy  effects  in,  249 
Chemical  basis  of  bond  type,  213 
Chemical  polarization  at  electrodes,  632 
Chemical    properties    and    polarization, 

122 
Chlorate  ion 

coordinating  ability  of,  29,  272 
structure  of,  580 
Chloride  ion,  donor  properties  of,  4-20 


Chlorides 
effect  on  electrodeposition  of  arsenic 

and  antimony,  648 
effect  on  electrodeposition  of  tin,  662 
in  nickel  electrodeposition,  656 
Chlorantine  Fast  dyes,  755 
Chloroamminebis(dimethylglyoximino) 

cobalt  (III),  284,  285,  313 
Chloroammineplatinum(II) ,      polymeri- 
zation isomers,  265 
Chloroaquo  -  octammine  -/x  -  amino  -  dico  - 

bait  (III)  chloride,  30 
Chloro  complexes 
absorption  of,  567 
bridged,  4,  8,  17,  81,  365,  462 
in  chromium  electrodeposition,  650 
of  beryllium,  evidence  for,  5 
of  cadmium,  405 
Chlorometallates,  three  coordinate,  385 
Chloropentaquochromium (III)    chloride 

monohydrate,  261 
Chloropentamminechromium (III)    chlo- 
ride, preparation  of,  18 
Chloropentamminecobalt(III)    chloride, 

preparation  of,  17,  153 
Chlorophyll,  739 

Chlorophyll  A,  structure  of,  74,  740 
Chlorophyll  X,  223 
a-Chloropropionic    acid,    resolution    by 

complex  formation,  33 
Chlorothallate(III)  ion,  6,  401 
Chlorotripyridyl  platinum(II)  ion,  288 
Choleic  acids,  559 
Chromate  ion,  structure  of,  580 
Chromate,  spectra  of,  568 
Chromium 
cathodic  reduction  of  cyano  complexes, 

629 
electrodeposition  of,  650 
tanning,  453,  454,  456,  461,  471 
Chromium  carbonyl,   double  bonds  in, 

192 
Chromium(II)-(III) 
ammine  couple,  valence  relations,  186 
aquated  couple,  valence  relations,  186 
cyanide  couple,  valence  relations,  186 
Chromium(I),  stabilization  by  2,2'  di- 

pyridyl,  68 
Chromium(II), 
complexes  of,  154,  411 


INDEX 


793 


cyclopentadienyl  compound  of,  196 
hydrazine  complexes  of,  111 
Chromium  (III) 
ammines,  explosive  character  of,  61 

chloride     hydrates,     equilibrium     be- 
tween, 458 
complexes  of,  154 
a-amino  acid  complexes,  37 
aquotisation,  rate  of,  574 
basic,  with  mixed  bridges,  463 
cyanide  complex,  double  bond-  in, 

194 
dye  complexes  of,  746,  749,  752-761, 

765 
halo  complexes  of,  10,  458 
hydroxo-aquo  complexes  of,  451 
lactate  complexes,  conductivity   of, 

596 
oxalato  complexes,  462 
polynuclear    complexes    with    fatty 

acids,  463 
spectra  of,  130 
hydroxide,  in  dyeing,  744 
number    of    unpaired    electrons    and 

structural  type,  209 
oxide,  hydrous,  decrease  in  chemical 

reactivity  on  aging  or  heating,  470 
oxide  solutions,  effect  of  aging  on  pH, 

465 
Baits,  olated  in  tanning,  471 
salt  solutions  investigated  by  ion  ex- 
change, 459 
salt  solutions,  neutralization  of,  453 
sulfate  solutions,  effect  on  pH  of,  by 
addition  of  neutral  salts,  458 
Chromium (IV)    and    (V),   existence   of, 

411,412 
Chromium (VI),  stereochemistry  of  tet- 

racovalent  complexes,  375 
Chromogen  Red  B,  751 
Chromotropic  acid  anion,  chelate  with 

iron  (III),  231 
( Shromoxanes,  751 

chxn,  see  1,2  frans-cyclohexanediamine 
Chymotrypsin,  metal  activation  of,  703 
ci,  see  citrate  ion 
Circular  dichroi>m.  337,  340 
Cz's-planar     configuration,     assignment 

from  chemical  behavior,  358 
Cis-tratts  isomerism 
infrared  stud}'  of,  577 


in  octahedral  complexes,  277   308 

in  olation,  1 19 

in  planar  complexes,  356,  360 
polarographic  b1  udy  of,  687 
potentiometric  determination  <>f,  594 

Hainan  sped  ra  of,  5M) 

\  ray  study  of,  610 

Citrate  ion,  36,  96 
( 'it  rates 
in  chromium  electrodeposition,  650 

in  copper  and  silver  electrodeposition, 

642 
Citrate  complexes,  ring  size  in,  232 
Citrate   group,    in    beryllate    hydrosols, 

469 
Claus'  theory  of  metal  ammines,  102 
Classification  of  complexes,  151-156 
Clathrates,  561 
Cleve's  salt,  97 
Cleve's  triammine,  97 
Cobalichrome,  738 

Cobalt(II)-(III),    aquated    couple,    va- 
lence relations,  185,  401 
Cobalt 

cationic  complex  in  chloride  solution, 

631 
configurations    of    cobalt  (-1)    and    co- 
balt (II)  complexes,  365 
dye  complexes  of,  746  et  seq. 
effect  of  coordinating  agents  on  elec- 
trodeposition of,  641 
electrodeposition  of,  650 

from  en  and  pn  complexes,  629 
glycyl  glycine      dipeptidase     complex, 

70  1 
in  vitamin  Bi2  ,  737 
stereochemistry    of    Co(C03)NO    and 
Co(C03)  (COH),375 
Cobalt  (0) 
in  carbonyls  and  oitrosyls,  610,  i 
in  cyanide  complex,  92 

Cobalt(I) 

cyano  complexes  o 

existence  of,  1 10 

nitrosyl  halides  of,  535 
Cobalt  (II),  complexi 

bis- (salicyl aldehyde)  ethylenediamine 

complex 

cyanide  complex,  isj 

lopentadieny]  compound 

halo  complex. 


794 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Cobalt(II)  —Cont. 

histidine  complex,  46,  735 

number    of    unpaired    electrons    and 
structural  type,  209 

peptide  complexes,  absorption  specta 
of,  704 

planar  configuration  of,  169 

stereochemistry  of  tetracovalent  com- 
plexes, 375 

thiocyanate  complexes,  688 
Cobalt(II)-(III) 

aquated  couple,  valence  relations  in, 
185 

hexammine  couple,   valence  relations 
in,  185 
Cobalt  (III) 

ammines 
explosive  character  of,  61 
polarographic  reduction  of,  629 

carbonatopentammine  complexes,  con- 
ductivity of,  596 

carbonatotetrammine,    dielectric    in- 
crement of,  599 

complexes 
a-amino  acid  complexes,  37 
binucleate,  448 
spectra  of,  130 

cyclopentadienyl  compound,  498 
cysteine  complex,  730 
ethylenediamine  complexes,  potenti- 

ometric  study  of,  594 
ethylenediaminetetraacetate      com- 
plex,    resolution     and    spectrum 
of,  235 
indinyl  complex  of,  499 
nitroammine  complexes,  spectra  of, 

568 
primary  amines,  complexes  of,  63 
tris-(biguanide)    complex,    stability 
of,  593 
Cobalt(III)  ion 

electronic  structure  of,  166 

hydrated,  184,  218 

oxidizing  power  of,  185 

relative  affinity  for  thioethers  and  for 
oxyethers,  51 

stabilization  by  coordination,  402 
Cobalt  (IV) 

existence  of,  410 

fluoride  complex  of,  10,  188 

peroxo  complexes  of,  26,  410 


reduction    potential    in    peroxo    com- 
plexes, 27 
Cobalt  carbonyl,  structure  of,  521 
Coerulein,  752 
Colloidal  behavior  of  hydrous  oxides 

adsorption  theory  of,  464 

complex  compound  theory  of,  463 
Colloidal  systems,   coordination  theory 

of,  471 
Color 

bond  function  of,  564 

magnetic  data  related  to,  605 

relation  to  complexing,  564 

relation  to  structure,  364 

relation  to  temperature,  564 
Color  of  complexes 

and  the  ionic  model,  130 

relation  to  temperature,  66 
Colors,  mineral,  743 

Colorimetric    methods    of   analysis,    co- 
ordination compounds  in,  688 
Columbium,  see  also  Niobium 

electrodeposition  of,  645,  665 

halo  complex  of,  16 

stereochemistry  of  Nb6Cli4-7H20,  375 
Complex  anions  as  precipitants,  682 
Complex  cations  as  precipitants,  681 
Complex  compound  theory  of  hydrous 

metal  oxides,  463 
Complexes 

classified  on  basis  of  molecular  vol- 
umes, 154 

classified  on  basis  of  chemical  proper- 
ties, 152 

of  zero  charge  in  hydrous  metal  oxides, 
470 
Complex  formation  in  polarography,  696 
Complexing  agents,  see  Chapter  1 

in  electrodeposited  metals,  630 

reducing  character  of,  412 
Complexing  tendencies  of  the  metal  ions 

according  to  periodic  groups,  3 
Complex  ions 

nomenclature  of,  93 

reduction  of,  586 

stabilities  of,  176-183;  221-252 
and  polarization,  125,  127 
and  second  ionization  potentials 
of  metals,  177 

stability  determination,  569 
Complex  species,  identification  of,  405 


INDEX 


795 


Complex  stability  and  ionic  radii,  177 
Compounds,  of  first  order,  168 
Compounds  of  second  order,  168 
Compressibility 
in  study  oi  complexes,  62 1 
structure  determination  by,  624 
( Soncentration  polarisation  at  elect  rodes, 

632 
Condensed  structures,  difficulties  arising 

from,  367 
Condensing  enzyme,  711 
Conductimetric  titrations,  595 
determination  of  degree  of  olation  by, 
455 
Conductivity  Btudies  on  complexes,  113 
Configurations,  see  also  Stereochemistry 
absolute,  581 

among  tetraeoordinate  complexes,  ob- 
served, 355 
assignment    through  olation,  449 
of  molecules  and  factors  determining, 

173 
of  molecules  and  electronic  constitu- 
tion, 174 
of     tetraeoordinate     complexes,     ob- 
served, 355 
of  tetraeoordinate  complexes   related 
to  chemical  reactions,  358 
table  of,  170 
Conjugate  theory  of  ammines,  100 
Continuous  variations,  569 
graphical  method,  574 
magnetic  data,  603 
pH  method,  571 
refractometry,  583 

-pectrophotometric  method,   570,  575 
BUrface  tension,  622 
Coordinate  bond,  1 

and  charge  distribution,  190  et  seq. 
Coordinate  covalent  bond,  157  et  seq. 
Coordinated  group 
inert,  214 
replacement    by   other  donor  groups, 

213-219,342-351,458 
nature  with  respecl  to  stability  of  com 
plex,  412 
Coordinating  ability.  Sec  also  individual 
coordinating  groups 
and  atomic  volume.  120 
and  base  strength  of  ligand,  I7(.»et  seq. 


and  localization  of  negative  charge  in 

ligand,  180 
of  acids,  relation  to  peptizing  ability, 

470 
Of  aliphatic  amine-.  62  67,    181 

of  alky]  substituted  hydrides  of  Group 

V   62  67,  78-84,  128 
of  alky]  substituted  hydrides  of ( rroup 

VI9  23  26,  19,  12s 

of  anions,  in  peptizing  hydrous  oxides, 

469 
Of  anion-,  relative,  459 
Of  cyclic  amine-.  67  60,  72  71,   181 
of  1,3-diketones,  41-45,  181 
influence  of  structure,  182 
of  malonic  acids,  influence  of  structure, 

35,  183 
of  pyridine  and  its  derivatives,  67-69, 
72,  181,400,677,686-691 
Coordination  as  an   acid  base  phenom- 
enon, 177 
Coordination  bond  electrons,  relation  to 

second  band,  565 
Coordination  bond,  homopolarity  of,  563 
Coordination  compounds 
as  factors  in  electrodeposition,  640 
stability  and  cation  charge  and  size, 
124 
Coordination,  energy  of,  174 
Coordination  groups,  relation  to  absorp 

tion  bands,  566 
Coordination  isomerism,  263 
Coordination  number 
abnormally  large,  145 
and  orbital  configurations  (Table)  ,170 
and  radius  ratio,  143 
determination   by   polarography,   584, 

586 
effect  on  electrodeposition  from  cya 

nide  complexes,  645 
estimation  by  electrostatic  methods, 

146 
definition  of,  111, 
fulfilment  of,  U3 
in  reference  to  structure  of  crystals, 

111 
of  cat  ions  for  w  ater  and  ammonia,  1 1 1 
of  metal  ions  in  poly  acids,  476,   177 

17'. » -484 
periodic    generalization    for.    1  13 

relationship  to  energy  terms,  143,  Ml 


796 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Coordination  number — Cont. 

role  of  anion  in  determining,  145 

role  of  ligand  in  determining,  145 
Coordination  number  eight,  170,  394 
Coordination  number  five,  387,  520 
Coordination  number  four 

configurations  observed,  354 

existence  of,  354 

stereochemistry  of,  354-381 
Coordination  number  nine,  397 
Coordination  number  seven,  8,  392 
Coordination  number  six,  165 

stereochemistry  of,  274-353 
Coordination  number  three,  384 
Coordination  number  two,  382 
Coordination  position  isomerism,  270 
Coordination  spheres,  the  possible  ex- 
istence of  several,  21 
Coordination  theory 

early  development  of,  100-118 

modern  developments  of,  119-219 
Coordination  theory  of  flocculation   of 

metal  oxide  sols.,  468 
Copaux's  formulation  of  poly-acids,  473 
Copper 

amine  complexes  of,  63 

configuration  of  complexes,  651 

dye  complexes  of,  747,  et  seq. 

electrode  polarization  in  complex  solu- 
tions, 637 

electrodeposition  of,  642,  651 
from  oxalato  complexes,  652 
from  thiosulfate  complexes,  630 

in  hemocyanin,  735 

in  phenol  oxidases,  723 

polarography  of  complexes  of,  403 

relation   of  oxidation  states   to  elec- 
tronic configuration,  185-186,  369-76 

requirement  in  hemoglobin  synthesis, 
736 

separation  from  bismuth,  666 
Copper(I) 

butadiene  complex,  495 

configurations  of  complexes,  364 

cyanide  complexes,  407 

electron  configuration,  167 

ethylene  complex,  494 

halo  complexes  of,  11 

number    of    unpaired    electrons    and 
structural  types,  209 


stabilization  of,  407 
three  coordinate,  385 
Copper (I)-(II)  couples 

ammine  couple,  valence  relationships, 
403 

aquo  couple,  valence  relationships  of, 
185 

cyano  couple,  valence  relationships  of, 
186 

iodo  couple,  valence  relationships  of, 
185 
Copper  (II)  complexes 

classification  of,  651 

configuration  of,  169,  364 

pentacoordinate,  390 

number    of    unpaired    electrons    and 
structural  types,  209 

relation  of  color  to  magnetic  moment, 
364 

stabilities  of,  181 

stereochemistry     of      tetracoordinate 
complexes,  371 

witha-amino  acids,  37 

with  ammonia,  conductivity  of,  595 

with  arsines,  magnetic  properties  of, 
604 

with  ethylenediaminebisacetylacetone , 
stability  of,  222 

with  glycine,  37 

with  halides,  11 
structure  of  CsCuCl3  ,  367 
structure    of     K2CuCl4-2H20     and 
CuCl2-2H20,  368 

with  peptides,  absorption  spectra  of, 
704 

with  substituted  /3-diketones,  stabili- 
ties of,  182 

with  substituted  malonic   acids,  sta- 
bilities of,  183 
Copper(III) 

complexes    of,    planar    configuration, 
169 

fluoride  complex  of,  188 

iodate  complex  of,  31 

preparation  of  complexes,  407 

stabilization  of,  407 

tellurate  complexes  of,  31 
Copper (II)  chloride  dihydrate,  structure 

of,  368 
Copper-gold     alloys,     electrodeposition 
from  cyanide  solutions,  637 


INDEX 


797 


Copper-olefin  compounds,  i 
Copper(II)  salicylate,  structure  of,  571 
Copper  (II) -5-6ulfo8aIicylaIdehyde,    for- 
mula of,  593 
Copper-tii)  alloys,  electrodepoeitioD  of, 

667 
Coprantine  dyes,  755 
38a'fl  First  Salt,  97 

•.'s  Second  Bait 
Cotton,  dyeing  of,  7 
Cotton  effect,  340,  5S1 

relation  to  structure,  364 

for  tetracoordinate  complexes,  356 
Covalent  bond,  207 

and  isolation  of  cis  and  trans  iosmers, 
211 

and  rate  of  exchange,  211 

and  resolution  of  optical  isomers,  211 

and  trans  effect,  196 

compared  with  ionic  bond,  136,  211,  21S 

early  treatments  of,  157 
Covalent  complexes,  137,  151,  208 
cptn,  see  1,2-  fra/is-Cyclopentanediamine 
Croceo  salts,  98 

Jorgensen's  structure  of,  107 
Cryoscopy-ia-study  of  complexes,  596 
Crystal  field 

effects  on  cobalt  (III) , 

splitting  theory,  218 

theory  of  magnetism, 

rys-baMfrtttce^  and  molecular  configura- 
tion, 173 
Crystal    orientation    and    structure    in 

electrodeposits,  640 
Crystallization  of  metal  in  electrodeposi- 

tion,  633 
Cupferron,  680 
cy,  see  Cyanide  ion 
Cyanide  complexes,  86  et  seq. 

alleviation  of,  87 

cobalt (II),  coordination  number  of,  87 

copper  (I),  88 

double  bonds  in,  193 

gold  (I),  infrared  study  of,  87 

in  electrodeposition,  645 
of  copper,  651 

mixed,  87 

silver  (I),  formation  constants,  88 
Cyanide  ion,  96 

as  bridging  group,  88,  365 


"L> 


completing  through  the  carbon  atom, 

displacement    of    other    coordinated 

groups  by,  ^7 
donor  properties  of,  75,  96 

effect  oo  bridged  structures,  365 
exchange  of,  in  cyanide  complexes,  88 
reaction   with   porphyrin    complexes, 
720,  721,  728,  729,  735 

reaction  with  vitamin  Bi2,  739 
Cyanide  solutions  for  alloy  plating,  668 
Cyanocobalamin,  737 

Cyanocobaltate(II)  ion,  composition  of, 
184 

Cyanonickelate  ions,  385 

1,2-Cycloheptanedionedioxime  in  deter- 
mination of  nickel,  674 

l,2-/rarcs-Cyclohexanediamine,  96 

Cyclohexanediamine   chelates,   228,   314 

1,2-Cyclohexanedionedioxime   in   deter- 
mination of  nickel,  674 

1 , 2-Cyclohexanediaminetetraacetate, 
calcium  complex 
chelate  effect  in  stability,  251 
steric  factors  in  stability,  243 

Cyclohexanone,  metal  complexes  of,  498 

Cyclohexene 
mercury  complex  of,  497 
palladium  complexes  of,  493 
platinum  complexes  of,  492 

Cyclooctatetraene,  543 

Cyclopentadiene  complexes,  207,  498 

1 , 2-/rans-Cyclopentanediamine,  96 

Cyclopentanediamine   chelates,   228 
stereochemistry  of,  314 

Cysteine,  oxidation  catalyzed  by  iron, 
731 

Cysteine  complexes,  730 

Cysteine-cystine  system,  730 

Cystine  complexes,  730 

Cytochrome-a,  729 

Cytochrome-b,  729 

Cytochrome-c,  7_'7 

Cytochrome  oxidase. 

Cytochrome  system,  7_'7 

Dacron,  dyeing  of, 

Decachloro-/i-o\odiruthenate(IV/)     ion, 
icture  of,  W7,  201,  - 

ilicylaldel 

copper  (II),  2 


708 


HIEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Decammine-/u-peroxocobalt(III)-cobalt- 

(IV)  ion,  203 
Decammine-ju-peroxodicobalt(III)      ion, 

203 
Decarboxylation 
mechanism  of,  707 
metal  activation  of,  706 
Decomposition  temperature,  relation  to 

ionic  volume,  621 
Decoordination  of  complex  ions,  633 
Deformation  of  cation,  role  in  coordina- 
tion, 125 
Dehydration  of  hydrates,  20,  454 
Dehydrogenases,  727 
Delta  bond,  201 
Delta  orbitals,  199 
Deolation 
effect  of  penetrating  power  of  anions 

on,  469 
in  dissolution  of  hydrous  metal  oxides, 
468 
Depolarization,  degree  of,  579 
Desolvation  of  complex  ions  in  electro- 
deposition,  633 
Desoxycholic  acid,  559 
Deoxolation,  rate  of,  457 
Diallylamine,  96,  491 

platinum  complexes  of,  491 
Dialysis 
in  study  of  complexes,  618 
of  basic  chromium  salt  solutions,  454 
Dialytic  constant,  619 
Diamagnetism,   absorption   as  criterion 

of,  171 
Diamagnetism  and  EAN  concept,  162 
Diamagnetism  of  complexes,  600-606 
1,2-Diamines,  as  complexing  agents,  63 
2,2'-Diaminobiphenyl,  cobalt  (III)  com- 
plexes of  67,  256 
a,7-Diaminobutyric    acid,   copper   com- 
plex of,  37 
Diaminocyclohexane-N,N'-tetraacetate, 

complexes  with  alkaline  earths,  230 
Diaminoglyoxime    in    determination    of 

nickel,  674 
l,2-Diaminopropane(propylenedia- 
mine) 
complexing  by,  63 
geometric  isomerism  due  to,  285 
optical  activity  of  its  complexes,  299, 
317-319 


stereospecific    reactions    of    its    com- 
plexes, 315 
Diamminecopper(II)  acetate,  77 
Diammine-ethylenediamine-bis    (acetyl  - 

acetone)cobalt(III),  320 
Diamond  Black  PV,  749 
Diamond  Flavine  G,  753 
Diaquobisoxalatochromate(III)  ion,elec- 

trodeposition  from,  650 
Diaquodiammineplatinum(II)   ion,   acid 

properties  of,  429 
Diastereoisomers,  332,  717 
Diazoamino   compounds,   chelation   by, 

74,  226 
Dibasic  acid  complexes,  254,  255 
dibenz,  see  Dibenzoylmethane 
Dibenzoylmethane,  96 

coordination  compounds  of,  41 
Dibenzoylsulfide,  complexes  with  gold, 

50 
Dibromobis(ethylenediamine)cobalt 

(III)  ion,  18 
1,3-Dicarbonyl  compounds,  see  1,3  dike- 
tones 
2,2'-Dicarboxyazobenzene,    copper  lake 

of,  757 
czs-Dichlorobis(ethylenediamine)chro- 
mium(III)  chloride,  preparation  of, 
18 
Dichlorobis(ethylenediamine)cobalt 
(III)  ion 
aquation  of,  302 
cis-trans  conversions  in  reactions  of, 

301 
isomers  of,  18 
reactions  of,  303-306 
Walden    inversions    in    reactions    of, 
344-347 
irans-Dichlorobis(ethylenediamine)plat- 
inum(IV)    ion,    preparation   of,  280 
Dichlorobis(oxalato)iridate(III)  ion,  301 

cis  and  trans  isomers  of,  281,  301 
Dichlorobis(oxalato)rhodiate(III)     ion, 

301 
Dichlorobis-  (phosphorus        trifluoride)  - 

platinum (II),  85,  205 
Dichlorobis  -  (propylenediamine)cobalt- 
(III)  ion 
diamagnetism  of,  211 
geometric  isomerism  of,  285 


INDEX 


799 


failure  to  exchange  with  radioactive 

Co(II),  211 
optical  activity  of,  299,  317-319 
stereospecific  reactions  of,  315 

Dichlorobis-(triethylphosphine)plati- 
DUm(II),  cist  rans  conversion  of,  205 

Dichlorodiamxnine  el  hylenediamine  co- 
l)alt(III)  ion,  isomers  of,  293 

Dichlorodiammine  platinum  (II) 
determination  of  configurations  of  iso- 
mers of,  360 
isomers  of,  356 
polymerization  isomers  of,  265 

Dichlorodiethylenetriamineplatinum- 
(II)  hydrochloride,  324 

frans-Dichloroethylene,    complex     with 
platinum (II),  492 

Dichloro-ethylenediamine-diammine 
cobalt  (III)  ion,  isomers  of,  293 

Dichlorotetraaquochromium(III)     chlo- 
ride, 458 

Dichlorotet  raaquochromium(III)     chlo- 
ride dihydrate,  261,  574 

Dichlorotet ramminecobalt (III)    ion,   cia 
and  trans  isomers  of,  279,  291 

t><,  //.^-Dichlorotet  rapyridylcobalt  (III) 
ion,  289 

I)ichlorotriethylenetetraminecobalt(III) 
ion,  320 
stereochemistry  of,  289 

Die  vanoamminenickel  (II) ,  benzene 

clathrate  compound  of,  498 

Dielectric    constant,   597 
relation  to  polarization,  597 

Dielectric  increments,  method,  599 

dien,  see  diethylenetriamine 

Diethylamine,  coordinating  ability  of, 
180 

Diethylenetriamine,  96 
chelation  by,  64 

Diethyl  gold  (III)  bromide,  structure  of, 
365 

Diethylsulfide,   compounds    with    plati- 
num. [I)chloride,  if) 

Diethylditbioethane,     complexes     with 
platinum  (II  .  60 

diffusion  coefficient-,  relation  to  polar 
ography,  586 

Diffusion  of  complex  ion-  to  electrode 
surfaces,  632 


1,3-Dike  tones 
cationic  complexes  of,  13 
chelation  as  a  result  of  enolization,  n 
coordinating  ability  of,  1 1 ,  182 
complexes  of  Cu (II),  stabilities  of,  L82 
mixed  complexes  of,  13 
resonance  in  stabilizing  of  chelates  of, 

246 
separation  of  metal  ions  by,   II 
•  sodium  complexes  of,  2,  182 

stability  of  complexes  of,  176 
Diket  onedithiosemicarbazone  c om j  ilexes 
with  copper (II)  and  nickel (II),  54 
3,3'-  dimethyl  -4,4'-  dicarbethoxydi i  >y  i 

romethene,  242 
Dimethyldithioethylene,    reaction    with 

copper(II)  and  gold(III),  48 
Dimethylglyoxime 
in  determination  of  nickel,  674 
monobasic  anion  of,  96 
spectra  of  complexes  of,  568 
2,9-Dimethyl-l,10-phenanthroline   com- 
plexes with  Cu(II)  and  Fe(II) ;  steric 
strain  in,  238 
Dinitrobis-(ethylenediamine)cobalt(III) 
ion, 
asymmetric  synthesis  of,  351 
chemical  behavior  of  cis  and  trans  iso- 
mers, 294 
preparation  of  cis  and  trans  isomers, 
280 
Dinitrobis-  (1-propylenediamine)     cobalt 
(III)  bromide,  configuration  of,  294 
Dinitrodiammineplatinum(II),  reduc- 
tion of,  659 
Dinitro-ethylenediamine-propyleiKM  1  i 
amine  cobalt  (III)   ion,  stereochem- 
istry of,  286,  318 
Dinitro(N -methyl  -N-ethylglycinato) 
platinate(II)  ion,  optical   resolution 
of,  38 
Dinitro-oxalato  diammine  cobaltate 

(III)  ion,  isomers  of,  292 
Dinitroresorcinol 
cobaltammine  complex  of,  7 17 
iron  complex  of,  717 
Dinitrotetrannninecoh.ilt  |  1 1 1     ion, 

chemical  behavior  of  cis-  and  trans 

isomers,  294 
preparation  and  properties  of  cis    and 

trans -isomers,  280 


800 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Dinitratotetramminecobalt(III)  nitrate, 

28 
Dinuclear  metal  carbonyls,  510 
Diolefins, 

coordination  of,  491 
separation  from  monoolefines,  500 
Dioximes 

cobalt(III)  complexes  of,  77 

iridium  complexes  of,  77 

isomers  of,  and  their  ability  to  form 

complexes,  77,  675 
rhodium,  complexes  of,  77 
use  in  analytical  chemistry,  673 
Dipentene,  platinum  complexes  of,  491, 

492 
Dipeptidases,  metal  activation  of,  704 
Diphenylcarbazide,  use  in  analysis,  692 
Diphenylcarbazone,  use  in  analysis,  692 
Diphenylethylene,   complex  with  plati- 
num (II),  492 
1,2-Diphenylethylenediamine,     96,     see 

also  Stilbenediamine 
Diphenylthiocarbazone,  691 
Dipole  moment,  596 

and  coordinating  ability,  123 

in  study  of  complexes,  596 

of  alkyl  derivatives  of  H20,H2S,  NH3,- 

PH3,  128,  129 
of  platinum (II)  complexes,  363 
of  tetracoordinate  complexes,  357 
use  in  distinguishing  cis-  trans  isomers, 
299 
Dipropylgold(III)  cyanide,  structure  of, 

88,  365 
dipy,  see  2,2'-Dipyridyl 
2,2'-dipyridyl,  96 

complexes   with  iron  (II),   steric   hin- 

rance  in,  237 
coordinating  ability  and  base  strength, 

205 
in  colorimetric  analysis,  689 
racemization  of  complexes  of,  328 
specificity  of  methyl  substituted,  690 
stability  of  complexes  of,  67 
substituted  derivatives,  effect  on  com- 
plexing  tendency  of,  67 
Directed  covalent  bonds,  356 
Direct  reduction  of  complex  ions,  628 
I  )i-sociation  of  complexes 
prior  to  electrodeposition,  626 
rate  of,  632 


Dissociation    constants    of    complexes, 
130,  402,  428 
calculation  of,  405 

Disalicaltriethylenetetramine,  as  a  do- 
nor, 321 

Disulfitotetramminecobaltate(III)     ion, 
configuration  of,  281 

Disk  method  for  infrared  measurements, 
577 

Dissimulation,  576 

Dissociation  constants,  see  also  Forma- 
tion constants,  Stability  constants 
and  Stability  constants 
graphical  method  for,  574 
polarographic  method  for,  584,  586 
tracer  method  for,  617 

Disulfides,  chelation  by,  50 

Disulfitotetramminecobaltate(III)     ion, 
293 

3 , 6-Dithia-l ,  8-bis-  (salicylideneamino) 
octane  complexes,  235 

Dithiane,  complexes  of,  48 

Dithiobenzoic  acid,  nickel  (IV)  complex 
of,  56 

Dithiocyanatodiethyldigold,      structure 
of,  53 

Dithiocyanatotetrapyridinenickel     (II) , 
stability  of,  67 

Dithio-j3-isoindigo,  metal  compounds  of, 
762 

Dithiooxamide,  (rubeanic  acid) 

complexes    with    nickel,    cobalt,    and 

copper,  56 
derivatives  of  diethyl  gold  bromide,  57 

Dithizone,  use  in  analysis,  691 

dim,  see  diallylamine 

DMG,  see  dimethylglyoximine  monbasic 
anion,  96 

Dodecammine  -jj.  -  hexol  -  tetracobalt  (III) 
ion, 
isomers  of,  266 
resolution  of,  277,  323 

Donor,  1 

Donor  groups,  abbreviations  for,  96 

Double  bond,  see  Ethylene,  Olefins,  and 
Unsaturated 

Double  salts,  early  theories  of,  107 

Drechsel's  chloride,  97 

dsp2  hybridization,  169 

dsp2  hybridization,   and   magnetic    mo- 
ments of  complexes,*  172 


INDEX 


s()| 


hybridization.   L66 

»and  magnetic  moments  of  complexes, 
172 
hybridisation,  109 
Durchdringungskomplexe,  151 
Dunrant'a  Salt.  23,  97 
Dye-metal-fiber-interactions,  763 
D' 
—COOH,— OH  substituted,  763  el  seq. 
coordination  compounds  as,  743  el  seq. 
o-dihydroxy-substituted,  749 
—NO,  —OH  substituted,  746 
sulfur-containing.  7t'>_, 
Dynel,  dyeing  of,  766 

EAN,  Bee  Effective  atomic  number 
Earnshaw'fl  theorem  of  electrostatics,  162 
Edge  displacement  in  substitution  reac- 
tions, 307 
EDTA,  Bee  Ethylenediaminetetraacetic 

acid 
Elective  atomic  number,  159 
as  -t ability  factor,  414 
in  carbonyls  151,  518,  519 
in  nitrosyls,  533 
Eight  coordinate  configurations,  394 
EUdit-membered  rings,  256,  260 
Elaidic    acid,   methyl   ester,    compound 

with  silver.  496 
Electrode  irreversibility  in  reactions  of 

complexes,  406 
Electrodeposition,  see  also  the  individual 
metal- 
coordination  compounds  in,  625-671 

theory  of,  625 
relation  to  electronic  configuration  of 

complex,  638 
relation  to  stability  of  complex,  642 
ectrodeposite 

chlorine  and  nitrogen  in.  630 
crystal  structure-  of.  640 
inclusions  in,  640 
Electrode  potential,  effect  on  character 

of  electrodeposits,  641 
Electrolytic  separation  of  metals  from 

complex  compounds,  666 
Electrometric  method-  in  study  of  com- 
plex 
eleetrometric  titration-.  600 
electrolytic  transference,  618 

force  measuremen- 


Electron    accepting    ability    of    donor 

atom-.   1,   194 

Electron  diffraction 

application  to  structure  of  complexes, 

nor, 

in  Btudy  of  bond  types,  213 

in  Btudy  of  complexes,  607 

in  study  of  betracoordinate  complexes, 

354 

Electron  distribution,  B,p,d-orbitals,  163 

Electronegativity    and    molecular    con 
figuration,  173 

Electronegativity   and   trans   effect,    1  '»«> 

Electronegativity    of    metals    in    com- 
plexes, 413 
and  stability  of  complex,  175 

Electroneutrality  principle,  190  et  seq. 

Electronic  configuration  and  electrode- 
position  mechanism,  638 

Electronic    constitution    and    molecular 
configuration,  174 

Electronic  effects  on  stability  of  chelate, 
244 

Electronic  isomerism,  272 

Electron  promotion,  167,  169,  1&4,  187 

Electron  quantization,  158 

Electrons,  shared  pair,  157 

Electrons,  stereochemical^'  active  pair, 
170 

Electronic  shifts,  relation  to  absorption, 
567 

Electronic  theory  of  acid  and  bases,  421 

Electronic  vibrations,  565 

Electron  transfer 
at  electrodes,  633 
in  reactions  of  complexes,  20,  406 

Electrophoresis,   investigation    of   basic 
chromium  -alt  solutions  by,  454 

Electrostatic  attraction  as  force  in  bind- 
ing of  complexes,  120 

Electrostatic     theory     of     coordination 
compounds,    modern    development-. 
119 
B  Et  hylenediamine 

enac,    Bee    Ethylenediamine-bis- (acetyl - 

acetone) 
enBigH,  see  Ethylenebiguanide 
Endopeptidases,  metal  activation  of,  703 
Energy  of  coordination,  137 
and  size  of  coordinated  group,  l-'7 

Lined  relation  to  AH,  224 


802 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Energy  of  coordination — Cont. 
of  ammines  of  zinc,  iron(II)  and  man- 
ganese (II),  142 
Enneachlorodithallate(III)     ion,    struc- 
ture of,  7 
Enneachloroditungstate    ion,    structure 

of,  16 
Enolase,  metal  requirement  of,  711 
Enterokinase,  metal  activation  of,  703 
Enthalpy  changes  in  chelation,  251 
Enthalpy   contribution    to    chelates    of 

copper  (II),  245 
Entropy  effects  in  chelation,  249,  251 
Entropy  effects  in  complex  stability,  224 
Enzyme  complexes,  metal  specificity  of, 

705 
Enzyme-like  action  of  complexes,  316 
Erdmann's  Salt,  97 
Erdmann's  "trinitrite",  113 
Erganil  dyes,  756 
Ergansoga  Brown  3R,  753 
Eriochrome  Azurol  B,  754 
Eriochrome  Black  T,  as  metal  indicator, 

685 
Eriochrome  Blue  Black  R,  755 

chromium  complex  of,  756 
Eriochrome  Flavine  A,  753 
Eriochrome  Red  B,  chromium  complex 

of,  755,  756 
Erythrochromic  ion,  271,  457 
Ethanolamine, 
complexes  of,  25 
coordinating  ability  of,  180 
Ether 
complexes  of,  25 
electrodeposition    from    solutions    in, 

670 
mixed  complexes  with  pyridine,  25 
Ethers,  coordinating  ability  of,  23,  123, 

129 
Ethylamine,  96 

coordinating  ability  of,  180 
Ethyl   bromide,   electrodeposition   from 

solutions  in,  670 
Ethyl  dithiocarbamate  nickel  (II)   com- 
plex; four  membered  ring,  227 
Ethylene,  see  also  Olefins  and  Unsatur- 
ated 
absorption  by  copper(I)  chloride,  494 
coordination  with  aluminum,  497 
palladium  complexes  of,  493 


platinum  complex  of,  488,  492 
trans  influence  of,  148,  149,  490,  491 
Ethylenebiguanide,  96 
Ethylenediamine 

as  bridging  group,  489 

complexing  by,  63 

C  substituted,  64 

coordinating  ability  and  base  strength, 

205 
formation  of  five  membered  rings  by, 

228 
monodentate,  65,  489 
N-substituted,  66 

stability   of   chelates   compared   with 
those    of   trimethylenediamine,    230 
Ethylenediaminediacetic-dipropionic 

acid,  41 
Ethylenediamine-bis  (acetylacetone) ,  96 
as  a  donor  molecule,  319 
copper (II)  complex  of,  390 
Ethylenediaminetetraacetato  cobaltate- 

(III),  stereochemistry  of,  320 
Ethylenediaminetetraacetic     acid,     96, 
223,  235,  777,  778 
complexes  of,  577 

stability,  uses,  39 
conductometric  titration  of,  780 
effect  of  calcium  salts  on  neutralization 

curve  of,  779 
hexadentate,  287 
homologs  of,  229 

relative  complexing  tendencies,  781 
in  dimeric  complexes,  253 
palladium  (II)  complex  of,  41 
pentadentate,  287 

rare  earth  complexes  of,  stability  con- 
stants, 179,  589 

use  in  iron  electrodeposition,  655 
use  in  water  softening,  777-782 
Ethylenediaminepropionates,      chelates 

of,  230 
Ethylenedibiguanide,  complexes  of,  70 
Ethylenethiocarbamide,  96 
Ethylenethiourea,  96 
complexes  of,  383,  385 
reaction  with  copper (II),  407 
Ethylidene  structure  of  platinum-olefin 

compounds,  503 
Ethylercaptan,  as  a  bridging  group,  83 
Ethylxanthogenate  nickel  complex;  four 
membered  ring  in,  227 


INDEX 


803 


etn,  Bee  Ethylamine 

etu,     su     Ethylenethiocarbamide     and 

ethylenethiourea 
Exchange     between     oxalate 
ami   trisoxalato   aluminum  (III)   ions, 

and    trisoxalato    chromium(III)    ions, 
326,  829 
and    trisoxalato   Lron(III)    ions,    326 

Exchange  of  functional   groups,   metal 

catalysed,  712 
Exchange  rate 

and  structural  features  of  complex  ion, 

213 
relation  to  bond  type,  615 
Exchange  reactions 

as  criterion  for  bond  type,  211 
of  complexes,  611-618 
of  ferrocyanides,  628 
of  platinum (II)    complexes,   relation- 
ship to  stability,  12 
of  trisoxalato  complexes,  326,  629 
results  compared  with  magnetic  sus- 
ceptibility data,  211 
results    compared    with    stability    of 
isomers,  211 
Exchange  resins,  use  of,  622 
Exopeptidases,  metal  activation  of,  703 
Expansion  of  crystal  lattice  as  size  of 

ligand  increases,  139 
Explosive  character  of  some  ammines,  61 

Fajans'  Quanticule  Theory,  132,  203 

Ferricyanide,  see  Hexacyanoferrate(III) 

Ferritin,  736 

Ferrocene,  494 

Ferro-  and  ferricyanide  pigments.  744 

structures  of,  90 
Ferrocyanide,  see  Hexacyanoferrate(II) 
Ferroin,  686 
Ferromagnetism,  600 
Fiber-metal-dye  interactions,  763 
First  absorption  band  of  complexes,  565 
First  order,  compounds  of,  158 
Fischer's  Salt,  97 

Five-coordinate  configurations,  387 
Five-membered  rings,  stability  of,  227 
Flash  electrodeposits,  639,  739 
Flavo  salts,  98 

Jorgensen's  structure  of,  107 
Flocculation  of  metal  oxide  sols, 


Fluoresence    of    complexes,    analytical 

uses  of,  604 
Fuoride  ion 
as  masking  agenl  for  molybdate  and 

tungBtate,  L6 
donor  properties  of,  I  20 
reaction  with  peroxidase,  726 
stabilisation  of  hi^h  oxidation  Btates 
by,  9 
Fluoro  complexes 

coordination  numbers  in,  144 

of  aluminum,  occurrence  and  proper 

ties  of,  6 
of  antimony,  configuration  of,  8 
of   tellurium,   electrodeposition    from, 

662 
rate  of  hydrolysis  of,  218 
stabilization  of  high  oxidation  states 

in,  9 
use  in  separation  of  niobium  and  tan- 
talum, 16 
use   in   separation   of   zinconium    and 
hafnium,  16 
Force    constants    of    coordinate    bonds 

from  Raman  spectra,  213 
Forced  configurations,  412 

for  tetracovalent  complexes,  354,  363 
Formamide,  electrodeposition  from  solu- 
tions in,  670 
Formate  as  bridging  group,  33 
Formation  constants,  see  also  Dissocia- 
tion    constants,     Instability     con- 
stants, and  Stability  constants 
determination  by  electrode  potentials, 

593 
determination  by  polarography,  405 
of  metal  chelates,  177 
Formato  complexes  of  chromium,  460 
Formatopentammine  cobalt(III)  ion,  33 
Formazyl  compounds,  metal  complexes 

of,  759 
F-strain  in  comple 

Four-coordinate,  see  Tetracoordinate 
Four  membered  ring 
evidence  for,  226 
in  bridged  molecule.  - 
Fourteen-member  ring,  260 
Fourth  absorption  band,  667 
Functional   groups,   blocking  by   metal 

tons,  71 1 
/3-Furfuraldoxime,  complexes  of,  78 


804 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Furildioxime  in  determination  of  nickel, 

674 
Fused  rings,  increased  stability  in,  221 

see  Chelate  effect 

Gallein,  752 
Gallium 

electrodeposition  of,  653 

halo  complexes  of,  6 

oxalato  complex,  claimed  resolution  of, 
212 

stereochemistry  of  tetracovalent  com- 
plexes, 374 

structure  of  halides  of,  365 
Gallocyanine,  752 
Gambine  Y,  746 
Geometrical  isomerism 

effect  on  acid  strength,  429 

in  hexacovalent  complexes,  277-308 

in  tetracovalent  complexes,  356  et  seq. 
Geometric  isomers 

absorption  spectra  of,  294-297,  364 

anionic  complexes,  281 

cationic  complexes,  279-281 

chemical  behavior  of,  294 

configuration  determination  by  chem- 
ical methods,  294,  360 

configuration  determination  by  physi- 
cal methods,  294,  361-364 

dipole  moments  of,  299,  363 

infrared  spectra  of,  300,  371-381 

interconversion  of,  301 

ion  exchange  separations  of,  300 

magnetic  susceptibilities  of,  300,  357, 
359,  364 

nomenclature  of,  94 

nonionic  complexes,  282-284 

polarographic   measurements    on,   299 

polydentate  donor  ligands  in,  286-289, 
318-329,  358 

Raman  spectra  studies  of,  300,  371-381 

rotatory  dispension  studies  of,  298,  338 

solubilities  of,  300 

substitution  reactions  of,  294,  299,  301, 
303-308,  327,  347,  348,  358 

x-ray  diffraction  studies  of,  297,  356, 
361,  367 
Gerard's  Salt,  97 

Germanium,  electrodeposition  of,  653 
Gibb's  Salt,  97 


Glutathione,  716 

complex  with  iron,  731 
gly,  see  Glycinate  anion 
Glycerol,  chelation  by,  24 
Glycinate  anion,  96 
Glycine  complexes,  structure  of,  578 
Glycine   in   silver   and   copper   electro- 
deposition, 642 
Glycol,  chelation  by,  24 
Glycollic  acid 
copper  complex  of,  36 
rare  earth  complexes  of,  36 
Glycylglycine  dipeptidase,  704 
Gmelin  reaction,  539 
Gold 
electrodeposition  of,  653 

colored,  744 
relation  of  oxidation  states  of  to  elec- 
tronic configuration,  369 
structure  of  Cs2Au2Cl6,  17,  368 
Gold  (I) 
cyanide,  structure  of,  89 
cyano-o-phenanthroline  complex, 

x-ray  structure  of,  609 
halo  complexes  of,  17 
number    of    unpaired    electrons    and 

structural  types,  209 
stereochemistry  of  tetracovalent  com- 
plexes, 371 
two  covalent,  383 
Gold (II),    non-existence   in    Cs2Au2Cl6, 

17,  368 
Gold  (III) 
bromide,  alkyl  derivatives  of,  19 
configuration  of  complexes,  169 
dipropyl gold  (III)    cyanide,    structure 

of,  88,  365 
halides,  structure,  365 
halo  complexes  of,  17 
stereochemistry  of  tetracovalent  com- 
plexes, 371 
thiocyanato  complex  of,  53 
Graham's  Ammonium  Theory,  101 
Grain  refining  agents  in  electrodeposi- 
tion, 642 
Grenzsauren,  473 
Grignand  reagent,  as  an  ether  complex, 

25 
Gro's  Salt,  97 
Group  IIIA,  halo  complexes  of,  6 


ixni-x 


so:, 


Group  VA  alky]  substituted  hydrides, 
coordinating  ability  of,  128 

Group  VIA  alky]  substituted  hydrides, 
coordinating  ability  of,  128 

Grunberg's  test  for  cis-trans  configura- 
tions, 35,  359 

Guany]  thiourea,  complexes  of,  55 

Hafnium,  halo  complexes  of,  16,  393 

Halt-cell  reaction,  399 
Half-wave  potentials,  402 
Halide  complexes,  5-20 

stabilities  of,  4 
Halide  coordinators,  3 
Halide  groups  as  less  abundant  donor 

species,  17-20 
Halide  ions 

analogy  to  hydroxy  ion,  4 

donor  properties  of,  4 
Halogen  bridges,  18,  527 
Halogen  in  metal  electrodeposited  from 

halide  solutions,  630 
Halogens  as  central  atoms,  384,  386 
Halometallates 

heptacoordinate,  393 

pentacoordinate,  388 
HD,   see  Dimethylglyoxime   monobasic 

anion 
Heisenberg's  Uncertainty  Principle,  162 
Heliogen  Blue,  761 
Hematin,  718 
Heme 

iron  in,  718 

magnetic  moment  of,  718 

reaction  with  monodentate  complexing 
agents,  720,  721 

reaction  with  oxygen,  719 
Hemichromes,  7_,n 
Hemiglobin,  734 
Hemin,  223,  718 

reaction  with  monodentate  coordinat- 
ing agent-.  I'll),  721 

structure  of,  74,  719 
Hemochromes, 
Hemocuprein,  726 
Hemocyanin,  45,  74, 
Hemoglobin,  15   }  _ 
icoordination, 

tluoro  anions  of   niobium  Y;    and 
tantalum  (V;,  16 


Heptafluorocobalta'-    l\     ion,  9 

structure  of,  393 
Beptafluorodiantimonatel  III)  ion, 

structure  of,  8 
Heptafluorohafniate,  structure  of, 
Heptafluorosirconate,  structure  of,  393 
Heterocyclic  amines,  coordinating  abil- 
ity of,  67,  677,  678,  686-691 
Heteropoly-acids,  472  et  seq.,  see  also 
Poly-acids 
central  atoms  in,  474 
definition  of,  472 
Hexaaquochromic  ion 
as  an  acid,  426 
isomers  of,  574 
Hexabromostannate(IV)   ion,  reduction 

of,  404 
Hexachlorogermanate  ion,  7 
Hexachloroiridate(IV)    ion, 
electron  configuration  of,  204 
paramagnetic  resonance  of,  204 
Hexachlororuthenate(IV)    ion,   electron 

configuration  of,  168 
Hexachlorostannate(IV)   ion,   reduction 

of,  404 
Hexachlorostibnate(V)     ion,     reduction 

of,  404 
Hexacovalent  atoms,  octahedral  struc- 
ture of,  274-277 
Hexacovalent  carbon,  165 
Hexacyanocobaltate(II)  ion,  401 
Hexacyanocobaltate(III)  ion,  91,  410 
Hexacyanoferrate(II)  ion,  90 
cathodic  reduction  of,  628 
dissociation  of,  628 
double  bonds  in,  193 
electron  configuration  of,  193 
stability  toward  oxidation,  400 
Hexacyanoferrate(III)  ion,  90 
alloy  electrodeposition  from,  667 
cathodic  reduction  of,  628,  632 
electron  configuration  of,  166 
negligible  exchange  with  labeled  cya- 
nide ion,  212 
reduction  of,  639 
Hexacyanomanganate (I)  ion,  409 

. •anomannanate(III)  ion,  409 
Hexafluorochromate(IV)  ion,  412 
Muorocobaltate(III)  ion,  10 
structure  of,  167 


806 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Hexafluorocuprate(III)  ion,  169,  407 

potassium  salt,  11 
Hexafluorogermanate  ion,  7 
Hexafluoromanganate(IV)  ion,  10 
Hexafluoronickelate(IV)  ion,  10,  409 
Hexafluoroniobate(V)  ion,  16 
Hexafluoroxybismuthate(V)   ion,  8 
Hexafiuoropalladate(IV)  ion,  13 
Hexafluoroplatinate(IV)  ion,  12 
Hexafluorosilicate  ion,  7 
Hexafluorotantalate(V)  ion,  16 
Hexahydroxystannate(IV)  ion,  7,  663 
Hexamethylenediamine,  complexing  by, 

64 
Hexamminecobalt(II)  ion,  properties  of, 

152 
Hexamminecobalt(III)  ion 

electronic  structure  of,  166 

exchange  of  hydrogen  atoms  in,  426 

failure  to  dissociate,  1 

in  detemination  of  ferrocyanide,  682 

in  determination  of  vanadium,  681 
Hexammineplatinum(IV)  ion,  acid  char- 
acter of,  121,  429 
Hexammine-/i-triol-dicobalt(III)  ion,  448 
Hexaquochromium(III)     chloride,     261, 

458 
Hexaquotin(IV)  ion,  reduction  of,  404 
Hexol  salt,  23,  310,  323 
Histidine,  cobalt (II)  complex  of,  46,  735 
Hittorf,  transference  numbers  of,  618 
Hofmann  bases,  complexing  by,  64 
Hund's  rule  of  maximum  multiplicity, 

166 
hx,  see  Hydroxy] amine 
H4Y,     see     Ethylenediaminetetraacetic 

acid 
Hybridization  of  orbitals,  164 

configurations  with  coordination  num- 
ber four  resulting  from,  359 

in  complex  formation,  415 
d2sp3  and  sp3d2  orbitals,  214 
Hydrated  ions 

as  acids,  425-431 

dehydration  of,  20,  454 

early  theories  of,  107 

exchange  reaction  with  solvent  water, 
21 
Hydrate  formation 

fractional   with    zirconyl   compounds, 
455 


nature  of  attractive  forces,  20 
relationship  to  size  and  charge  of  ca- 
tion, 21 
Hydrate  isomerism,  263 
Hydrate  isomers,  dehydration  of,  20 
Hydrate  isomers  of  chromium  (III)  chlo- 
ride, 262,  574 
Hydrate  isomers  of  rutherium(III)  chlo- 
ride, 14 
Hydration  effects  in  chelation,  252 
Hydration  of  poly-acids,  478,  480 
Hydrazine  complexes,  69 
chelation  in,  225 
with  chromium,  411 
with  palladium  and  platinum,  225 
Hydrocarbons,  unsaturated,  see  Ethyl- 
ene and  Olefin 
Hydrogen,  nascent,  in  silver  deposition, 

626 
Hydrogen  bonding 
in  aquo  and  hydroxo  complex  ions,  461 
relation  to  spectra,  567 
Hydrogen  cyanide 
electrodeposition    from    solutions    in, 
670 
Hydrogen  peroxide 
decomposition  by  metal  enzymes,  724 
"of  crystallization,"  26 
reaction    with    porphyrin    complexes, 
721,  724 
Hydrogen  sulfiide 
coordinating  ability  of,  123,  129 
solvent  properties  of,  48 
Hydrolysis 

accompanied  by  olation,  452,  468 
of  aluminum  salts,  451 

hydrous  aluminum  oxide  sols,  464 
of  aquo  compounds,  425-431 

factors     affecting,  451 
of  chromium  salts,  451 
Hydrosols,  463-471 
Hydroxide  ion,  coordinating  ability  of, 

22 
Hydroxo   bridge,    colloidal   oxides   and, 

22,  448-470 
Hydroxocobalamin,  738 
Hydroxo  complexes 
basicity  of,  424 

in  basic  chromium  salt  solutions,  454 
in  formation  of  ol  bridges,  449-455 


INDEX 


so; 


of  cerium .  401 
table  of.  442 
Bydroxo-complex     theory     of     ampho- 
tericin. 138 
Hydroxo  group 
as  bridge 
in  colloidal  oxides,  22,  148  470 
in  polynuclear  complexes,  22,   >s 
conversion  of  aquo  group  t<>.  161,  152, 

453,  465 
coordinating  ability  of.  22 
decrease  in  reactivity  by  olation.   166, 

470 
displacement  by  anion,  458,  405,  469, 

471 
distinction  from  ol  group,  448 
in  micelles   of  metal   oxide   hydrosols, 

464 
in  precipitated  hydrous  metal  oxides, 

470 
olated,  reaction  with  acid,  455 
Hydroxyacetone,       coordination       com- 
pounds, of,  41 
a -Hydroxy  acid  anions 
anion  penetration  by,  466 
boron  complexes  of,  36 
chelation  by,  35,  467 
copper(II)  complex,  36 
peptization  of  hydrous  zirconium  oxide 

by,  467 
solution  studies  of  complexes  of,  36 
1-Hydroxyanthraquinone,    metal     com- 
plexes of,  751 
o-Hydroxyazobenzene,    copper   lake    of, 

756 
o-Hydroxybenzeneazo-/S-naphthol 
aluminum  complex  of,  758 
chromium  complex  of,  757 
vanadium  complexes  of,  758 
Hydroxychlororuthenate   ion,   structure 

of  28,  167,201,  202 
2-Hydroxy-5,5'-dimethylazobenzene, 

copper  lake  of,  756 
Hydroxylamine,  96 
2-Hydroxy-5-methylazobenzene,    copper 

lake  of,  756 
2-Hydroxynaphthaldehyde- ;     resonance 

in  stability  of  chelates,  246 
7-Hydroxy-l,2-naphthoquinone-l-ox- 
ime,  cobalt  complexes  of,  717 


2  Ihdrow  .")  aitrobenseneaso  0  oapfa 
thol,  chromium  complex  of,  7.">; 
llydroxyoxinies,  678 
8  Bydroxj  quinaldine,  678 
8  1 1\  droxyquinoline 

in  coloriinet  lie  anal\  BIS,  690 

complexes  of,  72 
8  Hydroxyquinoline  derivat  h 
methyl   and   phenyl   buds titu ted,   238 
Bteric  hindrance  in  complexes  237  238 

ti77 
2-Hydroxy-5-sulfobenzeneazo-/3-naph- 

thol,  chromium  complex  of,  757 
2'-IIydroxy-5/-.sulfobenzeneazo-/3-na|>h 

thol,  copper  complex  of,  758 
2'-Hydroxy-4'-sulfobenzene-4-azo-l- 

l>licnyl-3-methyl-l-pyrazol-5-one, 

aluminum  complex  of,  758 
2'  Hydroxy-3'-sulfo-5'-methylbenzene- 

4-azo-l -phenyl -3 -met  hyl-1-pyrazol- 

5-one,  chromium  complex  of,  759 
II\  p-ochromic  effect,  565 
in  picrates,  553 

in  polynitro  molecular  compounds,  553 
in  quinhydrones,  551 
H4Y,     see     Ethylenediaminetetraa<  jetic 

acid 

Ibn,  see  Isobutanediamine 
Imidazole,  coordination  with  hemin,  721 
Iminodiacetic  acid,  complexes  of,  stabil- 
ity constants,  39 
Imino  group,  bridging  by,  62,  343 
Inclusions  in  electrodeposits,  effect    on 

brittleness,  643 
Indene.  platinum  complexes  of,  492 
Indicators, 

metal  ion,  221 

oxidimetric,  400 
Indicator  systems  involving  complexes, 

684 
Indigo,  metal  complexes  of,  762 
Indium 

complex  cyanides  <>\.  * 

electrodepositioD  of.  t;.")i 

halo  complexes  of,  6 

stereochemistry  of  tetracovalenl  com- 
plexes, 374 
Indium  halides,  -t  ructun 
Inductive  efTed  of  Quorine,  - 
Inert  complexes,  217 


808 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Inert  coordinated  groups,  214 
Inert  pair,  structure  of  compounds  con- 
taining, 370 
Infrared   absorption   in   study   of   com- 
plexes, 575-578 
Infrared  spectra 
acetylacetone  metal  complexes,  577 
cis-trans  isomerism  studied  by,  371— 

381,  577 
ethylenediaminetetraacetate  com- 
plexes, 577 
glycine  complexes,  578 
platinum-pentene   complexes,   504 
solvent  choice  for,  577 
use  in  structure  study,  300,  575-578 
Inner  complex,  definition  of,  672 
Inner  complexes 
insoluble,  672 

magnetic  moments,  table  of,  602 
value  in  analytical  chemistry,  673 
Inner  orbital  complexes,  207,  213,  217,  615 

electrodeposition  from,  639 
Inorganic  Maroon,  744 
Instability  constants  of  complexes,  130, 
402,  see  also  Dissociation  constants, 
Formation  constants  and   Stability 
constants 
Insulin,  reaction  with  metal  ions,  709 
Interatomic    distances    as   evidence   for 

multiple  bonds,  192 
Interaction  absorption,  in  bridged  halo- 
gen complexes,  19 
Interhalogens,  387 

Intermetallic  bonding,  522,  524,  534,  536 
Iodate, 
copper  (III)  complex  of,  31 
cobalt  (III)  complex  of,  29 
Iodide  ion,  donor  properties  of,  4-20 
Iodine (V)  fluoride,  387 
Iodine (VII)  fluoride,  structure  of,  394 
Iodocadmium  complexes,  405 
Iodomercurate(II)    complexes,    absorp- 
tion by,  566 
Ionic  bonds,  190,  195,  207 

compared  with  covalent,  136,  218 
Ionic  complexes,  137,  151,  208 
Ionic   mobility    of    complexes    in    basic 

chromium  salt  solutions,  454 
Ionic  model,   application  to  properties 
and  structures  of  complexes,  146 


Ionic  potential,  126,  423 

relation  to  complex  stability,  120 
Ionic  radii  and  stability  of  complexes, 

177 
Ionic  weight 
determination  of,  619 
of  aggregates  in  basic  chromium  (III) 
solutions,  451,  452,  459 
Ionization  isomerism,  267 
Ionization  of  hydrated  metal  ions,  22, 

425-431,  449-455 
Ionization  potential  of  metals,  and  sta- 
bility of  complexes,  177 
Ion  type,  see  also  Inner  orbital  complex, 
Outer  orbital  complex  and  Transi- 
tion metal  ion 
and  field  strength,  125 
importance  of  in  compound  stability, 
124 
Iridium 
electrodeposition  of,  660 
halo  complexes  of,  14 
olefin  compounds  of,  494 
Iridium (0)  pentammine,  151 
Iridium(II),  arsine  complexes  of,  80 
Iridium  (III), 
arsine  complexes  of,  80 
cyclopentadienyl  compound  of,  498 
thioether  complexes  of,  51 
thiourea  complexes  of,  53 
Iron  Blues,  744 

Iron  carbonyl,  509,  see  also  Iron  ennea- 
carbonyl    and     Iron    tetracarbonyl 
butylene  complex  of,  493 
configuration  of,  392,  520 
effect  of  light  on,  515 
reaction  with  phosphorus  (III)  halides, 
86 
Iron  carbonjd  hydride,  stereochemistry 

of,  375 
Iron  carbonyl  nitrosyl 
oxidation  state  and  configuration,  365 
stereochemistry  of,  375 
Iron     complexes, 
electronic  configurations  of,  638,  639 
oxidation-reduction  potentials  of,  188- 
190 
Iron  cyanide  complexes,  double  bonds 

in,  193 
Iron  cyanides,  use  of,  in  dyeing,  744 


INDEX 


SOU 


Iron,  1,3-diketone  complexes  of,  189 
Iron,  dye  complexes  of,  746,  747,  7  is.  7 19, 
766,  77.7.  768,  7(8 

Iron,  electrodeposition  of,  664 
Iron  enneacarbonyl.  see  also  Iron  car- 
bony]  and  Iron  tetracarbonvl 
iron-iron  interaction  in,  204 
ketonic  CO  in,  522 
structure  of,  621 
Iron  hydroxide,  in  dyeing,  71 1 
Iron  in  ferritin,  736 
Iron  nitrosyl,  161 
Iron-olefin  compounds,  493 
Iron  period  (first  long  period),  172 
Iron  porphyrin  complexes,  reactions  of, 

721 
Iron  a-pyridylhydrazine  complexes,  189 
Iron  a-pyridylpyrrole  complexes,  189 
Iron  tannage,  460,  471 
Iron  tetracarbonvl,  structure  of,  523 
Iron-tungsten    alloy,    electrodeposition 

of,  667 
Iron  (I) 
cyano  complexes  of,  628 
nitrosyl  halides  of,  535 
Iron  (II) 
ammines  of,  154 
chloride,  oxidation  of,  399 
complexes,     electronic     configuration 

of,  638 
a,a'-dipyridyl  complex 
covalent  bonding  in,  212 
diamagnetism  of,  212 
exchange  with  radioactive  iron  (II), 

212 
resolution   into   stable    optical    iso- 
omers,  212 
indenyl  complex  of,  499 
o-phenanthroline  complex  of 
covalent  bonding  in,  212 
diamagnetism  of,  212 
exchange    with    radioactive    iron- 
(II),  212 
phthalocyanine,    planar    configura- 
tion of,  365 
protoporphyrin,  718 
stabilization  of,  400 
stereochemistry     of     tetracovalent 
complexes,  375 
Iron(II)-(III)    aquo    couple,    standard 
potential  of,  400 


valence  relations  of  ISS,  399 

Iron(II)-(III)  dipyridy]  couple,  valence 

relations  of,   189 

Iron(II)-(III)    cyanide   couple,    valence 

relations  of,  188 
Iron(II)-(III)     fluoride    couple,    valence 

relations  of,  188 

Iron(II)-(III)     OXalato    COUple,     valence 

relations  of,  403 
Iron(II)-(III)   o-phenant  liroline  couple, 

valence  relations  of,  189 
Iron  (III) 
ammines  of,  154 

aquo  complexes,  neutralization  of,  157 
chloride  complex  with  phenol,  25 
chloro  complexes  of,  9 
citrate  complex  of,  structure,  570 
complexes,  ionic  and  covalent  charac- 
ter, 137 
cysteine  and  cystine  complexes,  731 
dimethyl   oxaloacetato  complex,  707 
dipyridyl    complex,    asymmetric   syn- 
thesis of,  583 
electron  configuration  of,  166 
ferriheme  complex;  failure  to  exchange 

with  radioactive  iron  (III),  213 
ferrihemoglobin    complex;    failure    to 

exchange  with  radioactive  iron  (III), 

213 
fluoro  complexes  of,  188 

analytical  importance  of,  9 
glutathione  complex  of,  731 
halides,  structure  of,  365 
in  catalase  and  peroxidase,  725 
malonate  complex, 

exchange  with    C-14   labeled   malo- 
nate ions,  212 

paramagnetism  of,  212 
oxalato  complexes 

dissociation  constants  of,  588 

exchange  with  labeled  oxalate  ions, 
212 

paramagnetism  of,  212 

reduction  of,  588 
oxide  hydrosols, 

dialysis  of,  463 

pure,  463 
oxo  complexes  of,  457 
salts,  basic,  citrate,  460 
stabilization  against  reduction,  400 


810 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Iron(III)— Cont. 

sulfate  solutions,  effect  on  pH  of  addi- 
tion of  neutral  salts,  458 
thiocyanate  complexes  of  76,  571,  688 
Irregular  tetrahedral  configuration,  orbi- 
tal hybridization  leading  to,  359 
Irreversibility  in  electrodeposition,  640 

and  brightness  of  deposits,  644 
Irreversibility  of  reduction  402,  406 
Irving-Williams  stability  series,  130 
Isobutanediamine,    96,    see   also    Isobu- 

tylenediamine 
Isobutylamine,  coordinating  ability  of, 

180 
Isobutylene 
absorption  by  CuCl,  495 
platinum  complexes  of,  492 
silver  complexes  of,  495 
Isobutylenediamine  complexes,  228 
Isomerism 
cis-trans 
infrared  study  of,  577 
in   hexacoordinate   complexes,   277- 

308 
in  olation,  449 
in  tetracoordinate  complexes,   357- 

364,  371-381 
polarographic  study  of,  587,  590 
Raman  study  of,  580 
effect  on  anion  penetration,  461 
types  of 

coordination,  263 
coordination  position,  270 
electronic,  272 

geometric,  277-308;  357-364;  371-381 
hydrate,  263,  574 
ionization,  267 
ligand,  271 

optical,  308-353;  357  et  seq. 
polymerization,  264 
position,  270 
ring  size,  272 
solvate,  261 
structural,  268 
summation,  272 
Isomer  pattern 
for  heptacoordinated   complexes,   392 
for  octacoordinated  complexes,  394 
for  octahedral  configuration,  276 
for  plane  hexagonal  configuration,  276 


for  polynuclear  octahedral  complexes, 

289-290 
for  pyramidal  tetracovalent  configura- 
tion, 358 
for  square  planar  configuration,  358 
for  three  coordinated  complexes,  385 
for   trigonal   prismatic    configuration, 

276 
for  tetrahedral  configuration,  358 
Isomorphism  of  poly-acids,  473,  477,  480 
Isonitriles 
complexes  of,  92,  533 
reaction  with  cobalt (II)  ion,  410 
reaction  with  metal  carbonyls,  93 
Isopoly-acids,   472,   477,   see   also   Poly- 
acids 
definition  of,  472 
Isoprene 
absorption  of,  by  CuCl,  495 
compound  with  copper(I)  chloride,  495 
Isothiocyanate  ion;  spectrum  of,  568 
Isotopic  exchange  studies  of  complexes, 
89,  211,  213,  326,  611-618 

Job,  method  of,  571 

Jorgensen,  theories  of  metal  ammines, 
103 

Keggin's  formulation  of  poly-acids,  481 

Keto  acids 

coordination  of,  707 
decarboxylation  by  metal  ions,  707 

iS-Keto  esters,  resonance  in  stability  of 
chelates,  246 

a-Ketoglutaric  acid,  decarboxylation  of, 
706 

Ketones,  coordination  compounds  of,  41 

Kinetics  of  electrodeposition  from  com- 
plex ions,  632 

Kopp's  rule  of  addition  volumes,  154 

Kurnakov's  test  for  cis-trans  isomers, 
53,  358 
failure  with  tertiary  phosphine  com- 
plexes, 79 

Labile  coordinated  groups,  214 
Lability  of  complexes,  215,  et  seq.,  see 

also  Mechanism  of  racemization  and 

Mechanism  of  reaction 
table,[215 


INDEX 


-II 


with  respect  to  charge  on  central  ion, 

217 
with  respect  to  substitution  reactions, 
214 
Laccase,  724 
Lactate  ion 
copper  complex  of,  ^t> 
rare  earth  complexes  of,  36 
Lanthanum    fluoride -potassium    fluoride 

mixture,  11 
Large  rings,  253  260 
Latimer's  convention,  399 
Lattice  energies  of  solid  complexes,  L38 
Lattice    types    in    occlusion    compounds, 
560 

Cage  lattice,  561 

channel  lattice,  560 

layer  lattice,  562 
Lead,  electrodeposition  of,  655 

Lead(II) 

acetato  complexes  of,  33 

acetylacetonate  of,  387 

iodo    complex,    dissociation    constant 

of,  588 
stereochemistry  of  tetracovalent  com- 
plexes, 374 
structure  of  complexes,  370 
Lead  (IV),  halo  complexes  of,  7 
Lead-silver  alloys,  electrodeposition  of, 

667 
Leveling  agents  in  electrodeposition,  642 
Ligand,  role  in  stability  of  complex,  175, 

177 
Ligand  effect,  249 
Ligand  isomerism,  271 
Ligands,  abbreviations  for,  96 
Limiting  poly-acids,  473,  474 
Lithium  halide  ammines,  stability  of,  140 
Lit  ton's  Salt,  97 

Logarithmic   method   of  studying  com- 
plex-    572 
Luteo  salts,  98 

Blomstrand's  structure  of,  104 
2,4-Lutidine,  coordinating  ability  of,  180 
•  .  copper  complex  of,  37 

M   gneaium 
double  fluoridf 
in  carboxylase,  706 

in  enolase,  711 


in  phosphorylation.  709 

.   planar  phthalocyanine  complex,  243 

Magnetic  criteria 

for  bond  type,  206,  759  % 

for  classification  of  complexes,  155 

for  planar  and   let  rahedral   Bt  fUCt  Hies, 

300,  359,  364 
for  structural  types  for  oontransition 

element-.  200 

for  study  of  complexes,  600 
Magnetic  moment 

affected   by  ligand  of  complex,   133 

anomalies  in,  603 

of  complexes  of  porphyrins,  718,  720, 
726,  728,  7:.t.  735 

of  complexes  of  transition  elements, 
172 

relation  to  configuration,  359,  364 

relation  to  susceptibility,  600 
Magnetic  susceptibility,  208 

relation  to  color,  605 
Magnetism  and  the  ionic  model,  132^/ 
Magnus'  Green  Salt,  97,  265 
Magnus'  Pink  Salt,  97 
Maleic  acid  complexes,  254 
Malonate  ion,  complexes  of,  35 

substituted,  complexes  of,  183 
Malonatotetrammine  cobalt  (III)  ion,  35 
Manganese 

cathodic  reduction  of  cyano  complexes, 
629 

dye  complexes  of,  755 

electrodeposition  of,  655 

in  arginase,  715 

nitrosyl  compounds  of,  539 

polarographic  determination  of,  696 
Manganese(I)  complexes,  preparation  of, 

409 
Manganese  (II) 

chloride  hydrates  of,  454 

cyclopentadienyl  compound  of,  198 

stereochemistry  of,  169 
Manganese  (III) 

cyano  complex  of,  88 

halo  complexes  of,  10 
Mangano(TV)-5-tungstic  acid,  177 
Mannitol  complex* 

mannito  boric  acid 

Maximum  multiplicity,  principle  of,  166, 

21  e 


812 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Mechanism  of  electro-deposition 
from  complex  ions,  632 
of  silver,  636,  641 
of  copper,  637 
Mechanism  of  racemization 
dichlorobis  (ethylenediamine)  cobalt- 

(III)  ion,  327 
dissociation,  325-329 
intramolecular  rearrangement,  329 
tris-(biguanidinium)    cobalt(III)    ion, 

330 
tris-dipyridyl  complexes,  327 
tris-orthophenanthroline      complexes, 

328,  330 
tris-oxalato  complexes,  326 
Mechanism  of  reaction 

cis-trans   interconversion    of    [Co   en2 

Cl2]+,  301-303 
displacement  (Sn2),  308 
dissociation  (SnI),  307 
"edge"  displacement,  307 
electron  transfer,  342,  353 
substitution    reactions,    146-149,    195, 

213-219,  305-308 
Walden  inversion,  307,  308,  347,  348 
Melano  chloride,  97 
Mercaptans,  coordinating  ability  of,  123, 

129 
Mercaptide  ion,   stability  of  complexes 

containing,  52 
o-Mercaptoazo  compounds,  metal  com- 
plexes of,  763 
Mercury 

complexes  with  primary  and  secondary 

amines,  63 
dye  complexes  of,  747 
electrode  polarization  in  cyanide  solu- 
tions, 638 
electrodeposition  of,  656 
olefin  complexes  of,  496,  579 
optically  active,  497 
Mercury  (II) 

amidochloride,  59 

chloro  complexes  of,  579 

complexes,  stabilities  of,  181 

cyanide,  structure  of,  89 

halo  complexes  of,  6 

iodo    complex,    dissociation    constant 

of,  588 
stereochemistry  of  tetracovalent  com- 
plexes, 372-4 


Mesitylene,  metal  complexes  of,  498 
Metachrome  Brown  B,  cobalt  complex 

of,  756 
Metal    acetylacetonates,    stabilities    of, 

41-45,  221,  231 
Metal-ammonia  bond   and  metal-water 
bond,  relative  strengths  of,  122-124 
Metal  ammonia  compounds,  150,  151 
Metal  carbonyl  halides,  513,  516 
mixed,  527 
stability  of,  517 
structure  of,  527 
table  of,  526 
Metal  carbonyl  hydrides 
formed  by  disproportionation,  516 
high  pressure  synthesis  of,  515 
hydrolysis  of,  515 
ionization  of,  516 
preparation  of,  515-517 
salts  of,  516 
structure  of,  530 
table  of,  510 
Metal  carbonyl  poisoning,  544,  545 
Metal  carbonyl  polymers,  510 
Metal  carbonyls,  151,  509 
bond  type  in,  517 
color  of,  529 
derivatives  of,  528 
formation  and  stability  of,  525 
heavy  metal  derivatives  of,  529 
industrial  significance  of,  540 
as  antiknock  agents,  540 
as  mordants,  745 
in  industrial  gases,  544 
interatomic  distances  in,  520 
mixed,  514,  517 
preparation  of,  511-515 
by  direct  union,  511 
by  disproportionation,  514 
by  Grignard  reagents,  511 
by  high  pressure  reactions,  512 
stability  in  relation  to  effective  atomic 

number,  160 
structure  of,  151,  517-537 
table  of,  510 

two  electron  bond  in,  521 
Metal  chelates,  formation  constants  of, 
178,  see  also  Dissociation  constants 
and   Ionization  constants 
Metal   crystallization   in   electrodeposi- 
tion, 633 


INDEX 


813 


Metal  ion  buffers,  221 
Metal  ion  indicators,  221 

Metal  ions 

absorption  by  wool.  76  ; 
biological  storage  <>t".  796 
shielding  of.  587 
Metal-ion   type   and   stability   of  com 

plexes,  177 
Metal-metal  bond.  522,  524.  534,  536 
Metal  oitrosyl  carbonyls 
bond  distances  in.  534 
reactions  of,  53S 
Metal  aitrosyls,  509,  531 
bond  distances  in,  533 
color  of.  52 

displacement  series,  535 
preparation  of,  534 
from  carbonyls,  537 
from  hydroxylamine,  539 
properties  of,  534 

stability  in  relation  to  effective  atomic 
number,  160 
Metal  nitrosyl  thio  compounds,  536 
Metal -olefin  compounds,  see  Olefin  com- 
plexes   and    the    individual    olefins 
Metal  oxide  hydrosols 
-    ay  salts.  464 
as  polymeric  ol  or  oxo  compounds,  464 
coordination    theory    of    flocculation, 

468 
effect  of  hydrolysis  on  charge  of  mi- 
celles, 467 
effect   of  oxolation   in   charge  of  mi- 
celles, 467 
flocculation  of,  468 
Metal   oxides,   coordination   with  metal 

ions,  28 
Metal  oxides,  hydrous 

•  rption  of  "foreign"  ions  by  dis- 
persed particles,  464 
adsorption  theory  of  colloidal  behav- 
ior, 464 
coordination  theory  of  colloidal  behav- 
ior, 463 
olation,  oxolation   and  anion   penet  ra- 
tion in  formation  of,  463 
properties  of,  17<i 

role  of  anion    penetration    and   deola- 
tion  in  dissolution  of, 
Metal-porphyrin    derivative- 
74U 


Metals,    purification    by    electrodeposi 

T  ion.  tiiiti 

Metaphosphoric  acid  in  copper  and  silver 

elect rodeposi! ion,  642 

Met  hemoglobin,  7:!  I 

(3  Met  ho\yeth\  lamine.  coordinating 

ability  of,  180 
Methylaminediacetic  acid,  i<> 
Methylbis-(3  -dimethylarsinopropyl)  - 

arsine,  complexes  with  iron,  cobalt. 

nickel,  and  copper.  ^1 

Methyl  elaidate,  Bilver  complex  of.  196 

Methyl  ferrocyanide,  isomer-  of.  Q2 

2-Mel  hyl-8-hydroxyquinoline,       coordi  - 
oating  ability  of,  678 

o-Methylmercaptobenzoic     acid,     com 
plexes  of,  51 

Methyloleate,  silver  complex  of,  496 

Methyl  tellurium  iodide,  isomerism  of, 
265 

Mineral  colors,  743 

Mineral  khaki,  743 

Mixed  carbonyls,  514,  517 

Molar  polarization,  507 

Molar  susceptibility,  600 

Molecular  asymmetry,  relation  to  polar- 
imetry,  580 

Molecular  beam,   dipole  moment   tech- 
nique, 598 

Molecular  compound,  definition  of,  547 

Molecular  compounds 
classes  of,  548 
properties  of,  548 
theories  of  structure,  554 
coordination  theory,  554 
ionization  theory,  556 
polarization  theory,  555 

Molecular  configuration  and  electronega- 
tivity, 173 

Molecular  orbitals,  198 

localized  and  non-localized,  199 
shapes  of,  200 

Molecular  orbital  theory,  197 
applied  to  complex  compounds,  201 

Molecular  symmetry,   determination   by 
infra!-  : .  576 

Molecular  vibrations,   determination  of, 
by  Raman  -[>• 

Molecular  volume  of  compl- 

criterion     for    classification    of    com- 
plexes, 104 


814 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Molecular  volume  of  complexes— Conf. 

relation    to    calcium    fluoride    lattice 
type,  155 
Molecular     weights, 

of  neutral  complexes  in  basic  chrom- 
ium salt  solutions,  454 

of  polyacids,  484 
Molybdate  ion,  structure  of,  8,  580 
Molybdenum 

Carbonyl-cyclopentadienyl  compound, 
499 

Cathodic    reduction    of    cyano    com- 
plexes, 629 

Cyanide  complexes  of,  395,  629 
double  bonds  in,  194 

cyclopentadienyl  compound  of,  498 

dye  complexes  of,  755 

electrodeposition  of,  665 

halo  complexes  of,  15 

oxyhalo  complexes  of,  16 

plating  of,  645 
Molybdenum  (II) 

stereochemistry  of,  375 

structure  of  [Mo6Cl8]4+,  366 
Molybdic  acid,  477 
Molybdic  acid-tartaric  acid  complexes, 

ionization  of,  595 
Monastral  Fast  Blue,  761 
Mond  process,  540 

Monohydroxytrifluoroberyllate,     resem- 
blance to  sulfate,  5 
Mononuclear    metal    carbonyls,    struc- 
ture of,  518 
Mordants 

inorganic  complexes  as,  743 

metal  carbonyls  as,  745 

phosphomolybic  acid  as,  745 

phosphotungstic  acid  as,  745 

potassium  ferrocyanide  as,  745 

tannin-tartar  emetic  as,  745 
Mordant  Yellow  O,  754 
Morin  in  fluorescence  analysis,  694 
Morland's  Salt,  97 

Morpholinc,  coordinating  ability  of,  180 
Multiple  bonds  between  metal  and  lig- 

and,  177,  189  et  seq.,  503,  521,  525 
Multiple  ring  systems,  234  et  seq. 
Murexide,  indicator  for  metal  ions,  684 
Mutarotation,  302,  303,  348 


Names    of    complexes,    see    Nomencla- 
ture 
Naphthalene,  metal  complexes  of,  498 
Naphthazarin,  beryllium  complex  of,  26 
Naphthol  Green  B,  746 
1 , 2-Naphthoquinone-l-oxime,         cobalt 

complex  of,  746 
Natural  products 
detection  of  coordination  compounds 

in,  698 
functional   relationships   between   co- 
ordination compounds  in,  700 
Negative     coordinating    groups,     suffix 

used  in  naming,  93 
Neocupferron,  680 
Neolan  Blue  B,  756 
Neolan  Red  B,  755 

Neopentanediamine,   coordinating  abil- 
ity of,  183,  229 
Nernst  equation,  399 

applied  to  alloy  deposition,  666 
Neutral   complexes   in  basic   chromium 

salt  solutions,  454 
Neutralization  of  aquo  compounds,  451 
Neutron  irradiation   of  complexes,   613 
Nickel,  diarsine  complexes  of,  187 
Nickel,  dye  complexes  of,  747,  755,  757, 

758,  759,  761,  762,  763 
Nickel,  electrodeposition  of,  656 

effect  of  coordinating  agents  on,  641 
from  thiosulfate  complexes,  630 
Nickel,  metallurgy  of, 
Mond  process,  540 
Orford  process,  48 
Nickel,  reaction  with  carbon  monoxide, 

509 
Nickel-tin   alloys,   electrodeposition  of, 

669 
Nickel-tungsten    allo3^s,    electrodeposi- 
tion of,  668 
Nickel (0)  complexes,  409 

configuration  of,  365 
Nickel  (0) 
cyanide  complex,  91 
nitrogen  sufide  complex,  151 
proposed  structure  for  K4Ni(CN).i,  370 
stereochemistry  of  tetracovalent  com- 
pounds, 375,  377 
Nickel  (I),  409,  538 
cyanide  complexes  of,  91,  514,  628 
nitrosyl  halides  of,  535 


INDEX 


815 


Nickel  (II) 
ammines,  resolution  Btudies,  212 
complexes 
eolor  related  to  structure,  364 
configuration  of,   L69  173,  365 
possible  electronic  distributions  for, 

360 
racemization  of,  328 
stereochemistry     of     tetracovalenl 

compounds,  375,  377 

coordination  number  of,  571 

cyanide  complex,  tracer  study  of,  616 

cyanide  complexes  with  amines,  137 

cyclopentadienyl  compound  of,  498 

dimethylglyoxime  complex- 
fused  rings  in,  232 
in  analytical  chemistry  of,  (171-677 
possible  structures  of,  233 

halo  complexes  of,  10 

magnetic  criterion  for  dsp2  bonding, 
360 

number    of    unpaired    electrons     and 
structural  types,  209 

porphyrin  complexes,  magnetic  prop- 
erties of,  718 

pyromethene  complexes,  magnetic 
properties,  604 

tris-(dipyridyl)    ion,   racemization   of, 
327,  328 

tris(o-phenanthroline)    ion,    racemiza- 
tion of,  328,  575 

triazine  complex  of,  254 
Nickel  (III),  187 
Nickel  (III)  complexes,  392 

with  cyclopentadiene,  498 

with  pho.-phines,  80,  187 

with  o-phenylenebis-(dimethylarsine), 
80,  187 
Nickel(IV),  187 
Nickel  (IV j  complexes,  409 

with  fluoride,  188 

with    o-phenylenebis(dimethylarsine), 
80,  187 

with  sulfur  compounds,  56 

with  triethylphosphine,  187 
Nickel  carbonyl,  171,  sec  also  Metal  car 
bonyls,  Chapter  16 

catalysis  by,  542 

configuration  of,  365 

discovery  of,  509 

electron  configuration  of,  191 


indust  rial  significance  of,  old 

preparation  of,  51  I  515 

reaction  with  antimony!  1 1 1     chloride, 
86 

react  ion  with  ph08phorU8<  Ml'  h.dides, 

86 

structure  of,  519,  608 
Nickel      carbonyl       hydride,       I'd 

Nine  membered  ring,  260 
Niobium.  84 1  also  ( lolumbium 

cyclopentadienyl  compound  of,  498 

electrodeposition  of,  665 

halo  complexes  of,  16 

structure  of  Nb6ClM-7H20,  366,  375 
Nitrate   ion,   complexing   ability  of,  28, 

280 
Nitrato  cerate  ions,  29 
Nitratopentamminecobalt(III)  ion,  28 

in  determination  of  phosphate,  682 
Nitrato  silver  (II)  complex,  408 
Nitric  oxide 

coordination  compounds  of,  531 

reaction  with  peroxidase,  726 

reaction    with    porphyrin    complexes, 
726,  735 

trans  influence  of,  537 
Nitrile  complexes  of  platinum,  75 
Nitriles,  donor  properties  of,  74 
Nitrilotriacetic  acid  (triglyeine),  stabil- 
ity of  complexes,  39,  777 
Nitrito-nitro  isomerism,  268,  280 
Nitroammine  cobalt  (III)  complexes 

absorption  spectra  of,  296,  565 

polymerization  isomers  of,  264 

relation  to  third  band,  296,  568 
Nitro  complexes  of  rhodium,  658 
Nitrogen  atom,  as  center  of  optical  activ- 
ity, 323 
Nitrogen,  donor  properties  of,  3,  59-78 
Nitrogen  in  copper-lead  alloys  deposited 
from  ethylenediamine  complexes,  630 
Nitrogen(II)oxide   pentammine   cobalt - 

(II)  ion,  273 
Nitro  group 

absorption  by,  567 

bridging  by,  I'd 
Nit  ro  oil  rito  isomerism,  268 
Nitroprussides 

Nitrosohydroxylamines  in  analysis,  680 

1  Nltro80-2-hydroxy-3-naphthoic       acid 

lamides,  iron  complexes  of,  748 


816 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Nitrosophenol  metal  complexes,  746,  747 

in  analysis,  680 
Nitroso-R  salt,  747 
Nitroso    Roussin    salts,    potentiometric 

titrations  of,  594 
Nitrosyl,  see  Metal  nitrosyl 
Nitrosyl  anion,  532 
Nitrosyl  cation,  533 
Nitrosylpentamminecobalt(III)  ion,  272, 

532 
Nomenclature  of  complexes,  93 
based  on  color,  98 
based  on  names  of  discoverers,  97 
I.U.C.  system,  93 
modifications  of,  94 
examples,  95 
Non-aqueous  solutions 
electrodeposition  from,  669 
solvent  systems,  418 
Nonbonding  orbitals,  199,  201 
Normal  (ionic)  complexes,  5,  18,  151,  207 
Nucleation  in  electrodeposition,  641 
Nullapon,  223,  777 

Number  of  coordinating  groups,  as  in- 
dicated by  prefix  in  name,  93 
Number  of  particles  and  entropy  of 

chelation,  252 
Nylon  type  fibers,  dyeing  of,  765 

Occlusion  compounds,  558 

choleic  acids,  559 

clathrates,  561 

4,4/-dinitrobiphenyl  adducts,  560 

thiourea  adducts,  560 

urea  adducts,  560 
Octacoordination,  configurations  of,  394 
Octacyanometallates,  395 
Octafluorometallate  ions,  395 
Octafluororuthenate(VI)  ion,  14 
Octafluorotantalate(V)  ion,  16 

structure  of,  396 
Octammines,  396 
Octammine  -/*  -  amino  -  ol  -  dicobalt  (III)  - 

chloride,  30 
Octammine-/x-diol-dicobalt(III)       salts, 

291,  448,  449 
Olation 

aging  in,  456 

"continued"  process  of,  451,  468,  470 

definition  of,  448 

degree  of,  455,  456,  466 


factors  promoting,  455 

formation  of  polymers  by,  451 

inflocculation  of  metal  oxide  sols,  468 

in  formation  of  hydrous  metal  oxides, 
468 

reversibility  of,  457 
Olefin  complexes,  487-508 

studied  by  Raman  spectra,  579 

structures  of,  501 
Olefins,  see  also  Unsaturated,  Ethylenes, 
and  the  specific  olefins 

coordination  in  terms  of  molecular  or- 
bitals, 207 

extraction  from  mixtures  with  satu- 
rated hydrocarbons,  500 

in  complex  formation,  159,  168,  487- 
508 

stabilization  of  lower  valences  by,  412 

trans  effect  of,  204,  490,  503 
Oleic  acid,  methyl  ester,  compound  with 

silver,  496 
Ol  group,  22,  448-471 

displacement  by  anions,  456,  463,  465 

distinction  from  hydroxo  group,  448 

formation  from  oxo  group,  464,  465 

in  micelles  of  metal  oxide  hydrosols, 
464 

in  precipitated  hydrous  metal  oxides, 
468,  470 
Optical  isomerism,  308-353 

as  criterion  for  bond  type,  209 

as  evidence  for  a  planar  configuration, 
361 

asymmetric  induction  of,  352,  581 

due  to  asymmetric  nitrogen  atom,  323 

due  to  asymmetric  sulfur  atom,  325 
Optical  isomers,  308-353 

absolute  configuration  of,  336,  337 

anionic  complexes,  312 

asymmetric  synthesis  of,  350,  351 

cationic  complexes,  309-311 

inorganic,  277,  322,  323 

mutarotation  of,  302,  304,  348,  349 

nomenclature  of,  94 

nonionic  complexes,  313 

optically  active  donor  ligands  in,  313- 
318 

oxidation-reduction  of,  353 

polydentate  ligands  in,  318-321 

polynuclear  complexes,  321-323 


INDEX 


817 


purely  inorganic  complexes,  277,  322, 

323 
racemisation  of.  325  330 
reactions   of   polynuclear   complexes, 

342,  343 
relative  configurations  of,  337-342 
resolution  of  raeeinie  complexes,  331- 

336 
substitution    reactions    with    retention 

of  configuration,  302,  308,  342,  344 
tables   of  optically    active   complexes, 

310,  312 
ptieally  active  donor  molecule-  in  com 

plexes,  313-3 In 
ptieally  active  unsymmetrical  ligands, 

317,  318 
ptical  methods 
as  criteria  for  bond  type,  213 
in  study  of  complexes,  5S0 
rbital  configurations  (table),  170 
rford  process,  4S 
rganic  anions 
donor  properties  of,  33 
penetration  by,  460 
rganic   dyestuffs,  metal  complexes  of, 

745  et  seq. 
Hon,  dyeing  of,  765,  766 
rnithine,  copper  complex  of,  37 
rthophosphates,  complexing  by,  770 
smium  carbonyl  halides,  stability  of, 

528 
smium,  electrodeposition  of,  660 
smium  enneacarbonyl,  523 
smium  nitrosyl  compounds,  539 
smium(II)-(III)  complexes 
electron  interchange  in,  20 
oxidation-reduction  reactions  of,  353 
smium  (III)  and   (IV),  halo  complexes 

of,  15 
smium  (IV),    hydroxo-halo    complexes 

of,  15 
smium (VI)  oxyhalo  complexes  of,  15 
smium  (VIII) 
fluoride,  396 
fluoro  complexes  of,  15 
oxide,  reaction  with  carbon  monoxide, 

513 
stereochemistry  of  tetracovalenf  com- 
plexes, 375,  379 
uter  orbital  complexes,  207-217 
verlap  integral,  165 


OX,  Bee  Oxalate  ion.  96 

<  bcalate  ion 

anion  penel ration  by,  150  162 

as  bridging  group,  35 

hydration  of,  35 

use    to    determine    configuration    of 
planar  complexes,  35,  359 
( bcalates 

in  chromium  electrodeposition,  650 

in  copper  elect  rodeposif  ion.  652 
in  lead  electrodeposition,  655 
Oxalate  complexes,  34 
basic  chromium,  450 

copper,  electrode  polarization  of,  637 

oxalatodileadl  II I  ion,  35 

oxalatobis  -  (ethylenediamine)cobalt 
(III)  ion,  30 

racemisation  of,  325 

resolution  of,  331,  333 
Oxalosuccinic  acid,  decarboxylation  of, 

706 
Oxazine  dyes,  752 
Oxidases,  722 
Oxidation-reduction,  biological,  role  of 

metal  ions  in.  722 
Oxidation-reduction  indicators,  400,  686 
Oxidation-reduction  potentials,  398 

of    hemochrome-hemiehrome    system, 
721 

use  in  formation  constant  determina- 
tion, 593 
Oxidation,  stability  toward,  398 
Oxidation  state  of  metallic  element 

degree  of  stabilization  of,  584 

indicated  by  name  of  compound,  93 

relation  to  bond  orientation,  171,  172 

relation  to  structure,  364,  365 

relation  to  thermal  stability,  620 
Oxides,  hydrated,  precipitation  of,  453 
Oximes,  coordinating  tendency  of,  76 
o-Oximinoketones,   metal  complexes  of, 

748 
Oxine,  complexes  of.  72 
Oxo  group,  448 

conversion  of  ol  group  to,  464,  465 

in  chromiumflll  I  conn 

in   micelles  of  metal   oxide   hyd: 
464 

in  precipitated  hydrous  metal  o: 

in  beryllium  complex,  28 


818 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Oxo  group — Cont. 

in  ruthenium  complex,  28,  167,  201 

reaction  with  hydrogen  ion,  465 

reactivity  of,  465 
Oxolation,  448,  456 

effect  of  on  charge  on  micelles  of  metal 
oxide  hydrosols,  467 

in  aluminum  complexes,  457 

in  basic  chromium  salt  solutions,  456 

in  flocculation  of  metal  oxide  sols,  468 

in  formation  of  hydrous  metal  oxides, 
463 

reversibility  of,  457 
Oxonium  theory  of  acids  and  bases,  416 
Oxo  process,  543 
Oxyacids,  classification  of,  438 
Oxyacid  theory  of  amphoterism,  436 
Oxy  anions 

donor  properties  of,  28 

electrodeposition  of  metals  from,  648 
Oxygen 

absorption  by  ammoniacal  cobalt  salt 
solutions,  45 

donor  properties  of,  3,  20 

reaction    with    porphyrin    complexes, 
719,  721,  728,  732 
Oxygen  carrying  chelates,  45,  735 
Oxygen  molecule,  paramagnetism  of,  203 
Oxyhemoglobin,  732 

Palatine    chrome   black   6B,    chromium 
complex  of,  756 

Palatine  Fast  dyes,  755 

Palladium 
dye  complexes  of,  747 
electrodeposition  of,  659 

Palladium  (0) 
cyanide  complex  of,  92 
proposed    structure    for   K4Pd(CN)4  , 
370 

Palladium  (II) 

arsine  complexes,  structure  of,  609 
configuration  of  complexes,  169 
cyanide  complex,  structure  of,  89 
cyclopentadienyl  compound  of,  498 
dimethylglyoxime  complex  of,  675 
halo  complexes  of,  13 
olefin  complexes  of,  493 
structure  of  ethylene  complex,  504 
structure  of  propylene  complex,  504 


resolution  of  a  tetracovalent  complex, 

361 
stereochemistry  of  tetracovalent  com- 
plexes, 375,  378 
structure  of  arsine  complexes,  365 

Palladium  (III)    in   M2PdCl5  ,   probable 
nonexistence  of,  13 

Palladium (IV),  halo  complexes  of,  13 

Palladium  period   (second  long  period), 
172 

Palladium  transition  elements;  electro- 
positive character  of,  195 

Para-isopoly-acids,  483 

Paramagnetic  resonance  of  IrCl6=,  204 

Partial    asymmetric   synthesis    of   com- 
plexes, 316 

Partial  ionic  structures,  165 

Paschen-Back  effect,  133 

Passivation  in  crystallization  in  silver 
electrodeposition,  641 

Pauling's  formulation  of  poly-acids,  479- 
481 

pc,  see  Phthalocyanine 

PDA,    see    o-Phenylenebis(dimethylar- 
sine) 

Penetration  (or  covalent)  complexes,  5, 
18,  151,  207 

Pentacarbonyls,  configuration  of,  392 

Pentachloronitrosylosmate  (III) ,    stabil- 
ity of,  414 

Pentachloronitrosylruthenate  (III) ,    sta- 
bility of,  414 

Pentacovalence,  determination  of,  616 

Pentacyanocobaltate(II)  ion,  391 

Pentadiene,      chelation     through      two 
double  bonds,  249 

1,3-Pentadiene,  copper  (I)  chloride  com- 
pound of,  495 

Pentafluoroantimonate(III)    ion,   struc- 
ture of,  8 

2,4-Pentanedione,  see  Acetylacetone 

2-Pentene,  platinum  complexes  of,  492, 
504 

2-Pentene(cis    and    trans),    silver    com- 
plexes of,  496 

Peptide  bonds,  metals  in  cleavage  of,  702 

Peptide  complexes,  structure  of,  704 

Peptide  group,  coordination  in  enzyme- 
substrate  complex,  704 


INDEX 


819 


Peptization 
of  hydrous  aluminum  oxide  effective 

nc--  of  various  acid-  in.   169 
oi  hydrous  metal  oxide-.   168 
of  hydrous  thorium  oxide  by  hydroxy 

arid  anions,   166 
of   hydrous    lirconium    oxide,    by    hy- 
droxy acid  anion-.    167 
theory  of  aniphoterism.   135 
rerchlorate  ion,  coordinating  ability  of, 

28 
Periodate  ion 
donor  properties  of,  407 

structure  of,  580 
Periodic  table,  front  end  paper 
and  electrodeposition,  646,  647 

and  electrode  reversibility,  646 
Perlon  Fast  color-.  765 
Peroxidase,  724 

k-Peroxo  complexes,  26,  47,  410 
ji-Peroxo  group  bridging,  451 

in  cobalt  complexes,  26,  47 
Perrhenate  ion.  structure  of,  580 
Peyrone's  Chloride.  97,  265 
Pfeiffer  effect,  581 

Pfeiffer's  formulation  of  poly-acids,  479 
ph.  see  1, 10-Phenanthroline 
pH,  measurements  of,  590,  592 
pH.  of  aluminum  oxide  sols,  465 
pH,  of  basic  chromium  salt  solutions,  ef- 
fect of  anion  penetration  on,  459 
I    Phase  changes,  in  continuous  variation- 
method,  621 
o-phen,  see  1, 10-Phenanthroline 
phenan,  see  1, 10-Phenanthroline 
1,10  Phenanthroline,  complexes  of,  69 

in  colorimetric  analysis,  688 

oxidation  and  reduction  of,  353 

racemisation  of,  328 

-t r  ric  factors  in  stability  of  iron  com- 
plexes, 244 
Phenol  oxidases,  722 
Phenol-,  test  for.  with  iron  (III),  25 
Phenylalanine  anion.  96 
Phenylbiguanide,  96 

complexes  of.  711 

o-Phenylenebis(dimethylarsine  . 

iron  complexes  of,  80 
nickel  complexes  of.  1^7 
palladium  complex 


o-Phenylenediamine,    mohodentate    be 

havior  of.  67 
Phosphate  complexes 

dissociation  constants  of,  77.") 
of  iron(IH),  ::i 
of  rhodium,  658 
Phosphates 

BSSOCiat  ion  with  cat  ion-.  77.") 

branched,  7t><» 
condensed,  T«  »* » 
glassy,  769 
ring,  769 
stability  of.  772 
ult  ra,  7ti!> 

used  in  water  Boftening,  7C>(.)-777 
Phosphine  (PH3) 
complexes  of,  127 
coordinating    ability,    compared    with 

ammonia,  127 
physical  properties  of,  127 
reaction  with  metal  sails,  78 
Phosphines,  complexes  of 
containing  two  different  metals,  83 
double  bonding  in,  81 
of  copper  (I),  79,  407 
of  gold  (I),  79 

of  group  VIII  metals,  80-82 
of    iron(0),    cobalt (0),    and    nickel  (0) 

carbonyls,  84 
of  mercury,  79 
of  nickel (II)  and  (III),  187 
of  platinum(II),  containing  ethyl  buI 

fide,      oxalate,      and      thiocyanate 

bridges,  83 
use  as  antiknocks,  79 
Phosphines,  organic 

coordinating  ability   of,  78,   123,   127- 

128,  129,  187,  392 
trans  influence  of,  79 
Phosphomolybdic  acid.  17:;.  i7»'>.  177 

as  a  mordant .  715 
Phosphoric    acid,    meta,    in    copper    and 
-ilver  elect  rodeposi  i  ion  ,  642 

Phosphorus  boron    bonds,    Btability   of, 

205 
Phosphorus  1 1 1 1    chloride,  complex* 

with  coppei  I 
with  gold  I  .  86 

with  iridium'  III  I  chloride,  -lability  of, 
85 


820 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Phosphorus  (III)  chloride— Cont. 

with  iron  (III)  chloride,  86 

with  nickel (0),  86,  151 

with  palladium (II)  chloride,  85 

with  platinum (II)  chloride,  84 
Phosphorus  (III)  fluoride 

complexes  of,  85,  148,  151,  205 

trans  influence  of,  148,  197,  205 
Phosphorus  (III)  halides 

donor  properties  of,  84 

reaction  with  nickel  tetracarbonyl,  86 
Phosphorus (V)  halides,  configuration  of, 

388 
Phosphorus (V)    oxychloride,    electrode- 
position  from  solutions  in,  670 
Phospho-12-tungstic  acid,  479-481 

as  a  mordant,  745 
Phosphorylation,  708 
Phthalic  acid,  chromium  complexes  of, 

461 
Phthalocyanine  metal  complexes,  73,  223 

as  dyes,  760 

magnesium  complex,  luminescence  of, 
741 

magnetic  moments  of,  243 

structure  of,  361,  370,  760 
Photosynthesis,  740 

Physical  methods,  application  to  coordi- 
nation compounds,  563-624 
Physical   properties    of   complexes,    130 
Phytates 

as  sequestering  agents,  769,  782 

calcium  complex,  782 
pi  bond,  201 
pi  orbitals,  199 

Picolines,  coordinating  ability  of,  180 
Picolinic  acid,  complexes  of,  73 

with  iron  (II),  38 
Picrates,  551 

bathochromic  effect  in,  554 

hypsochromic  effect  in,  553 

molecular,  553 

salt-like,  553 
Pigments,   coordination   compounds   as, 

743-763 
Pinene 

palladium (II)  complexes  of,  493 

platinum (II)  complexes  of,  492 
Piperidine,  coordinating  ability  of,  180 
Planar  configuration,  see  Square  planar 
configuration 


Platinum 
cathodic  reduction  of  cyano  complexes, 

629 
electrodeposition  of,  659 
Platinum  black,  electrodeposition  of,  658 
Platinum  complexes,  absorption  of,  567 
Platinum  period  (third  long  period),  172 
Platinum(O)  tetrammine,  151 
Platinum(II)  aquoammine  complex,  cis- 

trans  isomerism  of,  594 
Platinum(II)  chloride,  structure  of,  17 
Platinum  (II)  complexes 
ionic  model  of,  131 
dipole  moments  of,  363 
planar  configuration  of,  169 
resolution  of,  361,  369 
stereochemistry  of  tetracovalent,  375, 

379 
structure  of  K2PtCl4  ,  368 
trans  effect  in,  196 
trans  elimination  in,  360 
with  arsines,  phosphines,  and  stibines, 

598 
with  carbon  monoxide,  stability  of,  194 
with  halides,  12 
stability  series,  594 
radio  exchange  of  bromo  complexes, 
614 
with  olefins,   see   platinum (II)    olefin 

complexes 
with  primary  amines,  63 
with       /3,iS/,/3"-triaminotrieth3Tlamine; 
stereochemistry  of,  239 
Platinum  (II) -dinitro  (N-methyl-N-ethyl- 

glycine)platinate(II),  325,  334 
Platinum (II)  olefin  complexes,  487-492 
cationic,  490 

containing  amines,  stability  of,  489 
containing  anions,  stability  of,  489 
containing  two  moles  of  olefin,  491 
geometric  isomers  of,  490 
preparation  of,  489 
properties  of,  492 
stability  of,  489,  492 
structure  of,  501-506 
ethylene  complex,  502 
Platinum(III),    questionable    existence 

of,  13,  411 
Platinum  (IV) 
2-chlro-l ,  6-diammine-3 , 4 , 5-diethyl- 
enetriamineplatinum(IV),  279 


! 


INDEX 


S'Jl 


chloro-hydroxo  complexes  of,  439 
ethylenediamine  complexes,  potentio- 
metric  titrations  of,  595 

fluoro  complex  of,  lss 
halo  complexes  of,  11 

mixed  halo-hydroxo  complexes  of,  12 
tetramethyl,  165 
hexacovalenl  carbon  in,  132 

Platinum(V),   (VI)  and   .VIII),  possible 

existence  of,  41 1 
Platinum  group  metals 

cyano  complexes,  stability  of,  589 

electrodeposition  of,  657 

electropositive  character  of,  195 
Platosamminechloride,  116 
Platosemidiamminechlorid,  116 
Plutonium  (IV),  peroxo  complexes  of,  28 
pn,  see  Propylenediamine 
Pol  a  rime  try 

in  stud}'  of  complexes,  580-583 

structure  determination  by,  582 
Polarity,  induced,  597 
Polari  z  ability 

of     central     metal     ion,      126 

of  coordinated  molecules,  126 
Polarization 

at  electrodes,  632 

induced,  597 

molar,  597 

molecular,  related  to  absorption,  567 

of  ions,  nature  of,  121,  125 

orientation,  597 

permanent,  597 

relation  to  chemical  properties,  122 

relation  to  force  binding  electrons  to 
nucleus,  122 

relation  to  third  band,  568 
Polarographic  reduction 

of  antimony  complexes,  404 

of  cadmium  complexes,  405 

of  copper  complexes,  403 

of  cobalt  complexes,  629 

of  iron-oxalato  complexes.  403 

of  tin  complexes,  404 

of  uranium  complexes,  404 

of  vanadium  complexes,  405 

of  vitamin  B:_  .  7 
Polarography  in  Btudy  of  complexes,  102 
106,  5S4 

cis-trans  isomerism,  5s7 

polymetaphosphates,  588 


Polonium,  electrodeposition  of,  660 
Poly  acids,  172  L86 

aggregation  Btudiefl  of,  484   186 

analysis  of,  184  486 

Anderson's  formulation  of,  483  184 

anions  of,  with  chelate  containing  rat- 
ions, 486 
basicity  of,  478-480 
Blomatrand's  formulation  of,  473 
composition  of,  172  et  seq. 
central  atoms  in,  171 
coordination  number  of  metal  ions  in. 

475,  477,  479-484 
Copaux's  formulation  of,  473 
definition  of,  472 
hydration  of,  478,  480,  481 
isomorphism  of,  473,  477,  480 
Keggin's  formulation  of,  481-482 
limiting  series  of,  473-474 
molecular  weights  of,  484-485 
parent  acids  of,  474-475 
Pauling's  formulation  of,  479-481 
Pfeiffer's  formulation  of,  479 
physicochemical  studies  of,  484-486 
preparation  of,  485 
properties  of,  486 
Rosenheim-Miolati    classification    of, 

474-478 
salts  of,  474,  et  seq. 
structure  of,  472,  et  seq. 
unsaturated,  475-477 
Werner's  formulation  of,  473 
x-ray  studies  of,  479-483 
Polyamino  acids,  as  sequestering  agents, 

769,   777,   782,    see    also   individual 

polyamino  acids 
Polydentate  ligands,  234-236,  286,  318 
Polyhydric  alcohols,  chelation  by,  24 
Polymeric   complexes  of  beryllium,  42, 

466  et  seq. 
Polymeric  ions,  in  aluminum  oxychloride 

sols,  164 
Polymers,  see  also  Super  complexes 
cross-linked,    resulting   from   olation, 

i:»:; 
from  <lit bioxamide  complex* 
from  tetraketones,  12 
in  hydrous  metal  oxides,  170 

Polymerization  due  to  olation,   151 

Polymerization  isomerism,  264 


822 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Polymetaphosphates 
polarography  of,  588 
use  in  water  softening,  769  et  seq. 
Polymethylene    bis-a-amino    acid    com- 
plexes, 258 
Polymethylenediamines  donor  properties 

of,  253 
Polymolybdates,  477 
Polynuclear  complexes,  22,  321-323 
bridging  groups  in  formation  of,  462 
formation  of,  in  place  of  large  rings, 

230,  232,  244 
geometric  isomers  of,  289-291 
optical  activity  of,  321-323 
reactions  of,  342-343 
Polynuclear  metal  carbonyls,  521 
formation  rules  for,  525 
structure  of,  521 
Polyphosphates,  769 
free  alkalinity  of,  772 
natural  pH  of  solutions,  772 
properties  of  complexes,  773-775 
Polysulfide,  tin  electrodeposition  from, 

663 
Polytungstates,  478 
Poly  vanadates,  478 

Pontachrome  blue  black  R,  in  determina- 
tion of  aluminum,  695 
Porphin,  structure  and  derivatives,  73, 

717 
Porphyrin  complexes,  stability  of,  717 
Position  isomerism,  270,  317 
Positive-negative   theory   of   acids    and 

bases,  419 
Potassium  carbon}d,  509 
Potassium  chlorocuprate(II)  dihydrate, 

structure  of,  368 
Potassium  chloroplatinate(II),  structure 

of,  368 
Potassium  chlorostannate(II)  dihydrate, 

structure  of,  367 
Potassium    hexacyanoferrate(II),    as    a 

mordant,  745 
Potassium  hexacyanoferrate(III),  prop- 
erties and  crystal  field  theory,  134 
Potassium  hydroxychlororuthenate, 

structure  of,  167,  201,  202 
Potassium  tetrachloropalladate  (II) , 

structure  of,  356 


Potassium  tetrachloroplatinate  (II) 

crystal  field  splitting,  134 

structure  of,  356 
Potassium  tetracyanocuprate(I),  stabili- 
zation of  valence  in,  407 
Potassium  tetracyanonickelate  (0) ,  struc- 
ture of,  370 
Potassium  tetracyanopalladate  (0) , 

structure  of,  370 
Potential  energy  of  complex  ions,  634,  636 
Potential  energy  of  ions  at  electrodes, 

634 
Potentials,  signs  of,  399 
Potentiometry 

in  study  of  complexes,  590-596 

in  study  of  cis-trans  isomerism,  594 

used  to  determine  formula  of  complex, 
591 
Praseo  salts,  98 
Pressure,  effect  on  silver  chloride  am- 

mines,  622 
Primary  valence,  109,  157 
Principal  valence,  109 
Propylene 

absorption  by  copper  (I)  chloride,  494 

platinum  complexes  of,  488 
Propylenediamine,  complexing  b}^  63 
Propylenediamine  complexes,  228 

stereochemistry  of,  316-318 
Protein,  reaction  of  metal  ion  with,  703 
Proton  theory  of  acids,  421 
Protoporphyrin,  717 
Prussian  Blue 

as  a  pigment,  744 

as  a  super-complex,  75 

destroyed  by  2,2'-dipyridyl,  68 

structure  of,  90,  610 
ptn,  see  2,4-Diaminopentane,  96 
Purpureo  salts,  98 

Blomstrand's  structure  of,  104 
Purpurin,  751 
py,  see  Pyridine,  96 

Pyramidal  configuration  for  tetracova- 
lent  complexes,  354 

orbital  hybridization  leading  to,  359 
Pyridine 

complexes  of,  67,  128,  180 

electrodeposition    from    solutions    in, 
670 

mixed  complexes  with  ether,  25 
Pyridoxal,  712 


INDEX 


823 


Pyridoxal     complexes,     reaction     with 

alanine,  71 1 
Pyridoxamine  complexes,  reaction  with 

pyruvic  acid.  711 

a-Pyridylhydrasine,  complexes  of,  68 
ar-Pyridylpyrrole,  complexes  of,  87 
E>yrocatechindisulpho  acid,   Fe+++  che- 
lates oi,  231 
Pyrogens  Green,  762 
Pyrophosphate 

hydrolytic  degradation  of,  772 

in  copper  electrodeposition,  052 

sodium  hydrogen,  769 

tetrasodinm,  769 
Pyrophosphate  complexes 

in  water  Boftening,  771  et  seq. 

of  copper,  electrode  polarization  of,  637 

of   copper  II   .    nickel (II)    and    cobalt 
(II),  32 
Pyrromethene  derivatives,  structures  of, 

363 
Pyruvic  acid,  decarboxylation  of,  706 

Quanticule  theory  of  Fajans,  132,  203 
Quartz,  use  in  resolution,  333,  622 
Quinaldinic   acid,  complexes  with  iron- 

(II),  38 
Quinalizarin    in    colorimetric    analysis, 

693 
Quinhydrones,  519 

bathochromic  effect  in,  550 

color  in,  550 

hypochromic  effect  in,  551 

properties  of,  552 

table  of,  551 

used  as  a  half -cell,  550 
Quinoline 

donor  properties  of,  72 

platinum-ethylene  complex  of,  501 

platinum-styrene  complex  of,  501 
Quinolinic  acid 

complex  with  irondl 

complex  with  silver,  (06 

Racemates,  active,  563 

Racemic   comp  solution   of,  see 

Resolution 

Racemisation,  mechanism  of,  see 
Mechanism  of  racemisation 

Racemisation,  Btudied  by  tracer  tech- 
nique, til") 


Racemisation  of  amino  acids,  vitamin  B< 
and  metal  iona  in.  71  I 

Racemisation  rate  and  covalenl  binding, 

211 
Radioisotopes 
preparal ion  of.  til  I 
reparation  of,  613 

Radius  ratio 

and  coordination  number,  l  r> 
and  configuration,  356 

Hainan  effect .  563 
Haman  lines,  classification  of.  57fl 
Raman  spectra,  578 
ami    force    constants    of    coordinate 
bonds  ,243 

ran-  isomerism  >t  tidied  by.  300,  5*0 
in  study  of  metal -unsaturated  hydro- 
carbon complexes,  579 
in  study  of  oxyanions,  580 
interpretation  of,  563 
of  tetracoordinate  complexes,  356 
Rare  earth  acetylacetonates,  volatilities 

of,  42 
Hare     earth     complexes     of     dibenzoyl 
methane    and    benzoylacetone,    hy- 
drolysis of,  13 
Rare  earth  complexes 
stabilities  of,  176 
stability  constants  of,  179 
Rare     earth      ethylenediaminetetraace 

tates,  formation  constants  of,  589 
Rare    earths,     "anamalous"     oxidation 

states  of,  181 
Rare  gas  configuration  achieved   by  ac- 
ceptor atom,  111,  see  also   Effective 
atomic  number 
Elate   of  formation   and   dissociation   of 

complex  ion-.  627,  632 
Rate  of  oxidation  vs.  stability,  l"<> 
Rate  «>f  racemisation  and  ionic  bonds,  210 
Hate  of  racemisation  and  resolution  of 

complexes.  210 
Hate  of  reaction  and  bond  -trench.  213 
Hate  of  reduction  vs.  stability,  1<MI 

Hate  of  substitution   reaction-  and  bond 

stability,  213 
Ratio  of  radii 

And  coordination  number,  I  13 

And  tetrahedraJ  configuration, 
i:     oni     lyeinj 


824 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Reaction  mechanism,  see  Mechanism  of 

reaction 
Reaction  rate  constants  in  electrodeposi- 

tion,  638 
Recoura's  Sulfate,  97 
Rectangular  configuration  for  tetracova- 

lent  complexes,  354 
Redox  potential 

Of    complexes    of    chromium,    cobalt, 
copper,  iron  and  vanadium,  185-189 
Of    hemochrome-hemichrome    system, 
721 
Reduction,  stability  toward,  398 
Reduction  of  complex  ions,  628 

as  a  slow  process,  633 
Refractivity  and  polarizability  of  donor 
atoms  in  alkyl  derivatives  of  H20, 
H2S,  H3N,  H3P,  128,  129 
Refractrometry  in  study  of  complexes, 

583 
Reinecke's  Salt,  97 

Relative     configuration     of     analogous 
enantiomorphs 
active  racemate  method,  340,  341 
optically  active   quartz  method,  341, 

342 
rotation,  dispersion  method,  298,  337- 

340 
solubility  method,  337 
Replacement  series  of  donor  groups,  572 
Resolution  of  racemic  complexes 
by  circularly  polarized  light,  336 
by  crystallization  of  diastereoisomers, 

332 
by  differences  in  rate  of  reaction,  336 
by  equilibrium  method,  334 
by  method  of  active  racemates,  333 
by  method  of  "configuration  activity", 

335 
by  optically  active  quartz,  333,  622 
by  preferential  crystallization,  332 
by  spontaneous  crystallization,  331 
Resolution  of  optical  isomers,  uncertain- 
ties in,  368 
Resolution  of  tetracoordinate  complexes, 

356 
Resolvability  of  complexes  and  magnetic 

criterion  for  bond  type,  210 
Resonance  between  ionic   and    covalent 

types,  208 
Resonance,  conditions  for,  209 


Resonance  effects  in  stability  of  chelate 

rings,  245 
Resonance   stabilization   of   four  mem- 

bered  triazine  chelate  ring,  248 
Resonance  structures,  189-196 
Resonance  theory,  199 

and  trans  effect,  195 
Reversibility  in  electrodeposition,  640 
Rhenium 

electrodeposition  of,  660 

halo     complexes     of,      similarity     to 
ruthenium  complexes,  15 
Rhenium(III)  chloride,  structure  of,  18 
Rhodium 

electrodeposition  of,  658 

halo  complexes  of,  13 
Rhodium  complexes,  polarographic  anal- 
ysis of,  696 
Rhodium  (III) 

complexes  with  arsines,  80 

complexes  with  thioethers,  51 

complexes  with  thiourea,  53 

cyclopentadienyl  compound,  498 
Rhodochromic  ion,  271 
Rhombic  bisphenoidal  configuration  for 

tetracovalent  complexes,  354 
Rieset's  First  Chloride,  97 
Rieset's  Second  Chloride,  97 
Ring  phosphates,  in  water  softening,  775 
Ring  size  and  stability  of  chelates,  225- 
234 

of  diamines,  228 

of  dicarboxylic  acids,  230 

of  EDTA  homologs,  40,  229 
Ring  size  isomerism,  272 
"Robust"  complexes,  214 
Rochelle  salts  in  copper  plating,  652 
Rosenheim-Miolati  classification  of  poly- 
acids,  474-478 
Roseo  salts,  98 

Jorgensen's  structure  of,  105 
Rotatory  dispersion,  298,  337-340 

relation  to  magnetic  properties,  603 
Roussin  Salts,  534,  536,  594 

black,  97 

red,  97 
Rubeanic  acid,  see  Dithioxamide 
Ruthenium 

arsine  complexes  of,  80 

electrodeposition  of,  660 

halo  complexes  of,  14 


INDEX 


825 


Ruthenium (II),   cyclopentadieny]   com 

pound  of,  198 
Ruthenium(Il      III     dipyridyl  Bystem, 

valence  relations  in,  687 
Ruthenium  oitrosyl  compounds,  ">3!» 
Ruthenium  (II)-  (III)        phenanthroline 
stem,  687 

oxidation   and   reduction   reactions  of, 

Btereochemisl  ry,  353 

Salicylaldehyde  complexes,  41,  602 

resonance  in  stability  of  chelates,  264 
Salts 

effect  on  pH  of  aluminum  oxide  sols. 

465 
hydra  ted,  See  Hydrates  and  Hydrated 
ions 
Sarcosine  -  bis  (ethylenediamine)  cobalt 

[II)  ion,  324' 
Scandium 

acetylacetonate,  43 

complex  ions  of,  aggregation  from  ola- 

tion,  453 
halo  complexes  of,  11 
Scandium  perchlorate,  equilibria  in  solu- 
tions of,  453 
Schiff  bases,  metal  complexes  of,  713 
Second  absorption  band,  565 

stability  series  for,  566 
Secondary  amines,  coordination  by,  62 
Secondary  valence,  109,  157 
Second  order,  compounds  of,  158 
Selective    analytical    reagents;    role    of 
steric  factors,  237,  238,  672-681,  688 
et  seq. 
Selenato  group,  bridiging  by,  462 
Selenite  complexes  of  nickel,  copper  and 
cobalt,  58 
selenitopentamminecobalt(III)  ion,  58 
Selenium,  elect rodeposition  of,  660 
Selenomencaptides    and    ethers,    com 

plexes  oi 
Selenoxides.  coordination  by,  53 
Sequestering  ability 
of  phosphates,  factor-  affecting,  771 
-1  -  used,  768 
Sequestering  reaction,  nature  of,  77:: 
Sequestration,  definition  of,  768 
Sequestrene,  223.  777 

•i-coordinatc  bismuth   V  in 

[BiOF.1-,  8 


Seven  coordinate  configurations,  302 
Seven  inembered  rings,  254 
Sexidentate  ligand 

optical  activity  of  complexes,  235 
type-.  234 
Bigma  bond.  Jill 

Bigma  orbitals,  199 
Silicotungstic  acid,  173  476;  480  L81 
Silver 
configurations  of  silver(I)   and  Bilver 

(II)  complexes,  365 

deposition  from  the  cyanide  complex, 
alkali  metal  reduction  hypothesis, 
626 

electrodeposition  of,  642,  661 

higher  oxidation  states  of,  408 

mechanisms  of  electrodeposition  of, 
636 

relation  of  oxidation  states   to  elec- 
tronic configuration,  369 
Silver-lead  alloys,  electrodeposition  of, 

667 
Silver(I) 

cationic  complexes  in  iodide  or  cyanide 
solutions,  631 

complexes  with  dichlorotetrammine- 
cobalt(III)  ion,  19 

electrode  polarization  in  complex  solu- 
tions, 636 

halo  complexes  of,  17 

number  of  unpaired  electron-  and 
structural  type,  209 

stereochemistry  of  tetracovalent  com- 
plexes, 371 

three  coordinate,  385 

two  coordinate,  383 
Silver  (II),  190 

complexes  of,  408 

dipyridyl  complex  of,  190 

o-phenanl  broline  complex  of,  190 

planar  configuration  of  complexes,  169 

stabilization  by  2, 2* -dipyridyl,  68 

stereochemistry  of  tetracovalenl  cum 
plexes,  371 

impaired     electron-     and     structural 
type-.  200 
Silver    IN 

complexes  of,  108 

ethylenedibiguanide  complex  of,  71 
planar  configuration  of  complexes,  160 


826 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Silver  ammines 
effect  of  pressure  on,  622 
solvation  effects  and  stability,  224 
stabilities,  181 
Silver-benzene  complex,  structure  of,  507 
Silver-cadmium  alloys,  elect  rodeposition 

of,  667,  669 
Silver    complexes,     with    homologs     of 

ethylenediamine,  stability  of,  234 
Silver  cyanide,  structure  of,  89 
Silver  ferrocyanide,  structure  of,  91 
Silver-lead  alloys,  electrodeposition  of, 

667 
Silver-mercury(II)  iodide,  isomerism  of, 

263 
Silver-olefin  complexes,  495 
failure  to  isomerize,  504 
structure  of,  505 
Silver  tetranitrodiamminecobaltate 

(III),  x-ray  analysis  of,  298 
"Simple"  compounds  of  metals,  probable 

complexity  of,  11 
Simple  vs.  complex  ions,  670 
electrodeposition  from,  625 
Six-membered  rings,  preferential  forma- 
tion by  unsaturated  ligands,  231 
Size  of  coordinating  groups 
effect  on  nature  of  electrodeposits,  642 
effect  on  nickel  deposition,  657 
Sodium  chloride,  ammonates  of,  2 
Sodium  complexes  of  /3-diketones,  2,  182 
Sodium  salicylaldehyde  complexes,  2 
Sols,  see  Hydrosols 
Solubility,  study  of  complex  formation 

by,  623 
Solvate  isomerism,  261 
Solvation    effects,    role    of,    in    chelate 

stability,  224 
Solvation  of  ions,  20,  670 
Solvent  and  molecular  configuration,  173 
Solvent    distribution    studies    of    com- 
plexes, 44,  624 
Solvent  systems,  acid-base  reactions  in, 

420 
Sorbitol  complexes,  24 
Spatial  configurations    (table),    170,    see 

also  Isomer  patterns 
sp3  hybridization,  168 

and  magnetic  moments  of  complexes, 
172 
spd  hybridization,  165 


from 
358 


the 


sp2d  hybridization,  169 
sp8d2  hybridization,  167 
Specificity,  optical,  in  biochemical  proc 

esses,  716 
Spectra,  classification  of,  563 
Spectrophotometry 
study     of     complexes     by,     563 

cobalt  complexes,  575 
use  of  in  mechanism  studies,  573 
Spin  and  energy  of  molecule,  171 
Spongy  electrodeposits,  643 
Square    planar    configuration    of    com 
plexes,  169,  354  et  seq. 
determination    of    configuration 

reactions  with  bidentate  group, 
evidence  for,  356 
observed,  355 
of   platinum (II)    complexes,    and 

ionic  model,  131 
of  tetracovalent  complexes,  354 
orbital  hybridization  leading  to,  359 
planar  vs.  tetrahedral  configuration  for 
nickel  (II),  173 
Stability  and  bond  type,  3 
Stability    constants    of    complexes,    see 
Dissociation    constants,    Formation 
constants  and  Instability  constants 
and  ionic  potential,  121 
and    second   ionization    potentials    of 

metals,  177 
determination  of,  405,  569,  593 
of  diketone  complexes,  182 
of  ethylenediamine  complexes,  178 
of    ethylenediaminetetracetate    com- 
plexes, 179,  781 
of  phosphate  complexes,  775 
of  rare  earth  complexes,  179 
of  salicylaldehyde  complexes,  178 
of  silver  complexes,  180 
of  substituted  malonato  complexes,  183 
of  /3,/3', /3"-triaminotriethylamine  com- 
plexes, 178 
Stability  of  complexes  3,  180,  398 
and  atomic  orbital  theory,  174 
and    electronegativity    of   metal    ion, 

175,  413 
and  entropy,  130 
and  ionic  radii,  177 
and  metal -ion  type,  177 
and  polarization,  125,  127 
and  rate  of  reaction,  400 


INDEX 


827 


ami  ring  -  also  pages  2.">:;  260 

diamines,  228 

dicarboxylic  acids,  290 

EDTA  homologs,  10 
and  role  of  ligand,  17").  177 
and  role  of  metal  ion,  171 
and     second     ionization     potential     of 

metals,  177 
and  steric  factors,  130 
containing  five  and  six  members,  rea- 
son for.  250 
determination  of,  405,  569,  593 
effect  on  nature  of  electrodeposits,  642 
effect   on   throwing  power  in   electro 

deposition,  644 
oi  alkaline  earths,  181 
of  ammonia  and  water,  based  on  ionic 

model,  122-125,  126.  127 
of    ethylenediaminetetraacetate,    177, 

179 
of  ions  of  first  series  transition  metals, 

130 
relation  to  magnetic  moments,  602,  603 
Stability   series,    relation   of   to   second 

absorption  band,  566 
Stabilization 

resonance,     through     coplanarity     of 

complexing  agent,  718 
through  enveloping  cyclization,  718 
Stabilization  of  valence  by  coordination, 
91,  184-190,  398-115 
contributing  factors,  412 
Standard  state  of  ligand  in  definition  of 

chelate  effect,  252 
St  annate,  electrodeposition  from,  663 
Stannite,  electrodeposition  from,  663 
Statistical    effect    in    stability    of   com- 
plexes, 249 
Stereochemical!}-  active  electron  pairs, 
131 
in  K,SbF«  and  K2Sb2F7  ,  8 
Stereochemi-m 

and  nature  of  central  atom,  17! 
and  polarized  ionic  model,  131 
of  coordination  number  eight,  394-397 
of  coordination  number  five,  387-392 
of  coordination  number  four,  354-381 

annotated  bibliography,  370  381 
of  coordination  number  sev< 
of  coordination  number  six,  274 
of  coordination  number  three,  384-387 


of  i rdination  number  two,  382  384 

of  inorganic  complex  compounds  emu 
pared  to  that  of  organic  compounds, 
277 

Stereospecificity 
complexes  used  as  resolving  agents,  315 
limited  number  of  isomers,  313  315 
partial  asymmel ric  Bynl hesis,  316 

Steric  factors 
and  molecular  configuration,  17:! 

determined  by  the  metal   ion,  239 

in  complex  stability,  236 

Steric  inhibition  of  resonance,  248 
Stern-Gerlach.     molecular    beam     tech- 
nique, 598 
Stibine  complexes  of  platinum  and  palla- 
dium, 81 
Stibines,  complexing  ability  of,  78,  363 
stien,  see  1,2-Diphenylethylenediamine 

and  Stilbenediamine 
Stilbene,  ultraviolet  spectra  of  platinum 

complex,  504 
Stilbenediamine,  96 

chelates  of,  228 
Structural  isomerism,  268 
Styrene 
palladium  complex  of,  493,  501 
platinum  complex  of,  488,  489,  491,  492 
Substitution  rate 
and  covalent  character  of  bond,  217 
and  inner  d  orbitals,  214 
related  to  covalence,  616 
Substitution  reactions  in  complex  ions, 
213-219;  342-348 
relation  to  bond  type,  615 
Successive  formation  constants,  determi- 
nation of,  405,  593 
Sugars,  coordination  of,  2 1 
Sulfate  ion 

chleation  by,  29 
hydration  of,  31 
Sulfato-aqo     complexes      of      chromium 

Illi, 29 
Snlfatobis(ethylenediamine  coball   III 
bromide     l-hydrate,     questionable 
ucture  of,  30 
Sulfat.,  bridgei  162 

Sulfate  comple 
in  chromium  electrodeposition,  I 
of  iridium  111 
of  iron,  30 


828 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Sulfato  complexes — Cont. 

of  manganese,  electrodeposition  from, 

655 
of  rhodium,  658 
Sulfato-oxalato  complexes  of  iron  (III), 

30 
Sulf atopentamminecobalt (III)  bromide , 

29,  267 
Sulfide  ion  coordination,  in  qualitative 

analysis  and  metallurgy,  48 
Sulfide  ion,  reaction  with  porphyrin  com- 
plexes, 726,  728 
Sulfito  complexes 
of  chromium,  460 
of  cobalt  (III),  57 
of  iridium(III),  58 
of  platinum (II),  57 
of  rhodium  (III),  58 
of  ruthenium (II),  58 
Sulfito  group,  bridging,  463 
Sulfonyldiacetobisethylenediamineco- 

balt(III)  ion,  256 
Sulfur  and  nitrogen,   relative   affinities 

for  metals,  51;  front  end  paper 
Sulfur  atom,  as  center  of  optical  activity, 

325 
Sulfur-containing     dyes,     metal      com- 
pounds of,  762 
Sulfur  coordinators,  3 
Sulfur  dioxide 
complexes  with  ruthenium (II),  58 
electrodeposition    from    solutions    in, 
671 
Sulfur,  donor  properties  of,  47 
Sulfur  in  metal  deposits  from  thiosulfate 
solutions,  630 
in  nickel  deposits,  657 
Summation  isomerism,  272 
Super-complexes,  620 
electrostatic  treatment  of,  150 
ferro-  and  ferricyanides,  90 
olated,  451-455 
Surface  tension,  study  of  complexes  by, 

622 
Symbols  for  names  of  ligands,  96 
Synthetic  fibers,  dyeing  of,  765 
Szilard-Chalmers  process,  613 

Tannin-tartar  emetic  mordant,  745 
Tanning 
aluminum,  471 


chromium,  453,  454,  456,  461,  471 
iron,  460,  471 
zirconium,  471 
Tantalum 
electrodeposition  of,  645,  665 
halo  complexes  of,  16 
stereochemistry  of  Ta6Bri4-7H20  and 
Ta6Cli4-7H20,  366,  375 
Tartrates 
in  electrodeposition  of  copper,  642,  651, 

652 
in  electrodeposition  of  lead,  655 
in  electrodeposition  of  silver,  642 
Tartrate  complexes,  ring  size  in,  232 
TAS,  see  Methyl  bis(3-dimethylarseno- 

propyl)arsine 
Taube 

classification  of  complexes,  213 
substitution  theory  of,  615 
Tellurate 

as  donor  group,  407 
copper  (III)  complex  of,  31 
Tellurate  ion,  structure  of,  580 
Tellurium 
dye  complexes  of,  755 
electrodeposition  of,  662 
Telluromercaptides    and    ethers,    com- 
plexes of,  53 
Temperature,    relation    of    to    color    of 

complexes,  564 
Temperature   and  molecular  configura- 
tion, 173 
Ten-membered  rings,  258,  260 
Ternary  alloy,  electrodeposition  of,  667 
2,2',2"-Terpyridyl,  complexes  of,  68 
Tertiary  amines,  coordination  by,  62 
Tertiary  arsine  complex 

containing    both    copper  (I)    and  cop- 
per (II),  83 
containing  either  nickel  (II)  or  nickel 
(III),  187 
Tertiary  arsines,  donor  properties  of,  78 
Tertiary  phosphines,  donor  properties  of, 

78,  187 
Tetrabromoplatinate(II)   ion,   exchange 

reaction  of,  12 
Tetrachlorocuprate(II)  ion,  5 
Tetrachloro(0,j3'-diaminediethylsulfide)- 

platinum(IV),  385 
Tetrachlorodiammineplatinum(IV) ,    cis 
and  trans  isomers,  283 


INDEX 


829 


Tetrachlorodimethylphthalatotitanium 

IV 
Tetrachloropalladatei.il  >  ion 
dissociation  constant  of,  13 
reduction  of,  669 
Tetrachloroplatinate(II)  ion,  134,356 
Tetrachloro(thiodiethylenediamine- 

\.s  platinum (IV),  326 
Tetrachloro-/i-tris-benzidine  dinickel,  264 
Tetracoordinate  complexes 

annotated     bibliography     on     stereo- 
chemistry of,  370-S1 
configurations  of  (Uible),  355 
isomer  patterns  of,  357 
orbitals  involved  for  different  configu- 
rations, 359 
stereochemistry  of,  354-381 
Tetracyanocuprate(I)  ion,  stabilization 

of  valence  in,  407 
Tetracyanonickelate(O)  ion,  409 
Tetracyanonickelate(II)  ion 

exchange  of  cyanide   and  nickel  ions 
by,  89,  213 
Tetracyanopalladate(O)     ion,    structure 

off 270 
Tetradentate  ligands,  64,  65,  220,  320 
Tetraethyldigold  sulfate,  sulfato  bridge 

in,  31 
Tetraethylenepentamine,  cobalt  (II) 

complex  of,  65 
Tetrafluoroberyllate,  resemblance  to  sul- 
fate, 5 
Tetragonal    bisphenoidal     configuration 

for  tetracovalent  complexes,  354 
Tetrahedral  configuration,  168,  354,  355 
evidence  for,  354,  356 
for  nickel (II;  (vs.  planar),  173 
in  forced  planar  structure,  243 
in  nickel  carbony],  171 
orbital  hybridization  leading  to,  359 
relation  to  radius  ratio,  356 
Tetrahydroxydodecaquo-/i-decaolhe\a- 

chromium(III)  ion,  452 
Tetraketones,  polymeric  complexes  of,  42 
Tetrakis-(2-aminoethyl)ethylenedia- 
mine  complexes,  235 

kis-(ethylenediamine)-|*-amino-ni- 
tro-dicobalf  III  ion,  isomeric 
forms  of,  322 


Tel  ralris-(ethylenediamine  l-ji  amino 

peroxo -cobalt (III) -cobalt  (IV)      ion, 
isomeric  forma  of,  321 
Tetiakis-(ethylenediamine)-/i  diol  dico 

balt(III)  ion,  lis 
Te1  rakia    oxalato)-/i-diol-dichromate 

(III)  ion,  35,  462 
Tel  raids  -I  phosphorus   trichloride  I  nickel 

(0),  151 
Tetrakis- (phosphorus     trifluoride)  nickel 
(0),  151,205 
preparation  of,  86 
Tetrametaphosphate,  complexes  of,  769 
3,3',5,5'-Tetramethyl-4,4'-dicarbeth- 
oxypyrromethene,   steric    hindrance 
in  metal  complexes,  242 
Tetramethylenediamine,  complexing  by, 

64 
Tetramminecopper(II)  ion,  dissociation 

of,  2 
Tetranitrodiammine  cobaltate(III)  ion, 

configuration  of,  292,  298 
Tetraoxalato  uranate(IV)  ion,  structure 

of,  396 
2,2/,4,4'-Tetraphenylazadipyrrometh- 

ine,  metal  complexes  of,  762 
tetrapy,  see,  2,2',2",2'"-Tetrapyridyl 
a,a',a",a/"-Tetrapyridyl  complexes,  68, 

96,  690 
Thallium,  electrodeposition  of,  663 

from  cyanide  complexes,  664 
Thallium  (I) 

acetoacetic   ester  with   carbon   disul- 

phide,  258 
alcoholates,  386 
2,2-biphenol  complex,  255 
stereochemistry  of  tetracovalent  com- 
plexes, 374 
structure  of  complexes,  370 
three-coordinate,  387 
Thallium(I)-(III)  couple,   valence   rela- 
tionships of,  401 
Thallium(III) 
dimethylacetylacetone  complex,  H 
halo  complexes  of,  6 
questionable  planar  -tincture  for  com- 
plexes of,  370 
stereochemistry  of  tetracovalent  com- 
plexes, 374 
Thenoyltrifluoroacetone 

coordinating  ability  of,  41,  182 


830 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Thenoyltrifluoroacetone  —Cont. 

in    separation    of    zirconium    and    haf- 
nium,  1  I 
Thermal  measurements  in  study  of  com- 
plexes, 020 
Thermochemical  cycle 
in  complex  formation,  137 
quantitative  treatment,  142 
Thermochromic  complexes  of  copper(II) 
and  N -substituted  ethylenediamine, 
66 
Thermodynamic    activity,    relation    to 

Pfeiffer  effect,  582 
Thermodynamics  of  chelate  effect,  251 
Thiazine  dyes,  752 
Thiazone,    complexes    with    copper(II) 

and  nickel  (II)  ions,  54 
thio,  see  Thiosemicarbazide 
Thioarsenite,     electrodeposition     from, 

648 
Thiocarbanilide,  analytical  uses  of,  694 
Thiocarbonyl  compounds,  donor  proper- 
ties of,  53 
6,8-Thioctic  acid,  739 
Thioc3ranate 
as  a  bridge  group,  83 
cadmium  complex  of,  405 
cobalt (II)  complex  of,  76,  688 
iron  (III)  complex  of,  76 
Thiocyanate  ion,  donor  properties  of,  57 
Thiocyanato-isothiocyanato    isomerism, 

270 
Thiodicyandiamidine      complexes     with 
cobalt  (II),    cobalt  (III),    nickel  (II) 
and  palladium (II),  55 
Thioethers,  complexing  ability  of,  48,  49, 

123,  129 
Thioglycolic  acid,  complexes  of,  52,  731 
Thiohydrate  formation,  48 
Thiohydrolysis,  48 
a-Thiol  fatty  acids,  clinical  possibilities, 

52 
Thionyl  Purple  2B,  762 
Thiosemicarbazide,  96 
complexes    with   platinum(II),   palla- 
dium (II)  and  nickel(II),  53 
Thiosulfato  complexes 
electrode  polarization  of  copper  com- 
plex, 637 
electrodeposition  of  metals  from,  630 
electrodeposition  of  silver  from,  661 


isomerism  in,  270 

sulfur  in  electrodeposits  from,  630,  657 
Thiosulfate  ion,  donor  properties  of,  58 
Thiourea 

analytical  uses  of,  694 

complexing  by,  53,  96 

in  determination  of  cadmium,  682 

in     Kurnakov's     test     for     cis-trans- 
isomers  of  planar  complexes,  53,  358 

molecular  compounds  of,  560 
Thiourea  complexes 

in  electrodeposition  of  copper,  651 

in  electrodeposition  of  silver,  661 
Third  absorption  band,  567 
Thirteen-membered  rings,  258 
Thorium  oxide,  hydrous,  peptization  by 

salts  of  hydroxy  acids,  466 
Thorium  oxide,  hydrosols 

anionic,  466 

anion  penetration  in,  466 
Three-coordinate  configurations,  385 
Three -membered  chelate  rings,  evidence 

on,  225 
Threshold  treatment,  776 
Throwing  power  in  electrodeposition,  644 
Time  of  dissociation  of  complex  ion,  627, 

632 
Time  of  formation  of  complex  ion,  627 
Tin 

electrodeposition  of,  662 

valence  relations  in  complexes,  404 
Tin  (II) 

stereochemistry  of  tetracovalent  com- 
plexes, 374 

structure  of  K2SnCl4-2H20,  367 
Tin  (III),  existence  of,  373 
Tin  (IV)  chloride,  hydrolysis  of,  446 
Tin-copper  alloys,  electrodeposition  of, 

667 
Tin-nickel   alloys,   electrodeposition   of, 

669 
Titanium,  dye  complexes  of,  755 
Titanium,  electrodeposition  of,  665 
Titanium  (III),    cyclopentadienyl    com- 
pounds of,  499,  508 
Titanium  (IV),     cyclopentadienyl     com- 
pound of,  498 
Titanium(IV),  halo  complexes  of,  10 
Titanium  oxide  sols,  471 
Titration  of  metal  ions  with  complexing 
agents,  683 


INDEX 


831 


Tolidine,  complexes  of,  67 

Toluene,  metal  complexes  of,   198 

p  Toluidene,  platinum-ethylene  complex 

of,  501 
nt-Tolylazo-0-naphthol,  copper  lake  of, 

756 
Tracer   method,    in    study    of    racemiza- 

tion,  615 
Tracers 
artificial,  612 
natural,  612 
Tracer  techniques  in  study  of  complexes, 

611 
Transamination,  712 
Trans-diammine  diet  hylenet  riamine 

platinum (II)  ion,  259 
Trans  etYeet.  147,  196,  206,  358,  568 
application  to  acid-base  reactions,  42!) 
explanation    from    electrostatics    and 

polarisation,  147 
explanation     from      polarized      ionic 

model,  146 
explanation  from  resonance  theory,  195 
factors  promoting,  148 
of  coordinated  sulfur,  49 
of  double  bonds,  204 
of  easily  polarised  ligands,  146 
of  ethylene,  148,  149,  490,  491 
of  phosphorus  trifluoride,  197,  205 
rate  and  mechanism,  219 
relation  to  electronegativity,  196 
uncertainty    in    quantitative    evalua- 
tion, 149 
use  in  synthesis,  294 
Trans  elimination,  146,  198,  358 
Transfer  of  complex   ions  to  electrode 

surfaces,  632 
Transition  elements.  172 

electrodeposition  of,  640 
Transition  metal  ions,  relative  affinities 

for  halides  and  oxygen,  9 
Trans  planar  configuration,  assignment 

from  chemical  behavior,  358-9 
Transport    cumbers,   in   study  of  com 

plexes,  618 
Trans  positions,  spanning,  258 
t ren,  see  2.2' .2"-Triaminot riethylamine 

ar,/3,7-Triainino|>rop:me 

chelation  by,  65,  288 

preferential    formation    of    five  mem 

bered  ring,  227 


/3,/3',£"-Triamihoi  i  iet  h\  Limine 

complexes  of,  65,  96 

complex  u  ith  copper   II.  Bteric  factors 
in  stability,  211 

complex  with  nickeh  II  i,  stability  of. 

241 
complexes  with  nickel,  palladium,  and 

platinum,  363 

complex  with  plat  inuini  1 1 ),  2ii!) 
Triazene        complexes.       four-membcrcd 

rings  in,  226 
Triazine   rings,   resonance   stabilization 

of,  248 
Tri  1  iromobis- (t  lie t  hylphosphine) nickel 

(III),  392 
Tricyanonickelate(I),  reaction  with  ni- 
tric oxide  and  carbon  monoxide,  92 
Tridentate  ligands,  220,  288,  318 
trien,  see  Triethylenetetramine 
Triethanolamine  complexes,  25,  257 

in  cobalt  electrodeposition,  651 

in  iron  electrodeposition,  655 
Triethylenetetramine  complexes,  64,  220, 

320 
Triethylphosphite  complex  of  gold  (I),  86 
Triglycine,  777,  782 
Trigonal  bipyramid  structure,  520 
2,4,5-Trihydroxytoluene,     metal     com- 
plexes of,  749 
Trilon  A,  777 
Trilon  B,  777 

trim,  see  Trimethylenediamine 
Trimetaphosphate,  complexes  of,  769 
Trimethyldichloroantimonate,  389 
Trimethylenediamine 

coordinating  ability  of,  64,  96,  183,  228 

stability   of   chelates   compared   with 
those  of  ethylenediamine,  230 
Trimethylethylene,  palladium d  1 1   com- 
pound of,  493 
Trimethylphosphite 

complex  with  gold  (I),  86 

complex  wit  h  platinum!  II I  chloride,  s"> 
trin,  see  2, 2', 2"  Triaminotriethylamine 
Trinitrotriamminecobali   III.  28 

isomers  of,  282,  297 

polymerization  isomers  of,  264 
Triphenylmethane  derivatives,  lake  for 

mation  from.  754 
Triphosphate  complex,  with  calcium.  771 


832 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


Triphosphate   ion,   hydrolytic   degrada- 
tion of,  772 
Tripolyphosphate,  sodium  salt,  769 
tripy,  see  2,2',2"-Tripyridyl 
2,2',2"-Tripyridyl,  96 

in  determination  of  iron  and  cobalt, 
689,  690 
Tris(biguanidinium)cobalt(III),  330 
Tris(carbonato)cobaltate(III)  ion,  33 
Tris(a,a'-dipyridyl)iron(II)ion,  687 
Tris(dipyridyl)  osmium  (II)      and      (III) 

ions,  353 
Tris-(ethylenediamine)  complexes 
of  cadmium,  claimed  resolution  of,  212 
of  cobalt  (III),  resolvability  and  bond 
type,  210 

retention  of  optical  activity,  1 
of     rhodium  (III),     resolvability    and 

bond  type,  210 
of  zinc,  claimed  resolution  of,  212 
rotatory  dispersion  curves  of,  339 
Tris (glycine)  cobalt  (III),  cis  and  trans 

isomers  of,  284 
Tris-(N-hydroxethylethylenediamine)- 
cobalt(III)ion,  remarkable  stability 
of,  67 
Tris-(oxalato)aluminate  ion 
failure  to  resolve,  211 
rapid  exchange  with  labeled  oxalate, 
211 
Tris- (oxalato)chromate (III)  ion 
electrodeposition  from,  629,  650 
failure  to  exchange  with  labeled  oxa- 
late, 211 
mechanism  of  racemization  of,  413 
resolution  of,  211 
Tris (oxalato)cobaltate (III)  ion 
failure  to  exchange  with  labeled  oxa- 
late, 211 
diamagnetism  of,  211 
resolution  of,  211 
resolvability  and  bond  type,  210 
Tris(oxalato)      complexes,      absorption 
spectra  and  rotatory  dispersion  of, 
338 
Tris-  (oxalato)ferrate  (III)     ion,     failure 

to  resolve,  210 
Tris-(oxalato)gallate  ion,  claimed  resolu- 
tion of,  212 


Tris- (oxalato)manganate  (III)   ion,   fail- 
ure to  resolve,  210 
Tris-(oxalato)rhodiate(III)  ion,  resolva- 
bility and  bond  type,  210 
Tris-(o-phenanthroline)iron(II)  ion,  686 

potentials  of  methyl  substituted,  686 
Tris-(o-phenanthroline)      ruthenium  (II) 

and  (III)  ions,  353 
Trypsin,  metal  activation  of,  703 
TTA,  see  Thenoyltrifluoroacetone 
Tungstate  ion,  structure  of,  580 
Tungsten 

alloys,  electrodeposition  of,  667 

carbonyl   cyclopentadienyl   compound 
of,  499 

cathodic  reduction  of  cyano  complexes 
of,  629 

cyanide  complexes  of,  395 

cyanide  complexes  and  ionic  struc- 
tures, 194 

dye  complexes  of,  755 

electrodeposition  of,  665 

halo  complexes  of,  15 

oxyhalo  complexes  of,  16 

plating  of,  645 
Tungstic  acid,  477,  480,  482 
Turkey-Red  lake,  749 
Turnbull's  blue,  structure  of,  90,  610 
Twelve -membered  ring,  260 
Twenty-two  membered  ring,  260 
Two-shelled  complexes,  620 
Tyrosinase,  724 

Ultraphosphates,  structure  of,  769 
Ultrasonic  velocities,  study  of  complexes 

by,  624 
Ultraviolet    spectra,    interpretation    of, 

563 
Un,  see  Unsaturated  group,  492 
Unidentate  ligand,  220 
Unpaired  electrons 
and  stereochemical  configurations,  171 
and  structural  types,  209 
Unsaturated  acids 

compounds    with    copper(I)    chloride, 

495 
coordination  compounds  of,  488 
platinum  complexes  of,  488 
Unsaturated  alcohols 
compounds   with   copper(I)    chloride, 
495 


INDEX 


833 


silver  complexes  ol 
Unsaturated    compounds,    coordination 
compounds    with    metals,    187  506, 
S  ■    also  Ethylene,  Olefin,  and  the 

BPOcific     olefins 

[Jnsaturated  poly-acids,  17.~>  177 
rjnsymmetrica]   bidentate   donor  mole- 
cules in  complexes,  284 
Unusual  oxidation  Btates,  184  190,    i<>7 
U2,  573 

methods  for  characterising,   107 
Uranium  (V 

existence  of,  404 

reduction  of,  405 
Uranium  (VT) 

oxy chloride,  reduction  of,  405 

polarographic  behavior  of,  404 
Uranyl  ion,  polarography  of,  589 

I'rea 
in  copper  electrodeposition,  651 
molecular  compounds  of,  560 

A  alence  Btates 

and  electronegativity  of  bonding  atom 

in  ligand,  185 
and    stable    electronic    configurations, 

185 
generalizations     regarding     stabiliza- 
tion, 185-190 
stabilization  by  coordination,  184-100; 
398-415 
Vanadium 
coordination  with  oxygen,  406 
dipyridyl  complexes,  valence  relation- 
ships of,  186 
dye  complexes  of,  755,  758 
electrodeposition  of,  665 
halo  complexes,  of,  10 
quantitative    separation    of    V        and 
V+++,  187 
Vanadium(II)-(III), 

aquo  couple,  valence  relations  of.   186 
cyanide  couple,   valence  relations  of. 
186 
VanadiumflVi,    cyclopentadienyl    com 

pound  of,  498 
Vanadium (V  .  reduction  of.  406 
Vanadic  acid,  477 
Vanadium  complex) 
polarographic  characteristics  of,  105 
substitution  reactions  of,  616 


Vaquelin's  Salt.  :>7 

Versene,  2223,  777 

Vibration,  amplitude   related   to  color. 

564 
Vibrational  energy,  relation  of  to  first 

absorpl ion  band.  585 
Vicinal  influence.  .">s| 
Violeo  -alt-.  Hn 

Visible  Bpectra,  interpretation  of,  •">«'»:; 
Vitamin  H,  .  712 
Vitamin  Bu  ,  737 

polarographic   reduction   of,  739 
Volume  additivity,  study  of  complexes 

by,  623 
Volumetric  analysis,  complexes  in,  683 

et  seq. 
Vortmann's  Sulfate,  97 
as  peroxo  complex.  27 
Vulcan  Blue,  761 

Walden  inversion   type   reactions.   344 

348 
Water,  see  also  Hydrate,  etc. 

a-  a  bridge  group,  46,  391 

dipole  moment  and  polarizability,  124 
Water  softening,  through  complex  forma 

tion,  768-783 
Wave  mechanics,  162 
Werner 

biographical  sketches  of,  109 

configuration  rule,  582 

coordination  theory,  108 

formulation  of  poly-acids,  473 

system  of  nomenclature,  94 
Wolffram's  Red  Salt.  97 
Wool,  dyeing  of,  764 

X,  see  halide,  96 
Xanthene  dyes,  752.  7">t 
X-ray  analysis 
applications  to  structure  of  complexes, 
191,  356,  606 

criteria  for  bond  type,  213 
interpretation  of.  in  terms  of  bond-. 

368 
use  in  distinguishing  <-i-  and  Man-  iso- 
mers, -' »7 
validity  of  incomplete  studies,  307 

X-ray  Btudies 

of  ethylene  complexes,  •"><>! 

of  poly   acid- 


834 


CHEMISTRY  OF  THE  COORDINATION  COMPOUNDS 


X  ray  studies — Cant. 
of  Zeise's  salt,  504 
X-ray  structure  determinations,  coordi- 
nation number  from,  356,  361 
Xylene,  metal  complexes  of,  498 

Zeise's  acid,  488 
Zeise's  Salt,  66,  97,  488 
oxidation  of,  503 
x-ray  structure  of,  504 
Zinc(0)-(II)  couple,  401 
Zinc, 

electrodeposition  of,  664 
from  ammines,  638 
from  cyanide  solutions,  631,  638 
from  thiosulfate  solutions,  630 
from  zincate  solutions,  638 
in  carbonic  anhydrase,  708 
in  insulin,  709 

stabilization  by  hydroxyl  ion  coordi- 
nation, 402 
stereochemistry  of  tetracovalent  com- 
plexes, 372 
Zincate,  electrodeposition  from,  664 
Zinc  amide,  amphoterism  of,  418 
Zinc  chloride-amylene  compound,  497 
Zinc  complex  of  5-sulfo-8-hydroxyquino- 
line,  resolution  of,  72 


Zinc  cyanide,  structure  of,  89 
Zinc,  dye  complexes  of,  746,  757,  758,  760 
Zinc  halide  ammines,  stability  of,  139 
Zinc  halide  complexes,  5 

stability  series,  594 
Zinc-hydrazine  complexes,  225 
Zinc-o-phen    complexes,    ionization    of, 

596 
Zirconium 
dye  complexes  of,  755 
electrodeposition  of,  665 
halo  complexes  of,  16 

comparison  with  hafnium,  16 
Zirconate  hydrosols,  467,  468 
Zirconium    complexes    with    a-hydroxy 

acids,  468 
Zirconium,  cyclopentadienyl  compound 

of,  498 
Zirconium  oxide  hydrosols, 
anion  penetration  in,  467 
effect  of  aging  on  pH,  465 
effect  of  a-hydroxy  acid  anions  on  pH, 

467 
reversal  of  charge  of  micelles,  467 
Zirconium    oxide,    hydrous,   peptization 

by  salt  of  hydroxy,  acid,  468 
Zirconium,  plating  of,  645 
Zirconium  tanning,  471 
Zirconyl  chloride  hydrates,  454 


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