10 Coordination Chemistry 10.1. INTRODUCTION Tassaert made the beginning in 1798 with an orange coloured compound, CoCl,.6NH,, obtained on standing an ammoniacal solution of cobalt(II) chloride. This same. compound is now easily obtained in a pure state and in high yield by air-oxidation of ammoniacal cobalt (ID chloride in the presence of active charcoal as a catalyst. The structures of such compounds were the subject of prolonged controversies between two famous adversaries of the late nineteenth century, Jorgensen and Werner. Werner finally solved the mysteries and brought order into the area by announcing his famous coordination theory. In his coordination theory Werner introduced two great ideas : (1) In a coordination complex a metal ion is engaged in strong binding with a certain number of neutral and/or anionic groups in the first sphere of attraction, now called the coordination zone or sphere; (2) The metal ion is surrounded by the neutral and/or anionic groups in the coordination zone ina definite geometrical arrangement. Werner argued that metal ions possessed two kinds of valence : a primary valence and a secondary valence. Primary valence was satisfied by requisite number of anions or negative groups being present inside and/or outside the coordination sphere. Secondary valence indicated the capacity of the metal ion to accommodate certain number of groups around itself in the first sphere of attraction. Primary valence is now equated with the oxidation state and the secondary valence with the coordination number of the metal ion. C0,0,+NH,+KCI « * Co (NO,),.6 NH, + AgCl KOH (heat) AgNO, Co(SO,). 6NH, + HC! «£22 Go cl, 6NH, IS, Co(OH), . BNH, + AGC oS H,SO, ‘Ag,O KI conc. HCl 10 |, 6NH, + KCI No change Chart 10-1. Werner’s metathetie reactions on CoCl,,6NE). Werner’s ideas were supported by experimental findings on many compounds. It was shown by a large number of simple metathetic reactions on CoCl,. 6NH, in aqueous solution that changes were forced only on the chloride part leaving cobalt andl the six NH, groups intact.COORDINATION CHEMISTRY ‘Weiner redliaed that CoC, SNM, should Bete bo reyresentnd as w strongly bound unit (Co(NH,),]"", leaving the three chloride units as anions outside the square bracket, The unit within the square bracket was recognised as a complex entity, formed as a result of coordinaiton of the neutral NH molecules to the Co’ cation, both these being capable of independent existence. That such a formulation was indeed correct was amply verified by conductance measurements (“1024 - 432 ohms“cm? mole~). Werner named the zone within the square bracket as the first sphere of attraction, the anions being held in a second sphere of attraction. 283 According to Werner the geometrical arrangement around cobalt(III) was octahedral and hence there should be six groups around cobalt(III). The primary valence of +3 of cobalt AID in [Co(NH,),]CI, is satisfied by the three chloride ions while the secondary valence is satisfied by the six NH, groups. Werner formulated CoCl,.5NH, as [CoC\(NH,),ICI, indicating once again that the primary valence of cobalt(III) is satisfied by one CI inside the coordination zone and two CI” outside the coordination zone. In accordance with this formulation the compound is a bi-univatent electrolyte. The compound CoCl,.3NH, was given a non-electrolyte structure (CoCl,(NH,),]- NH,—HCI NH,—Cl I I oH. Co—NH,—NH,—NH,—NH,—Cl I NH,—HCI NH,—CI (a) (b) a cl \ Co—NH,— NE, NH—NE;—Cl Co—NH,—NH,—NH,—Cl \ NH,—Cl cl © @® It is now quite amusing to note the way Jorgensen tried to explain the structures of the cobalt(II) ammines. Following the structure of ethylenediamine chloride (a), Jorgensen represented the cobalt(III) ammines with chains of different lengths. He thought that the valences of cobalt in these compounds were different and therefore stabilised chains of different lengths. Since in CoCl,.6NH, (6) all the three chlorines were some distance away from cobalt these were, according to Jorgensen, readily precipitated by silver nitrate while in CoCl,.5NH, (c) two chlorine were expected to be readily precipitated. Such reasoning would predict that for the-then-unsynthesised CoCI,.3NH, (d) only one chlorine would be readily precipitated. This compound could not be prepared by Jorgensen but its iridium (III) analogue IrCl,.3NHL, was prepared and it failed to give any precipitate with AgNO,. This showed that all the three chlorines were strongly bound and that Jorgensen could not be right. Jorgensen’s problem was that he refused to think of these complicated molecules in any other way outside the chain theory. But characteristic of a genius Alfred Werner brought in a bundle of fresh thoughts in this area and that made all the difference.284 INORGANIC CHEMISTRY Alfred Werner played such a monolithic pivotal role in the development of coordination chemistry that the subject is now synonymous with his name. His researches on the coordination chemistry of metal ammines ure now classics, Metal ammines and related complexes are often lovingly called Werner complexes. 10.2, DEFINITIONS : COMPLEX, LIGAND AND COORDINATION NUMBER A metal ion (whether it is positively charged, neutral or even negatively charged) may combine with neutral molecules or anions to give a new reasonably identifiable entity called a complex. The groups that are bound to the metal ion in a complex are called ligands. The ligands are arranged M+ xL @ [ML,]"" around the metal ion inside the first sphere of attraction in preferred geometries, The number of the ligands bound around a metal ion by ¢ bonds is called the coordination number of the metal ion*. In the above reaction a metal ion M"* reacts with x moles of a neutral ligand L (each capable of taking one coordination position) to form a complex [ML,]"*. Here x represents the coordination number. In (Co(NH,),]Cl,, Co™ has a coordination number six. It will now be proper to discuss any difference between a double salt and a complex salt. A double salt such AIK(SO,),. 12H,O or 3CsCI.CoCl, is so very dissociated into its component ions (in, say, water) that there does not exist any reasonably identifiable complex entity. On the other hand, in complexes CoCl,, 6NH; or CuCl,. 4NH, we have identifiable complex entities [Co(NH,),]** or (Cu(NH,),]°*. A very weak complex wholly dissociates into its component ions. Therefore in a suitable solvent a double salt bréaks down into component ions, which can be detected by the usual tests of these ions. A stable complex ion, obviously, will not respond to the usual tests of its component units. 10.3. COMPLEXATION—A DONOR-ACCEPTOR INTERACTION Structure determinations have shown that most polar ligand molecules are so oriented in a complex that one unshared pair of electrons points directly to the metal ion, The ligand usually is a donor of electrons (Lewis base) and the metal ion an acceptor (Lewis acid). However this statement will be profoundly modified when the metal ion is in a low oxidation state and when the ligand possesses acceptor orbitals in addition to donor orbitals, (Chapter 24, Part II). 10.4. CLASSIFICATIONS OF LIGANDS Based on donor and acceptor properties ligands may be of the following types : * Ligands may be bound to the central atom both by sigma and pi-bonds but the pi-bonds are not considered to determine the coordination number.‘COORDINATION CHEMISTRY 285 Ligands (1) With one (or more) lone (2) Without a lone pair of electrons pair (s) of electrons but with _-bonding electrons e.g. : ethylene, benzene, cyclopenta- . dienyl ion. (@) No vacant orbitals to receive back donated electrons from metal e.g. : H,O, NH, F () Vacant orbitals that can be used to receive back donated x electrons from the low oxidation state metal e.g. : PR, CN”, CO Ligands may also be classified according as the number of coordination positions they occupy around the metal ion (Table 10.1). A ligand having only one lone pair of electrons can only be bonded to a metal ion at just one stereochemical point, that is, it occupies one coordination position only. This is called a unidentate or a monodentate ligand. When the ligand can occupy two positions around the metal ion it is di(bi)dentate; with three, four, five or six such links the ligands are called tri (ter) dentate, tetra (quadri) dentate, pentadentate or hexa (sexa) dentate respectively. Some authors prefer to classify ligands according to the number of donor atoms they can offer to metal ions. Thus, when a ligand 1s two atoms which can simultaneously serve as donors it is called didentate and so on. It follows that a didentate ligand has to have two lone pairs, a tridentate three lone pairs and so on, The Table 10.1. : Some Monodentate and Polydentate Ligands Ligand type Name Formula ve Ammonia NH Nu Monodentate Pyridine —1O> NH, f - 5 Didentate Ethylenediamine (en) . CH, Nay HANS oe Biguanide (Hbig) Ne Ne NSore INORGANIC CHEMISTRY Ligand type Name Formula J Non, Glycinate ion (gly) C=O SA Didentate Bipyridine (bpy) o-Phenanthroline (o-phen) (phen) Oxalate ion (ox) Dicthylenetriamine (dien) Tridentate oO. Ne —— CH, on Iminodiacetate ion (ida) C2 ae 1 Cc — CH, Z 1, o” (CH), — NH— (CH,), —> NH Triethylenctetramine “NH nH — (CH), (trien) TetradentateCOORDINATION CHEMISTRY 7 287 Ligand type Name Formula H c=N ra Ethylenebis (salicylaldimine) Or cH, anion (salen) 7 o-3 =n H 2 SS oes ‘cH NS Cn, Se Pentadentate ethylenediamine oO Ko . ie triacetic acid anion Span H oo Hexadentate ethylencdiaminetetra- (Sexadentate) acetic acid anion (edta) : / Q ‘> Go. p XN fie < ° 7 'N-— CH,—CH, —N B Soon,” Now —o% © ° \ f polydentate ligands may be further subdivided according to the nature of their donor centres. Thus ethylenediamine is a didentate ligand with two neutral donors (N) whereas oxalate ion is didentate with two acidic (anionic) donors (O°). Glycinate ion is again a didentate with one neutral donor (N) and an acidic donor (07). Such discussions may be extended to other polydentates as well. A chelating ligand is one that contains two or more donor centres so disposed that. they can simultaneously occupy more than one position around the same metal ion in the first sphere of coordination*. The resulting complexes are called metal chelates. Thus all TA chelating ligand offers more than one sigma-electron pair donor groups to the same metal ion.COORDINATION CHEMISTRY 293 above and below the pentagon. Some edta complexes are also found to be seven coordinate eg. [Fe(edta)(H,O)] [Mg(edta)(H,0)]*. Seven coordinated [NbF,|* and [TaF,]* have the six F- ions at the six vertices of a trigonal prism while the seventh F- is above the centre of a rectangular face of the prism. 10.5.6, Complexes with Coordination Number 8 : Complexes with this coordination number are usually found with lanthanides and actinides. The commonest structure is the cube —as for Na,(MF,] (M = Pa ; U, Np). Other geometries are hexagonal bipyramid and bicapped trigonal prism (with two ligands above and below the centres of the two triangular faces of the prism). Low oxidation state (with high electron density) of the metal and bulky ligands together favour low coordination number. Bulky ligands lead to overcrowding of the coordination sphere. On the other hand, high oxidation state (with low electron density) of the metal and small sized anionic ligands (eg. F-) favour high coordination number. 10.6. NOMENCLATURE OF COORDINATION COMPLEXES Every year many thousand new complexes are being synthesised and structurally identi- fied. Because of a very wide range of variations in respect of the oxidation number (state) of the central atom and donor atom, charge, ligational behaviour of the ligands etc. formula depicting and naming coordination complexes have been a real challenge. International Union of Pure and Applied Chemistry (IUPAC) has entrusted the work of systematising the literature to a Commission on the Nomenclature of Inorganic Chemistry (CNIC). CNIC has recom- mended rules for the purpose first in 1957, then in 1970 and lately in 1990. The 1990 recommendations are set out here in broad outline. For further details the reader is strongly urged to consult CNIC Recommendations (1990) as cited in Bibliography. Our task may be divided in four parts 1, Writing the formulae for mononuclear coordination complexes 2. Writing the names for mononuclear coordination complexes. 3. Writing the formulae and names of polynuclear complexes with ligand bridges or metal- metal bonds. 4. Writing the formulae and names of chelate complexes Interestingly formulae writing and naming do not follow the same sequence. 10.6.1. Writing the formulae for mononuclear coordination complexes Scquence of symbols within coordination zone : The central atom symbol is written first. Anionic ligands come next and then the neutral ones. Ligands are written in alphabetical order within each class (i.e. anionic and neutral) decided by the first symbol of their formulae. Thus H,O, NH,, SiH", NO,, SO, and OH" are cited at H, N, Si, N, S and O. Ligands having carbon and hydrogen only are treated under C. For organic ligands having heteroatoms other than carbon and hydrogen, C and then H precede other atoms put in alphabetical order. The alphabetical sequence of heteroatoms determines the position of such a ligand in the formula. If the coordination zone has two ligands with the same defining atoms then the one with fewer such atoms precede the one with more. Numbers of defining atoms being equal subsequent symbols decide the sequence. Thus P (C,H,), and C.H,N are cited under296, INORGANIC CHEMISTRY (ID, chromate(v), respectively. Neutral or cationic complexes simply end in the name of the Central atom with the oxidation number in parentheses, Oxidation number, charge number and ionic proportion : Oxidation number of the central atom is indicated by a roman numerical in parentheses a negative sign may be used before the number after its name. No positive sign is used but if necessary. Arabic zero is used for zero oxidation number. Alternatively the net charge on the complex may be written in arabic number with appropriate + or ~ sign after the number, all in parentheses, Tonic proportions of the coordinati Prefixes on the both the ions. Working examples : 1, Let us name the complex Na[PtBrCl(NO,)(NH,)] Here sodium is the counter ion and platinum (II) is the central atom of the anionic complex, So naming has to start with sodium and end in platinate(II) or platinate(1-). Unlike in formula, the ligands are to be named in alphabetical order irrespective of their charge so that we have ammine, bromo, chloro and nitrito. Assuming nitrito coordination occurs through N we have to use nitrito-W or nitrito-,NV. Thus we have : sodium amminebromochloronitrito-N-platinate(II) sodium amminebromochloronitrito-N-platinate(1-) 2. Take the case of the complex [PtCly(C,H,N)(NH,)] This is a neutral complex of platinum with oxidation number + The alphabetical order of the ligands is ammine, chloro, pyridine. So we have : amminedichloro(pyridine)platinum(11) amminedichloro(pyridine)platinum(0) 3. Now let us attempt naming the complex Na,[Fe(CN),NO} s in the anion so that the complex f the ligands will be cyanide and d then the name is : ion entities may be shown by using stoichiometric Sodium is the counter ion while the central atom iron i: anion is to be written a coordinated ferrate. The order of then nitrosyl. If the overall charge of anion is considere sodium pentacyanonitrosylferrate(2-) If we wish to specify the oxidation state of iron we have : sodium pentacyanonitrosylferrate(II) (assuming NO*) 10.6.3. Writing the formulae and names of polynuclear complexes with ligand bridges or metal-metal bonds : Complexes with ligand brid complexes apply here too. Bridging ligands before the ligand formula and separated by ges : In general the rules developed above for mononuclear are indicated by the Greek letter (mu) p appearing hyphen. If the bridging ligand appears more than once multiplicative prefixes are employed. The bridging index, the number of coordination centres connected by a bridging ligand, is indicated by a right subscript, 4.,, where n > 2. Bridging ligands are normally placed last in the formula, A bridging ligand is cited before a corresponding non-bridging ligand eg : di-.-chlorotetrachloro... Multiple bridging is listed in descending order of complexity eg : }13"X0di-,,-oxo-trioxo... Central atoms are listed inCOORDINATION CHEMISTRY 297 alphabetical order after the ligands. Tw i ands. Two or mi i a numerical prefix. For polynuclear anionie species the culfic we sag ees SHOuR BY . pecies the sulfix- i the charge on the ion are added after the central ator nn ine he sumber indicating Exampl UCKNH,),}, (x-OH)] Cl, UFeNO),}, (u-P(CoH,)p},) [CH(C,0,),}, (OH), --pchydroxo-bis (pentamminechromi- um) (5+) pentachloride «-bis(j.-diphenylphosphido)bis (di trosyliron) .di-y-hydroxo-bis(bis(oxalato) chromate(III) di-p-hydroxo-p-nitrito-,N-.O bis(triammine)cobalt(IID) bromide p-hydroxo-,.-imido-bis (bis (ethy- Ienediamine)cobalt(IIl) etrakis (,,-acetato-O : ;O') bis {(pyridine)copper(D)] hexakis (),-acetato-~O : ¢O")-py- oxo-trichromium(II) chloride (U{Co(NH,),}, (u-OH), (y-NO,)] Br, [{Cofen),}, {y-NH) (y-OH)}* ({(Cu(CH.N)}, (u-C,H,O,),] [Cry (-CH,COO), (,-O)] Cl Dinuclear complexes with metal-metal bonds : Symmetrical dinuclear complexes : For formula writing we may start with the metal with numerical subscript inside the square bracket. The ligands with appropriate numerical subscript come next being followed by the square bracket. Alternatively we may start with the ligands attached to one metal, then the symbols of the two metals and finally the ligands attached to the second metal. While naming two procedures are adopted : (1) the ligands with appropriate multiplicative prefix are first written followed by the name of the central atom with the affix di-. Thereafter the symbols of the central atom in italics separated by a long dash are enclosed in parentheses. (2) The name of half the molecule is enclosed in parentheses with the prefix ‘bis’. The rest is as in (1). Examples 1. [Mn,(CO),,] or [(CO), MnMn (CO),] decacarbonyldimanganese(Mn—Mn) bis(pentacarbonyl manganese) (Mn—Mn) 2. [Pb,(C)H,)g] or [(C,H,),PbPb (C,H,),] hexaethyldilead(Pb-Pb) or bis(triethyllead)(Ph—-Pb) Non-symmetrical dinuclear complexes : These may be of two types : (1) those having identical central atoms and (2) those with different central atoms. For the complexes of type (1) the central atom with greater number of ligands is numbered 1 and the other numbered 2. If both have the same number of ligands, that central atom which has the greater number In. Ch. I-38298 INORGANIC CHEMISTRY of alphabetically Preferred ligands is numbered 1, For the complexes belonging to type (2) ee ois assigned to the more metallic central atom in the formula even though they are in the alphabetical’ order in the name. While naming the complexes following order should be obeyed : names of the ligand, uber marking the central. atom, the Greek letter kappa with a right superscript denoting the number of such ligands and the italic capital symbol of the ligating atom in alphabetical order. For type (2) the prefix di-is added to the element name and for (2) the element names in alphabetical order. Finally the element symbols are given in italics with a long dash in between and enclosed in parentheses. Examples 2 1 1 2 [Co(CO),Re(CO),] or [(CO),Re~Co (CO),] nonacarbonyl-I ,5C, 2,4C-cobaltrhenium (Co-Re) 1 2 [[rCl, (CO) {P(C,H,),},] (HgC)] carbonyl-I x?C-trichloro-1 Cl, 2yCl- bis (triphenylphosphine-1 ,P) iridiummercury (Hg-Ir) 10.6.4. Writing the formulae and names of chelate complexes’: Formula for chelate complexes have the same order of symbols and the same conventions in respect of ligands, ionic charges and oxidation numbers as for mononuclear complexes cited above. Additionally care should be taken to specify, wherever possible and necessary, atoms of -the chelating ligands ligated to the central atom. Lower case letters should be used for all abbreviations, except for certain hydrocarbons like Me, methyl; Et, ethyl; Ph, phenyl etc. Examples : 1. [NiBr, {(CH,), PCH,CH,P (CH,),)}] dibromofethylenebis (dimethylphosphine)] nickel(I1) dibromo [1, 2-ethanediylbis (dimethylphosphine)-y2P nickel(II) 2. [PtCl, (edta)]** (edta being coordinated through one N and one 0). dichloro [ (1, 2-ethanediyldinitrilo-,N) tetraacetato-,O] platinate(II) dichloro (1, 2-ethylenediaminetetraacetato ,N, ,«O)platinate (4—) 3. [Co(edta) (H,O)]- (edta being coordinated through both N and three O) aqua (1, 2 ethanediyldinitrilo-,NN’) (tetraacetato-,30, O", O'") cobaltate(1—) aqua (ethylenediaminetetraacetato ;N, N'-°O, O", O'") cobaltate(II1) 10.6.5. Illustrations : Given below is a long list of examples of chemical formulae of -COordination complexes alongwith their naming as per recommendations (1990) of IUPAC. 1. [Co(NO;),(NH,),] triamminetrinitrito-~ N-cobalt(IID) | 2. [Co(N,)(NH,),]SO, pentaammineazidocobalt(III) sulphate 3. [Co(ONO)(NH,),]SO,COORDINATION CHEMISTRY 299 pentaamminenitrito-;O-cobalt(II1) sulphate 4, [Co(NH,)g] (Cr(CN)g] hexaamminecobalt(II1) hexacyanochromate(II1) 5. [CoCI(NH,),]* pentaamminechlorocobalt(2+) ion V pentaamminechlorocobalt(IIl) ion 6, [Co(NH,),JCISO, hexaamminecobalt(I1I) chloride sulphate 7. [CoCKNO,)(NH,),1C1 tetraamminechloronitrito- N-cobalt(II) chloride 8. [Co(H,0), (NH,),ICl, tetraamminediaquacobalt({I1) chloride 9. [Co(en),(bpy)}* bipyridinebis(ethylenediamine)cobalt(II1) ion bipyridinebis(cthylenediamine)cobalt(3") 10. Na[{Co(CO),] sodium tetracarbonylcobaltate(—1) [CoBr(NCS) (en),]* bromobis(ethylenediamine)thiocyanato- N-cobalt(II1) bromobis(ethylenediamine)thiocyanato- , N-cobalt(1+) 12. [CoCKNO,)(en),]” chlorobis(ethylenediamine) nitrito- ; N-cobalt(I1) 13. [Co(SO,)(NH,).] [Zn(OH),1 pentaamminesulphatocobalt(II1) tetrahydroxozincate(II) 14, [CoCI,(en),]NO, dichlorobis(ethylenediammine)cobalt(I11) nitrate 15, [Co(CO,)(NH,)4]s [Fe(CN),1 tetraamminecarbonatocobalt(II1) hexacyanoferrate(IIT) 16. [Cr(NCS),(NH,),1" diamminetetrathiocyanato- ; N-chromate(1-) diamminetetrathiocyanato- N-chromate(II) 17. K{CrF,O] potassium tetrafluorooxochromate(V) . NH, [Cr(NCS), (NH,),] ammonium diamminetetrathiocyanato- x N-chromate(II1) . [Cr(C,0,) (en),] [Cr-(C,0,), (H,O),] bis(ethylenediamine) (oxalato)(chromium) (IIT) diaquabis(oxalato) chromate(IIL) 20. [CuC1,(CH,NH,),] 2 “Ss300 2i x 8 23, 24. 25. 20 2 28. 29, 3 3 3: 8 3: 3. 35. 36. 3 3 % al s Ss ~ Fa s INORGANIC CHEMISTRY dichlorobis(methylamine)copper(II) - [Cu(NH,),] [CuCl,] tetraamminecopper({I) tetrachlorocuprate({1) . [Fe(H,0),)* hexaaquairon(III) / hexaaquairon(3+) [Fe(CO),]> tetracarbonylferrate(II) / tetracarbonylferrate(2-) K,[Fe(CN),] tetrapotassium hexacyanoferrate(I1) potassium hexacyanoferrate(I) potassium hexacyanoferrate (4) [Fe(CO),} pentacarbonyliron(0) H,[Fe(CN),] tetrahydrogen hexacyanoferrate(II) [Ru(NH,),(N,)] Cl, pentaammine(dinitrogen)ruthenium(I1) chloride K,[Ni(CN),] potassium tetracyanonickelate(0) K,[NiF,) potassium hexafluoronickelate(IV) (NiPCl,),] tetrakis(trichlorophosphine)nickel(0) - LifAIH,] Jithium tetrahydridoaluminate(II) . [BCIH, dichlorodihydroborate(1-) H[B(C,H,),] hydrogen tetraphenylborate(—) Na[B(NO,),] sodium tetranitratoborate(1—) / sodium tetranitratoborate(II]) Ba[BrF,], barium tetrafluorbromate(III) / barium tetrafluorobromate(1—) (PET hexafluorophosphate (1-) / hexafluorophosphate(V) K[PtCl,(C,H,)] potassium trichloro(ethylene)platinate(II) . [PCL (CH) trichloro(ethylene)platinate(II) ionCOORDINATION CHEMISTRY trichloro(ethylene)platinate(1—) 39. [PtCI(NH,CH,) (NH,),] Cl diamminechloro(methylamine)platinum({) chloride 40. K,[Pt(NO,),] Potassium tetranitrito-. N-platinate(2-) 41. (HE difluorohydrogenate(1-) 42. K[OsCI,N] potassium pentachloronitridoosmate(2~) - [WF,N(CH,),] Pentafluoro(dimethylamido)tungsten(V1) (GeF,{N(CH,),}] tetrafluoro(trimethylamine)germanium(IV) . [Hg(C.H,) (CHCI,)] ichloromethyl)(phenyl)mercury(II) . [U(C,H,0,),0,] bis(acetylacetonato)dioxouranium(VI) 2 [V(C,H,0,),0 (C5H,N)] / [VO.(C,H,0,), (CH,N)] bis(acetylacetonato)oxo(pyridine)vanadium(IV) bis(acetylacetonato)(pyridine)oxovanadium(LV) . (NH,), [V(C,0,),0] / (NH), [VO(C,O,),] ammonium bis(oxalato)oxovanadate(IV) . [ReC(CO),(py),] tricarbonylchlorobis(pyridine)rhenium(1) [Re(CO),(py),] _ tricarbonylbis(pyridine)rhenium(0) . Na, [Ag(S,0,),] sodium bis(thiosulphato)argentate(I) sodium bis(thiosulphato) argentate (3-) [Au(C,H,),(en)] Br diethyl(ethylenediamine)gold(III) bromide - K[Os(N)O,] potassium nitridotrioxoosmate(1—) potassium nitridotrioxoosmate(VIID) 54. y.-amido-,,-hydroxobis(tetramminecobalt (IID) chloride. 301 4 a 4 > 4 a 4 a 4 3 4 3% 4 Ny 5 S 5 3: Nn 5: o OH. (HN). COT CoN, .ICL ce NH;304 INORGANIC CHEMISTRY 10.9. TYPES OF ISOMERISM | ‘Two or more chemical compounds with same empirical composition but with different properties are called isomers. ‘Isos’ means ‘the same" and ‘meros’ means ‘the parts’. Isomers have the same number of same Kind of atoms but these are bonded in different manner. The phenomenon that gives rise to isomets is called isomerism. Isomerism in carbon compounds is restricted mainly to structural isomerism (ortho-, meta- and para-substituted compounds), geometrical isomerism (e.g. maleic and fumaric acids), optical isomerism (d-and /-tartaric acids) and conformation/rotational isomerism (chair and boat forms of cyclohexane and its derivatives). In the area of coordination complexes isomerism arises out of varied chemical linkages and complexity in stereochemical relationships, and are more diverse than isomerism in organic compounds. 10.9.1. Tonisation Isomerism : Ionisation isomers are produced by an interchange of position of coordinating anions (ligands) inside the complex zone and anions outside the complex zone. They yield therefore different ions in solution. Examples are : [CoBr(NH,),]SO, and [CoSO,(NH,),]Br [CoCI,(NH,),JNO, and [CoCI(NO,)(NH,),]CI [PtCL(NH,),JBr, and _—_[PtBr,(NH,),]Cl, Tonisation isomers are readily detected by quick determination of conductance in solution, and by chemical tests. Thus pentaamminebromocobalt(II1) sulphate will react with barium chloride to precipitate immediately barium sulphate but will not immediately precipitate silver bromide with silver nitrate. The conductance will be close to that of a bibivalent compound. On the other hand, pentaamminesulphatocobalt(III) bromide provides no immediate precipitate of barium sulphate on the addition of barium chloride. Besides, the complex ion will behave as a uni- univalent ion. Whereas in [CoBr(NH,),]SO, bromide ion is a ligand, in (CoSO,(NH,),] Br it is not. 10.9.2. Hydrate Isomerism : This isomerism arises out of different disposition of aqua molecules inside and outside the first sphere of attraction. Unfortunately the examples that have so far appeared in texts are not happy ones and can be considered merely as further examples of ionisation isomers only. The examples cited are those of chromium (III) chloride hydrates, The three hydrates are : (1) violet [Cr(H,O),]CI, which does not lose water over conc. H,SO, and the entire chloride is precipitated by silver ion; (2) blue-green [CrCI(H,0),]CI,-H,O loses one H,O over cone. H,SO, and two chloride ions are readily precipitated as AgCl. (3) [CrCL,(H,0),|C1.2H,0, dark green, loses two HO over conc. H,SO, and only one chloride gets immediately precipitated by silver nitrate. (1), (2) and (3) behave (on immediate conductance measurements) as tri-uni, bi-uni and uni-univalent electrolytes. ‘The three isomers can be separated from commercial CrCl,.6H,O via adsorption on cation exchange resin being followed by elution with dilute HCIO,. The major isomer in, the commercial product is trans [CrC1,(H,0),JCL.2H,O. Further examples covering both ionisation and hydrate isomerism are : . [Cr (en),(H,0),] Br, and [CrBr(en),(H,0)] Br,.H,O (CoCIH,O)(NH,),] G, —and-_—_[CoCI,(NHL),] CLH,OCOORDINATION CHEMISTRY aaa ‘A better example of hydrate isomerism is found in cream and olive green coloured varieties of bis(quinaldinato)oxovanadium(1V) : [VOQ,H,0] [VOQ,] H,O [HQ = quinaldinic = (CL) coon } olive green cream acid N ‘Thermal decomposition curves indicated that the olive-green variety lost the coordinated water around 170—180°C whereas the cream coloured isomer lost the water below 100°C. ‘The cream coloured variety was obtained by the reaction ‘of VOSO, and QH in water, and the olive green variety in methanol. 109.3. Coordination Isomerism : Such isomeric forms are exhibited when both the cation and the anion of a salt are complexes and there occurs a redistribution of the ligands between the two coordination zones : [Co[NH,)g] [Cr(on),] and [Cr(NH,),] [Color] [Co(NH,),] [CH(CN),] and [Cr(NH,)g] [Co(CN)g] [Cu(NH,),] [PtCh,] and [Pt(NH,),) [CuCl] The same metal ion may also function as a central unit in both cation and anion eg : (Cr(NH,)¢] [Cr(SCN),] and [Cr(SCN),(NH,),] (Cx(SCN), (NH),I- 10.9.4. Ligand Isomerism : [f two ligands are isomeric t0 each other their complexes with similar composition necessarily become isomeric. Thus trimethylenediamine NH, (CH,);—NH, (tn) and propylenediamine NH,—CH,—CH(NH,)—CH, (pn) form isomeric complexes, [CoCl,(pn) JCI and [CoCl, (tn),] Cl. Picolinic acid (pyridine o-carboxylic acid, Hpic) and nicotinic acid (pyridine p-carboxylic acid, Hnic) form, for example, isomeric inner metallic complexes with bivalent silver [Ag(pic),] and [Ag(nic),]- 10.9.5. Linkage Isomerism : Some coordinating groups possess two different donor atoms, Linkage isomerism arises out of two modes of attachment of the ligand to the central ‘metal ion through wo different donor atons, This NO, group may link via nitrogen (nitroisomer) or via oxygen (nitrito), e.g. : yellow brown pentaamminenitrocobalt(IMl) chloride [Co(NO,)(NH,),ICl, and the red _pentaamminenitritocobalt(II) chloride [Co(ONO)(NH,),]Cl,. Two isomeric forms involving the thiocyanate ion are also known : [Pd(SCN),(bpy)] and [PA(NCS),(bpy)]- K,[Pa(SCN),] + bpy —="S > (PAISCN), (bpy)] + 2KSCN orange yellow | solid heated to, 150°C + K,[PA(SCN),] + bpy —28°C_, [PACVCS),(bpy)] + 2KSCN light yellow Besides NO,” and SCN groups there are other ligands which are capable of playing double role, such as CN"; S,0,", (NH,),CO, (CH,),SO. However, actual isomeric compounds have not yet been fully established with any of these ligands. Pentaamminethiosulphatocobalt(IIT) salts obtained by Ray as purple crystals by air oxidation of cobalt(II) salts, thiosulphate and In. Ch. 1—3906 INORGANIC CHEMISTRY ammonia is now believed to be a mixture of 90% [(H,N),t [GN)CoS—SO,)" ar evidenced frome infrared studies (also see. Part I). Linkage isomers are generally identified by their distinguishing vibrational spectra ang electronic spectra. ‘The tess stable one of 2 pair of iekage tsomen often reverts to the more stable form, The leas stable formes hkrly to exist at bow aed with those metal ions which are known to form Lunctically inert compicacs (Co™, PY", AY), The above mentioned ligamts with to ponsshic sates for attachment to the metal ioa are called ambsdend (or sebadcrestc) lepamts Ht may be mentioned thst the okicr mammag of thi ioenctiem as aah jsamerism has now been replaced by the appropriate Lin see teencriten 10.9.4 Stervedsemmertum : This i1.6 form of itomerioo in which tw compteses of identical Airis rphare comguststacm bier om the tehative protitiome of the conondinating grams Such Gomart are ans termed coometiical isomer fot there incomend give tise to diferent rrangemandy of cium s grivert within the wume overall pecometry of the comptes teme Somme (mot all) of the promutsical nomen may develop optical actinty dur to distinct pone supertesponatlic metroy image form. That optical tours musy bie énalated froen a particular Sromctrical tomer wich in capable of cristing in ron. amprrieqonable mirror tmepe form Geometrical leemertum : Of the thene important and niche comenon sternochreitncs (etrabindeal, pile amd cctuahodral) prometrical iomertum camece attue ta 2 tetrabndeal serecterr. tecane en 3 teteahodron cach fegamd ie erpacicttane from dhe other theme sed at! bromd aemplios arr the ame (109°) Wet tm scgaare Plunae and conMndes Cormpieves advandae examples ate Laon, te squuce pune MAH, type complcucs the following gromwtrical homers ate observed © clvplaam vtractere cbexint when the two A gram of the two BL grep occupy neat and 10%, A NA A A 2p ard wu 0% Ne 0” No a’ “a a” Na rt mare ghee enperienponadie eevroe keag Suitter mage wapersmpenabte after fo cpeacad sctenity Fotaciom, mo optical activity neighbour Position, and a tnseplanse frometty when they oxxupy the distant most Ponitions. (PCI, (NH),),] caus in tao tomers fort Complcact of the type MA.BC can alto exist in two poometrical iscencra It fot pontible to obeais optical isomer of a square Planar complex with moncdcntate ligands, stutever the composition ix MA,B, > MA,BCOORDINATION CHEMISTRY 307 In octahedral stereochemistry cases of geometrical isomerism occur. When one of the A groups of [MA,] is replaced by a B group giving [MA,B] no geometrical isomerism can — wis . a = NH, cl NH, NH Co Co Nu — cl nu NH, [_ VIOLET NH, GREEN a (10-1X) (10-X) Cis-tetraamminedichloro Trans-tetraamminedichloro cobalt (III) (violet) cobalt (III) (green) occur. This is because whichever position of the octahedron is occupied by B it will be at 90° from four M—A groups and at 180° from the fifth M—A group. When a further replacement is made leading to [MA,B,] there occur two geometrical isomers, one cis and the other srans eg: (CoCI,(NH,), JCI. (10-IX and 10-X). A very large number of complexes of this type are known eg : cis{Co(NO,),(NH,),JX (yellow-brown) and trans [Co(NO,),(NH,),] X (yellow); cis [CoCI,(en),] Cl (purple) and trans [CoCl,(en),JCl (green); cis [Cr(NCS),(en),]NCS (orange red) and trans [Cr(NCS),(en),]NCS (yellow orange). Complexes of the type [MA,B,], [Co(NO,),(NH,),], [Co(NH,—CH,—CO0),} are also capable of exhibiting two geometrical isomers. Recall an octahedron has eight triangular faces A B B. A B, A B A A A B B (10 - Xl) (10 - XII) fac MAB, mer MAB, fac [ Co(NO,),(NH,),} mer { Co(NO,),(NH,)] and six vertices. In one geometrical isomer three A groups occupy the three corners (vertices) of one triangular face while the three B groups take up the remaining three vertices of another triangular face. This isomer is called facial or fac isomer. The second isomer has three A groups in one plane and the three B groups in a perpendicular plane. The A and the B groups lie along the meridian of a sphere (imagine our earth). Hence this isomer is named meridional or mer isomer. In the facial isomer the three A groups are cis to each other—the same is also true for the three B groups. In the meridional isomer two of the three A groups as also two of the three B groups are trans to each other. A tris308 INORGANIC CHEMISTRY i following two geometrical isomers, namely (didentate) chelate, [Co(NH,CH,COO),], has the fo! , facial and meridional (10-XIII ; 10-XIV). [ItCl, (PMe,),] has also two geometrical isomers : facial and meridional. A fy £ © [A NL | (10-XiN1) (10-xIV) facial meridional Complexes of the type [M(AA) B,C,] (eg. (CoCl, (NH,), (en)J*) admit of three geometri- cal isomers : (1) trans B, B; cis C, C (2) cis B, B: cis C, C and (3) rans C, C; cis B, B. Isomers (1) and (3) are symmetrical whereas isomer (2) is not. Determination of configuration of geometrical isomers : With the present state of knowledge about complexes, there are in fact several techniques by which configuration of geometrical isomers can be determined. The simplest technique is to study the action of didentate chelating ligands on the geometrical isomers, where isomerism is due to different disposition of two monodentate groups. Didentate ligands are capable of enforcing a S-membered or a 6-membered ring at the metal ions, and with this ring size it is possible only to span two nearest positions around the metal ion. Examination of models shows that usual didentate chelating ligands cannot span the furthest points, that is, trans positions. N [CoCO,(NH,),]" ee {CoCl(H,0)(NH,),* dil. | 1,80, cone. | H,S0, + HCI [Co(H,0),(NH,),]** [CoCl,(NH,),]* aq. | NOH green 3 [Co(OF)(H,O\(NH,),P* TRANS Series g heat J 100°C 5 on = INH), Com = Gg (NH,),)"* | on cone. ¥ HCI (12°C) > [CoCl(NH,),}* + {Co(H,0),(NH,),]* purple purpleCOORDINATION CHEMISTRY 10.11.1. Valence Bond Theor: and certain ligands is assumed t 317 at The formation of a complex between a given metal ion ‘0 follow the course enunciated below : (@ The metal M first loses requisite n: fe © number Ses requisite number of electrons to for i - rm the ion, thi of electrons thus lost being the valence of the resulting cation. " (6) The metal ion will make available number, for the formation of cov: certain definite stereochemistry, such stereochemical directions. : number of orbitals, equal to its coordination ‘lent bonds with the ligands. Since metal complexes attain feel ion orbitals must be so disposed that they point to letal ion orbitals (s, p or d) cannot ger i eae Is (s, generate on their own aa eee as tetrahedral, octahedral, square pyramidal and so on. secling eee invoked the hybridisation of suitable pure metal ion orbitals so that the 7 reochemistry is attained (Chapter 5). The types of hybridisation occurring in the irst row transition metal compounds are the following : 3 (gn? 4s 4p" (sp) : 4 coordinate, tetrahedral > 3dx*— y? 4s Ap*xly (dsp?) : 4 coordinate, square planar > 3dz"4sp*(dsp*) : 5 coordinate, trigonal bipyramidal 30-2Ix*y*4s4p°(’sp) >: 5 coordinate, square pyramidal 4s4p'aP7x7- y(sp'?) 6 coordinate, octahedral (outer orbital) 3d°2Ix?-y74s4p (asp) : 6 coordinate, octahedral (inner orbital) > (c) The hybridised orbitals are so directed as to facilitate their occupation by electron pairs coming from the donor ligands. Effectively there occurs an overlap of a filled ligand orbital and an empty hybrid orbital of the metal ion, (MeL). (@ The non-bonding electrons of the metal ion are then reorganised *o occupy the remaining metal orbitals (pure d, s or p as the case may be). The regrouping of electrons is achieved obeying Hund’s rule, that is, with maximum possible unpaired spins. (e) In addition to the g-bond a q-bond may be formed by overlap of a filled, suitable metal d-orbital with a vacant ligand orbital : M @ L. Such double bonding usually occurs in complexes of metal ions in low oxidation state (Chapter 24). ‘A given stereochemistry involves a particular hybridisation utilising certain d-orbitals. The orbitals available for non-bonding electrons are therefore the dictates of the stereochemistry. ‘Again the number of d-orbital electrons remaining indicates the oxidation state of the central ion. These points are illustrated in Chart 10-IV below with respect to iron (III). 3d 4s 4p 4d a ev [Tt] t {ttt 5 we (ttt) fe) belal= ~ welthtitit it! [=| ll=l=| [|= rs colt t=pe] bel babel , Chart, 10.1V, Valence Bond Theory as applied to iron (III). (7 or 4) represents the spin of a non-bonding electron, X denotes a bonding ligand electron and, n, number of unpaired electrons. Electrons in an incompletely filled shell give rise to a resultant spin angular momentum,318 INORGANIC CHEMISTRY The orbital moment is considered largely quenched by the surrounding ligands. Thus ignoring orbital contributions, the magnetic moment of a complex will depend on the actual number of unpaired electrons (Table 10.3). st Transition Series Moments of Complexes of Configu- Example Stereochemistry Hybridisa- Umpaired =, (B.M.) ration tion electrons __spin only Exptl. @ Kir) Octahedral asp 1 173 ‘1.70 [vO(acae),} Square pyramidal dsp” 1 1.73 1.70 & NH, (SO,), Octahedral bsp 2 2.83 2.80 12H,0 @ __(C(NH), Br, Octahedral asp 3 3883.77 a [Cr(H,0),]SO, Octahedral spd 4 4.90 4.80 & Na,[FeF,] Octahedral spa Ss 5.92 5.85 K,[Fe(CN),] Octahedral dsp? 1 1.73 225 & [Co(NH,),ICl, Octahedral d'sp* 0 0 0 [Et,N],[FeCl,] Tetrahedral sp 4 4.90 5.40 d’ Cs,[CoCl,} ‘Tetrahedral sp 3 3.88 4.60 K,Ba[Co(NO,),]__ Octahedral és 1 1.73 1.88 a [NiQNH,),ICI, Octahedral spd" B 2.83 3.32 (NiHbig), ICL, Square planar dsp? 0 0 oO @_ (Culbig), Cl, Square planar dsp? 1 1.73 1.79 In practice the magnetic susceptibility of a complex is determined whence the magnetic moment and the number of unpaired electrons are deduced (Chapter 11). The number of Unpairéd electrons is then checked with the number of unpaired electrons permitted by the different stercochemistries. Table 10.4 shows that although the correspondence between the number of unpaired electrons and the stereochemistry, bond type and oxidation state is not always unambiguous, a knowledge of the magnetic moment alongwith other chemical evidences may turn out to be of great practical value in understanding the nature of the complex. Outer orbital octahedral complexes are given by weak ligands while the inner orbital octahedral complexes are formed by strong ligands, Table 10.4, : Values of n (Unpaired Electrons) for different Stereochemistries and Configurations Number of d electrons 1203 4 5 6 7 8 9 Octahedral, outer orbital sp’ ——sdT~SC*«itSCS*«S 5 4 3 2 1 Octahedral, inner orbital dsp” 1-232 OT Tetrahedral, sp> 1 2,053 4 5 4 3 2 «271 Square planar dsp* 1 2,3 +4 #3 2 1 0 JCOORDINATION CHEMISTRY 321 10.11.2. Crystal Field Theory + A metal ion in a complex is interaction between a metal ion and the ligands is known as crystal field theory because it was originally discussed for ions in a crystal lattice (effect of chloride ions on K* in KCI lattice). 11 is important to record that the crystal field theory considers the ligand atoms as point charges or point dipoles, and does not consider any overlap between ligand orbitals» and metal ion orbitals. Crystal field theory examines the energetics of the d-orbitals, in particular, in given geometries. Placement of the donor atoms different stereochemical positions will perturb the d-orbitals of the metal ions to different extent. Several stereochemistries are now chosen for discussion, Octahedral Complexes : We recall that there exists only one s orbital (with a given 1) which is spherically symmetrical. ‘Therefore on applying a field of six negative charges located at octahedral points the energy of the s orbital will be raised compared to the field free situation. The three p orbitals have their lobes concentrated on either side of the nucleus along x(p,), y(p,) and z(p,) axes. Each of the p orbitals faces directly two ligands along the coordinate axes and since all the three p orbitals are also themselves degenerate, all p orbitals will be raised in their energy but there will be no separation (splitting) among the orbitals themselves (Fig. 10.6). It is important to note that instead of an octahedron if we take a square plane the p, and p, orbitals will face the four ligands but the p, orbital will face L none. Hence in a square planar environment there will Fig. 10.6. The three p orbitals in an octahedral surrounding of six ligands occur a splitting of p orbitals into two sets, higher one having p, and p, and the lower one having the p, orbital. The five d-orbitals are degenerate in a field-free ion. The five d-orbitals are not all alike. The three orbitals d,,, d_, d,_ have their four lobes concentrated in between the coordinate : for example d., orbital has its four lobes between the « and y axes. The ds» orbital axes has its lobes along the x and the y axes. The d_; orbital has two lobes along the z-axis and a concentric lobe all around the nucleus in the xy plane. When such a set of five d-orbitals is subjected to an octahedral crystal field the d-orbitals can no longer remain degenerate. Their overall energy is raised from that in the absence of any field, and they undergo splitting, Since the dy and the d_; orbitals are face to face with the six ligands they are now raised in energy whereas the d,,, d,. d,, orbitals are lowered in energy. The splitting is so adjusted that the centre of gravity (baricentre) of the five degenerate d-orbitals in a uniformly smeared out field of six negative charges is maintained (Fig. 10.7). The difference in energy of the orbitals d,,, d. d,. (tog Set) and the d.2—,2, d_, orbitals (e, set) is given as symbol 4 or 10 Dq, this difference being a measure of the strength of the crystal field operating on the metal ion. 4 or 10 Dg, is the separation between the 1,,(d,,) and the e, (d 7) and is known as the crystal field splitting parameter or as the crystal field strength. In. Ch. 141and the 1, set will go down by say y. If £ is the energy of each electron prior to splitting + if follows that m2 INORGANIC ‘CHEMISTRY Since the pre-splitting centre of gravity is to be maintained the ¢, set will go up by say i 1OE = {E+ 2) + 6 (E-y) $0 that 2x © 3y | 4 = 6Dq and y 2a 4Dq. Thus each f,, electron 3 Also since 4 y= q we have 1 3. 3 etn is stabilised by —4Dy and cach ¢, clectron is destabilined by + 6 Dg. with respect to the unypl baricentre energy (a) (b) 3) Fig. 107. Splitting of the d-cebitals in an cetahateat field (3) free ion (b) smesral out fickd (c) api The symbol't,, indicates a triplet orbital degenearcy ie. a three fold degeneracy comprising in this case the orbitals d,, dd, The symbol ¢ speaks of a nea fold degeneracy comprising in the present case the orbitals d+ + and d_; The subscript g stands for *perade’ of even. In an octahedral complethe metal ion is at the centre of inversion, and with respect {0 this inversion centre the d-orbitals maintain the same sign of theit wave function (i. same sign on the lobes) on inversion 1.23). But a tetrahedron has no centre of inversion and hence the g subscripts are dropped in a tetrahedral crystal field. The subscript 2 has its origin in group theory and refers to some particular symmetry operations. It may be noted that p orbitals change their sign on inversion in octahedral field and hence they are ‘ungerade’ or uneven and carry the symbol ‘u" The splitting of the d-orbitals has significant consequences on the colour and magnetic Properties of transition metal complexes. The d-orbital electrons in a crystal field will now have a choice of cither going to the 1, OF to the ¢,. There will be two opposing forces in this regard. The crystal field splitting (10 Dq) will tend to force as many electrons into the more stable 1,, sct whereas the electron pairing energy (energy required to cause pairingCOORDINATION CHEMISTRY 333 eof = SS hme ee a a a ty { == at me ae 7 ae ru ¢ ¢ : ? ge Fig. 10.8. The d', &, d?, a, d? and d configurations in octahedral field, of two electrons in the same orbital) P, will counter any move to induce spin pairing in-the 4,, set. In practice the relative magnitude of 4 (10 Dg), which for a given metal ion varies from ligand to ligand, and P * (which is fixed for a metal ion but varies from metal ion to metal ion) will decide whether or not spin pairing will take place, that is, will decide whether or not the ¢,, set will be occupied in preference to the e, set. The d', a’, d? systems will occupy the 1,, set with their electron spiris’parallel and this will be true irrespective of the strength of the crystal field. Again d°, d’ and d'° systems can be arranged in one way only with two, one and no unpaired electrons in the e, set (Fig. 10.8). : HIGH SPIN LOW SPIN HIGH SPIN LOW SPIN STATE STATE - Suir STATE {4 ipa t 4 — } 4 H t t dat, d= th, E=-2A+2P t tt t e pune tO pet ea, mm WE + T 4 d=tees =-2 Anse =-Bas+ap Es-ZA+2P Fig.’ 10.9. High-spin ‘and low-spin configurations and their energies for d', a’, d° and d' ions in octahedral field. * Pairing energy P is dictated by the principal quantum number n_of the clectrons. The. 4d. and the 5d orbitals extend more in space i.e. they are larger in size than the 3d orbitals. Therefore spin pairing i.e. double occupation of an orbital involves much-less interelectronic repulsion in 4d and Sd orbitals than in 3d orbitals. Thus P falls in the order : 5d < 4d < 3¢.324 INORGANIC CHEMISTRY But in the cases of d', d?, d® and d’ high-spin (maximum possible number of unpaired electrons) and /ow-spin (minimum possible number of unpaired electrons) may arise, According to Hund’s rule of maximum spin multiplicity (Chapter 11) » unpaired electrons will tend to occupy n orbitals to give rise to n unpaired spins. Thus pairing of spin is not normally a favourable process and the pairing energy P will have to be used to counter the electrostatic repulsion between two electrons in the same orbital. A d' high-spin system in a crystal field will have three electrons in the f,, set (energy = 3 x (-2/5 a) = -6/5 4 = — 12 Dg) and one in the ¢, set (energy = 3/5 4 = 6 Da), the total energy due to the crystal field being -6/5 4 + 3/5 4 = -3/5 4 = ~6Dg. On the contrary if the crystal field is strong enough to induce spin pairing, all four electrons will remain in the z,, set and taking the pairing energy into consideration the energy is 4 x (-2/5 a) + P =-8/S a+ P =-16 Dg + P. The energies of the a’, d°, d® and d’ systems in the two states are enumerated in (Fig. 10.9). These energies are with respect to the unsplit baricentre whose energy is conveniently taken to be zero. A perusal of the above energy values shows that on putting 4 = P, each of the a’, d’, and d’ systems provides the same energy for both the spin-states. With 4 > P the energy of the Iow spin state becomes smaller than that of the high spin state i.e. the low spin state becomes the preferred state. With < P the reverse holds. This condition for high spin d° is demonstrated below : For the high spin state , < Pie. P> a soP=yj+xa Energy of the high spin state UW yt+P=UWatatxa (h,¢,7) =3S atxa Energy of the low spin state. = -12/3. 4 + 3P = -12/5 4 +3, +3x.q 2," = 35 A +3xK4 Thus the energy of the low spin state is greater than the energy of the high spin state. Therefore the high spin state will result when < P and conversely the low spin state will be the ground state when > P. These two conditions are apparent on mere inspection of the energies of the two spin states of dl’ (Fig. 10.9). When 4 > P, the energy of the low spin state is negative while the energy of the high spin state is zero. Again when a < P, the energy of the low spin state is positive while the energy of the high spin state is still zero. ~ Thus the spin-state of an ion in a complex depends on whether the crystal field splitting is bigger or smaller than the pairing energy. When . is greater than P, electrons will tend to pair their spins and when q is smaller than P, electrons will tend to remain unpaired and will spread out in the d-orbitals. A weak-field (4 < P) produces high-spin complexes while a strong field (4 > P) gives low-spin complexes. Note also that the spin arrangement in weak field complexes is the same as in free ion, although the orbitals are no longer degenerate: For each of d‘, d’, d®, d’ ions there is a critical crystal field strength (4 = P) below which + all ligands will produce high spin state and above which all ligands will produce low spin state. The critical 10Dq (q) is sometimes referred to as the cross-over region (Fig. 10.10). A ligand, whose crystal field strength is close to the critical 10 Dq of a particular d" ion (n = 4, 5, 6, 7), can give rise to spin-state equilibrium. The difference in energies of the two spin-states around the cross-over region is in the range of thermal energy so that spin-INORGANIC CHEMISTRY The value of Dq in an octahedral geometry has been calculated with the following result : a a —— 5 electronic charge magnitude of the charge ‘on the ligand metal—ligand distance mean fourth power radius of the orbit of the d electrons of the central ion i.e. the distance of the electron from the nucleus. However this theoretical expression of Dq has little value as the: terms are difficult, to determine. In reality the value of Dq is evaluated from electronic spectral measurements and Tanabe-Sugario diagrams. Tetrahedral Complexes : A tetrahedron is best conceived by taking the alternate corners of a cube (Fig. 10.11). The centre of the cube is occupied by the metal ion. The dispositions Of the doy dn diy d,2_2 and dy orbitals relative to the tetrahedral points (where we place b D= day Gon he 2 At or 10Dq” a — re e L — : (@) eh). © Fig. 10.11. A tetrahedral ML, complex with the four ligands at the alternate Fig. 10.12. Splitting of the d orbitals in a tetrahedral corners of a cube, the centre of which field. is occupied by the metal ion. @ free ion (b) smeared out field (c) splitting four ligands) show that dy d,, and d,, orbitals lie just half an edge of a face of the cube ‘away from the nearest ligands whereas the d,._|2 and d_» orbitals lic half a diagonal of a face of the cube away. For example, each of the two lobes of the dy orbital along the z-axis will be placed in between the two ligands which occupy diagonally opposite corners . of a face of the cube. Thus in contrast to'the octahedral geometry, in a tetrahedral crystal field d,,, d,., d,, orbitals (s, set) are repelled more and hence destabilised comparéd to the da 2 and a. orbitals (e set) (Fig. 10.12). Further, calculations show that for a given metal ion and’a given set of ligands and the same internuclear distance the tetrahedral splitting At is about 4/9th of the octahedral splitting between the’ same metal ion and the same six ligands. In principle only a’, d', d° and a? may exist in both spin-states. The low spin-state will be favoured only if 4 tis greater than P.A tetrahedtal field produces only a weak splitting. No tetrahedral field is strong enough ‘to initiate spin pairing (4 ris always less than P).COORDINATION CHEMISTRY Splitting of d-orbitals in other Cryst ordering in several other fields. 327 Fields : In Fig. 10.13 we give. the d-orbital Ina square planar geometry the orbi e 2_,2 otbital alone directly faces the four li - ii Lay igands alon, the x and y axes. Bence da _y? orbital is the most destabilised. d,, orbital is’also rey Ned by the ligands but expectedly (0 a lower extent. The d.:and d. orbital cee nari ea vicinity of the ligands. In an ideali iar sterec Beene eee ee ised square planar stereochemistry there are no ligands and thus the d_, orbital is the most stabilised, Splitting i ° I L : sed. Splitting in a square Planar geometry may be likened to the spliting in a tatragonal elongation case (Fig. 10.16), In square pyramidal geometry the ¢ is fa i i i ee ee nety [2_,2 orbital is face to face with the four ligands in ce this is the most destabilised orbital. Next in order wi be the dz, d., and the doublet set d 4, : e the trigonal bipyramidal geometry the da orbital is the most destabilised as it faces wo ligands directly. However the equatorial orbitals dy, d.2_\2 do not face the three ligands — 22 —d22 moe “ — dy aay Oa ons Gye SQUARE PLANAR, SQUARE PYRAMIDAL TRIGONAL BIPYRAMIDAL Fig. 10.13. Splitting of the d-orbitals in square planar, square pyramidal and trigonal bipyramidal crystal fields. (in the xy plane) directly but are cquidisposed with respect to them. The d,. d,, orbitals, as expected, are the most stabilised. In the trigonal bipyramidal stereochemistry the energy difference (g,) between the lowest doublet set (d,,, d,,) and the upper doublet set (d.,, d. 2.2) is usually much smaller than the energy difference (5,) between the set (dz, d.?-*) and the highest energy d? orbital. It has been established that §, can never exceed pairing energy P of a metal ion. But 5, may be less than or greater than P, Therefore in this geometry a’, @,.d° and d* systems will always give the high-spin form while d°, d®, d’ and d* systems may be either high-spin or low-spin. Hence with favourable ligands ic. g,= P we may observe anomalous moments with the d°, d°, d’ and d® systems. In reality, however, for d® nickel (II) alone high spin and low-spin trigonal bipyramidal complexes are known. For all the other i.e. d°, d? and d’ systems all the known trigonal bipyramidal cémplexes are high~ spin. Academically, however, the spin arrangements will be say for d° as : high-spin B= (dy, d,)'s yy do d'2 low-spin d®—: (dy dy)", (dy da_2) da In trigonal planar geometry the two ligands along the z-axis are missing-and hence the d., orbital is now the most stabilised orbital. The energy order will be: daa. d,> diri INORGANIC CHEMISTRY e axi © ligand d..> d.;. Ina linear geometry the z-axis is taken as the unique axis and eee be Is her side of the metal ion. It can be guessed that the energy order 1S + d.> are placed on eit did > deed, Crystal Field Stabilisation Energy : The st of its d-orbitals by a crystal field and pre! ability gained by ad” ion due to the splitting occupation of the lower energy d-orbitaly is called its crystal field stabilisation ene The CFS! PO OES NN bfghe oe ns to form complexes with preferred g es. Thus CFSE of ion ih high-spin octahedral stereochemistry (15,°¢,") is S(-4 Dg) + 2 (+ 6 Da aD Sou ae heed not be considered for calculation of CFSE bn sare useful in interpreting several has to be cor behaviours of transition metals : , 2) variation of ionic radii (3) thermodynamic propertic. palt(HD) (3c) has the highest (1) stercochemic: (1) Ste + Octahedral low-spin col cobalt (IID) complexes ar weak crystal field strength ot High-spin [CoF,]* is the only exception. This is duc to the w : ; ind, which cannot exceed P (P energy of Co™). It is commonly chemi Jow-spin octahedral the Muoride fi “ suse these ions have high CFSE between octahedtal observed that d* and d* ions prefer octahedral = -12 Dg. A perusal of Table 10.5 reveals that the dif | ometries is smaller for cobalt(II) than for nickel(11). The difference in C and wtraheds of ad ion in octahedral and tetrahedral geometry is called octahedral site stabilisation energy Dg = Dq while that for Co(Il) (OSSE). Thus OSSE for Ni(ll) = -12 Dg + 8 Dy + 5.3 Dy = -2.7 Dg. Therefore a tetrahedral cobalt(I) complex has a better prospect to be formed in aqueous medium than a tetrahedral nickel(I]) complex. In fact addition of excess HCI to pale pink [Co(H,0),* changes the colour to blue due to the formation of tetrahedral [CoC1, while similar addition to [Ni(H,O),}°" has no eftect. The smaller the difference in the CFSE’s of the two complexes the greater is the prospect of the reaction to proceed from left to right under the mass action of the weaker chloride ligand : IMCL]? + 61,0 (M(H,O),)°" + 4CL (= -12 Dg) of [Ni(H,O),}* is 10.200 cnr! while that (= -8 Dg) of [NICH |’ is -3254 cm" so that the difference is about -7,000 cnr in favour of the octahedral ayia values of [Co(H,O),}°* (= -8 Dy) and of [CoC] (= -12 Dyn so that the difference is about -3700 em! m The CFS complex. The correspondin are — 7440 env! and -3774 em’! respective favour of the aqua complex. Thus nickel(II) favours octahedrat geometry to tetrahedial geometry much more than does cobalt(II). . of, d* (weak field) and a ions have no CFSE in either octahedral or tetrahedral felt! Hence such ions have no preference for any particular stereochemistry on crystal fick d-Ligaanne grounds, In such cases stereochemistry and coordination number are decided by li repulsion, size of the cation and cation charge. Zine(II) (3d'°) shows a preference for: tetrahedral geometry as this geometry minimises ligand-ligand repulsion by maximising |. M-—L angles to 109°.COORDINATION CHEMISTRY 329 Table 10.5. CFSE in Octahedral and Tetrahedral Fields. at High-spin Low-spin ‘Tetrahedral octahedral octahedral, (Pq) =¢ Dq (Da) (a) ° @ 4 a -6 Dat =-2.7Dq e aA a -12 =53 a _- 2 ‘e120 -8 ae a ~6 ~ 16 -4 a é 0 -20 0 2 é -4 -m4 = oa ? - on -12 =-53 é -12 12 ae 8 e ~6 " -6 : _ “° 0 ° 0 a That CFSE dictates the stereochemical preference of metal ions is also revealed by an inspection of the structures of normal spinels and inverted spinels. Spinels are mixed metal oxides, AB,O,, wheré A is a bivalent and B is a trivalent metal ion. Close packing of oxide ions gives both octahedral and tetrahedral holes (Chapter 6). In normal spinels the bivalent A ions occupy tetrahedral holes while the trivalent B ions occupy the octahedral holes, In inverted spinels the bivalent A ions have moved into the octahedral holes while half of the trivalent B ions have moved out of the octahedral holes to occupy the tetrahedral holes. Normal spinel Inverted spinel A[BBJO, BIAB]O, example : Ma" [Al,""JO, example ; Al" (Ni! Al"YO, (-12 Dg) of nickel(II) in octahedral geometry leads to the inverted spinel structure in NiAI,O,. Aluminium (III) is a 3d” system and has no CFSE. In low-spin octahedral state 3c cobalt(III) has a very high CFSE (-24 Dg) compared to that of high-spin cobalt(II) (-8Dq). Although oxide ion is weak it can still force spin pairing in cobalt(II). Thus Co,O, assumes a normal spinel structure Co" [Co,!"JO,. Iron (IIT) 3a’) in high-spin state has no CFSE but 3d’ iron(II) in high-spin octahedral state = eh Hence a stable system results if Fe,O, assumes an inverted spinel structure Fel" (Fe'Fe''] O, If Fe,O, had assumed a normal spinel structure the stability gain would have been only CFSE of 3d iron(I) in tetrahedral stereochemistry (-6 Dgt = - 3Dq) compared to ~4 Dq of the inverted structure. Similar reasoning would predict a normal spinel structure for Mn,O,. (2) Variation in Ionic radii : The bivalent metal ions in the high-spin metal oxides, MO, are octahedrally surrounded by the weak oxide ions. As we move from Ca”* to Zn’ there is a partial filling up of the 1,, level first, then partial filling of the e, level, then of the 1. Jevel again and so on to give finally £,,°e,' in Zn’. An e, electron repels ligand electrons more than does a 1,, electron. Since an e, electron faces the ligands directly it screens the The very favourable CF:COORDINATION CHinMistry oan for the case of univalent ani i 7 ‘alent anions. One lattice energies are plotted one nee again a double humped curve is obtained when the . Hi umbers © nbse cnerses should increase with dev Hau In the absence of any the lattice attice energies are corrected for the : ‘alula ent Hine is obtained when these compounds (Fig. 10.155), Mleulated from optical data of the MX, It is interesting to note , s that a double ed curve § in units of respective De's) ofwenk eee curve is also produced when the ifiaatianin @otney ae eee s are plotted against atomic number, C manganese(iD, Wr aa oo vanadium) and then drops to zero in h _ Semin down to zero again ett ean 8 A second maximum in d* nickel(II) and ~WVariation of 10 Dg : , 1. Given t eo ano Same rhetal ion and the same geometry it varies from ligand to ligand. This 8 the fo Seaper para Pera 7 field eligi lowing spectrochemical series of ligands in order of increasing crystal . I or bital will be destabilized. Again the d,, and the d,_ orbitals (with z components) will be stabilised while the d,, orbital willbe destabilised (Fig. 10.16). If the e two trans ligands (X) are pulled to infinity an idealised square planar geometry with the d-or- bital ordering da sil d,. #2 p>d2, 4. > d? The reverse will be true if the unpaired electron is in the d.2_,2 orbital. X-ray crystal pructure of CuCl, shows that each copper( tI) is octahedrally surrbuided by six chloride ions but with different copper-chlorine distance (4CI- at 230A and 2CI- at 2.95A) indicating that the unpaired electron is in the d orbital. The electron distribution then is most likely to be : ; d' aip> dato dy? > d2, d,, Hexaaquatitanium(I11), a 3d! system, is susceptible to Jahn-Teller distortion both in the Bround state and the excited state. Instead of a sharp, single absorption band (4, 5 ¢, broad asymmetric band centred around 20,3 LK covering several close transitions eg : (tetragonal elongation; * : deg Gye > dy 2 ide yoy 4.2, etc ; tetragonal compression : dy d ete.) is observed. Two S—O stretches (915 cm*! and 960 cm: ') in the infrared spectrum of {Mn(dmso),] (CIO,), (dmso = dimethylsulfoxide) indicate differently bonded dmeo molecules. (Free dmso has S—O stretch at 1055 cm), High-spin manganese(tI}) (3d*) has an uneven (unsymmetri- cal) distribution of electrons in the ¢, orbitals ' aa, 2° or ds 2°, ds"). Since the intensity of the 915 cm” band is almost double that of the 960 cm” bond, the nuclear Sercening is less in the square plane giving four strongly oxygen-coordinated dmso ligands. It has been concluded that the fourth electron is in the dz orbital. The electron distribution N doa (17 kK) anid d,., dy, dai (20 kk) High-spin [CoF,]* (d°) is Jahn-Teller sensitive in the ground state (é,'¢,°) and more so in the excited stage (t,,e,"). Instead of a single absorption band (57, 5E,; cf. Fig. 11.11.) a split spectrum (bands at ~ 700 nm and ~ 900 nm) is obtained. . Static Jahn-Teller effect results in a permanent distortion of the complex as in copper (II), high-spin manganese(II1), and high-spin cobalt(III), Howevet if the energy difference between the two possible distortions namely : (1) four long equatorial bonds and two short axial bonds and (2) four short equatorial bonds and two long’axial bonds is minimal, the final time average structure may be close to an undistorted octahedron. Such a case is said to arise from dynamic Jahn-Teller distortion and is difficult to establish experimentally. idy oda is:d.'.d,. COORDINATION CHEMISTRY Colour of Coordination Complexe: j sources : (1) d-d transition and (2) A complex absorbs the required red energy for a part ; white it and obviously transmits the remaining comple Teens suromseing oe entary colour is the colour that we observe (see sectign ane "ete ded transition : In its ‘ Serve (see section 23.2.5 Part II). irradiation ‘vith blue, een ee the d! electron of (TIGH,0),)°* is in the , set. On 3 en light of energy equal t ag, Sele :K) of the hexaaquatitani Tey equal to the 4 (493 my = 20,300 em’ mah ees quatitanium(I1) ion, the complex will absorb such energy to allow at of the f4, electron to the excited e. set, When we eae anne a a A ec the complex through transmitted li it appears purple violet. For multielectron system inter: Hw rough transmitted light to many clectronic stat : ‘actions between d-clectrons So that more than one transition is 7 a f nis Quantitative expressions havi connecting e1 335 + The cok lour of complexes ge S originates / charge transfer transition, ele sives rise possible (Chapter 11), pectral transitions on the crystal field model ; field strength and interelectronic repulsion ssions are also available connecting magnetic moments with re not forthcoming from VB theory. _ I can be Buessed that if 1,0 of [Ti(11,0),]"" is substituted by a stronger ligand, say by ‘pyridine giving [Ti(bpy),}", then light of much higher energy (presumably of the violet region ; ~ >, = 400 nm ; ~ 25,000 em! ; ~ 25 kK] will be required to effect the 1, -» ¢, electronic transition “since will now be much higher. The transmitted light will be correspondingly of lower energy and thus the observed colour will be yellow-green ot so. Thus we hi mple crystal field explanation of changing colour of complex with change in the coordinating ligand, assuming however the geometry remains the same, This is, in fact, how the spectrochemical series is obtained. It is also of interest to note that high-spin cobalt(III) (3¢) fluoro complex K,{CoF,] looks blue while all low-spin octahedral complexes of cobalt(III) (eg : [Co(NH,),|Cl,) are yellow to orange. Evidently K,[CoF,] absorbs low energy light and transmits high energy light while the diamagnetic complex does the re been derived for s : nergy of the transition, crystal Furthermore quantitative expres 10 Dg, which connections wer se i.e. absorbs high energy light and transmits low energy light. A simple undergraduate level explanation is that for the high-spin complex a, is less than the pairing energy P of cobalt(III) while for the low: larger than P i.e. a(1.s.) >> a(h.s.). Hence diff electronic transitions in the two spin stat -spin complex a is much rent amounts of energy are required for the , thereby showing varying colours. However this explanation is too simple. For an exact explanation one has t6 make use of the Tanabe-Sugano diagrams of his. d° (Fig. 11.10) and Ls.d° (Fig. 11D). - Charge transfer transition : There are many compounds where either the metal ion is highly oxidising and the ligands are reducing or that the metal ion is highly reducing and the ligands are oxidising. In such cases there occurs transfer of charge (j.e. electron) fom the reducing partner to the oxidising partner. The metal ions may or may C poses electrons. Common examples are intense purple permanganate eee er mercury(I) iodide (Hg(ID), Sd°), yellow chromate (Cr(VI), 3), dark re {Fe(bpy)5] 3c), blood red iron(II) thiocyanate (Fe(III),3d°), ete. : c i e: tronic transitions are : Liecaer chin Nata: mays he ground state and the excited state hav 1. No electronic transition is permitted where-the gro peta ( i i in multiplicities. Only those transitions ai different spin arrangements i.e. have different spinea which a $= 0 ice. transitions are permitted between states of the same spin aoe INORGANIC CHEMISTRY allowed f multiplicity. 2. Only those electronic transitions are allowed for which 4 l= due to Laporte points out that no electronic transitions ‘are allowed for which 4 1= 0 op , 2. Thus tran: ms such as 2s _, 2p, 2s 3p, 3p — 3d are allowed but 3d _, 3d, is = 2s or 3s 5 3d are not. No electronic transition can occur merely because of redistribution of electrons in similar orbitals i.e. d-d transitions (such as t,, > e, transition in {Ti(H,0),)+) are forbidden. For quantum mechanical proof of the two selection rules see Appendix (Par, ID. Although d-d transitions are forbidden yet the coordination complexes of transition metals are coloured. indicates that Laporte selection rule must be relaxed in practice. Whenever a complex lacks a centre of symmetry (e.g. : in a tetrahedron) or that odd vibrational effects lead to a change of bond lengths or bond angles there is scope for mixing of metal 3d-orbitals with metal 4p orbitals or with ligand p-orbitals. The pure d-character of the transition metal <-orbitals is then lost and Laporte selection rule is relaxed. A situation arises so that electron transfer from a d-orbital to a p-orbital can occur, Forbidden transitions have low molar extinction coeficients (10 to 200 J mol” cm’) for six-coordinate complexes and (100—1000 / mol"! cm") for tetrahedral complexes. In a non-centrosymmettic tetrahedral complex the selection rules are relaxed to a greater extent. Charge transfet transitions may be either from ligand to metal (L_—, M) or from metal to ligand (M _, L). In such complexes the ligand p-orbitals and metal d-orbitals interact to form M.O.’s.'When these M.O.’s are arranged in order of their energy it is found that some M.O.’s have predominantly ligand (i.c. p) character while some M.Q’s retain predominantly metal (i.e. d) character. Laporte selection rule allows p > d transition (4 [= +1) so that charge transfer transitions are allowed transitions with high molar extinction coefficients (10 to 10° ! mol"! cm”). Such transitions are usually high energy transitions and occur in the ultraviolet region. Charge transfer complexes appear coloured to our eyes only when the transitions occur in the near visible region or when the transition in the u.y. has a long tail in the visible. Common examples of LM charge transfer are mercury(I) iodide, MnO,, CrO,* etc. Examples of M -, L transition include [Fe(o-phen),]”*, [Fe(bpy),]** ete. While d-d transitions can occur only in complexes of transition metal ions with incomplete d-level, charge transfer can occur in both transition metals with incomplete d-level as also in @’, d'° and non-transition metal compounds, Charge transfer transitions are also known as redox selection rutg transitions. The more reducing M is and the more oxidising L is, the lower is the energy of the ML charge transfer transition. The absorption maximum of [Fe(bpy),]** occurs at 515 my while for (Ni(bpy),}*, [Co(bpy),]°* the maximum is at still lower wavelength range. Iron can be estimated down to 10° g/ml via formation of [Fe(bpy/o-phen),]**. The more oxidising M is and the more reducing L is the lower is the energy of the L _, M charge transfer transition. Thus iron(II) complexes give more L_,M charge transfer transition than does chromium(1D). Most colorimetric analyses of trace quantities of metal ions utilise charg? wansfer transitions. d-d transitions, being of low intensity, are of little value for such purposes: Charge transfer spectra need not be restricted to M., L or L_y M type. A third lessCOORDINATION CHEMISTRY common type is exhibi KPeFe(CN), One of a a So-called Prussian blue or Turnbull's blue precipitate possible from Fe(II) to Fe(tit). ease while the other is Fe(II). Electron transfer is are Fe(II) and thus there exists no sen K,FeFe (CN), is colourless because both the irons below are the salient dtinn are 4 coes £0" any (redox) charge transfer transition. Given S between d-d transitions and charge transfer transitions. 337 d-d transitions charge t itions 1. Presence of d electrons essential. 1 ee - Presence of d electron not essential, It may occur in d, d", d'° systems. . Laporte allowed. . Extremely high molar extinction coefficient (5,000 ~ 35,000 ! mot! cm’). 2. Laporte forbidden. 3. Low molar extinction coefficient 3 (5-200 I mol" cm), 4. No redox system involved, R 4, Redox system essential. 5. Not useful for spectrophotometric 5. Very useful for spectrophotometric estimation of metal ions. estimation of metal ions. Intensities of forbidden and allowed transitions are given below. type of transitions Epuax (I mol” cme) spin forbidden lan Laporte forbidden d — d 10 — 200 (octahedral complexes) * Laporte (partially) allowed d - d 100 - 1000 (tetrahedral complexes) charge transfer (Laporte allowed) 5,000 ~ 35,000 More on the electronic spectra from the viewpoint of the crystal field theory is given in Chapter 11. Passage of Crystal Fields to Ligand Fields : The crystal field model does not consider any covalent interaction between ligands and the metal ion. Experience shows that in several instances overlap of orbitals of the metal ion and the ligands does occur. (1) Ina low-spin hexachloroiridate(IV), [IrCl,]* (f,,°) the single unpaired electron has been shown from an analysis of the electron spin resonance spectrum to be delocalised into the six chloride ions. The multiband ESR spectrum convincingly proves that the iridium unpaired electron is definitely disturbed by the nuclei of the chloride ligands. This is not possible unless there occurs an overlap of metal—ligand orbitals. For details on the ESR evidence see Appendix V, Part II. (2) Intensities of the forbidden d-d transition cannot be adequately explained on the basis of mixing of metal d and metal p/s orbitals alone or on the basis of vibrations of the complex but that metal d-orbitals have to be assumed to be mixed up with ligand p-orbitals also. (3) Crystal field spectral bands cannot be fitted precisely if the interelectronic repulsion of metal electrons in a complex be assumed to be the same as in the free ion (section 11.10.5). Instead, good fits are obtained when the interelectronic repulsion is assumed to be 20-25% Jess in a complex than in a free ion. It follows that the d-orbitals have expanded in size and have penetrated into ligand orbitals so that repulsions between metal d-orbit s becomes smaller than in the free ion. The ligand electrons partially screen the orbital electrons from the nuclear charge of the metal ion. Thus the d-orbital electrons spend sometime on the ligands