harry b. gray and c. j. ballhusa- molecular orbital theory for square planar metal complexes

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260 HARRY . GRAY ND C . J. BALLHAUSEN VOI. 86

[CONTRIBUTIOXROM TH E DEPARTMENTF CHEMISTRY,OLUMBIANIVERSITY, EWYORK2i, N. Y.]

A Molecular Orbital Theory for Square PlanarMetal Complexes

BYHARRY . GRAY ND C. J . BALLHAUSEN'

RECEIVEDUL Y 25 , 1962

The bonding in square plana r metal complexes is described in terms of molewlar orbitals. The relative singleelectron molecular orbital energies are estimated for the different types of n-electron systems. Both th e d-dand charge transfer spectra of th e halide and cyanide complexes of NiZ f, Pd2 +,P t z + nd Au3+ are discussed interms of t he derived molecular orbital scheme. The magnetic susceptibilities of diamagnet ic d8-metal com-plexes are considered, and it is concluded th at a substantial "ring current" exists in KaNi(CN),.H20.

Introduction

The metal ions which form square planar complexeswith simple ligands have a d8-electronic configuration.Th e five degenerate d-orbitals of the uncomplexed metalion split into four different levels in a square planarcomplex. Thu s three orbital parameters are neededto describe the ligand field d-splittings.

Various efforts have been made to evaluate the d-orbital splittings in square planar metal complexes,and to assign the observed d-d spectral Th emost complete calculation has been made for Pt2+ om-plexes by Fenske, Ma rtin and Ruedenberg,B and con-siderable confidence may be placed in their ordering ofthe d-orbitals. However, for other complexes the d-

ordering may depend on the metal ion and the t ype ofligand under consideration.

It is the purpose of thi s paper to develop a molec-ular orbital theory for square planar metal complexes.Both the spectral and magnetic properties of typicalsquare planar complexes of Ni2+,Pd2+,Pt2+ nd A u 3 +will be considered in order to arrive a t consistent valuesfor the molecular orbital energies. However, in con-trast to previous workers, effort will be concentratedon the assignments of the charge transfer bands of rep-resentative square planar complexes.

Molecular Orbitals for Square Planar Complexes.-Figure 1 shows a square planar complex in a coordinatesystem with the central atom a t the origin, and the fourligands along the x- and y-axes. The orbital trans-formation scheme in the DQh symmetry is given inTable I .

TABLE

ORBITAL TRAXSFORMATIONCHEME IN D i h SYMMETRY

Repre- Metal

sentation orbitals Ligand orbitals

a 1 8 d 2 , s l / Z ( Q l + u2 + 6 3 + Q 4 )

azu P Z 1/2(X1,- + T Z Y + T3" + T 4 T )

blZ3 1/2(Ul - UZ + Q 3 - 4 )

bau . . '/Z(Tl. - TTZ" + T3 9 - T 4 " )

a28 . . ' /2 (Tlh + T2h + T 3 h f 4h )

d Z 2 y2

b2g d z , ' /z(Tlh - T z h + T3 h - 4h )

1 1

(1) Permanen t address: Inst itu te for Physical Chemistry, Cniversi ty of

Copenhagen, Denmark.

(2) (a) J. Chatt , G. A . Gamlen and L. E. Orgel, J . Chem . SOC.,86 (1958);(b) G. Maki, J . Chem. P h y s . , 28, 651 (1958); 29, 162 (1958); 29, 1129(1958).

(3 ) C. J. Ballhausen and A . D. Liehr, J . A m . Chem. Soc. , '71, 538 (1959).

(4 ) M. I. Ban, A d a Chin. c a d . Sci. H u n g . , 19, 59 (1959).(5) S. Kida, J. Fuji ta , K. Xakamoto and R. Tsuchida, Bull. Chem . SOC.

(Japan) , 8 1 , 79 (1958).

(6 ) R. F. Fenske, D . S. hlartin and K. Ruedenberg, I n o r g . Chem. , 1,44 1(1962).

(7) J. R. Perumareddi, A . D. Liehr and A. W. Adamson, J . A m . Chein.

Sac., 86, 249 (1963).(8) J. Ferguson, J . C h e m . P h y s . , 3 4, 611 (1961).

The single electron molecular orbitals are approxi-mated as

*(m.o.) = Cx(,)(@)(metal)+ C(L)j@(ligand)

where the C's are subject to the usual normalizationand orthogonality restrictions. The pure u-orbitals are

*[a~,(o)l = cM ( l ) @ [ f id r p l + CMW@(n+ 1)sI +

(1)

CL@[1/2(Ul + Q2 + u3 + Q 4 ) l ( 2 )

Q Z + Q3 - 1 (3 )

*[blg(Q)I = CM*[%dz'- r2 ] + cL@[1/2(Ul -

The form of the a-orbitals depends on the type of ligandunder consideration. Square planar complexes inwhich the ligands themselves have no n-orbital system

(Cl-, Br-, H20, "3 will be case I ; complexes inwhich the ligands have a n-system (CN-), and thusboth a-bonding (nb) and n-antibonding (n*) igandorbitals must be considered, will be case 2.

Case 1.-The pure a-orbitals are

*![aZdT) l = @['/Z(Tlh + T ?h + T 3h + T l h ) ]

*[@u(T)l = CM@[(n+ 1 ) p z l +( 4 )

c L @ [ 1 / 2 ( r l y + rZv T~ ~ + r4" j l ( 5 )

*[bZg(T)] = CM@[fidz,l + CL@ [ l /Z( TI h -XZh + T3 h - T4h)l (6 )

*r[e8(T)j = ~M@[nd, , l+ cL@ [ \/z ( m r - my) (8 )

*rr[e,(?r)~ = c M o [ ~ ~ , J

+cL@ [q n?,- inv)] ( 9 )

* [ b z u ( ~ ) I = @ [ ' / ~ ( X L V - xz v T B V - ~ Y ) ] (7)11

The mixed u- and n-orbitals are

* 1 1 [ e d u , ~ ) 1 C M @ [ ( ~ I)P,l +1 1

Case 2.-The pure a-orbitals are

C L ( l ) @[zuz - u4)] + CLW@[z



Feb. 5 , 1963 SQUARELANAR ETAL OMPLEXES 261

2

t

I A

1%. /I

L L - - y_

II

/

/

X

Fig. l.--B molecular orbital coordinate system for a square planar

The mixed u- and a-orbitals are

metal complex.

t - 4 1

The following general rules were adhered to in esti-mat ing the relative energies of t he single electron mo-

lecular orbitals:(I ) The order of the coulomb energies

is taken to be u (ligand), ab(ligand),nd(metal), (n+ 1)s(metal), n*(ligand), (n + l ) p (metal). (2) Th eamount of mixing of atomic orbi tals in the molecularorbitals is roughly proportional t o atomic orbital over-lap and inversely proportional to their coulomb energydifference. (3) Other things being approximatelyequal, a-bonding molecular orbitals are more stablethan a-bonding molecular orbitals, and a-antibondingmolecular orbitals correspondingly less stable than T -

antibonding molecular orbitals. (4) Interactionsamong the ligands themselves are expressed in termsof a ligand-ligand exchange integral p : the sign of thisinteraction for each molecular orbital is another factorin the final relative ordering. ( 5 ) The relative molec-ular orbital ordering is considered final only if it is

fully consistent with the available experimental re-sults: exact differences in the single electron molecularorbital levels can only be obtained from experiment.

The general molecular orbital energy level schemesarrived a t for square planar complexes are given in Fig.2 (case 1) and Fig. 3 (case 2). Group molecular orbita loverlap integrals (for Ni(CN)d2-) and ligand exchangeinteractions are summarized in Table 11.

Almost all square planar complexes with simpleligands are diamagnetic and contain a metal ion withthe de-electronic configuration. Thus the ground stat eis lAIg. The lowest energy excited stat es will be de-scribed separately for case l and case 2.

Case 1.-There are three spin-allowed d-d typetransitions, corresponding to the one electron transi-tions b*,(a*) -+ b lg (a*) ('AIg +. IA z g) , al,(a*) +blg(a*)

METALO R B TALS

MOLECULAR0 R B ITA LS

E .

LIGAND0RB ITALS

blg (ub) ,, g ( c b ) , e , ( u b )

Fig. 2.-Molecular orbital energy level scheme for square

planar metal complexes in which there is no intra-ligand r-orbital

system (case 1) .

('Alg +. IBlg) and e g ( T * ) +. blg(a*) (IAlg+. lEg). Allthese are pa rity forbidden as electric dipole transitions.Two allowed charge transfer transitions may be antic-ipated, corresponding to the one-electron transitionsbzu(n) + b ~ , ( u * ) lAlg+ ' A d and e,(+) +. blg(a*)

TABLE1

GROUP VERLAPSTEGRALSND LIGAND-LIGANDNTERACTIONSFOR SQUARE LANAR ETALCOMPLEXES

Ligand-

ligand

*'alp(.) 1 [?!d,Zl [L/a(ui + UP + ua + a)] 0.161 . . . . . .. . . .

* [ b i g ( u ) l [ n d z 2- yz l fl/z(ui - uz + ca - ua)] , 279 . , . , .

hlolecular Metal

orbital orbital Ligand orbitals G , j a interaction

[(n+ 1 ) ~ l [ ' / z ( ~ i + uz + ua + u4)l

['/Z(Tlh f s2h + T8h +

. 895

e [ u z g ( T ) ] . . . .

*Iazu(s)l I (n + l ) P z l [ ' / 2 ( 7 r l Y + 82 " + TliV +4h) 1 . . -2P(h,h)

T4") 0.464 +2@(v,v)

n4h) 1 0.166 +2,9(h,h)

* b g ( ~ ) l [ndzv l I1/2(T1h - ar h + sa h -P ( b r , ( n ) ] . . . 11/z(7riv - TZ V + ray -

Tar) 1 . . . -2P(v,v)

* I [ eg ( r ) l [ nd ,g ) L - m")] 0.117 . , . . , .

* ~ r [ e ~ ( s ) l Indyl l [,/i ( m ym.11

,117 . . . , . .

i 2

The group overlap integrals are for Ni(CNA)*-. The atomicwave functions are : 3d and 4s radial functions for nickel from ref.16 ; 2s carbon radial function from ref. 17 ; 2p carbon radial func-tion from ref. 18; the 4p nickel radial function is estimated as9(4p) = Rsp(l.40), where RsP (1.40) is a Slater orbital withexponent 1.40. The Ni-C distance is taken to be 1.90 A., aver-aged from the da ta given in ref. 21.



26 2 HARRY . GRAY ND C. J. BALLHAUSEN Yol. 85

M E T A L M O L E C U L A R L I G A N D

O R B I T A L S C R B T A L S 0RB IT A LS

e ( Q *,T *,

E.

b 1q(T , ,a 1q cb ,e (ub)

Fig. 3.-Molecular orbital energy level scheme for square

planar metal complexes in which th e ligands themselves have a

r-orbita: system (case 2:.

(IA?, + lEU). These are transitions from molecularorbitals essentially localized on the ligands to molec-ular orbitals essentially localized on the metal atom.An examination of the transition moment integrals forthese two transitions reveals that, for any reasonablemolecular orbitals, the 'Alg -f ]E, transition will bemore intense than 'AIg -+

Case 2.-There are three d-d transitions; 'AIg+ AZg, lAlg- Big and 'A1, + Eg, as in case 1. Th echarge transfer bands of lowest energy are expected toinvolve transitions from the highest filled metal orbitalsto the most stable empty "ligand" molecular orbital,th e U ~ ~ ( T * ) . he first charge transfer transition isb?,(T*) - ~ ~ ( T * )'Alg+ BlU). This transition is or-bitally forbidden and should have relatively little in-tensity. The second charge transfer band correspondsto t he transition u lg ( u* ) -+ u2,,(a*) ('AIg+ Azu). Th eenergy of this transition is calculated to be AE('A1, -+

'BlU) + A,, corrected for differences in interelectronic-repulsion energies in the lAzu and 'B1, excited states.The third charge transfer transition is eg(a*) -+ ~ 2 ~ -

( T * ) (lAlg+ lE,), and is calculated to be A.E(lA1,7IBl,) + A2 f A3, again corrected for repulsion dif-ferences in the 'EU nd lBIU xcited states . A summaryof the calculated orbital and interelectronic-repulsionenergies of all the excited states of interes t for case 2is given in Table 111.

Th e 'Al, + Azu and 'A1, -+ 'EU transitions are al-lowed, with the 'AI, -+ 'E, transition expected to haveconsiderably greater intensity. Thus a three-bandcharge transfer spectrum, with an intensity orderIAl, + E, > 'AIg * Azu> 'Al, + Blu, should betypical of square planar complexes with case 2 typeligands.

Spectral Properties of Square Planar Metal Com-plexes. A. Halid e Complexes (Case 1).-The com-

TABLE11

ORBITALAX D INTERELECTROSIC-REPULSION ENERGIESF SOME

COMPLEX

EXCITEDTATES OF ISTEREST FOR A4Nnd8 SQUARE PL.4NAR

Slater-CondonTerm designation Orbital energyn energy"-*

A. Singlet terms 'AI, (ground state) . . .'A?, AI -3534

d- d transitions 'BiS A I f Ar - 4Fn - 15F4:Eg Ai + A i f A: - 3F1 - 20Fc

Charge transfer 'BI, . . . . AE(iB8") . . .transitions 'AiU AE('BIu) + Ar -2QFn + 1OOF4

'E, A E ( 1 B i d f Az + Aa -15Fn f 75F4

d- d transitions JBig A i f i -12Fn - 45F4gEg Ai f Aa f A: - 9F z - 6OF4

B. Triplet terms AI - 05F4

Referred to the ground sta te as zero. * For the approxima-tions used to estimate the Slater-Condon energy for the chargetransfer transitions, see ref. 19.

plexes PdXd2-, PtX42- and AuX4- (X - = halide) aresquare planar and diamagnetic. The spectra of solidsamples of KZPdC14 and KzPd Br4have been measuredby Harris, Livingstone and R e e ~ e , ~nd by Yamada.12Three bands are observed which niay be assigned to d-dtransitions. In solution, PdC14?- and Pd Brd 2- showtwo charge transfer bands.

Th e interpretation of the spectrum of PtC1d2- iscomplicated by the presence of one or more "spin-for-

bidden" bands. However, PtCldz- seems to exhibitthe three spin-allowed d- d bands, and one charge trans-fer band. The spectra of th e AuC14- and AuBr4- com-plexes clearly show the two charge transfer bands.1°

Assignments of the spectra of these halide complexes,in terms of the derived molecular orbital level schemefor case I , are given in TabIe IV. The values of thesingle electron parameters AI , A2 and A3 also are cal-culated and are given in Table V for comparison pur-poses.

B. Cyanide Complexes (Case 2).-The spectra ofthe square planar Ni(CN)d2-, Pd(CN)4,- and Pt-(CN)42- complexes are given in Table IV . Each com-plex exhibits the expected three charge transfer bands,but the d-d bands for the most part are obscured bythese charge transfer bands. For Ni(CN)4?-, there

are two shoulders on the tail of the first charge trans-fer band . These shoulders probably represent at leasttwo d- d transitions. The charge transfer spectrum ofAu(CN)4- appears a t higher energies than for Pt-(CN)d2-. Thus only one of the expected three chargetransfer bandsis seen below 50,000 an.-'. The signifi-cance of this is discussed later.

Assignments of the spectra and the calculated valuesof A,, A2 and A3 for the cyanide complexes are given inTables IV and V, respectively.

Magnetic Properties of Square Planar Metal Com-plexes.-The magnetic susceptibility of the diamag-netic d8square planar metal complexes is given by

x = '/3(3X.L + XII) (22)where

x.L X A

+AH F(I) (23)

X I I = X A + XH . ( I I ) + XnW (24)

In eq. 23 and 24, X A s the atomic contribution for themolecule in question (the sum of the Pascal constants),Xringepresents the expected diamagnetic contributionof t he metal-ligand a-molecular orbital system (analo-gous to t he ring diamagnetism in benzene) and

(9) C.hl . Harris, S. E. Livingstone and I. H. Reece, J.Clievi. Soc., 1505

(1959).

(10) A. K. Gangopadhayay and A. Chakravorty, J . Chcm. P h y s . , 3 6 , 2 2 0 6( 1 9 6 1 ) .



Feb. 5, 1963 SQUARE PLANARMETALOMPLEXES 263

TABLEV

SPECTRALROPERTIESF SQUARE LANARETALCOMPLEXES

Reference

Y m X , to exptl.Complex (sample) cm. -1 c Assignment work

Case 1

PdClet- (solid K2PdClr) 16,700 . 'Aig + 1Ak 9 , 12 , 2021,500 . 'AI, * Big23,300 .... lAip"Eig

(aq. s o h , excess C1-) 36,000 12,00 0 'Aig -+ 'A-

44,900 30,000 'Alg * 'Eu

20 , 000 . . . . 'A?, + Big

26 , 0004 . . . . 'AI, -+ LE,

(aq. soln., excess Br-) 30,10 0 10,400 'AI , - AtU

40 , 500 30 , 400 'AI, -+ lEu

21,000 15 'Aig* lAzg25,500 59 'Aig -+ 'Big

30 , 200 64 ' A i g + 'E,

46 , 000 9 , 580 *Aig* A?,

19,700 15 'Atg24,300 100 'AI, + Big

28,200 120 'AI, + 'E,

37,300 7000 lAig + A-

44,200 29,500 'Aig -C 'Eu

39,400 38,900 'Ais + E,

PdBrcr- (sol id KrPdBrc) 16,000 . . . . 'AI, -+ 'Atp 9

PtC1,S- (as. soh.) 17 ,70 0 2 .6 1111 + Ai, 2a

(37,YOO) 250 b

PtBrdt- (aq. soln., 1 MBr-) 16,70 0 5 'AI. * Az, This work

AuCIc- (aq. soln.) 31,80 0 4,5 70 'Ai g+ 'As,, 10

AuBrc- (as . soln.) 2 6 , 3 0 0 5 , 0 1 0 ' A i g * l A t u 1 0

Case 2

Ni(CN)r'- (aq. s o h ) (22,500) 2 'Aig * Ai, This work

(30,500) 250 'AI, + Big32 , 300 700 'A * BIU

35 , 200 4 , 200 'Ai, * Aq,

37,6 00 10,60 0 'AI, -+ 'Eu

45,400 7 ,200 'Aip -+ 'Azu47 , 200 9 , 000 ' A is + E,

Pt(CN)2- (aq. s o h ) 35 , 720 1 , 590 'Ai, -+ 'Bin(38,680) 26,000 lAig + E,

39,180 29,500 'Aig * E,

41 , 320 1 , 850 ' A t , + Azu

37 , 850 331 'Ai, * Big46,080 2 ,400 'Aig .-c 'BI,

Pd(CN)ez- (aq. soln.) 41,6 00 1,200 lAil l -+ IBi, 20

Au(CN)a- (aq. so h ) 30,96 0 51 'Ai , * Ai, 22

this band is found a t 21,700 crn.-l in t h e K2PdBrd.2H20crys tal; see ref. 12. bPossibly a "spin-forbidden" charge

transfer transition.

TABLE

COMPLEXES

METAL&ORBITAL ENERGIESOR SELECTED SQUARE PLANAR

a,,- y t1'

d m a,,

Orbital energy differences (cm.-l) for Fz = 10F4 = 700 cm-1

Complex ion AI Ar AI

PdC1,'- 19,150 6,200 1450

PdBr42 - 18,450 5,400 5650 (1350)"

PtC142- 23,450 5,900 4350

PtBraa- 22,150 6,0 00 3550

AuC14- >20,000 . . . . . . . .AuBrr- >20,000 . . . . . . . .Ni( CN)4a- 24,950 9,900 650

Pd(CN)ra- >30,000 10,800 50

Pt( CN)da- >30,000 12,600 -4140

Au( CN)d- 33,410 10,620 . . . .a The value of 4 or KsPdBr4.2H10.

where g is the degeneracy of the ground state andAE0.i is the energy separation between th e ground sta te

$0 and excited state +i. For the d8 complexes underconsideration, the ground s ta te is lAlg, and eq. 25 and26 reduce to

Th e magnetic susceptibility of a single crysta l of Kz-Ni(CN)4.H20has been measured by Rogers." Thevalues he obtained were

XI1 = -127.8 X 10-B

X I = -146 X 10-6

(c.g.s.)

(c.g.s.)

Subtracting eq. 24 and eq. 23, the unknown XA iseliminated, yielding

XI - Xi 1 - XH . F . ( I )- Xt i .~ . t l i ) ) -Xr,,

Using the experimental values for the AE's , XH.F.(L)26 X and XH.F.(II)142 X are calculated;this gives X r i n g = -98 X loF6.

Discussion

The metal orbitals involved in u-bonding in squareplanar complexes are the nd,~, dxi-p, (n + l)s , (n +l)p, and (n + l )py . Th e ndX2-~2, n + lb, n + l ) P xand (n + l)p, orbitals account for most of the u-bond-ing, judging from the values of the overlap integrals inTable 11,and the ndz2makes only a minor contribution.The most important a-molecular orbital is the uzu, on-sisting of the (n+ l)p, metal orbital and a combinationof the four ligand a,-orbitals. This gives a very stabler-bonding orbital (with only a single node in the molec-ular plane) which may be called the "ring" 7-orbital.All the other pure n-molecular orbitals have an equalnumber of electrons in their bonding and antibondinglevels.

The tendency of nd8-metal ons to form square planarcomplexes increases in the order N iZT< Pd2f < Pt2+.Two features of the molecular orbital bonding schemeare consistent with this observation. Firs t, the avail-ability of the nd+ya metal orbital for u-bonding un-doubtedly increases in going from Ni2+ to Pt2+ (X -

c14'- is tetrahedral; PdC1d2- and PtC142- are squareplanar). Second, the square planar configuration is sta -bilized by the ring a-bonding, which is expected to in-crease from Ni2+ o Pt2+. The addition of a fifth groupabove the square plane drastically decreases the ring X-

bonding, by tying up th e (n+ l ) p , orbital in u-bonding.The assignments of t he electronic spectra of the

square planar halide and cyanide complexes of Ni2+ ,Pd2+, Pt2+and Au3+ are given in Table IV . Thehalide complexes will be considered first. Th e spectraof K2PdC14 nd K2PdBr4 show three bands which maybe assigned to d-d transitions. In the molecular orbitalenergy level scheme for case 1 ligands, the order of t hesingle electron molecular orbitals is reasonably ex-pected to be eg(7r*), ulg(u*), bzg(a*) and bl, (u*) . Thusthe bands are assigned 'AIg +. lAzg, lAlg + B1, andlAlg + lE, in order of increasing energy. The energyorder of the single electron molecular orbitals deducedhere is consistent with the results of a "complete" elec-trostatic calculation of square planar Pt2 complexesperformed by Fenske, et aL6

The spectra of PtC142- and P tBr2- are slightly morecomplicated; the first band is quite weak and is as-signed as the first spin-forbidden transition, 'A1, +

3A2g. The spin-allowed transitions are assigned in thesame way as for the PdXh2- complexes. The spectraof single crystals of K2PtC14 using polarized light showtha t the 25,500ern.-' band occurs in x ,y and the 30,200cm.-' band occurs in z polarization.12 These polariza-

(29)

(11) M. T. ogers, J. Am . Chem. SOC.. 9, 1506 (1947).(12) S. Yamada, ibid. , 73, 182 (1951).



VOl. 85G-L HARRY . GRAY ND C. J. BALLHAUSEN

tions are expected for the assignments given here, as-suming a vibronic intensity giving mechanism. Chat t,Gamlen and OrgelZapreviously assigned the spectrumof Pt CL- somewhat differently (17,700 cm.-l, 'AIg -+

3A2g;21,000 crn.-', 'Al, + 3Eg ; 25,500 cm.-l, 'Alg -+'Aza;30,200 cm.-', 'AIg + Eg). However, the 25,500cm.-l and the 30,200 cm.-' bands are both quite sensi-tive to axial perturbations (both bands shift to the redon changing the solvent from 12 d l HC1 to water).13This is evidence against the assignments of Chat t, etdjZaand is consistent with the assignments proposed here.

The AuC14- and A4uBr4- complexes do not exhibitany bands which can unambiguously be assigned to d-dtransitions. In these cases the two charge transferbands appear at lower energies (than for PtXd2- com-plexes) and mask the weaker d-d bands.

The only bands which can clearly be identified in thecyanide complexes are due to charge transfer transitions.A rough calculation may be performed to show that thecharge transfer bands in Ni(CN)i,- appear at aboutthe energy expected for transitions from metal d-orbitalsto the ligand uzU(n*) rbital. The ionization potentialof HCN is about 14 e.v.I4 The calculated separationof the single electron nb and x * levels in HCN is 9 e.v .l6Thus the coulomb energy of the x * level of complexed

CN- may be estimated a t -5 e.v . The coulomb en-ergy of the Ni2 +d-orbitals in Ni(CN)4*- may be esti-mated at about - 0 e.v. , using t he spectroscopic datacompiled by Mo ore,*4 nd assuming that the ne t chargeon the complexed Ni2+ is actually about +1. Thisleaves 5 e.v., or about 40,000 cm.-l, as the difference inenergy of the single electron d (Xi2+) and n*(CN-)orbitals. Since the uZu(n*) level is stabilized by theempty 4p, metal orbital (also by +2p(v,v,)) and sincethe 3dxy,3dxz,3d,, and 3dz2orbitals ar e actually weaklyantibonding, the appearance of the three charge trans-fer bands in Ni(CN)d*- between 32,900 cm.-' and 37,500cm.-l is in accord with theory.

The separation A1 for the cyanide complexes is by f arthe largest of the three orbital parameters, and illus-trates that the d,~- ,~s much more strongly antibondingtha n any of the other d-orbitals. The separation A3

decreases regularly in going from Xi2+ to Pt2+ (forNi(CN)42- and Pd(CN) J*- , the d,? is more unstablethan d,,, drs; for Pt(CN)4?-, the dZ?s more stable than

d,,, dy J. This is compatible with the idea that th eaxial interaction of the water molecules in water solu-tion decreases in going from Ni 2+ to Pt2+. The twocomponents observed for the intense charge transferband in Pt(CN)4*- are consistent with its assignmentas 'Al, + E,; th e 'E, excited state is expected to besplit by second-order spin-orbit effects.

The assignments given in this paper for the cyanidecomplexes differ considerably with assignments givenby other investigator^.^^^^^ The reason for this dis-agreement is found in our "band assignment philos-

ophy"-ussign only the bunds which can clearly be iden-t$ed as to energy and intensity. In th e case of t hecyanide complexes these bands are almost certainly dueto charge transfer transitions, as judged by their largeintensities.

A careful comparison of the charge transfer bands forth e PtC1d2-, AuCld-, Pt(CN)h2- an d Au(CN)4- com-plexes offers further proof of the assignments givenhere. Thu s the first charge transfer band appears a tlower energy for AuC14- than for PtCL2-, as expectedfor a ligand + metal type electronic transition, sinceAu3+ is a better electron acceptor than Pt2f . On the

(13) R. G. Pearson, H. B. Gray and F. Basolo, J.Am . Chem. SOL. ,2, 787

(14) F. H. Field and J. L. Franklin, "Electron Imp act Phenomena,"

(15) K. Iguchi, J . Chem. Phys., 23, 1983 (1955).

(1960).

Academic Press, Inc., New York, Ti. Y.,957, p. 273.

other hand, the first charge transfer band for Au(CN4-occurs a t higher energy than for Pt(CN)42-, expectedonly if th e charge transfer for these cyanide complexesis of t he metal+ igand type.

The charge transfer band systems observed for thehalide (case 1) and cyanide (case 2) complexes shouldserve as a guide in assigning the charge transfer in othersquare planar complexes as ligand (n)+metal (case 1)or metal+ igand [uzu(n*) (case2 ) . In the halide com-plexes the two charge transfer bands are separated by10,000 to 13,000 cm.-l. Thu s it can be estimated th atthe bZu(n) and eu(nb) molecular orbitals are separatedby a t least 10,000 cm.-' in a typical square planar com-plex. This is a reasonable value since the b2u orbitalis non-bonding with a ligand interaction of -2/3(v,v),and the e, orbital is w-bonding.?Z

The three charge transfer bands due to metal -+

ligand [uzU.*)I transitions and exhibited by the cyanidecomplexes are much more closely spaced. The twoorbitally allowed metal -+ ligand transitions (ulg(u*)+

U , ~ ( T * ) , eg(n*) -+ uzu(a*)) are only 2,000-3,000 cm.-'apart in all the cyanide complexes, since the u lg (u* )and eg(x*)metal orbitals are virtually non-bonding andthus very nearly equal in energy in these complexes.

Thus these two types of charge transfer show quite

differen t properties and should be distinguishable inother nda square planar complexes. The ligand (n ) ometal typ e will have two bands, spaced by about 10,000cm.-', the second more intense than t he first. Themetal + igand [usu(n*)]ype will have three closelyspaced bands with intensities bZg(n*)+uzU(n*) small],ul,(u*) +U ~ ~ ( T * )intermediate] and eg(n*)+ azU(n*)[large]. The exact positions of these bands give ofcourse the best clues to the relative ordering of the d-orbital energy levels. In all probability there will becases in which the bZg(n*) level is fairly close to alg(u*)

and eg(n*) (A , small), and thus only th e two orbitallyallowed bands will be observed. In any event thespacing of t he two orbital ly allowed charge transferbands will easily distinguish between ligand -+ metaland metal -+ ligand type transitions for the nds squareplanar complexes.

The "ring" diamagnetism apparently observed forK*Ni(CN) .HsO may be considered as experimentalevidence of electron delocalization through th e cyanidesystem. The electrons in th e uZu(nh)orbital, in par-ticular, would be expected to generate considerable"ring current."

Experimental.-The complex K2Ni(CN)4.H20 wasprepared as described in the 1iter atu1 -e.~~ he elec-tronic spectral measurements were made using a Cary14 spectrophotometer.

Acknowledgments.-The auth ors tha nk Dr . AndrewD. Liehr for making available a preprint of his paperwith Dr. Adamson a nd Dr . Perumareddi, and for stimu-

(16) J. W. Richardson, W. C. Nieuwpoort, R. R. Powell and W. .

(17) D. R. Hart ree, "The Calculation of Atomic Structures ," John Wiley

(18) C. Zener, Phys. Rev., 3 6, 51 (1930).(19) H. B. Gray and C. R. Hare, Inovg. Chem., 1, 363 (1962).

(20) C. K. Jp'rrgensen, "Absorption Spectra of Complexes of Heavy

(21) "Interatom ic Distances," T he Chemical Society, London, 1958, p. M

(22) A. Kiss, J. Csaszar and L. Lehotai, Acta Chim. Ac ad . Sci . Hung . , 14,

(23) W. C. Fernelius and J. J. Burhage, Inoug . S y n l h e s e s , 2 , p. 227.

(24) C. E. Moore, "Atomic Energy Levels," U. . Natl. Bur. Standards

Circular 467, 1949 and 1952.

(25) There is another assignment possible fo r the two charge transfer

bands in the square planar halide complexes. The first hand could contain

both th e b 2 U ( T )- l g ( u* ) an d a,(nb) * i , ( s* ) transitions; this would leave

the allowed eu(ub) - lg (c*) transition responsible for the second intense

band.

Edgell, i b id . , 36 , 1057 (1962).

an d Sons, Inc., New York, N. Y., 1957, pp. 169-171.

hfetals," Report to U. . Army, Frankfurt am Main, October, 1958.

178.

228 (1958).



Feb. 5, 1963 HEXACHLOROMETrZLLATES O F TRIVALENTHROMIUM, ANGASESE AN D IRON 26 5

lating discussions. Acknowledgment is made to thedonors of th e Petroleum Research Fund of the Ameri-

can Chemical Society for pa rtia l support of this re-search.

[CONTRIBUTIONROM THE WILLIAM LBERTNOYES ABORATORYF CHEMISTRY,NIVERSITYF ILLINOIS, RBANA,LLISOIS]

Hexachlorometallates of Trivalent Chromium, Manganese and Iron

BY ~ ' I L L I A ME. HATFIELD, OBERT . F.~Y, . E. PFLUGERND T. S. IPER

RECEIVEDUL Y 16, 1962

Compounds containing hexachlorometallate ions of t riva lent chromium, manganese, and iron have been pre-pared by stabilizing the anion in a crystal lattice with the tris-(1,2-propanediamine)-cobalt(11) cation. Theexistence of the complex anions has been proved by isomorphism with salts of known hexachlorometallate(III)ions and the same cation, and by X-ray dat a obtained from single crystals of [Co(C3Hl0X&] FeC&]. Magneticdat a showed that the ground st at e of these compounds is the expected high-spin state . Dq values of 1318, 1754and 920 cm.-' for [CrCl~l-, [MnC16]- and [ F ~ C ~ Q ]-, respectively, were determined from spectral data.

Introduction

Complex ions including those n ot found in appreciableconcentrations in solution often can be stabilized in acrystal lattice providing th at t he energetics of forma-tion of the complex ion from solvated species is not toounfavorable relative to the lattice energy. A goodguide to a suitable counter ion is the rule tha t th e lat-

tice energy usually is largest for ions of equal and oppo-site charge. Furthermore, a counter ion of similar sizeis best since disparate sizes promote solvation.

This approach has been used recently by Gill andNyholm, who have isolated stable compounds contain-ing such ions as the tetrahalonickellate(I1) ion. Simi-larly, Mori2 has prepared a number of chlorocupratesby stabilizing the complex anions in a lattice with thehexamminechromium(111) cation.

Although there are several references in the olderliterature to compounds having stoichiometries con-sistent with hexachlorometallate salts for trivalentfirst transition series elements, e.$., T & C T C ~ ~ , ~s3Fe-C ~ G . H ~ O , ~nd (h:H4)3CrC16,6he properties of the com-plex anions are unknown. In fact, evidence reportedhere casts doubt on the existence of any hexachloro-

metallate ions in these compounds. mre now wishto report the preparation and some physical propertiesof compounds which do contain hexachlorometallate

ions of trivalent chromium, manganese, an d iron. Upto now the position of the chloride ion in the neuphel-auxetic series has been based on spectral studies of an -hydrous chromium(II1) chloride and complexes ofsecond and thi rd transition series ions.6 Studies onoctahedral complexes of first transition series ions co-ordinated only by chloride should provide additionalinformation concerning the ligand field of the chlorideion.

ExperimentalTris-( 1,2-propanediamine)-cobalt(III) Hexachlorochromate-

(III).-A small amount of the pink compound crystallized froma concentrated hydrochloric acid solution (50ml.) of chromium-(111) chloride hexahydrate (5.4 g. , 0.020 mole) and tris-(1,2-pr0panediamine)-cobalt(II1) chloride (3.8 g., 0.010 mole)after standing for several days. The crystals were collectedon a Biichner funnel, washed well with water, then with alcoholand ether, an d dried in an oven at 110' for four days.

Anal. Calcd. for CoCgH&T6CrC16: C, 19.80; H, 5.53; N,

15.40. Found: C, 19.87;H, 5.61; N, 15.09.

Tris-(1,2-propanediamine)-cobaIt(111) Hexachloromanganate-(III).-This compound was precipitated from a hot solution ofmanganese(I1) sulfate monohydrate (3.4 g., 0.020 mole) and

(1) N . S. ill and R. S. yholm, J . Chem. Soc., 3997 (1959).(2) M. Mori, Bull. Chem. Soc. Japan . 33, 985 (1960).(3) G. Neuman, A n n . , 244, 329 (1888).(4) P. T. Walden. A m . J. Sci., 131 148, 283 (1894); 2. m r g . u. al lgem.

(5 ) H. I. Schlesinger, J . Am . Chem. Soc., 61 , 3520 (1929).

( 6 ) C. E. Schaffer and C. K . JWgensen, J . Ino ug . Nuclear Chcm., 8 , 14 3

(1958).

Chcm., 1, 332 (1894).

tris-(1,2-propanediamine)-cobalt(11) chloride (7.6 g., 0.020mole) in a minimum amount of hydrochloric acid by the addi-tion of small amounts of sodium chlorate . The dark brown com-pound was collected on a Biichner funnel, washed with alcohol,acetone, and ether, and dried in an evacuated desiccator overpotassium hydroxide.

Anal. Calcd. for COC&oN&fnCls: C, 19.70; H, 5.50; N,15.31; C1, 38.72. Found: C, 20.18; H, 5.61; N, 15.24;C1, 38.98.

The compound decomposed slo\vly upon standing presumablyevolving chlorine. The magnetic moment of a sample that hadbeen stored for three months was 0.3B.M . higher than th at of afreshly prepared sample.

Tris-(1,2-propanediamine)-cobalt(III) Hexachloroferrate(II1).---I olution of iron(II1) chloride hexahydrate ( 2 .7 g., 0.010mole) in 5 mi. of hot water was added to a solution of tris-(1,2-propanediamine)-cobalt(II1) chloride (3.8 g., 0.010 mole) in10 ml. of hot water. The solution was evaporated to about halfthe original volume. A few milliliters of concentrated hydro-chloric acid was added; the yellow crystals formed were col-lected on a Biichner funnel, washed with alcohol, acetone, andether, and air dried. The compound was recrystallized fromhot water with the addition of hydrochloric acid. The yieldwas 1.8 g. (29%).

Anal. Calcd. for CoCgHsoS+jFeClo: C, 19.66; H, 5.50; N,16.28; C1, 38.69. Found: C, 19.88; H, 5.14; N, 15.46; C1,38.69, 38.61.

Tris-(1,2-propanediamine)-cobalt(111) Hexachloroindate( 111).-This compound was prepared as was the corresponding ironcompound, The yield was 87%.

Anal. Calcd. for C O C ~ H ~ O N ~ ~ C ~ ~ :, 17.76; H, 4.96; N,13.81; C1, 34.96. Found: C, 17.68; H, 5.00; X, 13.08;C1, 35.00.

Tris-(1,2-propanediamine)-rhodium(111) Hexachlorometal-lates (II1) -Compounds with the tris-(1,2-propanediamine)-rho-dium(II1) cation were prepared by the methods described above.The analytical da ta for the new compounds are given in Table I .

Hexaamminecobalt(II1) Hexachloroferrate(III).--A solution offerric chloride hexahydrate (1.4 ., 0.005mole) in 10 ml. of hotwater was added to a hot solution of hexaammiriecobalt(II1)chloride (0.43 ., 0.0016mole) in 70 ml. of 3.4N hydrochloricacid. The resulting solution was heated for a few minutes untilan orange product crystallized. The crystals were collected on aBuchner funnel, washed with alcohol and ether and dried in anevacuated desiccator over potassium hydroxide. The yield was0.5g. (727,).

Anal. Calcd. for CoXcHlaFeCl6: N, 19.54; H, 4.18; C1,49.55. Found: N , 19.52; H,4.13;C1,48.42.

The preparation could not be made a t room temperaturebecause of the insolubility of hexaamminecobalt(I I1) chloride in asolution with a chloride ion concentration high enough to facili-ta te t he formation of the hexachloroferrate( 111) ion. Detailedstudies of this compound were not undertaken because we ex-perienced difficulties in preparing the corresponding manganesecompound.

Experiments with Other Cations and Anions.-Proper condi-tions could not be found for the preparation of hexachloro-ferrate(111) sal ts of the tris-(2,3-butanediamine)-cobalt(III) rtris-( 1,2-cyclohexanediamine)-cobalt(III) cation, An orangeproduct crystallized from a solution of tris-(ethylenediamine)-cobalt(II1) chloride and ferric chloride after a few days. Thecrystals were collected on a Biichner funnel, washed with alcoholand ether, and dried in air.

A n d . Calcd. for CoCsH?aSsFeCls.1.5H~O:C, 13.48; H,5.08; K , 15.74. Found: C, 13.21;H,4.68: N, 15.92.

The yield was 8.1 g. (75%).

harry b. gray and c. j. ballhusa- molecular orbital theory for square planar metal complexes

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