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Electronic structure, spectra and magnetic properties of oxycations—IV ligation effects on the infra-red spectrum of the vanadyl ion
J, lnorg. Nucl. Chem., 1963, Vol. 25, pp. 1359 to 1369. Pergamon Press Ltd. Printed in Northern Ireland
ELECTRONIC STRUCTURE, SPECTRA A N D MAGNETIC PROPERTIES OF O X Y C A T I O N S - - I V LIGATION
EFFECTS
ON THE INFRA-RED VANADYL ION
SPECTRUM
OF THE
J. SELBIN, L. H. HOLMES, JR.* a n d S. P. MCGLYNN Coates Chemical Laboratories, Louisiana State University Baton Rouge, 3, Louisiana (Received 11 April 1963)
Abstraet~Vanadium-oxygen multiple bond stretching frequencies have been tabulated for fifty-one oxovanadium (IV) and three oxovanadium (V) compounds. The spread of the values may be summarized by ~ = 985 ± 50 cm -1, which corresponds to a force constant spread, k = 7.0 ± 0"7 mdyne/A,. The frequency shifts are used to establish several short ligand series, and they are interpreted in terms of cr and ~r contributions to the vanadium-ligand bond. It is concluded that frequency decreases are directly and primarily caused by decreased p~r -- drr donation from oxygen to metal and to increased electrostatic repulsion of the vanadium and oxygen species. The effect may he written Af = --~r(L --~ M) -- 7rll(L ~ M) -- zrj(L ~ M) where it is estimated that the relative importance of terms is 8:2:1, and it is assumed that the M--~ L donation is negligible for the d 1 vanadium (IV). IN a c o n t i n u i n g study o f the spectroscopic p r o p e r t i e s o f m e t a l oxycations/1,2) we have been investigating the o x o v a n a d i u m (IV) (vanadyl) ion.~S,4) E x p e r i m e n t a l evidence for the existence o f the v a n a d y l ion, V O 2+, in a q u e o u s solutions is strong, b u t n o t una m b i g u o u s . <5-s) H o w e v e r , FELTHAM19) f o u n d t h a t the v a n a d i u m (IV) ion in a q u e o u s solutions has visible a n d electron spin resonance spectra which are very similar to those o f the solid f i v e - c o - o r d i n a t e d v a n a d y l complexes a n d quite different f r o m those o f the m o r e s y m m e t r i c VCI 4 a n d s i x - c o - o r d i n a t e d chelates. H e c o n c l u d e d t h a t the species present in a q u e o u s solutions is VO2+.nH20. F u r t h e r s u p p o r t for the o x y c a t i o n species in preference to a m o r e s y m m e t r i c a l species such as V(OH)z 2+ is s u m m a r i z e d by BALLHAUSEN a n d GRAY. (sI A g a i n the e x p e r i m e n t a l q u a n t i t y evaluated is the p a r a m a g n e t i c r e s o n a n c e (g) value. X - r a y diffraction studies o f the solid c o m p o u n d s V O S O 4. 5 H200°) an d VO(acac)2 ~l1) clearly d e m o n s t r a t e the existence a n d stability o f the V O 2+ entity in solid complexes. This is n o t surprising when it is realized that even in VO2, which crystallizes in a highly * Present address: Southeastern Louisiana College, Hammond, Louisiana. ~1~S. P. MCGLYNrq, J. K. SMITHand W. C. NELLY, ./7. Chem. Phys. 35, 105 (1961). ~2~B. SHAMaURGEg,M. S. Thesis, Louisiana State University (1960). (sl j. SELBIN, S. P. McGLYNN, J. K. SMITH,G. A. SIBILLE,and L. J. TaUEX. Paper presented at the 16th Southwest Regional ACS-Meeting, Oklahoma City, Okla., Dec. (1960). ~4) j. SELBIN,L. H. HOLMES,JR. and S. P. McGLYNN, Chem. andIndustr. 746 (1961). (~ L. P. DUCgET. Thesis, Paris University (1951). ~c,~F. J. C. RossovTI and H. S. ROSSOTTI,Acta Chem. Scand. 9, 1177 (1955). irt C. K. JORGENSEN,Acta Chem. Stand. 11, 73 (1957). 181 C. J. BALLHAUSENand H. B. GRAY, Inorg. Chem. 1,111 (1962). 19~R. D. FELTHAM. Thesis, University of California, UCRL-3867 (1957). I10) M . B. PALMA-VITTOI~ELLI, M . U . PALMA and F. SGARLATA, Nttol)o Cimento (Series 10) 3, 718 (1956); G. LUNDGaEN,Rec. Tray. Chim. 75, 585 (1956). (11~R. P. DODGE, D. H. TEMPLETONand A. ZALKIN, J. Chem. Phys., 35, 55 (1961). 1359
1360
J. S~LBIN,L. H. HOLMES,JR. and S. P. McGLYNN
distorted rutile structure, there is one short V--O bond in each VOe unit. ~12) Chemical analysis and molecular weight determination have also established that the formula of the acetylacetonate (acac) complex is VO(acac)2 and that it is monomeric in benzene.t13~ Indirect evidence for the mononuclear monoxy VO 2+ species may be deduced from the infra-red spectra of vanadyl compounds. BARRACLOUGHe t al. ~14) examined the infra-red spectra of several complex metal compounds which were presumed to contain metal-oxygen multiple bonds. They reported that the stretching frequencies of metal-oxygen double bonds may generally be found between 900 and 1100 cm -1. Specifically they reported probable V--O stretching vibrations for the following four compounds: VOSO4, 987, 1003, 1020; VO(acac)2, 995; (NHa)2[VO(ox)2].2HzO, 976; and (NH4)2[VO(mal)2].4H20, 967, 977 (see footnote to Table 1 for abbreviations used). Earlier, vanadium (IV) hydroxide had been formulated as VO(OH) 2 on the basis of its infra-red spectrum, ~15~which shows a strong band at 955 cm -1, attributed to the V--O stretching vibration. It is apparent from the above values for band maxima that the environment of the VO 2+ entity conditions a marked effect upon the stretching frequency, and hence upon the force constant and strength of the V--O bond. This is not unexpected and, in fact, MCGLYNN and co-workers tl~ have shown that the infra-red bands due to the dioxouranium (VI) symmetric and antisymmetric stretching vibrations may be used as a probe to establish a ligand series exhibiting a striking parallelism with the spectrochemical series. The persistence of the oxovanadium (IV) species in a number of compounds makes possible the utilization of the properties of the oxycation species in a study of the bonding structure of complexes of this d 1 electron system. Ideally, one might expect not only to be able to study the symmetry types available to this transition metal ion, but also to learn more about the nature of the bonding to various ligated species as they perturb the VO 2+ entity. Thus the stretching frequency of the metal-oxygen bond should serve as a sensitive probe from which certain structural and bonding features may be deduced. Furthermore, the oxymetal species will show electronic localized charge-transfer absorption bands (e.g., arising from electron transfer from orbitals located primarily on the oxygen to orbitals primarily located on the metal) in addition to those normally observed with d n ions. These additional bands may be expected to be independent upon the other ligands and the strength and symmetry of the ligand field. In this paper we shall be concerned only with the interpretation of the infra-red spectra of vanadyl compounds. A subsequent paper tle~ will deal in detail with the electronic spectra. The electronic energy levels in the vanadyl entity and in its complexes have been deduced from a combined crystal-field and molecular orbital approach and it is determined (see Appendix) that there is no basic disagreement with the molecular orbital scheme of BALLHAUSENand GRAY. (8) ~z) B. ANDERSON,Acta Chem. Seand. 10, 623 (1956). ~ls~M. M. JONESJ. Amer. Chem. Soc. 76, 5995 (1954). ~4~ C. G. BARRACLOUOH,J. LEWISand R. S. NYHOLM,J. Chem. Soc. 3552 (1959). tlS~ C . CABANNEs-OTT,C.R. Acad. Sci., Paris 242, 2825 (1956). tx6, S. P. McGLYNN,J. SELBIN,T. R. ORTOLAND,G. MAUS,and W. JORDAN,in preparation for publication.
Electronic structure, spectra and magnetic properties of oxycations--lV
1361
It appears that the v a n a d i u m in the vanadyl m a y be either five- or six-coordinated (13,17) and it is n o t always certain which obtains in a given c o m p o u n d . However, chemical analysis coupled with molecular weight and/or conductance data will generally suffice to settle this point. (18~ In sum, we have examined the infra-red spectra o f thirty-nine vanadyl complexes, thirty o f which are new c o m p o u n d s whose preparation and properties are reported elsewhere. (is) Stretching frequencies o f the perturbed VO2+ species have been assigned, certain ligand series established and discussed. EXPERIMENTAL The new compounds studied in this investigation were prepared and characterized according to procedures previously outlined. (18) Known compounds were prepared according to procedures found in the literature. Starting materials included reagent grade vanadyl sulphate, vanadyl chloride, ammonium metavanadate and vanadium (V) oxide. All other chemicals used were of reagent grade quality or better. Infrared spectra in the range 5000-650 cm -1 were obtained with a Perkin-Elmer Model 21 and a Beckman IR-7 recording spectrophotometer, both employing NaC1 optics. Polystyrene film was used to calibrate the instruments. Spectra were run with the samples in Nujol mulls since these gave better spectra than the alkali halide pellet technique. Interaction of some of the compounds with the alkali halide under the conditions necessary for pellet preparation was evident. Spectra of all ligands and/ or complexes of other transition metal ions with the ligands were obtained and using these, there was usually no difficulty in assigning the band(s) of the perturbed VO~+ stretching vibration. In those cases in which the ligands have bands in the region of the metal-oxygen stretch, comparisons were made with the spectra of analogous complexes of other divalent metal ions of the first transition series. Infrared spectra of many of the vanadyl complexes studied in this work were run in solution in a concurrent investigation to be reported elsewhere. (19) In all of those cases where solubility was effected in an inert solvent the value for the V--O stretch varied no more that ±5 cm -1 from the value found with the solid complex. RESULTS AND DISCUSSION Table 1 lists the frequencies assigned to the v a n a d i u m - o x y g e n multiple b o n d stretching vibration for fifty-one o x o v a n a d i u m (IV) and three o x o v a n a d i u m (V) c o m p o u n d s . In those cases where more than one b a n d was observed in the metaloxygen multiple b o n d stretch region, the strongest is listed first. The values enclosed by parentheses were taken f r o m the literature cited. The force constants, k, have been related to the frequency using a simple h a r m o n i c expression, and are f o u n d to vary f r o m 7.68 for VOCI 3 to 6.23 for [VO(DMSO)5 ] (C104) 2 in mdyne/A. The listing in Table 1 m a y then be approximately summarized by the single values k = 7.0 :k 0.7 and ~ = 985 ~ 50 cm-1; it behoves the authors to illustrate then that the spread of the effect being observed is larger than that which could arise f r o m (a) lattice forces or (b) simple mechanical coupling to the metal-ligand vibrations: (a) Certainly, some o f the shift observed to occur in the V - - O stretch frequency can undoubtedly be caused by the operation o f lattice forces which are expected to vary f r o m one solid c o m p o u n d to the next. However, the magnitude o f the crystallineenvironment effect is considered to be small. Specifically, some o f the c o m p o u n d s ~17)R. T. CLAUNCH,T. W. MARTIN and M. M. JONES,J. Amer. Chem. Soe. 83, 1073 (1961). (lS) j. SELBINand L. H. HOLMES,JR., J. Inorg. Nucl. Chem. 24, 1111 (1962). ,9) j. SELBIN,H. R. MANNINGand G. CESSAC,presented at the 17th Southwest Regional A.C.S. Meeting, Dallas, Texas, December, 1962, J. Inorg. Nucl. Chem. In Press.
1362
J. SELBIN,L. H. HOLMES,JR. and S. P. McGLYNN
TABLE 1.--VANADIUM-OXYGEN STRETCHING FREQUENCIES, OXOVANADIUM(IV) AND (V) COMPOUNDS
Compound* 1. 2. 3. 4. 5.
VOCIa VOBr3 V~O5 [VO(p-OCHs-TPP)] VOSO¢5H20
6. 7. 8. 9. 10. 11. 12. 13. 14.
[VO(phthal)] [VO(p-CI-TPP)] [VO(TPP)] [VO(p-CHa-TPP)] [VO(Benzac) 2] [VO(DBM)~] [VO(sal) 8] (Et,N)a[VOCNCS)6] [VO(acac)~]
15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41.
VOBrcH20 [VO(Etiopor)] VO(C104)2"xH20 (NH,)2(VO)2(SO4)2 VOClz.xH~O [VO(ophen) (ox)] (Me4N)s[VO(NCS)5] [VO(ophen)~](C104)~ VOF~-xH20 (Me,N)2[VOCI,(EtOH)] K3[VO(NCS)5] [VO(ophen)~C1]Cl [VO(ophen)2Br]Br.HzO [VO(en-bis-sal)] [VO(oxine)~] [VO(dipy)SO,] [VO(dipy) (ox)] [VO(dipy)~](C10~)~ WO(ophen)SOd VO(dipy)l.5(NCS)z (NH,)2VO(mal)2"4H~O (NH4),VO(ox),.2H,O VO(ophen)l.2(NCS), VO(ox)-H~ox'3H,O Na,VO(edta).5H,O [VO(ophen)F,] [VO(dipy)F,]
42.
VO(OH),
43. 44 45.
[VO(dipy)~C1]Cl [VO(acac)~py] VO(ophen)~SO~
V--O stretch, cm -x (1035) ~} (1025) c21} 1020 (1020) (14~ (1015) ~2~ 1003, 1017, 975 (1020, 1003, 987) t14~ (1008) 9, (980) c28~ 1003 (1002) ~22~ (1001) c2z~ (1001) c22' (1001) ta~ (1000) ~s~ (998) 19} 997, 979m 996 (995) 1~2~,(1005) 1°, (995) ~a, (I000) ~24~ 996 (995) {25} 995 991 br 990 989, 959m 987, 979 987 984 982, 969, 958sh 982, 969sh 982, 960sh, 908w 981, 984sh, 908w 990, (980) ~26~,(990) I°~ 979 979 979 979, 923m 978 977, 967 (977, 967) {m (976) I1.~ 973, 963sh 973, 988sh (970) t*" 968, 964, 928m 968, 957, 949sh 932w 968, 962 (955) ~ 965, 973sh (964) I ~ 960, 923sh 911,900
Electronic structure, spectra and magnetic properties of oxycations--IV
1363
Table 1 (cont.) Compound 46. 47. 48. 49. 50. 51. 52. 53. 54.
[VO(dipy)2Br]Br-N20 (Et4N)3[VO(CN)sI [VO(DMSO)3SO4] Cs3[VO(CN)s] (MeaN)3[VO(CN)5] [VO(DMSO)5]Br2 [VO(DMSO)3CI~] (NH4)3[VOFs] [VO(DMSO)sI(C10,)z
V - - O stretch, cm -~
958, 972sh 956, 965sh 954, 964sh 953, 968sh 952 950br, sh 948, 962sh, 936sh 947, 937 932, 955, 962sh
* ABBREVIATIONS" 7 P P = tetraphenylporphine; phthal = phthalocyanine; Benzac benzoylacetonate ion; D B M = dibenzoylmethanate ion; sal = salicylaldehyde; Et = ethyl; acac = acetylacetonate ion; Etiopor = etioporphyrin I; ophen = orthophenanthroline; ox = oxalateion; Me = methyl; en-bis-sal = ethylenediamine-bis-salicylaldehyde; oxine = oxinateion; mal = malonate; edta = ethylenediaminetetraacetate ion; DMSO = dimethylsulphoxide; sh = shoulder; m = medium; w = weak; br ~ broad. All bands not otherwise marked are strong or very strong in their respective spectrum. The values listed for compounds 4, 7, 8, 9, 28 and 39 were taken from the literature cited, although the authors did not assign the values to the V - - O stretch.
listed in Table 1 are soluble in inert solvents and the V---O stretch found for the solubilized complex varied no more than 4-5 cm -1 from the value found with the solid complex. (b) It is probably quite correct to assume that coupling of the V--O stretch with vanadium-ligand modes is insignificant for ligand masses which exceed, individually, 40 a.m.u. This assertion becomes tolerable even for ligands of low mass if the vanadiumligand vibrational modes are of low frequency, as they are generally believed to be. Nonetheless, this effect could become significant in certain cases, and cannot be dismissed too lightly. Our experience, here as elsewhere, suggests to us that such coupling is relatively unimportant, and this assumption is inherent in our future discussion. The effect of individual ligand to metal bonds on the V--O stretching mode is considered below. It is presumed that the appendix wilt be consulted for further detail. a-Bonding. The vanadium-oxygen bond in vanadyl compounds is a multiple covalent bond consisting of p~r--~ dTr donation of electrons by the oxygen to the vanadium superimposed upon the a-bond (see appendix). The amount (or degree) of p~--~ dTr donation depends both upon the tendency of oxygen to donate and ~20~H. J. EICHHOFF and F. WEIGEL, Z. Anorg. Chem. 275, 267 (1954); F. A. MILLER and L. R. Cous1NS, J. Chem. Phys. 26, 329 (1957). 121~F. A. MILLER and W. K. BAER, Spectrochim. Acta 17, 112 (1961). ~2j K. UENO and A. E. MARTELL, J. Phys. Chem. 60, 934 (1956). 1~3~K. NAKAMOTO, Y. MORIMOTO and A. E. MARTELL, J. Amer. Chem. Soc. 83, 4533 (1961). ~24) K. E. LAWSON, Spectroehim. Acta 17, 248 (1961). c25~j. G. EROMAN, V. G. RAMSEY, N. W. KALENDA and W. E. HANSON, J. Amer. Chem. Soc. 78, 5844 (1956). ~zn~K. UENO and A. E. MARTELL, J. Phys. Chem. 60, 1270 (1956). ~27~D. T. SAWYER and J. M. MCKINNIE, J. Amer. Chem. Soc. 82, 4191 (1960).
1364
J. SELBIN, L. H. HOLMES, JR. and S. P. McGLYNN
vanadium to accept electrons. The donating tendency of the oxygen, which may be associated with the high electron density in' the filled valence level of the small oxygen atom/28) should be very little affected by the presence of co-ordinated ligands. However, the electron-accepting capacity of the vanadium (IV), which has four empty d-orbitals and one half-filled d-orbital, should be markedly affected by coordinated ligands. Thus, co-ordination to ligands which owe their ligational strength primarily to their a-electron pair (Lewis Base) donating ability, will increase the electron density in the metal atom d-orbitals and effect a reduction in the amount of pzr-+ d~" donation from the multiply bonded oxygen. This lowering of the V--O bond order should be reflected in a lowering of the vanadium-oxygen stretching frequency. The greater the donation of ligand electrons to the vanadium, the greater will be the decrease expected in the V--O stretching frequency. Both the a 1 and eorbitals (see appendix) are stabilized in a fashion which increases with increasing adonor capacity of the ligand and which indicates a small shift of charge toward the vanadium atom. However, the bl(L)~ b~(M) is the most important contribution from a vibrational sense, since, despite no significant effect on the VO bond strength per se, it does lead to a considerable piling-up of charge in a bl-orbital of the metal which was previously unoccupied, and thus by secondary electrostatic repulsion to a destabilization of the V--O bond. It is undoubtedly this very effect which is manifested in the decrease of the vanadyl stretch upon going from V(V) compounds such as VOCI3 to V(IV) compounds such as [VO(phthal)] where a net change of 32 cm -1 occurs. In actuality VOCI3 should be compared with a compound K2VOCI4 or K3VOC15 in which a net decrease of some 70 or 80 cm -1 might be expected. Similarly a decrease ~la) of some 90 cm-1 is observed upon going from TiO(acac)2 to VO(acac)2; there is some uncertainty associated with the TiO(acac)2 frequency because of possible polymerization of titanyl systems. Despite this it seems reasonable to assume that the population of the w-group orbital of the MO ~+ species which occurs on going from V(V) to V(IV) oxycation species, or from TiO 2+ to VOz+ is the primary reason for the decrease of the M--O stretching frequency. It has now been suggested that population of the bl component of the original zr-orbital occurs in complexes, and to an extent roughly proportional to the a-donor capacity of the ligand. It seems reasonable therefore that the decrease of V--O stretching frequency which occurs in vanadyl complexes is a rough measure of this donor capacity. 7r±-bonding. Considerations here are exactly the same as those for a-bonding, except that all overlaps are less by a factor of ,-,8, and the importance of zr±(L ~ M) donation is accordingly minimized. No M--* L contribution is significant even in unsaturated ligands. zrjTbonding. There is a contribution of charge from the ligand to the original b~ AO. This effect may be expected to amount to some ½ of the a-effect since the ~b2(L)bz(V) dr overlap (see appendix) is of the order of ½-¼ of the corresponding b 1 overlap, and as such is of importance. In sum it is concluded that the frequency changes are directly related to decreased pzr ~ dzr donation from oxygen to metal and to increased electrostatic repulsion of the vanadium and oxygen species and that this effect may be written A~ = --tr(L --~ M) -- 7rlt(L-+ M) -- ~'±(Z --~ M) Izs) R. J. GILLESPIE, J..4mer. Chem. Soc. 82, 5978 (1960).
Electronic structure, spectra and magnetic properties of oxycations--IV
1365
It is impossible to estimate accurately the relative importance of changes in the bonding between the vanadium and oxygen since the energy calculations yield only the one-electron energies of the total complex, and not of one bond region. It is estimated that the terms in the above equation are of relative importance 8 : 2 : 1, and that the M - ~ L donation is entirely neglectable in complexes of d ~ vanadium (IV). It is further difficult to estimate the effects of the fifth ligand, and the only simple conclusion is that no direct electrostatic effect is expected (see appendix). We now proceed to discuss the results obtained on the basis of the attitudes presented in the foregoing. At least eleven different compound types are discernible among the vanadyl complexes listed in Table 1. They are, with the number of representatives of each type: [VOas], 10 (only 5 different a's); [VO(AA)2a ], 7; [VO(AAAA)], 6; [VO(AA)2 ], 4; [VO(AA)(BB)], 4; [VO(AB)2], 3; [VO(AA)az], 2; [VO(ABBA)], [VO(AA)a3]; [VOa4b] and [VOa3b2], 1 each. In addition there are four compounds of the type VOX~.xH20. Among the first type we observe the following order of ligands corresponding to decreasing frequency: H20 > NCSCN > DMSO ~ F-. This order should therefore correspond to increasing electron donation ability from left to right. The unexpected extreme position of fluoride may be explained by its strong tendency to delocalize itsp-electrons away from its compact filled valence level. It has only recently been proposed t2s) that multiple bonding by fluoride (and oxide) should be expected when it is bonded to an atom or grouping which can readily accept Tr-bonding electrons. The strong repulsions between the nonbonding electron pairs located on the small fluorine atoms will cause them to be partially delocalized into empty orbitals of a bonded atom. In this way the high electron density on the relatively small atom can be relieved. It is interesting to note that the fluoro complexes of VO z+ have been found (29) to be much more stable in aqueous solution than the chloro complexes. Furthermore, the fluoro complexes of VO 2+ are much stronger than those of the similarly sized Cu 2~, Zn ~+ and Ni 2+ ions, all of which have heavily populated d-orbitals. The position of dimethylsulphoxide is surprising but perhaps not very much so in view of the large number of very stable complexes it can form with a great many transition and nontransition metals. (3°) The intermediate position of cyanide compared with its extreme position in the spectrochemical series is perhaps to be traced to the fact that a large portion of its ligational strength derives from its capacity to accept ~-bonding electrons originating on the metal ions. As already stated, the vanadyl ion would be expected to be a very poor donor metal ion as compared with later transition metal ions having more populated d-levels. Finally, the oxygen of the neutral water molecule does not suffer from the high charge density of, for example, that of the oxide ion with the possible exception of NCS , it is surely the poorest Lewis base of the series. Both of these factors must contribute to its extreme position. Turning now to the second type of complex, [VO(AA)2a], we arrive at the series ophen > real ~ ox > dipy. The relative positions of the two nitrogen bases are reversed from those found in the spectrochemical series. However, the spectrochemical series has found to hold generally for the later transition elements which can partake (2a) S. AHRLANDand B. NOREN, Aeta Chem. Scand. 12, 1595 (1958). (3o~j. SELmN and L. H. HOLMES,JR., J. hlorg. Nucl. Chem. 16, 219 (1961); F. A. CoJrON and R. FRANCIS, J. Amer. Chem. Soe. 82, 2986 (1960); F. A. COTTON, R. FRANCISand W. D. HORROC~S,J. Phys. Chem. 64, 1534 (1960).
1366
J. SELBIN,L. H. HOLMES,JR. and S. P. McGLYNN
in back-donation ~r-bonding and in this respect it is not unusual that the more highly conjugated ophen is a "stronger" ligand than dipy. However, with vanadyl, where back donation is unimportant, it might be expected that the more flexible (spatially) molecule, would form the stronger bonds. It should be pointed out that the same relative positions of these two ligands is observed for the [VO(AA)z] compounds, where the order is DBM > acac > ophen > dipy. Again, with the [VO(AA)(BB)] compounds we find ophen > dipy. Finally, the same inversion of the normal order is found tl~ in UO22+ complexes of ophen and dipy. That the structurally similar and rigidly planar four-nitrogen quadridentate bases (TPP, its para-substituted derivatives and phthalocyanine) all give vanadyl complexes with similar V--O frequencies is expected, as are the rather high values for these frequencies (1001-1015 cm-1). Furthermore, from an electrostatic viewpoint it is not unexpected that the two-oxygen bidentate ligands with a formal minus one charge (benzac, DBM, sal, and acac) should lead to higher frequencies (996-1001 cm -1) than those with a formal minus two charge (mal, ox; 976-977). In conclusion, we might point out that while the V--O stretch frequencies for a large number of complexes have been reported here, the wide variety of complex and symmetry types reduces markedly the number of complexes that can be subjected to meaningful comparison. We shall have to wait for the results of studies now in progress to draw more meaningful and extensive conclusions concerning the bonding role of various ligands in oxycation species.
Acknowledgments--The authors wish to thank the National Science Foundation (Grant No. NSF15242) and The Research Corporation for financial support. APPENDIX It will be assumed that all vanadyl complexes are of C4~ symmetry. This is undoubtedly an incorrect assumption for some, if not all, of the complexes of Table 1 ; however, the use of this approximation should not affect in any way the rather general conclusions of this section. The vibrational frequencies tabulated in Table 1 are presumed to correspond to V--O stretching vibrations. Since the primary effort is to understand the manner in which various ligands influence this frequency, and conversely, to make use of the frequency variation to determine certain characteristics of the ligands, it would seem that a theoretical approach similar to the rather thorough one of BALLHAUSENand GRAYts~ would require considerable effort but would probably be infertile. One must rather consider the vanadyl complex in a stepped fashion: (1) The electronic structure of the V--O ~+ entity must first be elaborated; (2) the equatorial ligands must next be coupled to the V--O 2+ species, and the manner in which such ligands influence the strength of the polar V--O bond delineated, and finally (3) where appropriate the fifth (polar) ligand must be introduced, and the modifications consequent to such introduction prescribed. This procedure is similar to that previously used in a discussion of uranyl complexes. ~1~ The orbitals used were the 2p AO's of the vanadyl oxygen, the 3d, 4s, and 4p AO's of the vanadium atom and the tr, 7r± and 711 GO's of the nearest neighbour equatorial ligand atoms. These latter group orbitals are easily visualized by reference to Fig. 1,
Electronic structure, spectra and magnetic properties of oxycations--IV
1367
where the axes adopted are also specified; the equatorial group orbitals and their energies are given in Table 2. Only the np-AO's of the fifth (possible) ligand are considered, where n is the highest energy partially filled or filled orbital of the nearest neighbour atom of the fifth polar ligand. The one-electron levels of a VO 2+ entity are shown on the left of Fig. 2. This level diagram is principally qualitative and represents the solution of various secular equations the integrals in which were estimated by methods set out by FELTHAM(9) Z
]
0
l
A ; J ....
"Y
2 A
l"',
I \
\
×
FIG. l . - - T h e z=axis, as also the arrows representing the ~va.AO's are perpendicular to the plane of the paper. The A's represent the nearest neighbour atom of the ligand, L. Point group is C4v.
and by MCGLYNN and SMITH.(31) In such a species the ground state configuration (C
General orbital
1
1
1
or*
~rj t
~%++
Energy§
as
b2
(x + 2fl~
3
3
bl
a2
~ -- 2/4j
t ~' ~ =±~. + ~i (_l)~+~ll~. §/4 is the resonance integral and is negative; f o r j the substitution or, ~± and ~v~ is to be made where appropriate; the relative order of integrals is ; I/4~I> I/4~r~)[> lfl~±l (~1) S. P. McGLYNN and J. K. SMITH, J. Mol. Spect. 6, 164 (1961).
J. SELBIN,L. H. HOLMES,JR. and S. P. McGLYNN
1368
0"+
Ol
I
Tr
e
-'6e
0"
al
-~J
0" "
Ol ,
I"
~
l lev
=5e
m
0z
bz
C.v V O ÷÷
/ VOL~.*
o"
~T,
T~.
SCALE
FIG. 2.--An energy level diagram ofaVO 2+species (left), a VOL"+ species (centre) and a
D4n ligand arrangement (right). L signifies a general ligand, and n may be positive or negative. This diagram is of a semi-qualitative nature, and corresponds closely to that of BALLHAUSENand GRAYc8), The energy of the ligand orbitals is naturally variable and dependent on L; the VOL4 levels are also variable but to a lesser extent. Only a sufficient number of correlation lines are shown to identify the principal interactions discussed in the text. The equatorial ligand levels are shown schematically on the right of Fig. 2, and the composite orbitals are shown in the centre of the same figure. It is not intended to go into the details of this calculation, nor to elaborate all the spectral consequences o f this particular ordering of energy levels since this will be th e object of a further paper tie) on the electronic spectra of the vanadyl complexes considered in the present work. However, it is well to point out the following: (a) The ~(C~ov) level is split into two components, b 2 and bl in C4v. The b~ level remains unchanged (increase of 0.1 eV) energetically since it interacts only with the b 2 ligand group orbital of ~r type, and the overlap is small. However, the b 1 orbital overlaps strongly with one of the a-group orbitals of the ligands, and very weakly with one of the rr±-group orbitals. As a result the bl M O is increased in energy by an amount 1.3--+ 1.8 eV dependent on the nature of the ligand. In other words we conclude that the relative positions of the b I and e M O ' s are determined by the ligands and that the exact order is capable of inversion. This is a result which must be borne
Electronic structure, spectra and magnetic properties of oxycations--lV
1369
in mind when interpreting electronic spectra, and which may be considered an adjunct to the conclusions of BALLHAUSENand GRAY. (8) (b) The ground state of the vanadyl complex is given b y . . . . b 2, ~B 2, and there are now three intra-d-shell transitions possible: 2B2 --~ 2B, 2B2 ~ ~E and 2Bz -+ ZA1. The energies of the first two transitions may invert in complexes in which the ligands are strong a-donor species, while the third transition may overlap significantly with the 2B2 -~ 2A1(4S) transition, and may very well be completely occluded by chargetransfer bands. (c) Insertion of the fifth (polar) ligand would not be expected to have serious consequences for Fig. 2. It may lead to slight destabilization or stabilization of the original V - - O bond depending on whether it is more or less electronegative than oxygen. It will generally tend to raise the energy of the e MO relative to the b~ MO.
ELECTRONIC STRUCTURE, SPECTRA A N D MAGNETIC PROPERTIES OF O X Y C A T I O N S - - I V LIGATION
EFFECTS
ON THE INFRA-RED VANADYL ION
SPECTRUM
OF THE
J. SELBIN, L. H. HOLMES, JR.* a n d S. P. MCGLYNN Coates Chemical Laboratories, Louisiana State University Baton Rouge, 3, Louisiana (Received 11 April 1963)
Abstraet~Vanadium-oxygen multiple bond stretching frequencies have been tabulated for fifty-one oxovanadium (IV) and three oxovanadium (V) compounds. The spread of the values may be summarized by ~ = 985 ± 50 cm -1, which corresponds to a force constant spread, k = 7.0 ± 0"7 mdyne/A,. The frequency shifts are used to establish several short ligand series, and they are interpreted in terms of cr and ~r contributions to the vanadium-ligand bond. It is concluded that frequency decreases are directly and primarily caused by decreased p~r -- drr donation from oxygen to metal and to increased electrostatic repulsion of the vanadium and oxygen species. The effect may he written Af = --~r(L --~ M) -- 7rll(L ~ M) -- zrj(L ~ M) where it is estimated that the relative importance of terms is 8:2:1, and it is assumed that the M--~ L donation is negligible for the d 1 vanadium (IV). IN a c o n t i n u i n g study o f the spectroscopic p r o p e r t i e s o f m e t a l oxycations/1,2) we have been investigating the o x o v a n a d i u m (IV) (vanadyl) ion.~S,4) E x p e r i m e n t a l evidence for the existence o f the v a n a d y l ion, V O 2+, in a q u e o u s solutions is strong, b u t n o t una m b i g u o u s . <5-s) H o w e v e r , FELTHAM19) f o u n d t h a t the v a n a d i u m (IV) ion in a q u e o u s solutions has visible a n d electron spin resonance spectra which are very similar to those o f the solid f i v e - c o - o r d i n a t e d v a n a d y l complexes a n d quite different f r o m those o f the m o r e s y m m e t r i c VCI 4 a n d s i x - c o - o r d i n a t e d chelates. H e c o n c l u d e d t h a t the species present in a q u e o u s solutions is VO2+.nH20. F u r t h e r s u p p o r t for the o x y c a t i o n species in preference to a m o r e s y m m e t r i c a l species such as V(OH)z 2+ is s u m m a r i z e d by BALLHAUSEN a n d GRAY. (sI A g a i n the e x p e r i m e n t a l q u a n t i t y evaluated is the p a r a m a g n e t i c r e s o n a n c e (g) value. X - r a y diffraction studies o f the solid c o m p o u n d s V O S O 4. 5 H200°) an d VO(acac)2 ~l1) clearly d e m o n s t r a t e the existence a n d stability o f the V O 2+ entity in solid complexes. This is n o t surprising when it is realized that even in VO2, which crystallizes in a highly * Present address: Southeastern Louisiana College, Hammond, Louisiana. ~1~S. P. MCGLYNrq, J. K. SMITHand W. C. NELLY, ./7. Chem. Phys. 35, 105 (1961). ~2~B. SHAMaURGEg,M. S. Thesis, Louisiana State University (1960). (sl j. SELBIN, S. P. McGLYNN, J. K. SMITH,G. A. SIBILLE,and L. J. TaUEX. Paper presented at the 16th Southwest Regional ACS-Meeting, Oklahoma City, Okla., Dec. (1960). ~4) j. SELBIN,L. H. HOLMES,JR. and S. P. McGLYNN, Chem. andIndustr. 746 (1961). (~ L. P. DUCgET. Thesis, Paris University (1951). ~c,~F. J. C. RossovTI and H. S. ROSSOTTI,Acta Chem. Scand. 9, 1177 (1955). irt C. K. JORGENSEN,Acta Chem. Stand. 11, 73 (1957). 181 C. J. BALLHAUSENand H. B. GRAY, Inorg. Chem. 1,111 (1962). 19~R. D. FELTHAM. Thesis, University of California, UCRL-3867 (1957). I10) M . B. PALMA-VITTOI~ELLI, M . U . PALMA and F. SGARLATA, Nttol)o Cimento (Series 10) 3, 718 (1956); G. LUNDGaEN,Rec. Tray. Chim. 75, 585 (1956). (11~R. P. DODGE, D. H. TEMPLETONand A. ZALKIN, J. Chem. Phys., 35, 55 (1961). 1359
1360
J. S~LBIN,L. H. HOLMES,JR. and S. P. McGLYNN
distorted rutile structure, there is one short V--O bond in each VOe unit. ~12) Chemical analysis and molecular weight determination have also established that the formula of the acetylacetonate (acac) complex is VO(acac)2 and that it is monomeric in benzene.t13~ Indirect evidence for the mononuclear monoxy VO 2+ species may be deduced from the infra-red spectra of vanadyl compounds. BARRACLOUGHe t al. ~14) examined the infra-red spectra of several complex metal compounds which were presumed to contain metal-oxygen multiple bonds. They reported that the stretching frequencies of metal-oxygen double bonds may generally be found between 900 and 1100 cm -1. Specifically they reported probable V--O stretching vibrations for the following four compounds: VOSO4, 987, 1003, 1020; VO(acac)2, 995; (NHa)2[VO(ox)2].2HzO, 976; and (NH4)2[VO(mal)2].4H20, 967, 977 (see footnote to Table 1 for abbreviations used). Earlier, vanadium (IV) hydroxide had been formulated as VO(OH) 2 on the basis of its infra-red spectrum, ~15~which shows a strong band at 955 cm -1, attributed to the V--O stretching vibration. It is apparent from the above values for band maxima that the environment of the VO 2+ entity conditions a marked effect upon the stretching frequency, and hence upon the force constant and strength of the V--O bond. This is not unexpected and, in fact, MCGLYNN and co-workers tl~ have shown that the infra-red bands due to the dioxouranium (VI) symmetric and antisymmetric stretching vibrations may be used as a probe to establish a ligand series exhibiting a striking parallelism with the spectrochemical series. The persistence of the oxovanadium (IV) species in a number of compounds makes possible the utilization of the properties of the oxycation species in a study of the bonding structure of complexes of this d 1 electron system. Ideally, one might expect not only to be able to study the symmetry types available to this transition metal ion, but also to learn more about the nature of the bonding to various ligated species as they perturb the VO 2+ entity. Thus the stretching frequency of the metal-oxygen bond should serve as a sensitive probe from which certain structural and bonding features may be deduced. Furthermore, the oxymetal species will show electronic localized charge-transfer absorption bands (e.g., arising from electron transfer from orbitals located primarily on the oxygen to orbitals primarily located on the metal) in addition to those normally observed with d n ions. These additional bands may be expected to be independent upon the other ligands and the strength and symmetry of the ligand field. In this paper we shall be concerned only with the interpretation of the infra-red spectra of vanadyl compounds. A subsequent paper tle~ will deal in detail with the electronic spectra. The electronic energy levels in the vanadyl entity and in its complexes have been deduced from a combined crystal-field and molecular orbital approach and it is determined (see Appendix) that there is no basic disagreement with the molecular orbital scheme of BALLHAUSENand GRAY. (8) ~z) B. ANDERSON,Acta Chem. Seand. 10, 623 (1956). ~ls~M. M. JONESJ. Amer. Chem. Soc. 76, 5995 (1954). ~4~ C. G. BARRACLOUOH,J. LEWISand R. S. NYHOLM,J. Chem. Soc. 3552 (1959). tlS~ C . CABANNEs-OTT,C.R. Acad. Sci., Paris 242, 2825 (1956). tx6, S. P. McGLYNN,J. SELBIN,T. R. ORTOLAND,G. MAUS,and W. JORDAN,in preparation for publication.
Electronic structure, spectra and magnetic properties of oxycations--lV
1361
It appears that the v a n a d i u m in the vanadyl m a y be either five- or six-coordinated (13,17) and it is n o t always certain which obtains in a given c o m p o u n d . However, chemical analysis coupled with molecular weight and/or conductance data will generally suffice to settle this point. (18~ In sum, we have examined the infra-red spectra o f thirty-nine vanadyl complexes, thirty o f which are new c o m p o u n d s whose preparation and properties are reported elsewhere. (is) Stretching frequencies o f the perturbed VO2+ species have been assigned, certain ligand series established and discussed. EXPERIMENTAL The new compounds studied in this investigation were prepared and characterized according to procedures previously outlined. (18) Known compounds were prepared according to procedures found in the literature. Starting materials included reagent grade vanadyl sulphate, vanadyl chloride, ammonium metavanadate and vanadium (V) oxide. All other chemicals used were of reagent grade quality or better. Infrared spectra in the range 5000-650 cm -1 were obtained with a Perkin-Elmer Model 21 and a Beckman IR-7 recording spectrophotometer, both employing NaC1 optics. Polystyrene film was used to calibrate the instruments. Spectra were run with the samples in Nujol mulls since these gave better spectra than the alkali halide pellet technique. Interaction of some of the compounds with the alkali halide under the conditions necessary for pellet preparation was evident. Spectra of all ligands and/ or complexes of other transition metal ions with the ligands were obtained and using these, there was usually no difficulty in assigning the band(s) of the perturbed VO~+ stretching vibration. In those cases in which the ligands have bands in the region of the metal-oxygen stretch, comparisons were made with the spectra of analogous complexes of other divalent metal ions of the first transition series. Infrared spectra of many of the vanadyl complexes studied in this work were run in solution in a concurrent investigation to be reported elsewhere. (19) In all of those cases where solubility was effected in an inert solvent the value for the V--O stretch varied no more that ±5 cm -1 from the value found with the solid complex. RESULTS AND DISCUSSION Table 1 lists the frequencies assigned to the v a n a d i u m - o x y g e n multiple b o n d stretching vibration for fifty-one o x o v a n a d i u m (IV) and three o x o v a n a d i u m (V) c o m p o u n d s . In those cases where more than one b a n d was observed in the metaloxygen multiple b o n d stretch region, the strongest is listed first. The values enclosed by parentheses were taken f r o m the literature cited. The force constants, k, have been related to the frequency using a simple h a r m o n i c expression, and are f o u n d to vary f r o m 7.68 for VOCI 3 to 6.23 for [VO(DMSO)5 ] (C104) 2 in mdyne/A. The listing in Table 1 m a y then be approximately summarized by the single values k = 7.0 :k 0.7 and ~ = 985 ~ 50 cm-1; it behoves the authors to illustrate then that the spread of the effect being observed is larger than that which could arise f r o m (a) lattice forces or (b) simple mechanical coupling to the metal-ligand vibrations: (a) Certainly, some o f the shift observed to occur in the V - - O stretch frequency can undoubtedly be caused by the operation o f lattice forces which are expected to vary f r o m one solid c o m p o u n d to the next. However, the magnitude o f the crystallineenvironment effect is considered to be small. Specifically, some o f the c o m p o u n d s ~17)R. T. CLAUNCH,T. W. MARTIN and M. M. JONES,J. Amer. Chem. Soe. 83, 1073 (1961). (lS) j. SELBINand L. H. HOLMES,JR., J. Inorg. Nucl. Chem. 24, 1111 (1962). ,9) j. SELBIN,H. R. MANNINGand G. CESSAC,presented at the 17th Southwest Regional A.C.S. Meeting, Dallas, Texas, December, 1962, J. Inorg. Nucl. Chem. In Press.
1362
J. SELBIN,L. H. HOLMES,JR. and S. P. McGLYNN
TABLE 1.--VANADIUM-OXYGEN STRETCHING FREQUENCIES, OXOVANADIUM(IV) AND (V) COMPOUNDS
Compound* 1. 2. 3. 4. 5.
VOCIa VOBr3 V~O5 [VO(p-OCHs-TPP)] VOSO¢5H20
6. 7. 8. 9. 10. 11. 12. 13. 14.
[VO(phthal)] [VO(p-CI-TPP)] [VO(TPP)] [VO(p-CHa-TPP)] [VO(Benzac) 2] [VO(DBM)~] [VO(sal) 8] (Et,N)a[VOCNCS)6] [VO(acac)~]
15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41.
VOBrcH20 [VO(Etiopor)] VO(C104)2"xH20 (NH,)2(VO)2(SO4)2 VOClz.xH~O [VO(ophen) (ox)] (Me4N)s[VO(NCS)5] [VO(ophen)~](C104)~ VOF~-xH20 (Me,N)2[VOCI,(EtOH)] K3[VO(NCS)5] [VO(ophen)~C1]Cl [VO(ophen)2Br]Br.HzO [VO(en-bis-sal)] [VO(oxine)~] [VO(dipy)SO,] [VO(dipy) (ox)] [VO(dipy)~](C10~)~ WO(ophen)SOd VO(dipy)l.5(NCS)z (NH,)2VO(mal)2"4H~O (NH4),VO(ox),.2H,O VO(ophen)l.2(NCS), VO(ox)-H~ox'3H,O Na,VO(edta).5H,O [VO(ophen)F,] [VO(dipy)F,]
42.
VO(OH),
43. 44 45.
[VO(dipy)~C1]Cl [VO(acac)~py] VO(ophen)~SO~
V--O stretch, cm -x (1035) ~} (1025) c21} 1020 (1020) (14~ (1015) ~2~ 1003, 1017, 975 (1020, 1003, 987) t14~ (1008) 9, (980) c28~ 1003 (1002) ~22~ (1001) c2z~ (1001) c22' (1001) ta~ (1000) ~s~ (998) 19} 997, 979m 996 (995) 1~2~,(1005) 1°, (995) ~a, (I000) ~24~ 996 (995) {25} 995 991 br 990 989, 959m 987, 979 987 984 982, 969, 958sh 982, 969sh 982, 960sh, 908w 981, 984sh, 908w 990, (980) ~26~,(990) I°~ 979 979 979 979, 923m 978 977, 967 (977, 967) {m (976) I1.~ 973, 963sh 973, 988sh (970) t*" 968, 964, 928m 968, 957, 949sh 932w 968, 962 (955) ~ 965, 973sh (964) I ~ 960, 923sh 911,900
Electronic structure, spectra and magnetic properties of oxycations--IV
1363
Table 1 (cont.) Compound 46. 47. 48. 49. 50. 51. 52. 53. 54.
[VO(dipy)2Br]Br-N20 (Et4N)3[VO(CN)sI [VO(DMSO)3SO4] Cs3[VO(CN)s] (MeaN)3[VO(CN)5] [VO(DMSO)5]Br2 [VO(DMSO)3CI~] (NH4)3[VOFs] [VO(DMSO)sI(C10,)z
V - - O stretch, cm -~
958, 972sh 956, 965sh 954, 964sh 953, 968sh 952 950br, sh 948, 962sh, 936sh 947, 937 932, 955, 962sh
* ABBREVIATIONS" 7 P P = tetraphenylporphine; phthal = phthalocyanine; Benzac benzoylacetonate ion; D B M = dibenzoylmethanate ion; sal = salicylaldehyde; Et = ethyl; acac = acetylacetonate ion; Etiopor = etioporphyrin I; ophen = orthophenanthroline; ox = oxalateion; Me = methyl; en-bis-sal = ethylenediamine-bis-salicylaldehyde; oxine = oxinateion; mal = malonate; edta = ethylenediaminetetraacetate ion; DMSO = dimethylsulphoxide; sh = shoulder; m = medium; w = weak; br ~ broad. All bands not otherwise marked are strong or very strong in their respective spectrum. The values listed for compounds 4, 7, 8, 9, 28 and 39 were taken from the literature cited, although the authors did not assign the values to the V - - O stretch.
listed in Table 1 are soluble in inert solvents and the V---O stretch found for the solubilized complex varied no more than 4-5 cm -1 from the value found with the solid complex. (b) It is probably quite correct to assume that coupling of the V--O stretch with vanadium-ligand modes is insignificant for ligand masses which exceed, individually, 40 a.m.u. This assertion becomes tolerable even for ligands of low mass if the vanadiumligand vibrational modes are of low frequency, as they are generally believed to be. Nonetheless, this effect could become significant in certain cases, and cannot be dismissed too lightly. Our experience, here as elsewhere, suggests to us that such coupling is relatively unimportant, and this assumption is inherent in our future discussion. The effect of individual ligand to metal bonds on the V--O stretching mode is considered below. It is presumed that the appendix wilt be consulted for further detail. a-Bonding. The vanadium-oxygen bond in vanadyl compounds is a multiple covalent bond consisting of p~r--~ dTr donation of electrons by the oxygen to the vanadium superimposed upon the a-bond (see appendix). The amount (or degree) of p~--~ dTr donation depends both upon the tendency of oxygen to donate and ~20~H. J. EICHHOFF and F. WEIGEL, Z. Anorg. Chem. 275, 267 (1954); F. A. MILLER and L. R. Cous1NS, J. Chem. Phys. 26, 329 (1957). 121~F. A. MILLER and W. K. BAER, Spectrochim. Acta 17, 112 (1961). ~2j K. UENO and A. E. MARTELL, J. Phys. Chem. 60, 934 (1956). 1~3~K. NAKAMOTO, Y. MORIMOTO and A. E. MARTELL, J. Amer. Chem. Soc. 83, 4533 (1961). ~24) K. E. LAWSON, Spectroehim. Acta 17, 248 (1961). c25~j. G. EROMAN, V. G. RAMSEY, N. W. KALENDA and W. E. HANSON, J. Amer. Chem. Soc. 78, 5844 (1956). ~zn~K. UENO and A. E. MARTELL, J. Phys. Chem. 60, 1270 (1956). ~27~D. T. SAWYER and J. M. MCKINNIE, J. Amer. Chem. Soc. 82, 4191 (1960).
1364
J. SELBIN, L. H. HOLMES, JR. and S. P. McGLYNN
vanadium to accept electrons. The donating tendency of the oxygen, which may be associated with the high electron density in' the filled valence level of the small oxygen atom/28) should be very little affected by the presence of co-ordinated ligands. However, the electron-accepting capacity of the vanadium (IV), which has four empty d-orbitals and one half-filled d-orbital, should be markedly affected by coordinated ligands. Thus, co-ordination to ligands which owe their ligational strength primarily to their a-electron pair (Lewis Base) donating ability, will increase the electron density in the metal atom d-orbitals and effect a reduction in the amount of pzr-+ d~" donation from the multiply bonded oxygen. This lowering of the V--O bond order should be reflected in a lowering of the vanadium-oxygen stretching frequency. The greater the donation of ligand electrons to the vanadium, the greater will be the decrease expected in the V--O stretching frequency. Both the a 1 and eorbitals (see appendix) are stabilized in a fashion which increases with increasing adonor capacity of the ligand and which indicates a small shift of charge toward the vanadium atom. However, the bl(L)~ b~(M) is the most important contribution from a vibrational sense, since, despite no significant effect on the VO bond strength per se, it does lead to a considerable piling-up of charge in a bl-orbital of the metal which was previously unoccupied, and thus by secondary electrostatic repulsion to a destabilization of the V--O bond. It is undoubtedly this very effect which is manifested in the decrease of the vanadyl stretch upon going from V(V) compounds such as VOCI3 to V(IV) compounds such as [VO(phthal)] where a net change of 32 cm -1 occurs. In actuality VOCI3 should be compared with a compound K2VOCI4 or K3VOC15 in which a net decrease of some 70 or 80 cm -1 might be expected. Similarly a decrease ~la) of some 90 cm-1 is observed upon going from TiO(acac)2 to VO(acac)2; there is some uncertainty associated with the TiO(acac)2 frequency because of possible polymerization of titanyl systems. Despite this it seems reasonable to assume that the population of the w-group orbital of the MO ~+ species which occurs on going from V(V) to V(IV) oxycation species, or from TiO 2+ to VOz+ is the primary reason for the decrease of the M--O stretching frequency. It has now been suggested that population of the bl component of the original zr-orbital occurs in complexes, and to an extent roughly proportional to the a-donor capacity of the ligand. It seems reasonable therefore that the decrease of V--O stretching frequency which occurs in vanadyl complexes is a rough measure of this donor capacity. 7r±-bonding. Considerations here are exactly the same as those for a-bonding, except that all overlaps are less by a factor of ,-,8, and the importance of zr±(L ~ M) donation is accordingly minimized. No M--* L contribution is significant even in unsaturated ligands. zrjTbonding. There is a contribution of charge from the ligand to the original b~ AO. This effect may be expected to amount to some ½ of the a-effect since the ~b2(L)bz(V) dr overlap (see appendix) is of the order of ½-¼ of the corresponding b 1 overlap, and as such is of importance. In sum it is concluded that the frequency changes are directly related to decreased pzr ~ dzr donation from oxygen to metal and to increased electrostatic repulsion of the vanadium and oxygen species and that this effect may be written A~ = --tr(L --~ M) -- 7rlt(L-+ M) -- ~'±(Z --~ M) Izs) R. J. GILLESPIE, J..4mer. Chem. Soc. 82, 5978 (1960).
Electronic structure, spectra and magnetic properties of oxycations--IV
1365
It is impossible to estimate accurately the relative importance of changes in the bonding between the vanadium and oxygen since the energy calculations yield only the one-electron energies of the total complex, and not of one bond region. It is estimated that the terms in the above equation are of relative importance 8 : 2 : 1, and that the M - ~ L donation is entirely neglectable in complexes of d ~ vanadium (IV). It is further difficult to estimate the effects of the fifth ligand, and the only simple conclusion is that no direct electrostatic effect is expected (see appendix). We now proceed to discuss the results obtained on the basis of the attitudes presented in the foregoing. At least eleven different compound types are discernible among the vanadyl complexes listed in Table 1. They are, with the number of representatives of each type: [VOas], 10 (only 5 different a's); [VO(AA)2a ], 7; [VO(AAAA)], 6; [VO(AA)2 ], 4; [VO(AA)(BB)], 4; [VO(AB)2], 3; [VO(AA)az], 2; [VO(ABBA)], [VO(AA)a3]; [VOa4b] and [VOa3b2], 1 each. In addition there are four compounds of the type VOX~.xH20. Among the first type we observe the following order of ligands corresponding to decreasing frequency: H20 > NCSCN > DMSO ~ F-. This order should therefore correspond to increasing electron donation ability from left to right. The unexpected extreme position of fluoride may be explained by its strong tendency to delocalize itsp-electrons away from its compact filled valence level. It has only recently been proposed t2s) that multiple bonding by fluoride (and oxide) should be expected when it is bonded to an atom or grouping which can readily accept Tr-bonding electrons. The strong repulsions between the nonbonding electron pairs located on the small fluorine atoms will cause them to be partially delocalized into empty orbitals of a bonded atom. In this way the high electron density on the relatively small atom can be relieved. It is interesting to note that the fluoro complexes of VO z+ have been found (29) to be much more stable in aqueous solution than the chloro complexes. Furthermore, the fluoro complexes of VO 2+ are much stronger than those of the similarly sized Cu 2~, Zn ~+ and Ni 2+ ions, all of which have heavily populated d-orbitals. The position of dimethylsulphoxide is surprising but perhaps not very much so in view of the large number of very stable complexes it can form with a great many transition and nontransition metals. (3°) The intermediate position of cyanide compared with its extreme position in the spectrochemical series is perhaps to be traced to the fact that a large portion of its ligational strength derives from its capacity to accept ~-bonding electrons originating on the metal ions. As already stated, the vanadyl ion would be expected to be a very poor donor metal ion as compared with later transition metal ions having more populated d-levels. Finally, the oxygen of the neutral water molecule does not suffer from the high charge density of, for example, that of the oxide ion with the possible exception of NCS , it is surely the poorest Lewis base of the series. Both of these factors must contribute to its extreme position. Turning now to the second type of complex, [VO(AA)2a], we arrive at the series ophen > real ~ ox > dipy. The relative positions of the two nitrogen bases are reversed from those found in the spectrochemical series. However, the spectrochemical series has found to hold generally for the later transition elements which can partake (2a) S. AHRLANDand B. NOREN, Aeta Chem. Scand. 12, 1595 (1958). (3o~j. SELmN and L. H. HOLMES,JR., J. hlorg. Nucl. Chem. 16, 219 (1961); F. A. CoJrON and R. FRANCIS, J. Amer. Chem. Soe. 82, 2986 (1960); F. A. COTTON, R. FRANCISand W. D. HORROC~S,J. Phys. Chem. 64, 1534 (1960).
1366
J. SELBIN,L. H. HOLMES,JR. and S. P. McGLYNN
in back-donation ~r-bonding and in this respect it is not unusual that the more highly conjugated ophen is a "stronger" ligand than dipy. However, with vanadyl, where back donation is unimportant, it might be expected that the more flexible (spatially) molecule, would form the stronger bonds. It should be pointed out that the same relative positions of these two ligands is observed for the [VO(AA)z] compounds, where the order is DBM > acac > ophen > dipy. Again, with the [VO(AA)(BB)] compounds we find ophen > dipy. Finally, the same inversion of the normal order is found tl~ in UO22+ complexes of ophen and dipy. That the structurally similar and rigidly planar four-nitrogen quadridentate bases (TPP, its para-substituted derivatives and phthalocyanine) all give vanadyl complexes with similar V--O frequencies is expected, as are the rather high values for these frequencies (1001-1015 cm-1). Furthermore, from an electrostatic viewpoint it is not unexpected that the two-oxygen bidentate ligands with a formal minus one charge (benzac, DBM, sal, and acac) should lead to higher frequencies (996-1001 cm -1) than those with a formal minus two charge (mal, ox; 976-977). In conclusion, we might point out that while the V--O stretch frequencies for a large number of complexes have been reported here, the wide variety of complex and symmetry types reduces markedly the number of complexes that can be subjected to meaningful comparison. We shall have to wait for the results of studies now in progress to draw more meaningful and extensive conclusions concerning the bonding role of various ligands in oxycation species.
Acknowledgments--The authors wish to thank the National Science Foundation (Grant No. NSF15242) and The Research Corporation for financial support. APPENDIX It will be assumed that all vanadyl complexes are of C4~ symmetry. This is undoubtedly an incorrect assumption for some, if not all, of the complexes of Table 1 ; however, the use of this approximation should not affect in any way the rather general conclusions of this section. The vibrational frequencies tabulated in Table 1 are presumed to correspond to V--O stretching vibrations. Since the primary effort is to understand the manner in which various ligands influence this frequency, and conversely, to make use of the frequency variation to determine certain characteristics of the ligands, it would seem that a theoretical approach similar to the rather thorough one of BALLHAUSENand GRAYts~ would require considerable effort but would probably be infertile. One must rather consider the vanadyl complex in a stepped fashion: (1) The electronic structure of the V--O ~+ entity must first be elaborated; (2) the equatorial ligands must next be coupled to the V--O 2+ species, and the manner in which such ligands influence the strength of the polar V--O bond delineated, and finally (3) where appropriate the fifth (polar) ligand must be introduced, and the modifications consequent to such introduction prescribed. This procedure is similar to that previously used in a discussion of uranyl complexes. ~1~ The orbitals used were the 2p AO's of the vanadyl oxygen, the 3d, 4s, and 4p AO's of the vanadium atom and the tr, 7r± and 711 GO's of the nearest neighbour equatorial ligand atoms. These latter group orbitals are easily visualized by reference to Fig. 1,
Electronic structure, spectra and magnetic properties of oxycations--IV
1367
where the axes adopted are also specified; the equatorial group orbitals and their energies are given in Table 2. Only the np-AO's of the fifth (possible) ligand are considered, where n is the highest energy partially filled or filled orbital of the nearest neighbour atom of the fifth polar ligand. The one-electron levels of a VO 2+ entity are shown on the left of Fig. 2. This level diagram is principally qualitative and represents the solution of various secular equations the integrals in which were estimated by methods set out by FELTHAM(9) Z
]
0
l
A ; J ....
"Y
2 A
l"',
I \
\
×
FIG. l . - - T h e z=axis, as also the arrows representing the ~va.AO's are perpendicular to the plane of the paper. The A's represent the nearest neighbour atom of the ligand, L. Point group is C4v.
and by MCGLYNN and SMITH.(31) In such a species the ground state configuration (C
General orbital
1
1
1
or*
~rj t
~%++
Energy§
as
b2
(x + 2fl~
3
3
bl
a2
~ -- 2/4j
t ~' ~ =±~. + ~i (_l)~+~ll~. §/4 is the resonance integral and is negative; f o r j the substitution or, ~± and ~v~ is to be made where appropriate; the relative order of integrals is ; I/4~I> I/4~r~)[> lfl~±l (~1) S. P. McGLYNN and J. K. SMITH, J. Mol. Spect. 6, 164 (1961).
J. SELBIN,L. H. HOLMES,JR. and S. P. McGLYNN
1368
0"+
Ol
I
Tr
e
-'6e
0"
al
-~J
0" "
Ol ,
I"
~
l lev
=5e
m
0z
bz
C.v V O ÷÷
/ VOL~.*
o"
~T,
T~.
SCALE
FIG. 2.--An energy level diagram ofaVO 2+species (left), a VOL"+ species (centre) and a
D4n ligand arrangement (right). L signifies a general ligand, and n may be positive or negative. This diagram is of a semi-qualitative nature, and corresponds closely to that of BALLHAUSENand GRAYc8), The energy of the ligand orbitals is naturally variable and dependent on L; the VOL4 levels are also variable but to a lesser extent. Only a sufficient number of correlation lines are shown to identify the principal interactions discussed in the text. The equatorial ligand levels are shown schematically on the right of Fig. 2, and the composite orbitals are shown in the centre of the same figure. It is not intended to go into the details of this calculation, nor to elaborate all the spectral consequences o f this particular ordering of energy levels since this will be th e object of a further paper tie) on the electronic spectra of the vanadyl complexes considered in the present work. However, it is well to point out the following: (a) The ~(C~ov) level is split into two components, b 2 and bl in C4v. The b~ level remains unchanged (increase of 0.1 eV) energetically since it interacts only with the b 2 ligand group orbital of ~r type, and the overlap is small. However, the b 1 orbital overlaps strongly with one of the a-group orbitals of the ligands, and very weakly with one of the rr±-group orbitals. As a result the bl M O is increased in energy by an amount 1.3--+ 1.8 eV dependent on the nature of the ligand. In other words we conclude that the relative positions of the b I and e M O ' s are determined by the ligands and that the exact order is capable of inversion. This is a result which must be borne
Electronic structure, spectra and magnetic properties of oxycations--lV
1369
in mind when interpreting electronic spectra, and which may be considered an adjunct to the conclusions of BALLHAUSENand GRAY. (8) (b) The ground state of the vanadyl complex is given b y . . . . b 2, ~B 2, and there are now three intra-d-shell transitions possible: 2B2 --~ 2B, 2B2 ~ ~E and 2Bz -+ ZA1. The energies of the first two transitions may invert in complexes in which the ligands are strong a-donor species, while the third transition may overlap significantly with the 2B2 -~ 2A1(4S) transition, and may very well be completely occluded by chargetransfer bands. (c) Insertion of the fifth (polar) ligand would not be expected to have serious consequences for Fig. 2. It may lead to slight destabilization or stabilization of the original V - - O bond depending on whether it is more or less electronegative than oxygen. It will generally tend to raise the energy of the e MO relative to the b~ MO.