Jan., 1957
ELECTRON DEFICIENT COMPOUNDS
observed when one compares the methyl affinities of quinone and ch loroq~inone,~ naphthoquinone and 2-chloronaphth o q u i n ~ n e , maleic ~ anhydride and chloromaleic anhydride16 and so forth. On the other hand, the electron donating groups seem t o decrease the reactivity, compare, e.g., methyl affinities of toluquinone and metho~yquinone.~ At this stage i t is necessary to remark that steric hindrance is also a factor determining the reactivity. Blocking of “reactive points” decreases the reactivity. This effect is seen when one compares the reactivities of styrene, the diphenyl ethylenes, triphenyl ethylene, and tetraphenyl ethylenea6 The steric effect is even more dramatically manifested in the series of chlorinated derivatives. Whereas one chlorine activates the double bond of a quinone, and the activation is even greater when each of the double bonds contains chlorine (e.g., 2,5- or 2,6-dichloroquinones), the tetrachloroquinone is unreactive.a The same observations hold in other series. For ‘example, vinyl chloride is more reactive than ethylene, vinylidene chloride is even more reactive, but tetrachloroethylene is unreactive. That this phenomenon is due to steric hindrance is shown well when one compares the very high methyl affinity of tetrafluoroethylene (a small substituent) with the low reactivity of tetrachloroethylene (a large substituent). Finally, let us consider how the changes in the slope of the repulsion curve affect the intrinsic re(15) Unpublished data of Leavitt from our laboratories.
t
45 EXCITED MOLECULE
ENERGY OF THE SYSTEM
W DISTANCE A
-R
_.
Fig. 9.
activity of the radical. Inspection of Fig. 9 shows that for a flatter repulsion curve the coefficient a! in the relation A E a c t = aAEe,is lower. Since the reciprocal of a! measures the intrinsic reactivity of a radical, this means that a flatter repulsion curve corresponds to a higher intrinsic reactivity. In conclusion, it is hoped that the studies of methyl affinities and of similar reactions will be helpful in understanding of factors influencing the rates of radical addition processes.
ELECTRON DEFICIENT COMPOUNDS1v2 BY R. E. RUNLLE Contribution No. 478 from the Department of Chemistry and Institute for Atomic Research, Iowa State College, Ames, Iowa Received Auoust IS. 1966
There are a number of com ounds for which the number of chemical bonds exceeds the number of valence electron pairs, presenting a serious problem g r classical valence theory. Historically, the boron hydrides provided the first example, but the problem is general t o chemistry. Indeed, the first “electron deficient compound” to be given adequate structural study WBS the tetramethylplatinum tetramer, in which a quarter of the methyl groups are bonded equally t o three platinums. The “cause” of electron deficient bonding is easily stated. When atoms, inevitably metallic, with fewer valence electrons than low energy bonding orbitals form compounds with atoms or groups containing no unshared electron pairs, bonding is delocalized to make use of all of the low energy metallic orbitals. This arises from the same quantum mechanical principles that lead to chemical bonding, and has perhaps been best treated for the Hs+ ion, which serves as a prototype for electron deficient compounds. It has been possible to predict where new electron deficient compounds are to be sought, but understanding of structure and compositions, both in the bpron hydrides and generally, has come from structural work. The geometry of methyl bridges is singled out for special discussion. Delocalization of bonding arising from “excess” orbitals, always from metal atoms, provides an interesting correlation running from organic compounds, through organometallic compounds and interstitial compounds, to metals themselves, delocalization increasing with fraction of “excess” orbitals. Certain special features of the generalization giving rise to weak metal-metal bonds in dspa-square transition metal complexes will be noted. One might term this latter bonding ‘‘configuration interact,ion bonding.”
Introduction The formulation of the electron pair bond3 and the discovery of the boron hydrides4 occurred almost simultaneously, so that some of the first attempts to write bond structures for BzHe were uninhibited by ideas of what a bond was. The correct geometry was, thereby, included in the many early (1) Work was performed in t h e Ames Laboratory of the Atomic Energy Commission, (2) Presented a t the Symposium on Valence and Chemical Bonding, Madison, Wisconsin, June 22, 1956. (3) G. N. Lewis, J. A m . Chem. Soc., 38, 7 6 2 (1918). (4) A. Stock and K. Friederici, Ber., 46, 1959 (1913).
proposal^.^ But it soon became evident that the boron hydrides presented a difficult problem for the electron pair theory of valence, since in these cnmpounds the number of bonds is required to exceed the number of valence electron pairs, hence the term “electron deficient compounds.” Though there are some who still write as though boron hydrides were unique in this excess of bonds to electron pairs, the electron deficient bonding problem was extended to the rest of the third group elements by the discovery that trimethylalu(5) W.Dilthey, 2. angew. Chum., 34,696 (1921).
R. E. RUNDLE
46
minum was a dimer.6 The problem was further extended t o chemistry generally when tetramethylplatinum was found t o be a tetramer.’ It would be unfair t o review the work of early theorists in the field of the boron hydrides. They were misled by wrong structures “determined” by over optimistic experimenters. Nor would it be fair to review the work of early experimenters who were misled by theorists into believing they’ could determine n parameters from n - m pieces of data if they applied a little theory to reduce the number of parameters. Experiment and theory thereby confirmed each other and, with all problems apparently solved, progress in the field ceased until a reevaluation was initiated many years later.8 One historical footnote might be added, however. The field of electron deficient compounds is one to make the chemical theorists blush. For the most part they followed the experimentalist, reconciling whatever the experimentalist found or thought he found with theory. Nevertheless, the-
@=Pr @ Fig. 1.-Structure
- CL IN
ME IN
P T M ECL~
ME,
of tetramers of tetramethylplatinum and trimethylplatinum chloride.
P 121.5”
Fig. 2.-Structure
of BzH6.
(6) A. W. Laubengayer and W. F. Gilliam, J . Am. Chem. Soc., 68, 477 (1941). Earlier work, indicative of dimers, is referred to in the above article. (7) R. E.Rundle and J. H. Sturdivant. ibid., 69, 1561 (1947). (8)B. V. Nekrasov, J. Oen. Chem. (U.R.S.S.), 10, 1021, 1156
(1940).
Vol. 61
ory was sufficiently developed that it could have led instead of following. It was suggested that BzHs should be similar to isoelectronic ethylene. Group theoretical arguments strongly reinforce this suggestion, since a BzH6 molecule with ethanesymmetry (D3h or D3d) should almost surely have a triplet ground state, while with ethylene-symmetry (h) it should have a singlet ground state as observed. This argument was, indeed, made early enough t o have changed the history of the boron hydrides,1° but was so well qualified as to be considered less significant than doubtful experimental data. Some theorists were overly cautious and some experimentalists undercautious during a critical period for this subject. The first electron deficient compound whose structure was surely known was the tetramethylplatinum tetramer (Fig. l),’ though some shrewd examinations of spectral and other data made the bridge structure for BzH6 (Fig. 2) seem probab1e.l’ This model for BZH6 was finally established in 1948 by W. C. Price by an infrared study,12 but bond angles and distances were not obtained by the much maligned electron diffraction method until 1951.13 Bond Delocalization Cause and Consequences.-The structure of tetramethylplatinum, in which a single methyl group is bonded equally to three platinum atoms, does violence to the time-honored tetravalence of carbon. Another look at the structure reveals, however, the familiar octahedral configuration about platinum(IV), and this feature suggests the principle (below) which leads to electron deficient, or better, delocalized bonding. For example, the common feature of B2Ha, the trimethylaluminum dimer and the tetramethylplatinum tetramer is a n atom (usually metallic) with more low energy orbitals than valence electrons combined with atoms or groups containing n o unshared electron pairs, i.e., four sp3-orbital us. three valence electrons for boron and aluminum, six d2sp3-orbitals us. four valence electrons for platinum (IV). It appears that under these circumstances delocalixation of bonding occurs so as to use all of the low energy bonding orbitals of the metallic atom.14 Such a generalization is easily justified by quantum arguments. 14b The valence bond approach had already been given by Hirschfelder, Rosen and Eyring in considering H3+,16where one electron pair serves to bond three atoms together very strongly. H3+ serves as a prototype for electron deficient bonding, just as Hz furnishes a model of the electron pair bond, both valence bond and molecular orbital approaches (9) E. Wiberg, Bet-., 69, 2821 (1936). Also reviews all other Buggestions very completely. (10) R. S. Mulliken, Phys. Rev., 43,765 (1933). (11) Ref. 8. See also (a) Y. K. Syrkin and M. E. Dyatkinrt, A c t o Physicochim. (U.R.S.S.), 14, 547 (1941); Compt. Tend. acad. aci. (U.R.S.S.), 86,180 (1942); (b) H.C. Longuet-Higgins and R. P. Bell, J. Chem. Soc., 250 (1943); (c) K.9. Pitzer, J. Am. Chem. SOC.,67, 1126
(1945). (12) W.C. Price, J. Chem. Plrys., 16, 894 (1948). (13) I
