Jack Edwin Bisseyl
California State College LOS Angeles, 90032
Some Aspects of d-Orbital Participation in Phosphorous and Silicon Chemistry
It is conventional to compare the bonding of phosphorus with that of nitrogen because they are in the same group of the periodic table (1). They have similar outer electronic configurations and, therefore, we expect similarities in their chemical behavior. There are frequently striking diierences in the chemistry of third row elements when compared with their second row analogs, however, and phosphorus fulfills expectations in this regard. Thus, there are no phosphorus compounds analogous to the nitro, nitroso, azo, diazo, azido, nitrile, and imino compounds, and phosphorus does not give stable heterocycles like pyridine or pyrrole. On the other hand, there are compounds of phosphorus-phosphoric acid, phosphonie acids, phosphinic acids, and phosphoranes-whose analogs do not appear in nitrogen chemistry. And, molecular nitrogen is a triply bonded diatomic species, whereas, elemental white phosphorus is a tetrahedral species with each atom connected by three single bonds to the other atoms in the molecule. I t has been suggested that the large size of phosphorus relative to nitrogen and, therefore, the larger interatomic distances of its compounds cause less effective overlap of phospboms 3p-orbitals with 3p- or 2porbitals of other atoms ( 2 ) . Alternative suggestions are repulsion of bonding orbitals and innemhell electrons
bonds like the phosphoryl group, P-0, d,-p, bonding is more important than p,-p, bonding since electron density is concentrated nearer the more electronegative atom, oxygen. Evidence for the double bond character of the phosphoryl group is derived from bond lengths and dipole moment data for the phosphine oxides (paP=O). For the P-0 bond length a value of 1.76 A was calculated, whereas the measured length for several compounds is 1.45-1.48 A (6). Similarly, the bond moments of amine oxides are greater than those of the corresponding phosphine oxides and lead to estimates of charge distribution in the P-0 and N-0 bonds as shown in Figure 2 (7). The reverse order
Figure 2.
Charge dirtribvlion
in P-0
and N-0
bonds.
would he expected on the basis of inductive effects alone since nitrogen is more eleetronegative (3.0) than phosphorus (2.1). The character of d,-p, bonding in phosphorus compounds on the basis of energy eonsiderations, symmetry, bond lengths, spectra, and chemical properties has been extensively reviewed by Hudson (8). Phosphazenes
Figure 1.
d,-p,
Bond [left) and p , p ,
bond.
is greater with large atoms and, also, single bonds to phosphorus have partial double bond character r e sulting from the use of d-orbitals, so there is less to gain in forming a double (p,p,) bond than is the ease with nitrogen ( 3 , 4 ) . However, the calculation of overlap integrals for the d-orbitals on phosphorus with the p-orbitals on a second atom led to the conclusion that the overlap increased with increasing charge on phosphorus, thus contracting and stabilizing d-orbitals normally too diffuse for bonding (5). I n contrast to p,-p, bonding where maximum overlap occurs in a region equidistant between the nuclei, d;p, bonding is polar with electron density roneentrated near the ligand atom (Fig. 1). Thus, for polar Present address: Rio Hondo Junior college, Whittier, Calif. 90661.
In 1834, Rose described the reaction between PC& and NH3 (9). This was the beginning of a long history for a series of compounds which continues to evince interest to this day, that is the phosphazenes, or phosphonitrilie, compounds. It was proposed by Stokes, in 1895, that these compounds possessed a eyelie structure and formed a homologous series (PNC13, (Fig. 3) (10). He isolated and identified the members for n = 3-7. The main features of these compounds
I PNC12)3 Figure 3.
(PNC1214
Some phorpharenes.
Volume 44, Number 2, February 1967 / 95
Molecular Geometry of Phosphazenes (PNX&,r
Toble 1.
(PNCl.). ~.~ NPN angle XPX an le PNP an$e P-N length (A) P-X length (A) P-X in X;P=O IA)
Toble 2.
(PNClsh
(PN(NM4e)d
(PNFzh planar
119.8' 104 132 1.60 1.80 1.81
121.Z0 102.8 132 1.58 1.99 1.99
120" 104 133 1.58 1.68
122.6" 99.9 147 1.51 1.51 1.52
119.6" 101.9 120.0 1.59 1.99 1.99
Ligmd electronegativity
Table 3.
Me
Br
2.5
2.8
NMen CI
3.0
3.0
OMe
3.5
F
3.9
Representation: n = 3 n = 4
Nitrogen Ps
-
13
B2
W l c r e .4 ie totally w n m . c t r w rrwrwntatmo. t.' la ~niirymmatr:; revrea n l li rr 12ublg nec?nerrre r r ~ r ~ a r n r a r i n n
.mtattoo.
are comparative chemical inertness and thermal stability. Many derivatives have been obtained, but reaction is often slow and incomplete, and the products have not always been completely characterized. Although they are acid chlorides, they are not hydrolyzed rapidly; the trimeric chloride can be steam-distilled without serious loss (11). The molecular geometry of many of the bomologs has been determined, principally by X-ray or electron diiraction (1.2). The values for some compounds are shown in Table 1. The planar fluoride has such a low electron density on nitrogen, and is consequently so weak a base, that it barely dissolves in sulfuric acid. The ring bond angles a t phosphorus are nearly constant and agree with other molecules with doublebonded phosphorus, e.g., C1-P4 in C 4 P 4 is 118", and &P-0 in tetrametaphosphate ion is 121'. The P-N bond is sensitive to ligands on phosphorus. One way of measuring this effect is to compare the P-N stretching frequencies in the infrared for various ligands (13). I n Table 2 correlations for a variety of ligands are shown. The use of group electronegativities in correlations of this type, however, may involve a circular argument, since one of the sources of group electronegativities is from the assignments of infrared stretching frequencies. The fact that these compounds undergo substitution reactions to a limited extent, without scission of the ring, has led to some interesting controversiea conceming their aromatic character. 96
/
Journal o f Chemical Education
Craig and Paddock, in 1958, published a short article purporting to introduce a new theory of aromaticity using the phosphazenes as their model (14). This was a molecular orbital description, its novelty lying in the fact that the lower filled energy levels are degenerate, whereas in benzene the highest filled level is degenerate (Fig. 4). Thus, where aromatic hydrocarbons r e quire (412 2) ?r-electrons in order to maintain their aromatic character (Hiickel's rule), Craig's model merely requires an even number increase in r-electrons. This arises from the gerade inversion symmetry of the d-orbitals, where porbitds are ungerade to inversion (Fig. 5). Dewar, Lucken, and Whitehead attacked Craig's model in 1960, proposing an alternative model with an allylic or three-center structure of a-bonds (15). Using as a first approximation D3 point group symmetry, Dewar outlined the symmetries of the orbitals involved, ae shown in Table 3, emphasizing that symmetry favors the dVrorbitd fot interaction with the pa-orbital, since the d,,orbital cannot interact with the p,orbital in the totally symmetric state (Fig. 6). Craig had pointed out that dimensions and energies of phosphorus d,-orbitals will be affected by adjacent ligands; since the ligands are not all the same, the sizes and energies of dZr and d,,-orbitals may differ. He had used Dasymmetry and believed the change in coulomb and resonance integrals would be enough to make the d,-p, interactions much more inportant than the d,,orbital interactions. He had, therefore, neglected the &-orbital in his treatment. Having cast doubt upon Craig's choice of d-orbitals,
+
Orbital Symmetries in (PNCld3.r Qumtum No.
...
Aromaticity Arguments
Infrared Stretch Frequencies of the P-N Bond
X in (PNXd.
Orbital
(PNMenh
4-
I f - + "J
-tl-
+-It
CeH6
IPNCIIII
u
++ X
Figure 4.
Benzene and phorphazene 7-eleshon energy levels (not to
sca1eJ.
Figure 5.
IPNC1214
Inversion rymrnelry of
d,- and p,.orbitalr.
Figure 6. Relatiomhip of the d = v , dvr, and prorbitah in the phor phorenes (of* Craig and Pod-
~i~~~~ 7. Linearly combined orbitals d,' and dR (after Dowar, ~"cken,ond Whitehead).
dock).
Dewar proceeded to replace the d-orbitals with a pair of linear combinations d,' and d," (Fig. 7), where "d,
1/.\12(dzz - dv*)
in order to attain efficient overlap with the nitrogen porbitals, and combined the atomic orbitals into threecenter allylic-type molecular orbitals (Fig. 8).
The three-center orbitals overlap only weakly with one another, the break a t each phosphorus atom implying that the adjacent three-center set need not be coplanar. This was consistent with certain puckered compounds and showed that ring electron delocalization (as in benzene) was an unnecessary hypothesis for the phosphazenes. Dewar's most cogent argument concerned the ultraviolet spectra of the phosphazenes. The spectra show little resemblance to organic aromatic systems (although this is probably to be expected), their interesting feature being the fact that the spectra of all the cyclic polymers are almost the same. They are broad based peaks with no fine structure, X maximum = 200 mp (e.g., (PNC1J3 = 199 mp, (PNCI,), = 203 mp), and extinction coefficients which increase with increasing molecular weight. This is to be expected if the size of the basic conjugated unit is unchanged throughout the series, but not if its size were expanded by conjugation through the phosphorus atom. I n a remarkable display of one-upmanship, Craig and Paddock (perhaps it was two-upmanship) in 1962, brought out their theory in detail in 15 pages of the Journal of the Chemical Societv (18). Craig systematically attacked the symmetry species of all five of the phosphorus d-orbitals, considering each possibility with a correspondmg calculation of overlap integrals, and (for good measure) used s p hybridized orbitals on nitrogen. Also, to confuse the innocent, he turned his coordinate system 90" in the midst of his discussion. He presented a very impressive table, not reproduced here, except for the colorful headings:
Craig discussed the bond strengths in the phosphazenes and concluded that, in the trimeric chloride, for example, the P-N bond is G10 Kcal stronger than a single P-N bond. He rationalized the fact that this is so much smaller than the energy in C-C bonds (cu. 38 Kcal) in benzene by referring to the diffuseness of the d-orbitals and the fact that maximum overlap in a d;p, bond occurs in a region of low nuclear field. One cannot argue with Craig's conclusions because they are so general. He states that the results indicate the distribution of d-electrons is sensitive to the detailed environment and justify the assumption of dorbital participation. The electron distribution is controlled by symmetry and by interactions which vary with the molecule and can have large effects on the molecular structure. Silicon Multiple Bonds
Silicon, like carbon, has a normal valence of four and forms stable bonds to other silicon atoms, to carbon, hydrogen, halogens, oxygen, nitrogen, etc. However, a compound containing a double bond to silicon has never been isolated. Attempts have been made to form such bonds, for example, by the treatment of PhzSiC12 with sodium (16). No formation of compounds of the type Ph2Si=SiPhz took place although ring stabilization by conjugation with such a bond would be expected to promote its formation. This inability to form double bonds is a general characteristic of second row elements, probably resulting from the d i u s e nature of the 3p- and 3d-orbitals compared to the 2p-orbitals. There is, however, a great deal of evidence for partial double bond character in many compounds of silicon (17, 18). On the basis of electronegativity we would expect trisilylamine (HsSi)3Nto be a stronger base or electron donor than trimethylamine (H&)3N, since silicon is less electronegative than carbon. I t is, in fact, a poorer donor and forms no complexes with diborane or trimethylboron, although an unstable complex was formed with boron trifluoride (19). I n addition, the molecule is planar, whereas trimethylamine is tetrahedral. One explanation is the bonding of the nitrogen lone pair with silicon through its d-orbitals,- or contributions from resonance structures such as H3Si= N(s~H,),. Similar arguments have been used to explain the linear structure of silyl isothiocyanate H8iNCS (SO), the more protonic character of trimethylsilanol compared to tertiary butanol @I), the high acidity of triphenylsilanol (approaching that of phenols) (22), and the anti-Markownikov addition of hydrogen halides to trialkylvinylsilanes (83). I n all these cases, the predicted inductive effect of the silyl group is overcome by the resonance effect of the d,-p, interaction between silicon and the adjacent atom. This is shown by the fact that when the resonance effect is inoperable, as in allyltrialkylsilanes, the Markownikov rule is followed in HBr addition (83) : EtSiCHLXI=CH?
+ HBr
-
Et$iCHDHBrCHa
Electron spin resonance experiments were performed by Allred and Curtis, enabling the first calculation of a a-bond order (0.18) for a silicon-carbon bond (24). substituted biphenyl radicals were used (Fig. 9) beVolume 44, Number 2, February 1967
/
97
cause their symmetry makes the spectra simpler and easier to interpret and the radical ions are particularly stable (a tetrahydrofuran solution of b (Fig. 9) is stable for six months at room temperature). A Test of d-Orbital Participation
Probably the most significant recent development toward a proof of &orbital participation by silicon is an experiment by Goodman, Sommer, and Konstam (25). This is a spectroscopic experiment which is relatively easy to perform, and it is of interest to consider it in some detail. The electronic energy levels of benzene and substituted benzenes have been studied theoretically and experimentally more than those of any other group of
c (Fig. 10) cannot &ect the intensity of the 2600 band since they are symmetric about the C2, axis, but b (Fig. 10) and d (Fig. 10) may do so. The intensity of I this band comes from two sources: (1) borrowing from the ben; zenoid band a t 1850 A (Fig. 12), and (2) conR tributions from structures like b (Fig. 10) and d (Fig. 10) to a transition moment. Source (1) is more important provided the perturbations from the substituents are not too strong. I Goodman related F ~ Q W 1I . ~,.,~.iti~~ moment ( M L ~ the transition moment polorirotion for tho 2600h transition in MLbfor the forbidden dirubstituted benzenes. transition to the cnmponent transition moment of appropriate symmetry M,, of the borrowed 1850 A transition, and to the moment BBzU of appropriate symmetry for structures b (Fig. 10) and d (Fig. 10): M u = eosAMm + sinAM~a (1) where the quantity A is determined by the amount of mixing of the benzene Bz. and El, wave functions caused by the substituent perturbation. For substituents like -CH3 or -C1 which do not conjugate strongly, the ratio MBb/MBZu is greater than lo2, so that the second term in eqn. (I), the "borrowing" term, is greater. As sin A becomes greater, the "borrowing" term predominates and the intensity of the absorption band a t 2600 A increases. To a first approximation:
j: s
Figure 9. Substituted biphenyl radicol ions for ESR rhldier.
molecules. Goodman had published a theory with Shull ($6) using A t 0 theory and first order configuration interaction theory to account for the observed effects of substitution upon the electronic spectrum of benzene. When a silicon atom is attached to a benzene ring, two effects should become apparent-the inductive and the resonance. The inductive effect is expected to release electrons into the ring since carbon is more electronegative than silicon, whereas the resonance effect can either withdraw or release r-electrons depending upon the s-bonding of the substituent. Some of the resonance structures for phenylsilane using the
sin A =
Figure
10. Some resonance rtructurer for phenylsilane.
3d-orbitals of silicon are shown in Figure 10. Contributions from b (Fig. lo), a type of hyperconjugation, would tend to reinforce the inductive effect represented by a (Fig. 10). Structures c (Fig. 10) and d (Fig. 10) oppose the inductive and hyperconjugative effects by withdrawing ring s-electrons into silicon dorbitals. For substituted benzenes of C2. or higher symmetry, the probability of the 2600 A transition in the ultraviolet spectrum is related to the electronic distributions of the ring alone. This transition is formally forbidden in benzene itself, but upon mono- or paradisubstitution, it becomes formally allowed, with the transition moment ML"erpendicularly polarized to the symmetry axis (Fig. 11). Because of this, one need only consider electronic distortions in the ring with components perpendicular to the C2. axis. Hence, structures like a (Fig. 10) and 98 / Journal o f Chemical Educafion
(+.mu
IH I ICmu)/(smu- r s w )
(2
and the matrix element in this equation reduces to a difference in energy between configurations entering into the symmetry-determined benzene B2"and Eluwave functions: (#a% IHI + a d = '/n(Eiz - El$) (3) where Eu is the energy of the promotion i -+ j, and i
=SO
Figure 12.
i
2600
A-
The ultraviolet spectrum of benzene.
-
aud for para-disubstitution, the resulting intensity of the 2600 A band in terms of dipole strength D resulting from each group alone (Da D9)is
+
D = Dg
-
H3
a
b
BENZENE
L
Figure 13. Perturbation of benzene MO'r b y (a) electron-donoting rubrtituents, and (b) electron-withdrawing rvbrtituentr (after Goodman, Somrner, and Konrtoml.
refers to a node in the i molecular orbital on the symmetry axis (Fig. 13). Figure 13b represents the effect of an interaction of the benzene orbitals with a low lying " - 3d-orbital which has the effect of lowering the energies of the la- and 2norbitals, where 2~ is lowered more than la. The Iaand %-orbitals are unaffected because they have a node on the C2"axis. This situation requires that the matrix element and, therefore, sin A be less thau zero. I n Figure 13a we see the effects of either hyperconjugation or a group which donates electrons to thering, such as methoxy or methyl, which interacts to raise the energy of the T-orbitals, requiring that sin A be greater than zero.
Figure 14. Substituted phenylsiloner wed in spectroscopic experiment%
OMe Goodman extended equ. (2) to disubstituted benzenes in terms of El., and B",,matrix elements for each group alone (R and-s in Fig.-11), to obtain sin A = sin(Az
+ As)
Figure 15. U.V. absorption specha for phenyirilme (solid line) and p-methylphenylrilone (broken linel.
+ DS + 6in Ax sin hs(Dsi.)
(5)
The interference term in eqn. ( 5 ) is important. If both sin A's are in phase, the interference is constructive and the band intensity will be greater thau for either group alone. Conversely, destructive interference requires a decrease in intensity. The reason for using disubstituted benzenes and eqn. (5) is that the dipole strength D is determined by the square of the transition moment and no useful information can be obtained from the intensity in phenylsilane itself, since: DU trr ' / z (sin2A)D~,,, (6 and eqn. (6) has no interference terms. Both methoxy and methyl groups fulfill the condition that sin A be greater than zero, and for the compounds shown in Figure 14 a phase difference will exist in the spectroscopic moments if the 3d-orbitals of silicon interact with the ring a system. The experimental spectra for the ultraviolet absorption of the 2600 A band for the con~poundsin Figure 14 are shown in Figures 15 and 16, where they are compared with the monosubstituted derivatives. The pronounced decrease in absorption intensity upon disubstitution with the silyl group definitely confirms that the -SiHa group is withdrawing electron density into the 3d-orbitals of silicon from the benzene ring. Goodman and co-workers have obtained like results with some lower halogens (27) and the thiol group (28) as substituents on the benzene ring, and Bissey and Goldwhite have utilized his method to demonstrate the electron acceptor properties of phosphorus (29). Acknowledgment
The author wishes to thank Prof. Harold Goldwhite for suggesting that this manuscript be written and for critical reading of it. Errors of omission or commission are, of course, solely those of the writer. Literature Cited ( 1 ) See, for exsmple, ROBERTS, J. D., AND C z < s ~ ~id. ~ oC., , "Basic Principles of Organic Chemistry," W. A. Benjamin, Inc., New York, 1964, p. 1177.
Figure 16. U.V. absorption spectra for anirole (solid line) ond p-mothoiyphenylriione (broken line).
Volume 44, Number 2, February 1967
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99
MULLIKEN, R. S., J . Am. Chem. Soc., 72,4493 (1950). PITZER. K. S.. J . Am.. Chem. Soc.. 70, 2140 (19481. . . C O U L S ~ N , C. A,, "Valence," oxford university Press, Inc., London, 1952, p. 178. CRAIG,D. P., MACCOLL, A., NYEOLM, R. S., ORGEL,L. E. AND SUTTON. L. E..J . C h . Soe... 332 (19541. . , PADDOCK, N. L., "Structure and Reactions in Phosphorus Chemistry," The Royal In~nstituteof Chemistry, Lecture Series 1962, No. 2, London, 1962, p. 32. PHILLIPS,G. M., HUNTER,J. S., AND SUTTON, L. E., J. Chem. Soe., 146 (1945). HUDSON,R. F., "Structure and Mechanism in OrganoPhos~harus Chemistrv." Academic Press. Inc.. New ~ o r k ;1965, Chapter 3: ' Snaw. R. A.. FITZSIMMONS. B. W.. AND SMITH.B. C.. AUDRIETH, L. F., STEINMAN, R., AND TOY,A. D. F., Chem. Revs., 32,109 (1943). CRMG,D. P., AND PADDOCK, N. L., J . C h a . Soc., 4118 (1962). SHAW,R. A., FITZSIMMONS, B. W., AND SMITH,B. C., Ref. ( 9 ) , p. 272. CRAIG,D. P., AND PADDOCK, N. L., Natwe, 181, 1052 (19581. . . DEWAR,M. J. S., LUCKEN,E. A. C., AND WHITEHEAD, M. A,, J . Chem. Soc., 2423 (1960).
100 / Journol of Chemkcrl Education
(16) KIPPING,F. S., MURRAY, A. G., A N D MALTBY,J. G,, J . Chem. Sac., 1180 (1929). D., J . Ino~g.& N u d . C h m . (17) STONE,F. G. A., AND SEYFERTE, 1,112 (1955). L., "The Nature of the Chemical Bond," Cornell 118) PAULING, University Press, Ithaca, New York, 3rd ed., 1960, p. 310. (19) BURG,A. B., AND KUIJIAN,E. S., J . Am. Chem. Soe., 72, 3103 (1950). A. G., AND MADDOCK, A. G., J. Inorg. & (20) MACDIARMW, N w l . Chem. 1, 411 (1955). E. G. AND ALLRED,L., J . 1nwg. (21) STONE,F. G. A,, ROCEOW, & Nuel. Chem. 2,416 (1956). (22) WEST, R., AND BANEY,R. H., J. Inmg. & Nucl. C h m . 7, 297 (19581. (23) SOMME'R, L.'H., et al., J . Am. C k m . Soe. 76, 1613 (1954). (24) CURTIS,M. D., AND ALLRED, A. L., J . Am. Chem. Soc. 87, 2554 (1965). L. H., GOODMAN, L., AND KONSTAM, A. H.,J. (25) SOMMER, Amw. Chem. Soe., 87, 1012 (1965). L., AND SKULL,H., J. Chem. Phys. 27, 1388 (26) GOODMAN, (1957). L., AND FROLEN,L. J., J . C h a . Phys. 30, (27) GOODMAN, 1361 (19591. ~
~
