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J . Am. Chem. SOC.1985, 107, 1922-1930
Electronic Structure of the Phosphoryl and Thiophosphoryl Bonds? Michael W. Schmidt* and Mark S. Gordon* Contribution from the Department of Chemistry, North Dakota State University, Fargo, North Dakota 58105. Received August 13, 1984
Abstract: The phosphoryl bond is presented as a resonance hybrid between singly and triply bound structures. The types of electronically excited states and a number of their potential surfaces can be understood within the confines of a singly bound model. However, a detailed analysis of the electronic distribution within the ground state shows significant contribution from the triply bound structure. In fact, increasing the electronegativity of the ligands at phosphorus increases the contribution of the triply bound structure, so that the PO bond order approaches that of a double bond. The thiophosphoryl bond may be viewed in much the same manner but is rather weaker and contains less multiple bond character than does the PO group.
As noted in a recent paper,’ the phosphoryl group in phosphine oxide can be written in several ways,
I
I1
I11
The singly bonded structure I contains only u bonds and three lone pairs on oxygen. The d orbitals of phosphorus are not required to explain the bonding in structure I, although their energy is low enough that they may be. invoked in other bonding schemes, such as I1 and 111. These multiply bonded structures indicate the possibility of additional bonding usually ascribed to oxygenphosphorus p,-d, overlap, leading to donation of the oxygen lone pairs back to the phosphorus d levels. The customarily written structure I1 can be rejected as inconsistent with the threefold symmetry of the molecule. After all, both oxygen e symmetry lone pairs should participate in this back-bonding to an equivalent extent, leaving one al oxygen lone pair. Although this symmetry requirement for triple rather than double bonding was recognized as early as 1961,2the incorrect resonance structure I1 is in nearly ubiquitous use by chemists. Recently, localized orbital calculations3 have been used to argue for full triple bonding in phosphine oxide. Asymmetric substitution, as in XYZPO, removes the degeneracy of the oxygen lone pairs in structure I but does not prevent both from participating in back-bonding, so that even in these compounds multiple bonding would be better described by structure 111 than by 11. The full triple bonding implied by structure I11 overstates the extent of the actual a-bonding interaction. The true bonding between phosphorus and oxygen must fall somewhere between the two resonance structures I and 111, and this may well be the origin of the frequent use of structure 11. The aforementioned investigation’ established that the phosphoryl bond is more closely represented by the dative structure I, with some admixture of the resonance structure 111. The evidence’ we presented for predominantly single binding in phosphine oxide, supplemented by some additional a backbonding may be summarized as follows. The bond formation can be thought of as an acid-base reaction in the Lewis sense, with the phosphine lone pair forming a donor bond that completes the oxygen octet. Such a dative bond gives a formally ionic Lewis structure, and indeed the Mulliken charges on P and 0 were found to be +0.9 and -0.7, respectively. This charge separation is mirrored in the 3.7-D dipole moment. Energy-localized S C F orbitals were found to consist of one strong a bond and three lone pair orbitals on oxygen. These lone pairs were found by Mulliken population analysis to participate in a noticeable amount of a back-bonding to P, to the extent of some 0.2 e- per oxygen lone The group notation is being changed in accord with recent actions by IUPAC and ACS nomenclature committees. A and B notation is being eliminated because of wide confusion. Group I becomes groups 1 and 11, group I1 becomes groups 2 and 12, group 111 becomes groups 3 and 13, etc.
0002-7863/85/1507-1922$01.50/0
pair. It is this back-bonding which causes the phosphoryl bond length to be about the same as that in the doubly bound HPO, rather than the singly bound H2POH. However, the primary source of P O binding is due to the u bond. No drastic change in the localized orbitals occurs as H,PO isomerizes to the singly bound H2POH, except for a loss of most of the back-bonding. Finally, the P O stretching frequency in H 3 P 0 is intermediate between those in H P O and H2POH. There is a great deal of experimental and theoretical evidence for the presence of multiple binding in phosphine oxides. References to experimental evidence for the Occurrence of this binding have been given by Carlson and Meek4 and Albright et aL5 However there is much disagreement as to the magnitude of this a bonding. For example, Carlson and Meek4 state that the bonding is best described as a “a bond possessing a character intermediate between a single and a double bond”. At the opposite extreme, Wilkins et aL6 somewhat arbitrarily obtain a P O bond order of 2.68 in trimethylphosphine oxide. References to previous ab initio theoretical calculations on phosphine oxides were given in our previous work.’ All these calculations stress the importance of d orbitals on phosphorus in describing structure and bonding in phosphine oxides. Hudson’ has given an extensive review of bonding in phosphoryl and other related compounds. Thermochemical data for phosphorus compounds have been compiled by Hartley et aL8 There is some recent experimental evidence for the interpretation of single bonding with some degree of a reinforcement. X-ray PES studies9 show a large stabilization (3.95 eV) of the trimethylphosphine lone pair upon P O bond formation. There is a much smaller shift of the H O M O e symmetry oxygen lone pairs as they participate in some a back-bonding. X-ray fluorescence spectral0 of various phosphine oxides have been interpreted as electronic transitions occurring from orbitals responsible for a weak p,-d, interaction. Very recently, several workers”J2 have suggested that p orbitals are responsible for at least part of the a-accepter capability of (1) Schmidt, M. W.; Yabushita, S.; Gordon, M. S. J . Phys. Chem. 1984, 88, 382-389.
(2) Lucken, E. A. C.; Whitehead, M. A. J . Chem. SOC.1961, 2459-2463. (3) Wallmeier, H.; Kutzelnigg, W. J . Am. Chem. SOC.1979, 101, 2804-28 14. (4) Carlson, R. R.; Meek, D. W. Inorg. Chem. 1974, 13, 1741-1747. ( 5 ) Albright, T. A,; Freeman, W. J.; Schweizr, E. E. J . Org. Chem. 1975, 40, 3437-3441. (6) Wilkins, C. J.; Hagen, K.; Hedberg, L.; Shen, Q.; Hedberg, K. J . Am. Chem. SOC.1975, 97, 6352-6358. (7). Hudson, R. F. “Structure and Mechanism in Ogranophosphorus Chemistry”; Academic Press: New York, 1965; Chapter 3. (8) Hartley, S. B.; Holmes, W. S.; Jacques, J. K.; Mole, M. F.; McCoubrey, J. C. Quart. Rev. 1963, 17, 204-223. (9) Elbel, S.; tom Dieck, H. J . Chem. Soc., Dalton Trans. 1976, 1757-1762. (10) Mazalov, L. N.; Yumatov, V. D.; Dolenko, G. N. Zh. Strukt. Khim. 1980, 21, 21-29. (1 1) Xiao, S.-X.; Trogler, W. C.; Ellis, D. E.; Berkovitch-Yellin, 2. J . Am. Chem. SOC.1983, 105, 7033-7037. (12) Marynick, D. S. J . Am. Chem. SOC.1984, 106, 4064-4065.
0 1985 American Chemical Society
Phosphoryl and Thiophosphoryl Bonds
J . Am. Chem. SOC.,Vol. 107, No. 7 , 1985
1923
phosphines. T h i s conclusion was reached on t h e basis of X a calculations on t h e free phosphines” and PRDDO calculations on phosphine-metal complexes, using bases without d orbitals.I2 This conclusion is partially supported by our earlier finding’ that t h e back-bonding in phosphine oxide is predominantly into t h e p orbitals of phosphine. However, our preliminary results were obtained by using a basis of only moderate quality (containing a single phosphorus d orbital). More refined calculations a r e clearly necessary to settle the relative importance of the phosphorus p a n d d orbitals in phosphine *-acceptor behavior. In t h e present work w e expand our investigation] of t h e phosphoryl bond in H 3 P 0 a t its equilibrium geometry in several directions. The strength of t h e bond and several of its dissociative potentials are presented for H,PO in various electronic states. These ground- and excited-state surfaces a r e rationalized in terms of t h e reinforced PO single bond model presented above. Electronegative substitution a t phosphorus (a sine qua non for t h e experimental existence of these compounds) is considered with regard to its influence on the bond strength and electronic structure of t h e PO moeity. For a few of t h e molecules we have carried out calculations in extended bases designed t o probe the influence of t h e basis set on t h e back-bonding chemistry of phosphines. Finally, we briefly discuss t h e thiophosphoryl bond found in t h e phosphine sulfides.
exponents are taken from Francl et a1.16a The only exception to this prescription is F,PO, where the large basis size forced the omission of d orbitals from the fluorines. As the phosphine oxides are well represented within the closed shell S C F model, restricted Hartree-Fock (RHF) calculations are used to predict structures and electron distributions. The interpretation of R H F results is aided by a Mulliken population a n a l y s i ~ and ’ ~ by the determination of energy-localized MOS.~O d orbital populations are taken from the 3-21G* R H F wave functions, since these exclude the 3s contaminant in the d orbitals. The most accurate total charge populations are derived from 6-31G* bases. For selected molecules, more accurate electronic distributions are obtained from the ETGC basis. As already demonstratedl for the phosphine oxides, the energy-localized MOs are quite different from and clearly superior to those obtained by the commonly used Boys2I localization procedure. Electron correlation effects are treated with two different methods. First, Mdler-Plesset perturbation theory at third order22(MP3) is used to recover the major portion of the correlation energy of all the valence electrons at the equilibrium structures. We will use the following notation to summarize these calculations; for example, MP3/6-31G*// STO-2G* means a MP3 calculation using a 6-31G* basis at the STO2G* R H F geometry. Because MP3 calculations with the 6-31G* basis are quite time-consuming, we often estimate energy differences by the additivity relationship AE(MP3 /6-3 1G*) = AE(RHF/6-31G*) [AE(MP3/3-21G*) - AE(RHF/3-21G*)]
Computational Methods
that is, the correlation correction is found by using the more economical 3-21G* basis. This approximate equality is usually accurate to about 5 kcal/mol, as will be demonstrated below for a few cases (see also ref 12). Second, multiconfiguration S C F (MCSCF) calculations within the full optimized reaction space (FORS) are used to describe the dissociation of the PO bond in phosphine oxide and to characterize its excited states. Computer codes used include A L I S , ~for ~ MCSCF and CI calculations and orbital localizations, and GAUSSIAh’80,25 for structure determination26 and perturbation theory computations.
Five atomic basis sets of increasing accuracy are employed. Equilibrium geometries are found by using the minimal STO-2G* basis.’) This is a very economical and fairly reliable means of predicting structures in phosphine oxides.14 The split-valence 3-21G* basisl5 gives improved accuracy for energy quantities, while remaining of tractable size for correlation and some geometry calculations. The addition of a Rydberg set of s and p orbitals centered at phosphorus (with common exponent 0.039) yields a third basis set denoted 3-21G*+ that will afford better accuracy for excited states that may possess Rydberg character. The 6-31G’ basisI6 is used to obtain the most accurate results for energies for all the molecules considered. The first three bases contain d orbitals only for third-row atoms (phosphorus exponent 0.5516”in all bases) to permit a description of the PO pn-d, interaction at all basis levels. The 6-3IG* basis set contains polarization functions on second-row atoms as well. For the first two bases only, the 3s contaminant of the Cartesian d orbitals is excluded. The fifth basis set, referred to as ETGC below, strives to approach the Hartree-Fock limit for the electronic density between phosphorus and oxygen as closely as is practical for these large molecules. The sp basis for phosphorus and third-row substituents is a 16s,l l p even-tempered basis, generally contracted to a split-valence set 4s,3p. The phosphorus basis is supplemented by two d sets of five functions with exponents 0.22 and 0.77, since calculations with the above four bases show a large contribution from a single d function. The use of an extra set of d orbitals will of course act to increase back-bonding. The sp basis for oxygen and second-row substituents is a 12s,6p even-tempered basis generally contracted to the split-valence set 3s,2p. The basis for oxygen and all heavy atom substituents at P is complemented by a single set of five d functions and a diffuse s and three p functions to allow these atoms to accommodate substantial negative charge. The exponents for these diffuse functions are the appropriate even-tempered parameters a. The diffuse functions on oxygen should reduce back-bonding by permitting oxygen to retain a larger negative charge. The diffuse functions on the ligand atoms will increase PO back-bonding, by causing the phosphorus to be more positively charged due to the more negative ligand. The basis for hydrogen atoms is an unpolarized, generally contracted even-tempered basis 4s/2s, scaled by 1.2. The even-tempered portions of this basis are taken from Schmidt and Ruedenberg.” Second-row polarization exponents are taken from Dunning and Hay,’* and third-row polarization (13) Collins, J. B.,; Schleyer, P. v. R.; BinMey, J. S.;Pople, J. A. J . Chem. Phys. 1976, 64, 5142-5151. (14) Gordon, M. S.; Boatz, J. A.; Schmidt, M. W. J . Phys. Chem. 1984, 88, 2998-3002. (15) Pietro, W. J.; Francl, M. M.; Hehre, W. J.; DeFrees, D. J.; Pople, J. A.; Binkley, J. S. J . Am. Chem. SOC.1982, 104, 5039-5048. (16) (a) Francl, M. M.; Pietro, W. J.; Hehre, W. J.; Binkley, J. S.;Gordon, M. S.; DeFrees, D. J.; Pople, J. A. J . Am. Chem. Phys. 1982, 77, 3654-3665. (b) Hariharan, P. C.; Pople, J. A. Theor. Chim.Acta 1973, 28, 213-222. (17) Schmidt, M. W.; Ruedenberg, K. R. J . Chem. Phys. 1979, 71, 3951-3962. (18) Dunning, T. H.; Hay, P. J. “Methods of Electronic Structure Theory”; Schaefer, H. F., 111, Ed.; Plenum Press: New York, 1977; pp 1-27.
+
Results and Discussion Dissociative Potentials for Phosphine Oxide. Angular momentum coupling rules require that when a ground-state phosphine oxide eliminates a closed-shell phosphine moiety, t h e atomic oxygen cofragment must be in a singlet state. A t first glance it might appear that the removal of oxygen from phosphine oxide might be accurately described within t h e RHF framework, since no electron pairs are broken for this process. However, t h e RHF wave function dissociates t o a singlet valence s t a t e of oxygen, namely 2s22pX22p;, t h a t is a linear combination of t h e ID and ‘S states of oxygen. This incorrect RHF dissociation h a s been noted p r e v i ~ u s l y . ~Dissociation t o a correct ‘D oxygen s t a t e requires the addition of t h e oxygen pz orbital or, in MO language, t h e u* orbital correlating the u dative bond, which itself becomes t h e phosphine lone pair upon dissociation. We have performed MCSCF calculations within t h e FORS model. These MCSCF calculations contain all possible configurations in which t h e eight reacting electrons (six from oxygen, two from phosphine) a r e distributed among t h e five active orbitals: the three oxygen lone pairs and t h e u and u* orbitals of t h e oxide molecule. These calculations correctly describe the dissociation of phosphine oxide into a n y valence state of t h e oxygen a t o m or ion and the ground (19) Mulliken, R. S. J . Chem. Phys. 1955, 23, 1833-1846. (20) (a) Edmiston, C.; Ruedenberg, K. Reo. Mod. Phys. 1963, 35, 457-465. (b) Edmiston, C.; Ruedenberg, K. J . Chem. Phys. 1965,43, S97S115. (21) Foster, J. M.; Boys, S. F. Reo. Mod. Phys. 1960, 32, 300-302. (22) Pople, J. A.; Binkley, J. S.; Seeger, R. Int. J . Quantum Chem., Quantum Chem. Symp. 1976, I O , 1-19. (23) (a) Ruedenberg, K.; Schmidt, M. W.; Gilbert, M. M.; Elbert, S. T. Chem. Phys. 1982, 71,41-49. (b) Ruedenberg, K.; Schmidt, M. W.; Gilbert, M. M. Chem. Phys. 1982, 71, 51-64. (c) Ruedenberg, K.; Schmidt, M. W.; Gilbert, M. M.; Elbert, S. T. Chem. Phys. 1982, 71, 65-78. (24) Elbert, S. T.; Cheung, L. M.; Ruedenberg, K. National Resource for Computers in Chemical Software Catalog 1, Program QMO1, University of California-Berkeley, 1980. (25) Binkley, J. S.; Whiteside, R. A,; Krishnan, R.; Seeger, R.; DeFrees, D. J.; Schlegel, H. B.; Topiol, S.; Kahn, L. R.; Pople, J. A. Quantum Chem. Program Exchange Catalog 13, Program 406, Indiana University, Bloomington, 1981. (26) Schlegel, H. B. J . Comput. Chem. 1982, 3, 214-218.
1924 J. Am. Chem. SOC.,Vol. 107, No. 7, 1985
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