3080
J. Phys. Chem. 1983,87,3080-3082
Lone-Pair and Core Ionization Potentials of Planar Ammonia and Phosphine. The Use of Core and Valence Ionization Potentials To Quantify the Bonding and Antibonding Character of Molecular Orbitals of Compounds of Nitrogen and Phosphorus Charles J. Eyermann and Wllllam L. Jolly' Department of Chemistry, University of Californk at Berkeley and the Materkls and Molecular Research Division, Lawrence Berkeley Laboratory, Berkeley, Calfomla 94720 (Received:December 8, 1982; I n Final Form: February 3, 1983)
For both ammonia and phosphine, the differences between the ionization potentials of the pyramidal and planar forms were calculated by an ab initio SCF method. The lone-pair ionization potential differences, IP(p1anar) - IP(pyramidal),are -1.0 and -2.3 eV for NH3 and PH3,respectively. The core binding energy shifts, .&(planar) - EB(pyramidal),are -0.3 and -0.4 eV for NH3 and PH3,respectively. By combining these data with appropriate experimental valence and core ionization potentials, it is possible to quantify the bonding or antibonding character of molecular orbitals of nitrogen and phosphorus compounds. The method is illustrated by using data for several nitrogen- and phosphorus-containingspecies.
Introduction There is ample evidence, both experimental and theoretical, that shifts in strictly nonbonding valence orbital ionization potential are approximately eight-tenths of the corresponding shifts in core binding energy.'v2 On the basis of this approximation, core binding energies can be used to correct valence ionization potential shifts for changes in potential and relaxation energy. Thus, one can determine the part of a change in ionization potential that is due to a change in the bonding or antibonding character of the MO. By considering ionization potential changes relative to an MO which by symmetry must be strictly nonbonding, one can quantitatively determine the absolute bonding or antibonding character of MOs. The only molecules with strictly nonbonding valence MOs for which experimental core binding energies are available are H20, H2S, H2Se, and the hydrogen halides. (In an H2A molecule the nonbonding MO is the p a lonepair orbital on atom A, which is perpendicular to the molecular plane. In an HX molecule, the nonbonding MO comprises the two pa lone pairs on atom X, which are perpendicular to the bond axis.) Thus, it is possible to determine, using appropriate core binding energies, the hypothetical localized orbital ionization potentials (LOPS) for the oxygen 2p orbitals of oxygen compounds, for the chlorine 3p orbitals of chlorine compounds, etc. Comparison of actual ionization potentials with LOIPs can be very instructive. For example, if a chlorine "lone-pair" ionization potential is lower than the corresponding chlorine 3p LOIP, the lone pairs must be undergoing repulsive interaction with other electrons in the molecule. If the lone-pair IP is higher than the LOIP, the lone pairs must be undergoing a stabilizing (bonding) interaction with a higher-energy orbital in the molecule. In this paper we calculate the core and lone-pair ionization potentials of two hypothetical molecules: planar NH3 and planar PH3. (Each of these molecules possesses a strictly nonbonding pa MO on the central atom.) The calculated ionization potentials, in combination with other available data, allow the calculation of LOIP values for nitrogen and phosphorus compounds. Thus,quantification (1)Jolly, W.L. J . Phys. Chem. 1981,85,3792. (2)Jolly, W.L.;Eyermann, C. J. J. Phys. Chem.1982,86,4834.LOIPS are calculated by using the relation LOIP = LOIP(ref) + 0 . 8 [ E ~-.EB(reo], where ref refers to the reference molecule whose lone pair is stnctly non bonding. 0022-365418312087-3080$01.50/0
of the bonding or antibonding character of MOs can be extended to a much wider variety of compounds than heretofore. We shall show how the results can aid the interpretation of the ionization potentials of several nitrogen- and phosphorus-containing species.
Computational Details Quantum-mechanical calculations for pyramidal and planar NH33and PH3475have been carried out by other workers using a variety of basis sets. Unfortunately, these studies to not report all of the data necessary to accurately determine the lone-pair ionization potentials and core binding energies of planar NH3 and PH,. To obtain these data we have performed our own calculations using previously reported basis seta and optimized geometries which are of near Hartree-Fock accuracy. We used the HONDO program of King, Dupuis, and Rys6 to calculate the ab initio SCF total energies required. For our calculations on ammonia we have used an optimized Gaussian orbital double-zeta plus d-polarization basis set3because it yields an N-H bond length and HNH bond angle for pyramidal ammonia in excellent agreement with the experimental values.'~~For planar NH3 we assumed that the N-H bond length of 0.991 8, predicted by this basis set is reliable. The phosphorus d-polarization functions and hydrogen basis set used for our calculations on PH3were taken from the work of Lehn and Munsch4 on the inversion barrier of phosphine. To complete our phosphorus basis set we have chosen Dunning and Hay's9 (lls7p/6s4p) contracted bash set. Using this bash set and the optimized geometries for pyramidal and planar phosphine reported by Lehn and Munsch, we obtained for these species total energies within 0.0003 au of those calculated by Lehn and Munsch. (3) Carlsen, N. R.; Radom, L.; Riggs, N. V.; Rodwell, W. R. J. Am. Chem. SOC.1979,101,2233. (4)Lehn, J. M.; Munsch B. Mol. Phys. 1972,23,91. (5) Marynick, D. S.; Dixon, D. A. J. Phys. Chem. 1982,86,914. (6)King, H.;Dupuis, M.; Rys, J. " R C C Software Catalog"; Lawrence Berkeley Laboratory: Berkeley, CA, 1980;Vol. 1, Lawrence Berkeley Laboratory Report No. 10811. (7)Experimental* N-H bond length is 1.012 A; HNH bond angle is 106.7'. Calculated N-H bond length is 1.003 A; HNH bond angle is 107.3'. (8)Kuchitsu, K.; Guillory, J. P.; Bartell, L. S. J . Chem. Phys. 1968, 46,2488. (9)Dunning, T.H.;Hay, P. J., In "Modern Theoretical Chemistry"; Schaeffer, H., Ed.; Plenum Press: New York, 1977;Vol. 3.
@ 1983 American Chemical Society
The Journal of Physical Chemistry, Voi. 87, No. 16, 1983 3081
IPS of Planar NH, and PH,
TABLE I: Calculated Total Energies (au) for NH,, PH,, and Related Ionized Molecules
TABLE 111: Ionization Potentials ( e V ) for NH,, NMe,, PH,, and PMe, ~~
D3h
c3u
NH, NH,' OH,' PH, PH,' SH,'
- 56.200 22 -55.854 86 -76.307 96 -342.456 22 -342.108 37 -398.946 19
-56.190 -55.879 -76.309 -342.128 -342.128 -398.898
14 75 17 13 13 27
TABLE 11: Lone-Pair Ionization Potentials ( e V ) of NH, and PH, (c,uand D 3 h )
KT
A(SCF)
exptl
9.40 10.8' 11.49 C," NH, 10.44 8.45 NH3 D J h NH3 (cor) 9.75 9.85 c3u PH, 10.50 9.46 10.6b 8.05 7.27 D3h pH, 8.14 8.40 Dah PH3 (cor) a Turner, D. W.; Baker, C.; Baker, A. D.; Brundle, C. R. "Molecular Photoelectron Spectroscopy"; WileyInterscience: New York, 1970; p 356. Branton, G. R.; Frost D. C.; McDowell, C. A.; Stenhouse, I. A. C h e m . Phys. L e t t . 1970, 5 , 1. D3h
Therefore, we believe that our choice of basis set also yielded near Hartree-Fock values. The differences in core binding energy between the pyramidal and planar forms of NH, and PH3 were calculated by using the equivalent cores approximation.1° The oxygen basis set for OH3+(the equivalent cores ionization product of NH3) was Dunning's" (Qs5p/4s2p) contracted Gaussian basis set. The sulfur (lls7p/6s4p) contracted Gaussian basis setgwas used for SH3+(the equivalent core ionization product of PH,). The geometries of OH3+and SH3+ were assumed to be identical with those of the neutral NH, and PH3 molecules, respectively, in conformance with the Born-Oppenheimer approximation.
Planar Ammonia and Planar Phosphine Calculations The calculated total energies of the pyramidal and planar forms of ammonia and phosphine are given in Table I. The differences in the total energies of NH3 and NH,', and of PH, and PH3+, are the SCF lone-pair ionization potentials of NH3 and PH3, respectively. These A(SCF) values, along with ionization potentials obtained via Koopmans' theorem (KT), are presented in Table I1 for both the pyramidal (C3J and planar (D3h)forms of NH3 and PH,. The calculated absolute ionization potentials are in poor agreement with the experimental data. However, we note that the shifts in ionization potential between the CBVand D3hforms as calculated from the KT results are in good agreement with those calculated from the A(SCF) results. We believe that these shifts are reliable because of the cancellation of systematic errors in taking the differences and that more sophisticated calculations (e.g., those accounting for electron correlation) would not significantly alter the shifts. In Table I1 we have tabulated the ionization potentials for the D3hspecies, calculated by adding the calculated shifts between the C* and D* forms to the experimental ionization potential of the C,,,form. Thus,we obtain 9.8 eV as the lone-pair ionization potential for planar ammonia, and 8.3 eV as that for planar phosphine. (10) Jolly, W.L. In "Electron Spectroscopy: Theory, Techniques, and Applications";Brundle, C. R., Baker, A D., Eds.; Academic Press: London, 1977; Vol. 1, p 119. (11) Dunning, T. H.J. Chem. Phys. 1970,53, 2823.
-
~
molecule
core EBa
lonepair IP
LOIP
IP LOIP
NH, NMe, PH, PMe,
405.60 404.80 137.02 135.93
10.84b 8.54' 10.5gd 8.6'
10.0 9.4 8.6 7.8
0.8 -0.9 2.0 0.8
'Reference 14. Aue, D. H.; Webb, H. M.; Bowers, M. T. J. A m . C h e m . SOC. 1975, 9 7 , 4136. Elbel, S . ; Bergmann, H.; Ensslin, W. J. C h e m . SOC., Faraday Trans. 2 1974, 7 0 , 555. Branton, G. R.; Frost, D. C.; McDowell, C. A , ; Stenhouse, I. A. Chem. Phys. L e t t . 1970, 5 , 1. Differences in core binding energy may be calculated by using the equivalent cores approximation.1° For example, the difference between the N is binding energies of pyramidal and planar NH3 can be expressed as the energy of the following reaction: NH,(pyramidal)
-
+ OH3+(planar)
OH3+(pyramidal) + NH3(planar)
Because core binding energies are reported as uertical ionization potentials, the OH3+ species in this reaction should have the same geometries as the corresponding NH3 species. Using the total energies for these species from Table I, we calculate that the N 1s binding energy of planar NH, is 0.3 eV lower than that of pyramidal ammonia (405.60 eV12);i.e., the N 1s binding energy of planar NH, is predicted to be 405.3 eV. Similarly, we calculate that the P 2p3I2binding energy of planar PH3 is 0.4 eV lower than that of pyramidal PH, (137.02 eV12). Therefore, the P 2p3 binding energy of planar PH, is assigned a value of 13d.6 eV.
Ammonia, Trimethylamine, Phosphine, and Trimethylphosphine The lone-pair ionization potentials and valence p LOIP values for ammonia, trimethylamine, phosphine, and trimethylphosphine are listed in Table 111. The lone-pair ionization potential of ammonia is 0.8 eV greater than the corresponding LOIP, corresponding to a significant amount of N-H bonding character, as well as some s character, in this orbital. The lone-pair ionization potential of trimethylamine is 0.9 eV lower than the LOIP, indicative of repulsion between the lone pair and the methyl groups. In the case of phosphine and trimethylphosphine, the IP - LOIP values are considerably higher than in the case of ammonia and trimethylamine, probably reflecting an increase in the valence s orbital character of the lone pairs. The IP - L O P value of trimethylphosphine is only 1.2 eV less than that of phosphine, in contrast to a corresponding decrease of 1.7 eV on going from ammonia to trimethylamine. This decrease in lone pair-methyl group repulsion is probably due to the greater size of the phosphorus atom. Bis(trimethylsily1)amine The nitrogen lone-pair ionization potential and nitrogen 1s binding energy of bis(trimethylsily1)amine are 8.66 and 403.32 eV, respectively.lSJ4 From these data and the corresponding ionization potentials of planar ammonia, we calculate a nitrogen p LOIP value of 8.2 eV. Obviously (12) Perry, W.B.;Schaaf, T. F.; Jolly, W. L. J. Am. Chem. SOC.1975,
97,4899.
(13) Starzewski, K.A. 0.;tom Dieck, H.;Bock, H. J . Organomet. Chem. 1974,65, 311. (14) Bakke. A. A.: Chen, H.W.: Jolly, W. L. J. Electron Spectrosc. Relat. Phenom. 1980,20, 333.
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T A B L E IV: I o n i z a t i o n Potentials ( e V ) f o r Pyridine and P h o s p h a b e n z e n e molecule
EBa
b,(n) IPb
LOIP
IP LOIP
pyridine phosphabenzene
404.88 13 5.8
10.5 9.2
9.5 8.9
1.0 0.3
core
a
Reference 14.
R e f e r e n c e 15.
the nitrogen lone pair of bis(trimethylsily1)amine has undergone a net stabilization of about 0.5 eV, which we believe is concrete evidence for delocalization of the lone pair onto the valence d a orbitals of the silicon atoms.
Pyridine and Phosphabenzene On going from pyridine to phosphabenzene, the highest-occupied bl(a) orbital (derived from the p7~orbital of the heteroatom) changes from a next-to-highest occupied MO to a highest occupied M0.16 Although this change suggests considerably reduced involvement of the heteroatom with the carbon A orbitals in phosphabenzene, the valence ionization potentials alone give little quantitative information about such a interactions. The bl(a) ionization potentials and LOIP values for pyridine and phosphabenzene are listed in Table IV. Clearly the b1(r) orbital of pyridine, stabilized by 1.0 eV, is moderately strongly bonding, whereas that of phosphabenzene, stabilized by 0.3 eV, is relatively weakly bonding. It would be of interest to determine the core binding energies of arsabenzene and stibabenzene, to allow a similar quantitative interpretation of the valence ionization potentials and heteroatom a bonding of these analogous molecules. Atomic Nitrogen and the Amino Radical The core binding energy of atomic nitrogen can be estimated from the known heats of formation16of the species N and O+ and the "core replacement energy", A,, empirically determined from data for various nitrogen compounds." It has been shown that the value of 4 corresponding to nitrogen compounds with a valence lone pair (15) Batich, C.; Heilbronner E.; Homung, V.; Aahe, A. J.; Clark, D. T.; Cobley, U. T.; Kilcast, D.; Scanlan, I. J.Am. Chem. SOC.1973,95,928. (16) Rosenstock, H. M.; Draxl, K.; Steiner, B. W.; Herron, J. T. J. Phys. Chem. Ref. Data. 1977, 6, Supplement No. 1. (17) Jolly, W. L.; Gin, C.; Adams, D. B. Chem. Phys. Lett. 1977, 46, 220.
Eyermann and Jolly
(and hence relatively high s electron density at the nitrogen nucleus) is about 0.6 eV higher than that corresponding to nitrogen compounds without a valence lone pair (and relatively low s electron density at the nitrogen nucleus). In the case of atomic nitrogen the 2s orbital is completely nonbonding, and the s electron density at the nucleus is higher than in nitrogen compounds.18 Using the 4 value corresponding to relatively high nuclear s electron density, we calculate EB = 410.9 eV for atomic nitrogen. By combining this with the data for planar ammonia, we estimate a 2p LOIP of 14.3 eV for atomic nitrogen. Analogous calculations for the NH2radical (using experimental heats of formation16for NH, and OH2+)lead to EB = 408.0 eV and LOIP (2p) = 12.0 eV. The experimental ionization potentials of atomic nitrogen and the amino radicallg are 14.53 and 12.45 eV, respectively, values which are slightly higher than the corresponding estimates. This close agreement between the experimental and estimated values is unexpected, because both of these valence ionizations involve loss of an electron from a half-filled 2p orbital. Such ionization potentials would be expected to be 1-2 eV higher than those for filled orbitals because of the lack of Coulombic repulsion associated with two electrons sharing the same orbitdm In the case of atomic nitrogen, the discrepancy would be expected to be even higher because of the loss of exchange stabilization. We have no explanation for the close agreement of the experimental values with the values estimated without taking account of intraorbital repulsion and exchange stabilization. If we were to use a A,, value more appropriate for a very high nuclear electron density, the estimated LOIP values would be higher, in even greater disagreement with expectations.
Acknowledgment. This work was supported by the Director, Office of Energy Research, Office of Basic Energy Sciences, Chemical Sciences Division of the U.S. Department of Energy under Contract Number DE-AOO376SF00098. We are grateful to Dr. Michel Dupuis for his assistance with the HONDO program. Registry No. NH,, 7664-41-7; PH,,7803-51-2. (18) Snyder, L. C.; Basch, H. "Molecular Wave Functions and Properties"; Wiley: New York, 1972. (19) Dunlavey, S.J.; Dyke, J. M.; Jonathan, N.; Morris, A. Mol. Phys. 1980,39,1121. (20) Blake, A. B. J. Chem. Educ. 1981,58, 393.
