J. Phys. Chem. 1981, 85,3792-3797
3792
drogen, from Figures 6 and 7 it is apparent that there may be a transition from a molecular H2 phase to a molecular (HA3 phase. Transition pressures can be derived from Figure 7 by finding the slope of the line that is a common tangent for the curves describing the phases of interest. Table I11 lists the rough ranges for transition pressures thus estimated. Note that the pressure required for transition to an atomic metallic state would not be substantially changed by the existence of an (HJ3 phase. The existence of an (H& phase is more consistent with a more gradual transition8J2rather than a direct first-order change to the metallic ~ t a t e Ramaker . ~ ~ ~ et a1.,22in an ab initio study of solid hydrogen at high pressures (see curves 4 and 4m in Figure 7), found that indeed there is a gradual weakening of the H2 bonds as an extended cluster of H2 molecules is formed. Other workers have suggested various clusters of the metal that include planes or strings of H atoms. Brovman et suggest that a filamentary structure may be a stable form of the metal and may be metastable at zero pressure. Harris and Delhalle'l point out that at zero pressures there would be an explosion to form H2 molecules and that such a chain would need a minimum of 700 kbar to maintain hydrostatic stability. Brovman et also suggest that the metal contains planes of H atoms. The work of Dixon et al.: however, indicates that, of the possible clusters of hydrogen molecules, H6is the most stable. This suggests that, in the process of extended cluster formation, a phase involving termolecular clusters would be a first step. The forces driving such a process are displayed in Figure 2. As the H6 complex dissociates, the molecular self-energy drops, but, as the volume expands, the repulsive forces with neighboring molecules increase. Of course, this is a greatly oversimplified view, since all bonding between the center molecule and its neighbors is neglected. At some pressure, the neighboring molecules will be driven together with such force that bonding may occur. N
The view that there is a gradual transition to a metallic phase is qualitatively consistent with a recent experimental study which looked at the Raman scattering from solid H2 at room temperature to 630 kbar.18 The Hz stretching frequency increased as the pressure was increased to 350 kbar, leveled off nem 400 kbar, and then decreased slightly. The initial increase is expected from compression of the Hz bond.20 Since the formation of clusters of H2molecules at higher pressures is expected to lead to a decrease in frequency (with longer bonds in the clusters), the theoretical pictures is at least qualitatively compatible with this experiment. Perhaps the most useful aspect of the present calculation is the suggestion that species related to transition states may be trapped in compressed solids. The most stable form of the termolecular complex expected in the solid is not the most stable transition state, but rather a distorted, partially dissociated form. Configuration B, in which the molecule is really a bonded cluster of Hz molecules, is a compromise between the loss of energy through dissociation of the molecule and the gain of energy through expansion of the lattice. Since we have not examined other configurations, including all nonplanar ones, we cannot say that this configuration is the most stable; however, it should be more stable than the undissociated molecule. This implies not only a pressure-induced reaction (such as the disproportionation reactions given above) but a pressure-stabilized reaction for which the products could not exist in the gas phase. By performing the reaction in the solid state, the molecules are fixed close together and do not have the entropic handicap suffered by termolecular processes in the gas phase.
Acknowledgment. We have enjoyed discussions of several aspects of ultrahigh-pressure chemistry with R. L. Mills. We are grateful for support of this work by the Department of Energy under Contract DE-ACOB79ER10470.
Use of Core Electron Binding Energies for the Comparison of Valence-Shell Ionization Potentials and the Quantification of the Bonding and Antibonding Character of Molecular Orbitals William
L. Jolly
Department of Chemistry, University of California, and the Materials and Molecular Research Division, Lawrence Berkeley Laboratory, Berkeley, California 94720 (Received:July 27, 198 1)
A localized-orbitalionization potential (LOIP) is defined as the ionization potential that an atomic orbital would have if it were in a hypothetical nonbonding state and were localized on a particular atom in a molecule. The LOIP includes the entire molecular potential at the atom and the electronic relaxation energy associated with the creation of a positive charge on the atom. It is shown that differences in core electron binding energy between compounds of the same element are proportional to the corresponding difference in LOIP. Thus, it is possible to measure quantitatively the relative bonding or antibonding character of actual molecular orbitals, relative to hypothetical nonbonding atomic orbitals, by using experimental valence and core ionization potentials.
The traditional interpretation of a valence-shell photoelectron spectrum consists of the assignment of the various peaks to molecular orbitals (MOs), a procedure which is usually aided by comparison of the measured ionization
potentials with values estimated by quantum-mechanical methods. Knowledge of the ionization potential corresponding to a particular MO is of limited value to a chemist because the absolute magnitude of an ionization
0022-3654/81/2085-3792$01.25/00 1981 American Chemical Society
Bonding and Antibonding of Molecular Orbitals
potential is usually not directly related to a chemically significant quantity such as the bonding or antibonding character of the MO. If we consider an MO to be a linear combination of atomic orbtals, it is logical to measure the bonding or antibonding character of the MO in terms of the difference between the ionization potential of the MO and that of an atomic orbital which constitutes a significant part of the M0.l However, it is important to use for the atomic orbital a special ionization potential, corresponding to a localized, nonbonding orbital, which is the potential a t the atom in the actual molecule less the electronic relaxation energya2 Of course such a “localized-orbital ionization potential” (LOIP) cannot generally be directly measured and would be very difficult to calculate theoretically. However, it is the thesis of this paper that differences in LOIP between compounds are proportional to the corresponding differences in core electron binding energy and that therefore core electron binding energies can be used to establish bench marks for the comparison of valence-shell ionization potentials. Recent studies have shown that lone-pair ionization potentials and the corresponding core binding energies are linearly correlated only for sets of very similar molecules, in which there is little change in the hybridization or delocalization of the lone pair on going from one molecule to anothera3-’ The data indicate that, for such sets of similar molecules, both the lone-pair ionization potential and the core binding energy are affected similarly by changes in potential (atomic charges) and by substituent changes which alter the electronic relaxation accompanying electron emission. Now, whereas core binding energies are only affected by changes in potential and relaxation energy, “lone-pair” ionization potentials are also affected by changes in lone-pair delocalization (i.e., changes in bonding or antibonding character). Thus, it has been observed that, on going from one compound to another in which there is a marked increase in lone-pair bonding character (or a decrease in antibonding character), the lone-pair ionization potential increases much more than the core binding energy. These results strongly suggest that shifts in core binding energy can be used to estimate the corresponding hypothetical LOIP shifts. The well-known observation that core binding energy shifts are approximately independent of the core level can be rationalized by using a crude model of the atom in which each atomic orbital is a spherical shell centered on the nucleus.* If this model were strictly valid, shifts in (1) In simple molecules there is usually no difficulty in identifying the atomic orbitals associated with particular MOs. If an atomic orbital does not constitute a significant part of an MO, the magnitude of this energy difference will have little significance. Obviouslyone must clearly specify the atomic orbital one uses as a reference when measuring the bonding or antibonding character of an MO, and one must somehow know that that atomic orbital is involved in the MO. These points will be clarified later in the discussion of particular compounds. (2) By equating the ionization potential of a localized orbital to the potential less the relaxation energy, we are, in effect, combining several terms (electron flow relaxation, orbital contraction relaxation, rehybridization relaxation etc.) in the relaxation energy term. (3) Lee, T. H.; Jolly, W. L.; Bakke, A. A.; Weiss, R.; Verkade, J. G. J. Am. Chem. SOC.1980,102, 2631. (4) Davis, D. W.; Rabalais, J. W. J. Am. Chem. SOC. 1974, 96, 5305. (5) It has been shown that there is a good eneral correlation of lone-pair ionization potential with proton affinity?$ Hence, evidence that core binding energy is not generally correlated with proton affinity’ is essentially evidence that core binding energy is not generally correlated with lone-pair ionization potential. (6) DeKock, R. L.; Barbachyn, M. R. J. Am. Chem. SOC.1979, 101, 6516. (7) Mills, B. E.; Martin, R. L.; Shirley, D. A. J. Am. Chem. SOC.1976, 98, 2380. Carroll, T. X.; Smith, S. R.; Thomas, T. D. Ibid. 1975,97,659. (8) Siegbahn, K. et al. “ESCA. Atomic, Molecular and Solid State Structure Studied by Means of Electron Spectroscopy”; Almqvist and Wiksells: Uppsala, 1967.
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LOIP would also be exactly equal to the various core binding energy shifts. However, ASCF Hartree-FockSlater calculations have shown that ionization of the valence shells of C, N, 0, F, and Tm2+causes valence ionization potential shifts that are -80% of the corresponding 1s binding energy ~ h i f t s . ~Rough ? ~ calculations from this laboratory indicate that, for a second-row atom such as sulfur, the 3p ionization potential shift is also -80% of the 1s binding energy shift.1° Therefore, we shall assume for all atoms, as a reasonable approximation, that A(L0IP) = 0.8Ah’~ The utility of this approximation can ‘be briefly illustrated as follows. The p a lone-pair orbital of H 2 0 is a strictly nonbonding MO. Its ionization potential (12.61 eV) may therefore be taken as the LOIP of the oxygen 2p orbital in HzO. The 0 1s binding energy of OF2is 5.53 eV greater than that of HzO. We therefore estimate that the LOIP of the oxygen 2p orbital in OF2 is (0.8)(5.53) or 4.4 eV greater than 12.61eV, or 17.0 eV. The observed oxygen lone-pair ionization potential of OFz is 13.13 eV, Le., 3.9 eV lower than the LOIP of the oxygen 2p orbital. Hence, it is clear that the oxygen lone-pair orbital of OF2 is strongly antibonding, and not bonding, as might be naively concluded from the increase in ionization potential on going from H20 to OF2. It should be pointed out that meaningful quantification of bonding and antibonding character from experimental ionization potentials has not heretofore been accomplished. Comparisons of MO ionization potentials with, for example, atom valence-state ioniza ion potentials (VSIP) have lacked significance because IP values do not take account of either the potential due to the other atoms in the molecule or the electronic relaxation energy associated with ionization.” On the other hand, LOIP values do include the effects of these important factors. In the remainder of this paper, we shall demonstrate how core and valence ionization potentials for some relatively simple compounds may be used to evaluate quantitatively the bonding or antibonding character of MOs. It will be seen that the results are in good accord with modern views of the bonding in the compounds. We hope the results will give chemists the confidence to apply the method to more complicated or unusual molecules, in which the bonding is poorly understood. The method should also aid photoelectron spectroscopists in the assignment of spectra.
&
Bonding and Antibonding Character of MOs We shall use the difference between the ionization potential (IP) of an MO and the LOIP of an atomic orbital as a measure of the bonding character of the MO relative to the atomic orbital. The quantity A is defied as follows: A = IP - LOIP If ,!&(ref) and IP(ref) are the core binding energy and the ionization potential, respectively, of a reference compound for which we assume IP(ref) = LOIP(ref), then we may write A = IP - IP(ref) - 0 . 8 [ E ~- E ~ ( r e f ) ]
A positive value of A corresponds to bonding character in (9) Carlson, T. A. “Photoelectron and Auger Spectroscopy”; Plenum Press: New York, 1975; pp 167-70. (10) Bakke, A. A., unpublished data. (11)McGlynn, S. P.; Vanquickenborne, L. G.; Kinoshita, M.; Carroll, D. G. “Introduction to Applied Quantum Chemistry”; Holt, Reinhart and Winston: New York, 1972. On pp 112-13, the shortcomings of VSIPs are discussed.
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Jolly
TABLE I: “Lone-Pair” Valence-Shell and Nitrogen Is Ionization Potentialsa species
adiabatic IP, eV
N “2
“3
CH3NH, (CH,),” (CH3)3N C,H,”, NF3 C5HJ N3P3F6
-
orbital symmetry
14.12‘ 11.3 10.16 8.97 8.24 7.82 10.45e 13.00 9.27 10.8f 15.58 13.59 13.1ph 12.7$ 13.83 16.39
PA A U
U U
0 U U U
u and
A
U
EB ( N Is), eV
LOIP, eV
410.gd 408.0d 405.60 405.17 404.93 404.81 405.31 414.2 404.88 405.29 409.93 406.8 405.55 404.95 406.0 408.6’
13.6 11.3 9.4 9.0 8.8 8.7 9.1 16.3 8.8 9.1 12.8 10.3 9.3 8.9 9.7 11.8
eV 0.5 0.0 0.8 0.0 -0.6 - 0.9 1.3 -3.3 0.5 1.7 2.7 3.3 3.8 3.8 4.1 4.6
%N= Llg U CH,CN U C,H,CN NSF, U U N2 0 a Unless otherwise stated, ionization potentials are from Rosenstock e t al. (Rosenstock, H. M.; Draxl, K.; Steiner, B. W.; Herron, J. T. J. Phys. Chem. Ref. Data 1977, 6, Supplement No. 1 and core binding energies are from Bakke e t al. (Bakke, A = I P - 0.8E,(N I s ) + 315.1. A. A.; Chen, H.-W.; Jolly, W. L. J. Electron Spectrosc. Relat. Phenom. 1980, 20, 333. Valence-state ionization potential for 2p level of nitrogen is given by Hinze and Jaffg.“ Reference compound: NH,. Calculated by the equivalent cores method. See: Jolly, W. L.; Gin, C. Int. J. Mass Spectrom. Ion Phys. 1977, 25, 27. e Reference 15, p 309. f Reference 18. g Reference 23. Reference 30. Griebel, R.; Hohlneicher, G.; Dorr, F . J. Electron Spectrosc. Relat. Phenom. 1 9 7 4 , 4 , 185. Neijzen, B. 3. M.; De Lange, C. A. Ibid. 1978, 14, 187. j Cowan, D. 0.; Gleiter, R.; Glemser, 0 . ;Heilbronner, E. Helv. Chin. Acta 1972, 55, 2418. k Terminal nitrogen atom.
--
TABLE 11: “Lone-Pair’’ Valence-Shell and Oxygen 1 s Ionization Potentialsa species
0
HZO
CH,OH (CH3)20 C,H,OH OF2 0 2
co 0 3
(cH3)2c0 CF,O ONF, OPF, OPCI,
adiabatic IP, eV
orbital symmetry
17.9‘ 12.61 10.84 9.96 12.61e 8.37@ 11.22e 13.13 12.07 19.69g 12.56h 13.02h 13.5Yh 12.9Sk 13.77 12.89 11.18
n n
10.88
0
9.70 13.2 13.36: 12.77 11.58‘
U
E , (0 Is), eV
A
545.4d 539.80 538.88 538.7 538.9 538.9 538.9 545.33 543.75f 542.57 546.2’ 541.51:
U
541.5’
u and n
539.84 541.30 541.42 540.3 539.38 537.9 540.74 541.93 539.2 538.0
2P n A TI A
n U
n TI
U
0
A
0 ?I A
A
LOIP, eV 17.1 12.6 11.9 11.7 11.9 11.9 11.9 17.0 15.8 14.8 17.7 14.0 14.0 12.6 13.8 13.9 13.0 12.3 11.1 13.4 14.3 12.1 11.2
eV 0.8 0.0 - 1.0 - 1.8 0.7 -3.5 - 0.7 -3.9 -3.7 4.9 -5.2 - 1.0 - 0.4 0.3 0.0 - 1.0 - 1.8 - 1.4 - 1.4 -0.2 - 1.0 0.6 0.4
Average of valence-state Reference compound: H 0. A = IP - 0.8EB (0 1s) i- 419.23. a See footnote a , Table I. Calculated by using the equivalent cores method and the heats ionization potentials for 0 2p given by Hinze and Jaffe‘. ,a of formation of O(g) and F+(g). See: Jolly, W. L.; Gin, C. Int. J. Mass Spectrom. Ion Phys. 1 9 7 7 , 2 5 , , 2 7 . e Reference 1 5 , Reference 24. Middle atom. J Terp 309. f Weighted average of the two peaks of the spin doublet. g Reference 23. minal atoms. Reference 30, p p 84-6. Reference 27.
the MO, and a negative value of A corresponds to antibonding character. In Tables I-IV we have listed adiabatic12 IPSfor lone-pair13orbitals (derived from valence p orbitals), the corresponding core binding energies, LOIP values, and A values for compounds of nitrogen, oxygen, fluorine and sulfur. In Table V we have listed vertical12 IPSof orbitals derived from metal 3d orbitals, the corresponding 2p3I2binding energies, LOIP values, and A values (12) The use of adiabatic IPSis perhaps preferable because then one compares transitions which include a consistent (zero) amount of vibrational excitation. However, vertical IPS can be used as long as they are consistently used when making comparisons. (13) We are permissive in our use of the term “lone pair”. Although all of the orbitals in the tables have significant density on the indicated atoms, some of them are quite delocalized and have strong bonding or antibonding character
.
for compounds of iron, cobalt, and nickel. Nitrogen Compounds. The half-occupied ?r orbital of the NH2 radical is strictly nonbonding, and we have used it as our reference zero for the 2p orbitals of the nitrogen compounds. The corresponding A value for atomic nitrogen (calculated by using the valence-state ionization potential14 for the N 2 p i ~orbital) is higher than the expected value of zero, but probably not by more than the uncertainty in the estimated EB(N 1s) value. The lone-pair orbital of ammonia has a weak bonding interaction with the hydrogen 1s orbitals, and therefore the small A value of 0.8 eV is reasonable. The gradual decrease in A on going from ammonia to trimethylamine is evidence for the re(14)Hinze, J.; Jaff6, H. H. J. Chem. SOC.1962, 84, 540.
Bonding and Antibonding
of Molecular Orbitals
The Journal of Physical Chemistiy, Vol. 85, No. 25, 198 I 3795
TABLE 111: “Lone-Pair’’ Valence-Shell and Fluorine 1s Ionization Potentialsa
A values of the analogous orbitals of HCN, CH3CN, c g H5CN, NSFBand N 2 0 show greater bonding character. This increase in bonding character is probably due to orbital E, adiabatic sym(F Is), LOIP, better overlap caused by rehybridization.21 In N2, the 2s compd IP, eV metry eV eV, ~ , eV b orbitals essentially make up two nonbonding orbitals and the HOMO is weakly bonding because it is mainly made HF 16.01 IT 694.22 16.0 0.0 -1.3 691.3 13.7 up of the relatively poorly overlapping 2p orbitah. In NzO, nu* XeF, 12.35c 13.7 0.3 ng 691.3 14.00‘ the nitrogen atom attached to the oxygen atom forms two 13.7 1.6 nu 691.3 15.25‘ u bonds using orbitals which are approximately sp hybrids OF, 13.11d n* 695.07 16.7 -3.6 and which therefore give good overlap and in turn cause 15.74d U* 695.07 16.7 -0.9 the terminal nitrogen atom to provide an orbital with more 16.44d n(u) 695.07 16.7 -0.2 s character. Similarly, the u orbitals offered by the groups ONF, 14.83e n 695.27 16.8 -2.0 CH, CCH3, CC6H5,and SF3 in HCN, CH3CN, C6H&N, CH,F 12.53 IT 692.92 15.0 -2.4 n 695.6 17.1 -1.3 CF, 15.8f and NSF,, respectively, have more s character and give OPF, 16.6ge n 695.6 17.1 -0.4 better overlap than the corresponding orbital of a nitrogen u 694.8 16.5 -0.9 BF, 15.56 atom in N2. 14.9 -1.1 13.85g u 692.88 C,H,F Oxygen Compounds. We have used the nonbonding a 18.0 -2.3 15.69 n 696.66 F, orbital of water as the A = 0 reference for the 2p orbitals a See footnote a, Table I. Reference compound: HF. of the oxygen compounds in Table 11. The A value for Reference 29. RefA = IP - 0.8EB ( F I s ) + 539.37. atomic oxygen is, perhaps within the probable error of the erence 22. e Reference 27. f Reference 30, p p 221, 237. estimated 0 1s binding energy, also equal to zero. The g Reference 1 5 , p 309. ionization potentials listed for methanol and dimethyl ether correspond to the R lone-pair orbitals. The decrease pulsive interaction between the nitrogen lone pair and the in A with increasing number of methyl groups can be exC-H bonding electrons. Presumably the A value of aniline plained, as with the methyl amines, in terms of increasing would be lower than it is (perhaps near that of direpulsions between the lone pair and the C-H bonding methylamine) because of repulsions by u bonds of the electrons. In the case of phenol, interaction of the outphenyl ring if it were not for a-bonding interaction with of-plane lone pair with the highest occupied phenyl ?r the highest occupied a orbital of the ring. The effect of orbital yields a bonding MO (IP = 12.61 eV; A = 0.7 eV) this a bonding is to raise the A value and to delocalize the and an antibonding MO (IP = 8.37 eV; A = -3.5 eV). The lone-pair electron density onto the ring.15 more nearly nonbonding in-plane lone pair (IP = 11.22 eV) If one considered only the high ionization potential has a A value of -0.7, corresponding to a small amount of (13.00 eV) and the low basicity of nitrogen trifluoride, one bonding pair-lone pair repulsion. might conclude that the nitrogen lone pair is in a bonding In oxygen difluoride, the ionization potential correorbital. However, the nitrogen atom of NF, has a high sponds to the antibonding bl(a) orbital.22 Thus, even positive charge. When one eliminates the effect of this though the ionization potential is fairly high (13.13 eV), high charge by calculating A, it becomes clear that the the A value is low (-3.9 eV). Similarly, the ionization nitrogen lone pair is essentially in an antibonding orbital potential of molecular oxygen corresponds to the antiand that the extra electron binding energy is due to the orbital, and the A value is, as expected, very bonding rTg charge on the nitrogen atom. This antibonding character negative. is a consequence of repulsions by the fluorine lone pairs. There is no truly nonbonding pair of electrons on the The same repulsive interaction has been used to explain oxygen of carbon monoxide. However, the third band in the remarkable weakness of the N-F bonds in NF3.16 the photoelectron spectrum of carbon monoxide (IP = The A value of pyridine is about the same as that of 19.69 eV)23corresponds to a bonding (r orbital on carbon ammonia, apparently a consequence of a balance between and oxygen.20 The A value of 4.9 eV is consistent with this increased bonding pair-lone pair repulsion and increased assignment and interpretation. s character. The bonding pair-lone pair repulsion cannot The first ionization potential of ozone (12.56 eV) corbe very great in pyridine, and therefore we conclude that responds to a strongly antibonding al(u) orbital on the the pyridine lone pair does riot have much s character. middle oxygen atom; the A value of -5.2 eV is entirely Indeed, CNDO/2 calculations show that the hybridization consistent. The second ionization potential (13.02 eV) has changes from 6.8% s character in ammonia to only 15.4% been assigned to the nonbonding la2(a) orbital on the s character in pyridine.l’ The latter precentage is much terminal atoms, and the third ionization potential (13.57 less than the 33% corresponding to sp2hybridization. In eV) has been assigned to the weakly antibonding 4b2(u) the structurally related (NPF2)3molecule, the A value of orbital on the terminal atoms.24 Although the A value of 1.7 eV indicates significant a bonding between the nitrogen -0.4 eV seems reasonable for the weakly antibonding 4b2 and phosphorus atoms.ls orbital, the A value of -1.0 eV for the nonbonding l a 2 The A value of 2.7 eV for the HOMO of molecular niorbital seems too negative. Indeed, BrundleZ4has indicated trogen corresponds to the usual description of this MO as that the data do not strongly support the reported 4b2,1a2 a weakly bonding u orbital mainly derived from the 2pa order, and perhaps the reverse assignment should be seorbitals, but with a small amount of 2s c h a r a ~ t e r . ’ The ~ ~ ~ ~ riously considered. The 12.98 eV ionization potential of SOz corresponds to the essentially nonbonding la2 orbital ~~~
(15) Rabalais, J. W. “Principles of Ultraviolet Photoelectron Spectroscopy”; Wiley-Interscience: New York, 1977; pp 306-8. (16) Jolly, W. L. Inorg. Chem. 1964, 3, 459. (17) We are grateful to Mr. A. A. Bakke for these calculations. (18) Branton, G. R.; Brion, C. E.; Frost, D. C.; Mitchell, K. A. R.; Paddock, N. L. J . Chem. SOC.A . 1970, 151. (19) Purcell, K. F.; Kotz, J. C. “Inorganic Chemistry”; Saunders: Philadelphia, 1977; pp 143-4. (20) Streitwieser, A., Jr.; Owens, P. H. “Orbital and Electron Density Diagrams”; Macmillan: New York, 1973.
(21) For a discussion of the relation between overlap and hybridization, see: Maccoll, A. Trans. Faraday SOC.1950,46, 369. Bent, H. A. Chem. Reu. 1961,61, 275. (22) Brundle, C. R.; Robin, M. B.; Kuebler, N. A. Basch, H. J. Am. Chem. SOC. 1972,94, 1451. (23) Potts, A. W.; Williams, T. A. J.Electron Spectrosc. Relat. Phenom. 1974,3,3. (24) Brundle, C. R. Chem. Phys. Lett. 1974,26, 25.
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Jolly
TABLE IV: “Lone-Pair’’ Valence-Shell and Sulfur 2p,, Ionization Potentialsa adiabatic IP, eV
compd
orbital symmetry
10.47 9.44 8.68 11.11 9.04 9.4d 10.08 11.18 9.63e 12.34 12.4f 9.3g 11.3g 12.6g
EB
(S 2 ~ , , , ) eV ,
170.20 169.4 169.09 172.51 170.1‘ 171.57 169.80 170.60 168.87 174.80 177.8‘ 171.91 174.53 176.20
LOIP, eV
A , &e v
10.5 9.8 9.6 12.3 10.4
0.0 -0.4 -0.9 -1.2 - 1.4 -2.2 -0.1 0.4 0.2 - 1.8 -4.2 -2.5 -2.6 -2.7
11.6 10.2 10.8 9.4 14.2 16.6 11.8 13.9 15.3
‘
a See f o o t n o t e a, Table I. Reference compound: H,S. A = IP - 0 . 8 E ~(S 2p,,,) i- 125.69. Calculated from E , (S 2p) by subtracting 0.4 eV. Colton, R. J.; Rabalais, J. W. J. Electron Spectrosc. Relat. Phenom. 1974, 3, 345. e Bunzli, J. C.; Frost, D. C.; McDowell, C. A. J. Electron Spectrosc. Relat. Phenom. 197213, 1 , 481. f Reference 31. Chadwick, D.; Cornford, A. B.; Frost, D. C.; Herring, F. G.; Katrib, A , ; McDowell, C. A.; McLean, R. A. N. In “Electron Spectroscopy”; Shirley, D. A., Ed.; North-Holland Publishing Co.: Amsterdam, 1 9 7 2 ; p 453.
TABLE V: Metal 3 d and 2p,,, Ionization Potentialsa
compd Fe( co 1 5 FeCp, HCo(CO), Co(CO),NO COCP, Ni(CO), W P F , 14 CpNiNO NiCp,
orbital E, vertical sym- (M 2 p ) LOIP, IP, eV metry ev3’* ’ ev
A,eV ~~
8.6 6.86 8.90 8.90 5.66 8.90 9.69 8.48 6.51
e’ eZg e e eig
t, t, a, elg
715.79 713.05 786.86 786.85 784.83 861.15 862.00 861.36 859.87
8.6 6.4 8.9 8.9 7.3 8.9 9.6 9.1 7.9
O.Ob 0.5 0. O b 0.0 -1.7 O.Ob 0.1 -0.6 -1.4
a Ionization potentials are from Cowley (Cowley, A. H. Prog. Inorg. Chem. 1979, 26, 4 5 ) and core binding energies are from Bakke et al. (Bakke, A. A,; Chen, H.-W.; Jolly, W. L. J. Electron Spectrosc. Relat. Phenom. 1980, 20, 333). Reference 3d level,
on the terminal oxygem atoms of this molecule. The small positive A value of 0.3 eV is consistent with a proposed small amount of d t - p r bonding character in this orbita].2426 The ionization potential of COz corresponds to removal of an electron from the nonbonding tgorbital, and the A value of 0.0 eV is reasonable. In the case of N20 and OCS, the analogous oxygen lone pairs have A values of -1.0 and -1.8 eV, respectively, suggesting appreciable antibonding character. This result is a consequence of the fact that the nonbonding ‘IT orbitals of N20 and OCS have some density on the opposite terminal atoms as well as the oxygen atoms. Because of the lower electronegativities of nitrogen and sulfur, the orbital energies are higher than they would be if the orbitals were localized completely on the oxygen atoms. In formaldehyde, acetone, and carbonyl fluoride, the ionization potentials correspond to the in-plane nonbonding p orbitals of the oxygen atoms. In formaldehyde and acetone, bonding pair-lone pair repulsions give some antibonding character to the orbitals and cause the rather low A value, -1.4 eV. In carbonyl fluoride, the C-F bonding electrons are held so tightly by the fluorine atoms that they do not strongly repel the oxygen lone pair. An equivalent explanation is that the C-F CJ bonding orbitals (25) Linnett, J. W. Discuss. Faraday Soe. 1963, 35, 226. (26) Hillier, I. H.; Saunders, V. R. Mol. Phys. 1971, 22, 193.
have such low energies that they do not interact appreciably with the oxygen lone-pair orbital. The ionization potentials of ONF,, OPF,, and OPCl, have been assigned to the degenerate nonbonding t orbitals on the oxygen atoms.27 It is generally believed that the electrons in these orbitals are engaged in some t bonding to the central atom either through hyperconjugative interaction with the u* orbitals of the N-F, P-F, and P-C1 bonds or, in the case of OPF3 and OPCl,, through the use of the empty phosphorus 3d orbitals.28 Consequently the slightly positive A values of OPF, and 0PCl3 are quite reasonable. However, the -1.0 eV value of ONF, is incredible. We believe that the 0 1s binding energy of ONF3 is in error by at least 2 eV or that the ultraviolet photoelectron spectrum has been misassigned. Bassett and Lloyd themselves point out several inconsistencies in their assignment of the UPS spectrum of ONF3.27 In order to have the A value consistent with the x bonding in the N-0 bond of ONF,, it would be necessary to assign the 5e T orbital on the oxygen atom to one of the higher ionization-potential bands, such as that at IP = 16.34 eV. Fluorine Compounds. We have used the nonbonding P‘ITorbital of HF as the reference level for the 2p orbitals of fluorine compounds. In XeF,, the antibonding, nonbonding, and bonding t orbitals have A values of -1.3,0.3, and 1.6 eV, r e s p e c t i ~ e l y . The ~ ~ relative magnitudes and signs of these are all quite reasonable. The first and second ionization potentials of OFz correspond to orbitals which are T and u antibonding, respectively, and the A values of -3.6 and -0.9 eV are appropriate for such The third ionization potential corresponds to the nonbonding u orbital on the fluorines; the A value of -0.2 eV is quite reasonable. In ONF3, CH,F, and CF4,the fluorine lone pairs acquire antibonding character by repulsive interaction with the three CJ bonding orbitals corresponding to the bonds between the central atom and the other three ligand atoms. Thus, the A values are quite negative (-2.0; -2.4; and -1.3eV, respectively). In the case of O P F , the central atom is relatively big, the lone pair-bonding pair interaction is (27) Bassett, P. J.; Lloyd, D. R. J . Chem. Soc., Dalton Trans. 1972, 248. (28) Frost, D. C.; Herring, F. G.; Mitchell, K. A. R.; Stenhouse, I. A. J. Am. Chem. SOC.1971, 93, 1596. (29) Brundle, C. R.; Robin, M. B.; Jones, G. R. J. Chem. Phys. 1970, 52, 3383.
Bonding and Antibonding of Molecular Orbitals
less intense, and A is less negative. In BF3 and C6H5F,the in-plane fluorine lone pairs interact repulsively with the two adjacent bonding pairs; however, because of the greater bond angle and the presence of only two repulsive bonding pairs, the overall repulsion is significantly less than in the case of the tetrahedral molecules of first-row elements. The A values of -0.9 and -1.1 eV, respectively, are therefore reasonable. The ionization potential of elemental fluorine corresponds to the a antibonding orbitals, and thus the very negative A value (-2.3 eV) is not unexpected. Sulfur Compounds. The A values of H2S, CH3SH, (CHJ2S, and (CF2)2Sare consistent with A lone pairs having no repulsive interactions in H2S (our reference compound) and repulsions with bonding pairs in the other molecules. As expected, the magnitudes of the A values of CH3SH and (CH3)2Sare significantly less than those for the corresponding oxygen compounds. The sulfur lone pairs of S8and S2C12occupy antibonding orbitals, with significant lone pair-lone pair repulsion, and this is antibonding character is reflected in the A values of -1.4 and -2.2 eV, respectively. The a nonbonding orbital of CS2,situated on the sulfur atoms, does not interact significantly with any other orbital and therefore has a A value near zero. The corresponding orbital of OCS is situated on both the oxygen and sulfur atoms, and, because of the greater electronegativity of oxygen, the orbital energy is lower than it would be if it were localized on the sulfur atom. Thus, we can rationalize the slightly positive A value. The A value of the degenerate a orbitals on the sulfur atom of SPC13is near zero, indicating that these orbitals do not interact strongly with other orbitals. Calculations have shown that the sulfw lone pairs of SO2 and SF4 are antib~nding,~OBl and one would expect the (30) Turner, D. W.; Baker, C.; Baker, A. D.; Brundle, C. R. “Molecular Photoelectron Spectroscopy”; Wiley-Interscience: London, 1970. (31) Cowley, A. H.; Lattman, M.; Walker, M. L. J. Am. Chem. SOC. 1979,101, 4074.
The Journal of Physical Chemistry, Vol. 85,No. 25, 198 1 3797
same to be true for the closely related compounds (CH3)2S0, SOCl2, and SOF2. Thus, the very negative A values for all of these compounds are understandable. Transition-Metal Compounds. The A values in Table V measure the bonding or antibonding character of the highest occupied orbitals with mainly metal 3d character. We have taken the HOMOS of Fe(C0)5, HCO(CO)~, and Ni(C0)4 as the reference 3d levels for iron, cobalt, and nickel, respectively. In Fe(C0)5and HCO(CO)~, the HOMOS are essentially the metal d,zxYz and d,, orbitals, and in Ni(C0)4 they are mainly the metal dxy, d,,, and d,, orbitals. These orbitals undergo both u interactions (energy-raising) and a interactions (energy-lowering)with the carbonyl groups, and therefore we believe it is reasonable to use these orbitals as the nonbonding references for other compounds of iron, cobalt, and nickel. On this basis we find that the A value for the e2 orbital of ferrocene is 0.5 eV, corresponding to weak boding character, and that the A values for the elg orbitals of cobaltocene and nickelocene correspond to antibonding character. These results agree with the usual ligand-field energy-level scheme for the metall~cenes.~~ We also find that A for both CO(CO)~NO and Ni(PF3)4is essentially zero, in agreement with the general opinion that the a-acceptor abilities of NO and PF3 are quite similar to that of CO. On the other hand, the HOMO of CpNiNO has significant antibonding character, consistent with the weak a-acceptor ability of the cyclopentadienyl group.
Acknowledgment. This work was supported by the National Science Foundation (Grant CHE-7926097) and the Director, Office of Energy Research, Office of Basic Energy Sciences, Chemical Sciences Division of the U S . Department of Energy, under Contract No. W-7405-Eng48. I thank Professor T. D. Thomas of Oregon State University for a helpful conversation and Mr. Albert A. Bakke for certain potential calculations. (32) Warren, K. D. Struct. Bonding (Berlin) 1976,27, 45.
