1432
J. Phys. Chem. 1995, 99, 1432-1435
Theoretical Calculation of Intrinsic Acidity and Basicity of FOH Peeter Burk and Ilmar A. Koppel" Institute of Chemical Physics, Tartu University, Jakobi 2, EE2400 Tartu, Estonia
Alar Rummel and Aleksander Trummal Institute of Chemical Physics and Biophysics, Ravala 10, EEOOOI Tallinn, Estonia Received: July 5, 1994; In Final Form: October 31, 1994@
A b initio calculations have been used to study the intrinsic basicity and acidity of hypofluorous acid (FOH), which plays a key role as an intermediate in the interaction of F2 with water. At the G2 level of theory, the proton affinity and deprotonation enthalpy of FOH were predicted to be 135.3 and 360.3 kcal/mol at 298 K, respectively. In particular, the latter value is in excellent agreement with the best available thermodynamically calculated one. The computational results show that FOH is a significantly stronger acid (by ca. 24 kcaV mol) and a weaker base (by ca. 30 kcaYmol) than H20. The influence of the replacement of one hydrogen atom in water molecule with F atom on its acid-base properties is rationalized in terms of field-inductive, polarizability, resonance, and adjacent lone pair- lone pair interaction effects.
Introduction It has been known since the time of Moissan' that fluorine reacts vigorously with water. The principal products of this reaction are usually HF, 0 2 , H202, and O F I . ~In the early 1930s, it was claimed by Dennis and R o ~ h o w that ~ . ~they succeeded in liberating hypofluorous acid from the solutions of oxy salts of fluorine made by passing F2 into aqueous alkali followed by the neutralization of the excess alkali. However, the unambiguous confirmation of the existence of FOH came only in 1968 in the work of Noble and Pimentel, who isolated it in a solid N2 matrix at 14-20 K.5 Later it was discovered that FOH plays the key role of an intermediate in the interaction of FZ with water.6 Besides that, the electronic structure and reactivity of this molecule have attracted the attention of theoretical chemists because of the presence of the highly electronegative F atom and the effect of the lone pair-lone pair interactions of adjacent oxygen and fluorine atoms in the relatively small (one hydrogen and two heavy atoms) molecule. During the last 2 decades, FOH was experimentally characterized by its microwave ~ p e c t r aIR , ~ spectra in the gas phase,* in the nitrogen and argon m a t r i ~ e sand , ~ in the solid state,'O Raman spectra," photoelectron spectra,I2and reactivity.6 An excellent review of the earlier studies of the chemistry and physics of this molecule can be found in ref 13. There are numerous papers reporting the quantum chemical investigation of FOH. So the equilibrium structure and anharmonic force-field constant^,'^^'^ the energy of its ground state and homolytical d i s s ~ c i a t i o n , ' rotational ~~'~ and vibrational spectra,I7and gas-phase acidity and basicity'8-zo have been calculated. However, no high-level calculations or experimental work on the acid-base properties of FOH have been performed so far. As a rule, the existing empirical predictions of these quantities are at variance with each other as well as with those calculated from the existing thermochemical data. So the gasphase acidity of FOH could be predicted on the basis of the multiparameter correlation equation to be 362 kcaVmol,*I from which a proton affinity of around 369 kcaVmol could be estimated. Thermochemical estimates of the proton affinity of FO- on the basis of the existing thermochemical data on the @
Abstract published in Advance ACS Abstracts, January 1 , 1995.
0022-3654/95/2099- 1432$09.00/0
FOH molecule [AHdFOH) (-23.0 kcal/mo122 or -19.9 kcaY moll6), D(F0-H) = 98.5 kcal/m01~~] and suggested electron affinities of FO' (1.4-2.15 eV23,24)lead to values in a wide range from 362 to 380 kcal/mol, whereas the most probable value, using EA(FO') = 2.15,24seems to be in the range 361.8364.8 kcal/mol. Unfortunately, no experimental data on AHf(FOH2+) are available for the thermodynamic estimation of the proton affinity of the neutral FOH molecule. In the present work, high-level ab initio calculations (up to G1 and G2 levels of theory) of the acidity and basicity of FOH were performed to investigate the influence of the abovementioned effects (the presence of the highly electronegative F atom and the lone pair-lone pair interactions of the adjacent oxygen and fluorine atoms) on the acidity and basicity of the title compound. Methods Standard ab initio molecular orbital c a l c ~ l a t i o n sof~ ~the neutral, protonated, and deprotonated FOH molecule were performed with full optimization of geometry using the STO3G, 3-21G, 3-21+G, 6-31G, 6-31G*, 6-31G**, and 6-31++G** basis sets as implemented in the Gaussian92 system of programs.26 With the 6-3lG* basis set, post-Hartree-Fock level (MP2, second-order Moller-Plesset perturbation theory; CISD, configuration interactions with inclusion of singles and doubles) calculations with geometry optimizations were also carried out. G12' and G2**calculations of energies of the above-mentioned species were performed. G1 theoryz7is based on MP2=FuIl/6-3 lG* (second-order MP perturbation using all electrons) geometries. The energies are calculated at the MP4DSTQ/6-31 lG(d,p) level (complete fourthorder MP perturbation) with corrections due to diffuse-sp functions on non-hydrogen atoms and higher polarization functions on non-hydrogen atoms and for correlations beyond the fourth-order perturbation theory using quadratic configuration interaction. A final high-level correction, to make Ee exact for the hydrogen atom and hydrogen molecule, is added. Finally, the total energy (EO)is obtained by adding a zeropoint energy (from scaled HF/6-31G* frequencies) to E,. In the G2 theory,28the EO,calculated as given above, is further corrected for nonadditivity caused by the assumption of separate
0 1995 American Chemical Society
Intrinsic Acidity and Basicity of FOH
J. Phys. Chem., Vol. 99, No. 5, 1995 1433
TABLE 1: Calculated Geometries, Total Energies, and Dipole Moments for Neutral FOH" STO-3G 3-21G 3-21+G 6-31G 6-31G* 6-31G** 6-31++G** MP2/6-3lG* CISD/6-31G* GI G2 e~ptl~.~'
r(0-H)
r(F-0)
1.006 0.976 0.978 0.960 0.952 0.949 0.950 0.979 0.971 0.979 0.979 0.966
1.355 1.444 1.447 1.440 1.376 1.376 1.377 1.444 1.426 1.444 1.444 1.442
A,
"Bond lengths in moments in D.
LH-0-F 101.38 98.96 99.71 100.04 99.80 100.06 100.24 97.18 98.05 97.8 97.8 96.8
E
P
-172.374 209 -173.800673 -173.844'800 -174.684 918 -174.729 582 -174.735 917 -174.744 110 -175.092 856 -175.071 968 -175.352024 -175.353 387
1.468 2.429 4.433 2.595 2.155 2.150 2.220 2.262 2.235 2.230
angles in deg, energies in au, and dipole
TABLE 2: Calculated Geometries and Total Energies for the FO- Anion" E
r(F-0) STO-3G 3-21G 3-21+G 6-31G 6-31G* 6-31G** 6-31++G** MP2/6-3 1G* CISD/6-3lG* G1 G2 Bond lengths in
1.400 1.525 1.559 1.564 1.490 1.490 1.497 1.498 1.507 1.498 1.498
-171.552 631 -173.189 786 -173.305 672 -174.114 384 -174.126 312 -174.126 312 -174.161 524 - 174.478 805 - 174.460902 -174.780 963 -174.781 120
A and energies in au.
TABLE 3: Calculated Geometries and Total Energies for the Protonated Form of FOHn r(F-0) r(0-H) STO-3G 3-21G 6-31G 6-31G* 6-31G** 6-31++G** MP2/6-31G* CISD/6-31G* G1 G2
1.370 1.428 1.413 1.354 1.352 1.351 1.427 1.408 1.427 1.427
Bond lengths in
1.010 0.995 0.975 0.978 0.973 0.973 1.008 0.998 1.008 1.008
LH-0-F 110.24 109.73 111.52 106.12 106.78 106.69 101.89 103.39 101.89 101.89
LH20FH'
E
124.34 136.33 143.75 120.62 122.69 122.24 114.30 115.85 114.30 114.30
- 172.687846 -174.031 301 -174.903 446 -174.954 602 - 174.968 308 -174.971 431 -175.319047 -175.299 008 -175.564 340 -175.566 680
A, angles in deg, and energies in au.
basis set extensions for diffuse-sp functions and higher polarization functions in G1 theory and for the addition of a third d function to the non-hydrogen atoms and a second p function to the hydrogen atoms. Finally, the high-level correction used in G1 theory is modified to give the zero-mean deviation from the experiment of the calculated atomization energies of 55 molecules having well-established experimental values.
Results and Discussion The results of the calculations of FOH, FO-, and FOH2+ are presented in Tables 1-3. In the latter case, both altematives, 0 protonation and F protonation, were considered. In accordance with earlier literature findings,20 the protonation on the 0 atom is more favorable by 18 kcdmol at the G2 level of theory. Only the results for the 0-protonated form are given in Table 3. The electron correlation is known to be important for molecules with 0 - F bonds,I5so it must be explicitly accounted for for reliable predictions for these systems. In the case of FOH, the inclusion of electron correlation (within the 6-31G*
basis set) expands the calculated 0-F bond length (around 0.04 A) and 0-H bond length, while the F-0-H angle decreases (see Table l), resulting in a good overall agreement with experiment. It is noteworthy that the 3-21G basis set also gave the geometry of FOH in good agreement with experiment. Although for FOH2+ the experimental geometry is not known, the above trend of bond lengthening and angle decreasing by inclusion of electron correlation holds (Table 3) and probably also leads to the better agreement with experimentalvalues. The bond lengthening in FO- is much smaller (Table 2) and is possibly caused by much stronger electrostatic repulsion between the oxygen and fluorine atoms. When comparing the deprotonation enthalpies calculated at different levels of theory (see Table 4),one can see that the results cover a very wide interval from 331 up to 508 kcaY mol. The span for the calculated proton affinities for FOH is much smaller, especially when excluding the very strongly deviating point from the STO-3G calculations. In the latter case, the scatter of the calculated proton affinities covers only a 10 kcdmol wide region in the PA scale. Thus, it can be concluded that the above-mentioned scatter of the deprotonation enthalpy is mainly due to the difficulties of the calculations of the anion. Hence, it becomes clear that the calculations of FO- depend very significantly on the choice of the basis set. This sensitivity can be attributed to the very high electrons per orbital ratio in these species. So it can be assumed that for the correct calculation of the deprotonation enthalpy of hypofluorous acid, very large basis sets, incorporatingboth polarization and diffuse functions, are needed. An attempt has been made to correct the obtained results by using an empirical linear correlation relationship between the calculated and experimental proton affinity values for a large variety of bases (both neutral and However, as can be seen from Table 4,the results did not improve significantly. This is probably caused by the above-mentioned high electrons per orbitals ratio in the considered compounds. As the final total energies obtained in the G2 theory are effectively at the QCISD(T)/6-311+G(3df,2p) leve128-30 and thus correspond to the highest level of theory applied presently, we expect that the corresponding values of the proton affinity and deprotonationenthalpy (133.8 and 359.1 kcdmol, respectively) of FOH are the best available estimates of those values at 0 K. This assumption is further c o n f i i e d by the extensive comparisons of the experimental proton affinities and those calculated at the G2 We note also that both the G1 and the G2 levels of theory gave relatively consistent results (the difference is less than 1.0 kcdmol). The results given above were corrected to 298 K (using the standard procedure31)where the experimentalresults are usually reported, giving the values for the proton affinity and deprotonation enthalpy, 135.3 and 360.3 kcdmol, respectively. The latter value is rather close to the best thermodynamical value (363 f 2 kcaYmo1) estimated in this work (vide supra) and to the quantity obtained for FOH using a multiparameter correlation analysis of the contributions from field-inductive, polarizability, and resonance effects (AGacid= 362 k c a Y m 0 1 , ~DPE ~ ~ ~=~ 369 kcdmol). One might assume that the change of one H atom in the H20 molecule for fluorine should, due to the extremely high electronegativity (field-inductive effect) of the latter, result in a compound which must be significantly more acidic than water. The counteracting effects on the acidity are expected to be displayed by the increased lone pair-lone pair repulsion in the FO- anion as compared to the neutral acid, FOH, and by the slightly smaller polarizability of the F atom as compared to the
Burk et al.
1434 J. Phys. Chem., Vol. 99, No. 5, 1995
TABLE 4: Comparison of Proton Affinities and Deprotonation Enthalpies for FOH (kcaVmol) Calculated at Different Levels of Theorv STO-3G 3-21G 3-21+G 6-31G 6-31G* 6-3 1G** 6-31 ++G** MP2/6-31G* CISD/6-31G* G1 G2
DPE
DPE"
PA
PA"
515.5 (384.0)b 383.3 (372.2)b 338.3 (352.4)b 358.0 378.8 (367.6)b 382.5 365.5 385.3 383.4
507.9 375.7 330.7 350.4 371.1 374.9 358.0 377.7 375.8 358.3 359.1 (360.3)c
196.8 (172.3)b 144.7 (150.2)b
189.1 137.0
137.1 141.2 (149.7)b 145.8 142.6 141.9 142.5
129.4 133.5 137.1 134.9 132.2 134.8 133.2 133.8 (135.3)'
a Corrected to include AZPVE calculated as 0.8929AZPVE/6-31G*. Corrected using linear correlation between experimental and calculated proton a f f i n i t i e ~ . Corrected ~~ at 298 K.
hydrogen atom. Since, in fact, the calculated acidity is much higher for FOH than the corresponding experimental quantity for H20 (DPEFoH= 360.3 kcavmol and D P E H , = ~ 390.8 kcav mol33),the two latter effects must have much less influence on the acidity of FOH than the first one (Le., field-inductive effect). Similarily, the basicity of FOH should be much less than that of H20 due to the stronger field-inductive effect and also to the slightly lower polarizability of the F atom. Simultaneously, due to the protonization of FOH, the counteracting basestrengthening effect of the elimination of the adjacent lone pair repulsion should be present. As can be seen from comparison of the proton affinities for H20 (166.5 kcavm01~~) and FOH (135.3 kcavmol, this work), the first two effects seem to be dominant. Since there are no acceptor orbitals for either FOH, FO-, or FOH2+, the resonance effects on the acidity or the basicity of the FOH molecule could be considered inoperative. The comparison of the calculated acidity of FOH with the suggested experimental value for CF3OH (AGacld= 341 k c d DPE x 348 kcavmol) shows that the latter is more acidic than FOH by ca. 12 kcal/mol. On the assumption of the roughly similar field-inductive effects of the F atom and CF3 group2'-32,34 and contribution (ca. 2.3 kcal/mo13*) to the stronger acidity of CF30H due to the higher polarizability of the CF3 group as compared to the fluorine atom, the remaining gap, 12 - 2.3 = 9.7 kcavmol, is believed to be caused by the above-mentioned destabilizing lone pair-lone pair repulsion effect, since in CF3OH and in its anion there are no lone pairs in the a-position to the deprotonation center, and by some contribution due to the acid-strengtheningresonance (anionic hyperconj~gation~~) effect which accompanies the acidic dissociation of CF30H. A somewhat higher contribution of the polarizability effect and, consequently, less significant lone pair-lone pair repulsion effect could be estimated on the basis of the approach given in ref 21. A similar comparison of the calculated G2 (without ZPVE correction) proton affinity of FOH (141.5 kcdmol) with the calculated proton affinity of CF3OH at the 6-31+G* level (161 kcal/moP6) so far is not strictly justified. However, the rough comparison of these two quantities shows that the latter compound is expected to be up to 20 kcdmol more basic than FOH. By using once more the assumption given above of roughly similar field-inductive effects of the F atom and CF3 group, one must assume that the contribution from the higher polarizability of the CF3 group must be rather significant ('20 kcdmol), as there are no more base-strengthening effects in CF30H, while in the case of FOH there is also the abovementioned elimination of the adjacent lone pair repulsion which must have some base-strengthening effect. Evidently, for a more correct comparison of the proton affinities of FOH and CF3-
OH, further high-level (including G2) calculations of the latter molecule and its protonated form are needed. Conclusions High-level ab initio calculations have been used to study the acid-base properties of the FOH molecule, as well as the energetics, geometry, and electronic structure of the latter, and its deprotonated (FO-) and protonated (FOH2+) forms. It was shown that for current calculation of the deprotonation enthalpy of hypofluorous acid, very large basis sets, incorporating both polarization and diffuse functions, are needed. Also, the need for explicit accounting of the electron correlation for the correct prediction of the geometries of the studied species was confirmed. FOH was found to be a significantly stronger acid and weaker base than H20. The rationalization of these findings is given in terms of field-inductive, polarizability, and (counteracting) lone pair-lone pair repulsion effects. References and Notes (1) Moissan, H. Le Fluor et ces Composes; Masson: Paris, 1900. (2) Cady, G. H. J . Am. Chem. SOC.1935, 57, 246. (3) Dennis, L. M.; Rochow, E. G. J . Am. Chem. Soc. 1932, 54, 832. (4) Dennis, L. M.; Rochow, E. G. J . Am. Chem. SOC.1933,55, 2431. ( 5 ) Noble, P. N.; Pimentel, G. C. Spectrochim. Acta, Part A 1968, 24, 797. (6) Appelman, E. H.; Thompson, R. C. J . Am. Chem. SOC.1984, 106, 4167. (7) Kim Hyunyong; Pearson, E. F.; Appelman, E. H. J. Chem. Phys. 1972, 56, 1. ( 8 ) Pearson, E. F.; Kim Hyunyong, J. Chem. Phys. 1972, 57, 4230. (9) Goleb, J. A.; Claassen, H. H.; Studier, M. H.; Appelman, E. H. Spectrochim. Acta, Part A 1972, 28, 65. (10) Kim Hyunyong; Appelman, E. H. J. Chem. Phys. 1982, 76, 1664. (11) Christe, K. 0. J . Fluorine Chem. 1987, 35, 621. (12) Appelman, E. H.; Kim Hyunyong. J . Chem. Phys. 1972,57, 1972. (13) Appelman, E. H.; Acc. Chem. Res. 1973, 6, 113. (14) Halonen, L.; Ha, T. K. J. Chem. Phys. 1988, 89, 4885. (15) Thiel, W.; Scuseria, G.; Schaefer, H. F.; Allen, W. D. J . Chem. Phys. 1988, 89, 4965. (16) Pople, J. A.; Curtiss, L. A. J . Chem. Phys. 1989, 90, 2833. (17) Peterson, K. A,; Woods, R. C. J. Chem. Phys. 1990, 92, 7412. (18) Koppel, I. A.; Comisarow, M. B. Org. React. (Tartu) 1980, 17, 495. (19) Koppel, I. A.; Molder, U. H. Org. React. (Tartu) 1981, 18, 42. (20) Alcami, M.; Mo, 0.; Yanez, M.; Abbound, J.-L. M.; Elguero, J. Chem. Phys. Lett. 1990, 172, 471. (21) Taft, R. W.; Koppel, I. A,; Topsom, R. D.; Anvia, F. J . Am. Chem. SOC. 1991, 112, 2047. (22) Baulch, D. L.; Cox, R. A.; Crutzen, P. J.; Hampson, R. F., Jr.; Ken, J. A.; Troe, Watson, R. T. J . Phys. Chem. Ref. Data 1982, 11, 327. (23) Thynne, J. C. J.; MacNeil, K. A. G. J. Mass Spectrom. Ion Phys. 1970, 5, 95. (24) Alekseev, V. I.; Volkov, V. M.; Fedorova, L. I.; Baluev, A. V. Izv. Akad. Nauk SSSR Ser. Khim. 1984, 1302.
Intrinsic Acidity and Basicity of FOH (25) Hehre, W. J.; Radom, L.; Schleyer, P. v. R.; Pople, J. A. Ab Initio Molecular Orbital Theory; Wiley: New York, 1986. (26) Frisch, M. J.; Trucks, G. W.; Head-Gordon, M.; Gill, P. M. W.; Wong, M. W.; Foresman, J. B.; Johnson, B. G.; Schlegel, H. B.; Robb, M. A.; Replogle, E. S.; Gomperts, R.; Andres, J. L.; Raghavachari, K.; Binkley, J. S.; Gonzalez, C.; Martin, R. L.; Fox, D. J.; Defrees, D. J.; Baker, J.; Stewart, J. J. P.; Pople, J. A. Gaussian 92, Revision C; Gaussian Inc.: Pittsburgh, 1992. (27) Pople, J. A.; Head-Gordon, M.; Fox, D. J.; Raghavachari, K.; Curtiss, L. A. J . Chem. Phys. 1989, 90, 5622. (28) Curtiss. L. A.: Raehavachari. K.: Trucks, G. W.: Poule, J. A. J. Chem.’Phys. 1991,94, 7251. (29) Koppel, I. A,; Mblder, U. H.; Palm, V. A. Org. React. (Tartu) 1985, -21 - , 3. (30) Smith,B. J.; Radom, L. J. Am. Chem. SOC. 1993, 115, 4885.
J. Phys. Chem., Vol. 99, No. 5, 1995 1435 (31) Foresman, J. B.; Frisch, A. Exploring Chemistry with Electronic Srructure Methods: A Guide to Using Gaussian; Gaussian Inc.: Pittsburgh, 1993. (32) Koppel, I. A.; Molder, U. H. Org. React. (Tartu) 1983, 20, 3. (33) Lias, S. G.; Bartmess, J. E.; Liebman, J. F.; Holmes, J. L.; Levin, R. D.; Mallard, W. G. J. Phys. Chem. Ref. Data 1988, 17 (Suppl. 1). (34) Hansch, C.; Leo, A,; Taft, R. W. Chem. Rev. 1991, 91, 165. (35) Koppel, I. A.; Pihl, V.; Koppel, J.; Anvia, F.; Taft, R. W. J. Am. Chem. Soc. 1994, 116, 8654. (36) Burk, P.; Koppel, I. A. To be published. (37) Rock, S. L.; Pearson, E. F.; Appelman, E. H.; Noms, C. L.; Flygare, W. H. J. Chem. Phys. 1973, 59, 3440. JF94 16679
