J. Phys. Chem. 1996, 100, 16137-16140
16137
Theoretical Study of Dimethyl Sulfoxide-Anion Clusters Peeter Burk,*,† Uldo Mo1 lder,† Ilmar A. Koppel,† Alar Rummel,‡ and Aleksander Trummal‡ Institute of Chemical Physics, Tartu UniVersity, EE2400 Tartu, Estonia, and AS Eltex, Va¨ ike-Karja 10, EE0001 Tallinn, Estonia ReceiVed: May 3, 1996; In Final Form: July 17, 1996X
Complexes of dimethyl sulfoxide with F-, Cl-, OH-, CH3O-, HCOO-, and NO2- anions were studied at MP2/6-31+G* and B3LYP/6-31+G* levels of theory. The energies, geometries, NBO charges, and topological charge densities were analyzed. The formation of bidentate hydrogen-bonded complexes was verified. The combined hydrogen bonding and electrostatic charge-dipole interaction was proposed to be the base of complex formation.
DMSO + X- a DMSO‚‚‚X-
Introduction Dimethyl sulfoxide (DMSO) is a widely used solvent capable of dissolving both hydrophobic and hydrophilic solutes. For the modeling of chemical reactions in DMSO it is helpful to understand the role of the solvent in complexing solute molecule or ion. Therefore, it is important to understand the structural chemistry of DMSO and the way it bonds to the other molecules and ions. More than ten years ago, it was confirmed by Kebarle’s group that DMSO forms stable complexes with anions in the gas phase.1 This interaction was explained by the authors1,2 as purely electrostatic, i.e., charge (anion)-dipole (DMSO, mainly S-O bond) interaction. Similar conclusion was obtained by Koppel3,4 based on STO-3G* and 3-21G* ab initio calculations. Later, Brown, on the basis of solid state X-ray spectroscopy data, proposed5 that DMSO can bond to anions by forming hydrogen bond between its methyl hydrogen and free electron pair in anion. The co-existence of ion-dipole interactions and hydrogen bonding in similar systems has been discussed and quantitatively analyzed by means of correlation analysis techniques by Abboud et al.6 DMSO has been extensively studied by high-level ab initio calculations.7 Unfortunately, no high-level calculation results are available for complexes of DMSO with anions. As the above given hypotheses on bonding in DMSO-anion complexes are somewhat contradicting, current study was undertaken to clarify the situation. Another aim of this work was to get some structural information about such complexes as practically nothing is known about them. Methods Complexes of DMSO with F-, Cl-, OH-, CH3O-, HCOO-, and NO2- anions were calculated at MP2/6-31+G* and B3LYP/ 6-31+G* (the hybrid method includes a mixture of HartreeFock exchange with DFT exchange-correlation by the Becke’s three-parameter functional,8 and the nonlocal correlation is provided by the Lee, Yang, and Parr functional9) levels of theory using the Gaussian94 program package.10 All geometries were fully optimized, and stationary points found were verified by vibrational analysis to be minima on the potential energy surface (no imaginary frequencies). The theoretically calculated complexation enthalphy ∆H° was defined by analogy of proton affinities as the negative of the heat of the following reaction: X
Abstract published in AdVance ACS Abstracts, September 15, 1996.
S0022-3654(96)01269-5 CCC: $12.00
∆H° ) E°(DMSO) + E°(X-) - E°(DMSO‚‚‚X-) where E°(DMSO), E°(X-), and E(DMSO‚‚‚X-) are total energies of the DMSO molecule, anion, and DMSO-anion complex, respectively.11 The results of frequency calculations were used to evaluate ∆H298 for comparison with the experiment. Topological charge density analysis12 was performed using Bader’s PROAIM13 program package. Results and Discussion Calculated dissociation energies of complexes are presented in Table 1. Comparison with the available experimental results for the DMSO-Cl- complex (∆H298 was found to be 18.6, 16.3, and 18.6 kcal/mol by MP2, B3LYP, and experiment,1 respectively) and the DMSO-NO2- complex (19.5, 17.2, and 19.2 kcal/mol by MP2, B3LYP, and experiment,14 respectively indicates good agreement between experiment and calculations. As the primary experimental data are Gibbs energy changes we have also evaluated these quantities for the two complexes above, and once more the agreement between experiment and calculations is good: for the DMSO-Clcomplex, the reported1 experimental ∆G298 is 12.5 kcal/mol while our calculated values are 12.3 and 10.9 kcal/mol for MP2 and B3LYP. For the DMSO-NO2- complex, the reported14 experimental ∆G420 is 8.7 and our calculations give 9.6 and 7.5 kcal/mol for MP2 and B3LYP, respectively. Both methods used give similar results, with B3LYP energies being systematically somewhat lower. It has been pointed out in earlier publications15 that the strength of the hydrogen bond (dissociation energy of complex) in the B--HX complexes increases with the basicity of anion. This is true also for data in Table 1. The stability order of complexes matches exactly the proton affinity order of anions. Moreover, the correlation between proton affinity and “DMSO affinity” of the anions is remarkable (R2 ) 0.967 for MP2 results and R2 ) 0.816 for B3LYP). This suggests considerable contribution from hydrogen-bond type interaction between DMSO and anion. Obtained geometries of complexes are given in Tables 2 and 3 (numbering of hydrogens is given on Figure 1). In all cases complexes preferred Cs symmetry. Differently from what is reported in earlier studies,1-4 the anions tended not to be on the same line with the O-S bond. They preferred positions in the plane defined by carbons and hydrogens participating in © 1996 American Chemical Society
16138 J. Phys. Chem., Vol. 100, No. 40, 1996
Burk et al.
TABLE 1: Calculated Total Energies of Monomers E0X and Complexes E0C (au) and Complexation Enthalpies ∆H° and ∆H298 (kcal/mol) MP2/6-31+G* DMSO FClHOCH3OHCOONO2-
Becke3LYP/6-31+G*
EX0
EC0
∆H°
∆H298
EX0
EC0
∆H°
∆H298
-552.13005 -99.62385 -459.67115 -75.58836 -114.74453 -188.71138 -204.64827
-651.79998 -1011.83196 -627.76871 -666.92092 -740.87539 -756.81020
28.9 19.3 31.6 29.1 21.3 20.0
28.5 18.6 31.2 28.0 21.0 19.5
-553.20011 -99.85970 -460.27472 -75.79668 -115.11531 -189.21950 -205.16963
-653.10622 -1013.50165 -629.04428 -668.35294 -742.45036 -758.39814
29.1 16.8 29.8 23.5 19.3 17.8
28.9 16.3 29.6 23.1 19.1 17.2
TABLE 2: Selected Geometrical Parameters of DMSO and Its Complexes with Monodentate Ligands (Bond Lengths in Angstroms and Angles in Degrees) MP2/6-31+G* r(S-O) r(S-C) r(C-H1) r(C-H2) r(C-H3) r(X-S) r(X-H2) R(O-S-C) R(C-S-C) R(S-C-H1) R(S-C-H2) R(S-C-H3) R(O-S-X) R(C-H2-X)
1.522 1.809 1.094 1.094 1.092 106.7 96.3 108.9 110.1 107.3
Becke3LYP/6-31+G*
F-
Cl-
HO-
CH3O-
-
1.545 1.798 1.096 1.108 1.093 3.434 1.901 108.1 95.9 107.8 106.2 107.8 156.3 152.4
1.540 1.800 1.095 1.098 1.093 4.090 2.520 107.5 96.2 108.1 108.1 107.7 158.7 157.1
1.548 1.797 1.096 1.107 1.094 3.419 1.950 107.6 96.4 107.8 106.9 107.5 158.0 149.3
1.546 1.798 1.096 1.106 1.094 3.425 1.962 107.6 96.4 107.9 106.9 107.4 158.3 149.2
1.519 1.837 1.094 1.095 1.093
-
ligand
106.8 96.5 109.2 109.9 106.8
F-
Cl-
HO-
CH3O-
1.543 1.825 1.097 1.114 1.095 3.466 1.850 108.6 95.7 108.0 105.8 107.7 153.4 154.8
1.537 1.828 1.096 1.100 1.094 4.162 2.568 107.6 96.3 108.3 108.3 107.2 158.0 157.2
1.545 1.825 1.097 1.112 1.096 3.459 1.922 108.1 96.2 107.9 106.6 107.3 156.0 151.2
1.543 1.826 1.097 1.109 1.095 3.494 1.959 108.0 96.2 108.0 106.8 107.3 156.2 151.5
TABLE 3: Selected Geometrical Parameters of DMSO and Its Complexes with Bidentate Ligands (Bond Lengths in Angstroms and Angles in Degrees) MP2/6-31+G* ligand r(S-O) r(S-C) r(C-H1) r(C-H2) r(C-H3) r(X-S) r(X-H2) r(O2-H2) r(X-O2) R(O-S-C) R(C-S-C) R(S-C-H1) R(S-C-H2) R(S-C-H3) R(O-S-X) R(C-H2-X) R(C-H2-O) R(O2-X-O3) R(H2-O2-X)
Becke3LYP/6-31+G* -
-
HCOO-
NO2
1.522 1.809 1.094 1.094 1.092
1.541 1.798 1.095 1.101 1.093 4.115 2.941 2.104 1.269 107.3 96.6 108.1 109.4 107.2 154.7 156.7 175.9 129.0 119.3
1.539 1.799 1.095 1.099 1.093 4.146 3.102 2.158 1.278 107.3 96.4 108.1 109.1 107.3 155.4 153.8 175.2 115.4 127.1
106.7 96.3 108.9 110.1 107.3
-
HCOO-
NO2-
1.519 1.837 1.094 1.095 1.093
1.539 1.825 1.096 1.104 1.094 4.526 2.938 2.100 1.260 107.6 95.7 108.4 109.3 106.8 160.2 156.8 178.7 129.2 119.9
1.537 1.826 1.096 1.102 1.094 4.727 3.089 2.156 1.260 106.1 96.3 108.4 108.9 107.1 159.4 160.1 179.1 116.1 127.3
106.8 96.5 109.2 109.9 106.8
hydrogen bonding (see Figures 1 and 2). It allows much shorter distance between the anion and hydrogens of DMSO methyl groups while the distance with S-O bond dipole is increased. The distances between the anion and adjacent hydrogens in the methyl groups of DMSO are usually between 1.9 and 2.1 Å, which is fairly typical for hydrogen-bonded complexes. The changes in geometry of DMSO upon complex formation are quite small. Most pronounced is the lengthening of S-O bond by 0.02 A. The C-H bond, which should participate in hydrogen bonding, lengthens only very slightly (about 0.01 Å at the MP2 level of theory). The results of NBO charge distribution analysis for all investigated species are presented in Tables 4 and 5. It can be seen that charge transfer from anions to DMSO is small but constant (ca. 0.04 electrons). Following the small charge
Figure 1. Structure of the complex between DMSO and monodentate anion (OH-) and the numbering of hydrogens in the DMSO methyl groups.
transfer, the changes in charge distribution within DMSO molecule are also small. There are some common trends which should be noted. Upon complex formation, the negative charge on oxygen increases while the positive charge on sulfur decreases. The charge distributions, obtained by us, clearly indicate that the interaction between DMSO and anions cannot be explained only as interaction between the S-O bond dipole and a point charge. Our results show that alongside with the increase of the negative charge on oxygen there is similar rise in the negative charge on carbon atoms and positive charge on potentially hydrogen-bonded hydrogens. It once more suggests considerable contribution from hydrogen-bonding to the complexation of anions by DMSO. To confirm the existence of a hydrogen bond between the anion and hydrogens of the methyl groups of DMSO the Bader’s topological charge density analysis12 was carried out. In all cases except for the presence of a bond between the anion and hydrogens of methyl groups of DMSO was verified by finding bond critical point between these atoms. Inspection of the changes in the frequencies of the C-H2 bond valence vibrations also indicates the presence of the hydrogen bonding in studied complexes. The unscaled MP2 frequency of that vibration changes from 3100 cm-1 (free
Theoretical Study of Dimethyl Sulfoxide-Anion Clusters
J. Phys. Chem., Vol. 100, No. 40, 1996 16139
TABLE 4: NBO Charges of DMSO and Its Complexes with Monodentate Ligands MP2/6-31+G*
Becke3LYP/6-31+G*
ligand
-
F-
Cl-
HO-
CH3O-
-
F-
Cl-
HO-
CH3O-
S O C H1 H2 H3 X ligand
1.307 -1.074 -0.869 0.251 0.241 0.261
1.269 -1.124 -0.888 0.222 0.347 0.221 -0.950 -0.950
1.281 -1.113 -0.874 0.232 0.306 0.233 -0.962 -0.962
1.272 -1.128 -0.882 0.220 0.347 0.219 -1.382 -0.953
1.274 -1.125 -0.881 0.222 0.337 0.221 -1.092 -0.945
1.210 -0.972 -0.895 0.257 0.251 0.269
1.168 -1.034 -0.912 0.230 0.341 0.230 -0.910 -0.910
1.188 -1.021 -0.897 0.240 0.306 0.242 -0.948 -0.948
1.173 -1.039 -0.904 0.228 0.340 0.229 -1.356 -0.920
1.176 -1.033 -0.903 0.231 0.327 0.232 -0.996 -0.917
Figure 2. Two possible structures of the complex between DMSO and bidentate anion (HCOO-).
TABLE 5: NBO Charges of DMSO and Its Complexes with Bidentate Ligands MP2/6-31+G*
Becke3LYP/6-31+G*
ligand
-
HCOO-
NO2-
-
HCOO-
NO2-
S O C H1 H2 H3 X OX ligand
1.307 -1.074 -0.869 0.251 0.241 0.261
1.277 -1.154 -0.878 0.227 0.319 0.232 0.809 -0.907 -0.960
1.281 -1.113 -0.878 0.229 0.311 0.234 0.411 -0.685 -0.960
1.210 -0.972 -0.895 0.257 0.251 0.269
1.182 -1.025 -0.900 0.236 0.312 0.241 0.619 -0.797 -0.935
1.184 -1.021 -0.901 0.238 0.302 0.242 0.231 -0.579 -0.926
DMSO) to 2849 cm-1 (DMSO-MeO- complex). The smallest change (by 53 cm-1) was found for complex between DMSO and chloride anion. Similar changes were found for B3LYP frequencies. Comparison of energies and geometries of different complexes between DMSO and monodentate anions reveals that the interaction energy depends on both S-X- and H2-X- distances. So for F-, CH3O-, and Cl- the energy order follows the lengths of the H2-X- bond. The OH- anion is exception from this rule, and even as CH3O- obeys it, the difference in energy between the F- and CH3O- complexes is small (0.5 kcal/mol). For these two complexes somewhat shorter S-X- distances should be considered, which can indicate the enhanced dipolecharge interaction. Inspection of NBO charges on the oxygen atom in these anions reveals that this may be the case as the charges are considerably higher (-1.38 and -1.10 in OH- and CH3O-, respectively) than in other two anions (-0.95 and -0.96 in F- and Cl-, respectively). Further support of the hypothesis of combined hydrogen bonding and electrostatic interaction between anion and DMSO dipole can be obtained from the study of potentially bidentate
anions, NO2- and HCOO-. In both cases the complexes in which both oxygens participate in hydrogen bonding were found to be slightly more stable (by 0.32 and 0.30 kcal/mol for NO2and HCOO-, respectively, at the MP2/6-31+G* level of theory) than monodentate complexes (see Figure 2) despite the fact that in latter case the anion can approach the S-O bond dipole considerably more close (S-Oanion distance is 4.11 Å in the bidentate complex and 3.25 Å in the monodentate complex, MP2/6-31+G* results), while the Oanion-H2 distance increases only slightly (from 2.104 to 2.370 Å, MP2/6-31+G* results). However, as the energy difference between these two conformers is small despite the loss of one hydrogen bond interaction, one should consider the possibility of combined interactions between anions and DMSO, i.e., both hydrogen bonding and charge dipole interaction. While the hydrogen bonding should be considerably weakened in monodentate complex, the enhanced charge-dipole interaction (much shorter distance between dipole and charge) seems to compensate for most of this weakening. Conclusions All of the arguments given above lead to the conclusion that the interaction between DMSO and anion cannot be characterized only as electrostatic interaction between the anion and S-O bond dipole. The formation of bidentate hydrogen-bonded complexes was verified in current work, and combined hydrogen bonding and electrostatic charge-dipole interaction should be considered as a more realistic alternative. Acknowledgment. This work was supported in part by the Estonian Science Foundation (Grants No. 81 and 1231) and the International Science Foundation (Grants LCP000 and LLJ 100). References and Notes (1) Magnera, T. F.; Caldwell, G.; Sunner, J.; Ikuta, S.; Kebarle, P. J. Am. Chem. Soc. 1984, 106, 6140. (2) Kebarle, P.; Caldwell, G.; Magnera, T.; Sunner, J. Pure Appl. Chem. 1985, 57, 339. (3) Koppel, I. A., Mo¨lder, U. A. Org. React. 1983, 20, 3. (4) Koppel, I. A. Org. React. 1987, 24, 263. (5) Brown, I. D. J. Sol. Chem. 1987, 16, 205. (6) Abboud, J.-L.; Notario, R.; Botella, V. In Theoretical Computational Chemistry; Politzer, P.; Murray, J. S., Eds.; Elsevier: Amsterdam, 1994; Vol. 1, pp 135-182. (7) Rao, B. G.; Singh, U. C. J. Amer. Chem. Soc. 1990, 112, 3803. Wolfe, S.; Schlegel, H. B. Gazz. Chim. Ital. 1990, 120, 285. Leroy, G.; Riffi Temsamani, D.; Sana, M.; Wilante, C. J. Mol. Struct. 1993, 300, 373. Fueno, H.; Ikuta, S.; Matsuyama, H.; Kamigata, N. J. Chem. Soc., Perkin Trans. 2, 1992, 1925. (8) Becke, A. D. J. Chem. Phys. 1993, 98, 5648. (9) Lee, C.; Yang, W.; Parr, R. G. Phys. ReV. B 1988, 37, 785. Miehlich, B.; Savin, A.; Stoll, H.; Preuss, H. Chem. Phys. Lett. 1989, 157, 200. (10) Gaussian 94, ReVision B.2; Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Gill, P. M. W.; Johnson, B. G.; Robb, M. A.; Cheeseman, J. R.; Keith, T. A.; Petersson, G. A.; Montgomery, J. A.; Raghavachari, K.; AlLaham, M. A.; Zakrzewski, V. G.; Ortiz, J. V.; Foresman, J. B.; Cioslowski, J.; Stefanov, B. B.; Nanayakkara, A.; Challacombe, M.; Peng, C. Y.; Ayala, P. J.; Chen, W.; Wong, M. W.; Andres, J. L.; Replogle, E. S.; Gomperts,
16140 J. Phys. Chem., Vol. 100, No. 40, 1996 R.; Martin, R. L.; Fox, D. J.; Binkley, J. S.; Defrees, D. J.; Baker, J.; Stewart, J. J. P.; Head-Gordon, M.; Gonzales, C.; Pople, J. A. Gaussian, Inc.: Pittsburgh, PA, 1995. (11) Foresman, J. B.; Frisch, A. Exploring Chemistry with Electronic Structure Methods: A Guide to Using Gaussian; Gaussian Inc.: Pittsburgh, PA, 1993. (12) Bader, R. F. W. Atoms in Molecules: A Quantum Theory; Oxford University Press: Oxford, 1990.
Burk et al. (13) Biegler-Ko¨nig, F. W.; Bader, R. F. W.; Ting-Hua, T. J. Comput. Chem. 1982, 3, 317. (14) Sieck, L. W. J. Phys. Chem. 1985, 89, 5552. (15) Arshadi, M.; Kebarle, P. J. Phys. Chem. 1970, 74, 1483. Payzant, J. D.; Yamdagni, R.; Kebarle, P. Can. J. Chem. 1971, 49, 3308. Yamdagni, R.; Kebarle, P. J. Am. Chem. Soc. 1971, 93, 7139.
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