Why are dimethyl sulfoxide and dimethyl sulfone such good solvents?
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- Clark, T., Murray, J.S., Lane, P. et al. J Mol Model (2008) 14: 689. doi:10.1007/s00894-008-0279-y
We have carried out B3PW91 and MP2-FC computational studies of dimethyl sulfoxide, (CH3)2SO, and dimethyl sulfone, (CH3)2SO2. The objective was to establish quantitatively the basis for their high polarities and boiling points, and their strong solvent powers for a variety of solutes. Natural bond order analyses show that the sulfur–oxygen linkages are not double bonds, as widely believed, but rather are coordinate covalent single S+→O− bonds. The calculated electrostatic potentials on the molecular surfaces reveal several strongly positive and negative sites (the former including σ-holes on the sulfurs) through which a variety of simultaneous intermolecular electrostatic interactions can occur. A series of examples is given. In terms of these features the striking properties of dimethyl sulfoxide and dimethyl sulfone, their large dipole moments and dielectric constants, their high boiling points and why they are such good solvents, can readily be understood.
KeywordsDimethyl sulfoxideDimethyl sulfoneElectrostatic potentialsσ-Hole bondingNoncovalent interactions
Experimental physical properties of some organic solventsa
Dipole moment, D
Melting point, °C
Boiling point, °C
Dimethyl sulfoxide, (CH3)2SO
Dimethyl sulfone, (CH3)2SO2
Dimethyl sulfide, (CH3)2S
Carbon disulfide, CS2
The dipole moments and especially the dielectric constants of DMSO and DMSO2 indicate that they are quite polar, which suggests strong intermolecular interactions in the liquid phase. This can be used to account for the very high boiling points of DMSO and DMSO2 (compare DMSO and DMSO2 to the other three compounds in Table 1).
DMSO and DMSO2 are important and widely used solvents [2, 4], the latter being especially valuable for high-temperature reactions. They can dissolve a wide range of solutes and are miscible with many other solvents; this is true not only for polar compounds but also for some of low polarity, e.g., naphthalene and toluene. Being aprotic, DMSO and DMSO2 can tolerate relatively strong bases.
Why are the DMSO and DMSO2 molecules so polar? Oxygens certainly have the capacity to become highly negative, but sulfur (which is of intermediate electronegativity) and methyl groups are not normally expected to become highly positive. Acetone, (CH3)2CO, is similar to DMSO, except that carbon (also of intermediate electronegativity) replaces sulfur; however, the dipole moments and dielectric constants in Table 1 show acetone to be much less polar than DMSO, and its boiling point accordingly much lower.
In this paper, we will explore the issue of DMSO and DMSO2 polarities in terms of their computed structures, electronic properties and electrostatic potentials. We will also look at some of their intermolecular interactions, as a means of addressing their very high boiling points and their notable solvent capabilities. Since both DMSO and DMSO2 can be prepared by the oxidation of dimethyl sulfide, DMS [2, 4], we shall include the latter in our computational analysis as a reference point.
Electrostatic potential: definitions
The electrostatic potential has been found to be a particularly effective tool for analyzing and predicting noncovalent interactions. For this purpose, we generally compute V(r) on the surface of the molecule, labeling it VS(r). We take the surface to be the 0.001 au (electrons/Bohr3) contour of the electronic density, as suggested by Bader et al. . The most positive and most negative values of VS(r) on a given molecular surface are designated as VS,max and VS,min, respectively; there may be several such local maxima and minima. The magnitudes of VS,max and VS,min have been shown to correlate well with empirical measures of hydrogen bond donating and accepting tendencies .
VS(r) can be characterized further by means of several statistically defined quantities, such as its average positive and negative values, and its positive and negative variances. In terms of these and related quantities, it has been found possible to develop analytical expressions for a variety of condensed phase physical properties that depend upon noncovalent interactions: heats of phase transitions, solubilities, boiling points and critical constants, viscosities, surface tensions, diffusion constants, etc. For reviews, see Murray and Politzer [9, 10].
The computational procedures were the same as in our earlier studies of σ-hole-bonded systems [19–21, 35]. To obtain the electrostatic potentials VS(r) on the surfaces of the molecules of interest, we used the density functional B3PW91/6–31G(d,p)//B3PW91/6–31G(d,p) method. For natural bond orbital (NBO) analyses , and to compute interaction energies ΔE, we used higher computational levels. We optimized geometries at B3PW91/6–311G(3df,2p) and used these for NBO and ΔE at B3PW91/6–311G(3df,2p) and for just ΔE at MP2-FC/6–311++G(3df,2p). Energy minima were confirmed by the absence of imaginary vibration frequencies. With such large basis sets, any errors in ΔE due to basis set superposition should be minimal  and were accordingly not evaluated. The ΔE are the differences between the energy minima at 0 K, products minus reactants.
Structures and NBO analyses
Some optimized bond lengths and bond angles in (CH3)2S, (CH3)2SO and (CH3)2SO2, at the B3PW91/6–311G(3df,2p) level. When two or more bond lengths or bond angles in the molecule have the same magnitudes, this is indicated in parentheses. Experimental values are given in bracketsa
Bond length, A
Bond angle, degree
S−C: 1.801 (2) [1.802 (2)]
C−S−C: 100.0 [98.9]
S−C: 1.807 (2) [1.799 (2)]
C−S−C: 96.4 [96.6]
S−O: 1.481 [1.485]
C−S−O: 106.9 (2) [106.5 (2)]
S−C: 1.778 (2) [1.777 (2)]
C−S−C: 103.8 [103.3]
S−O: 1.437 (2) [1.431 (2)]
C−S−O: 107.9 (4) [107.8 (4)]
O−S−O: 120.3 
The nature of the sulfur–oxygen bonding in DMSO and DMSO2 has, in the past, been a matter of some disagreement [13–16]. It is now commonly described as involving S=O double bonds, with sulfur 3d orbitals having important roles. Indeed, the sulfur–oxygen bond lengths in Table 2 are very similar to our calculated value for the SO molecule (1.484 Å), which is certainly expected to have a double bond. However, it has also been argued that DMSO and DMSO2 have coordinate covalent single bonds between the sulfur and oxygen, in which both electrons are provided by the sulfur: S+→O−.
Natural bond orbital (NBO) analyses of (CH3)2S, (CH3)2SO and (CH3)2SO2, at the B3PW91/6–311G(3df,2p) level. When the results are the same for two bonds or two lone pairs in a molecule, this is indicated in parentheses
Bond (BD) or lone pair (LP)
Atom, percent contribution
BD: S−C (2)
BD: S−C (2)
BD: S−C (2)
BD: S−O (2)
LP: O (2)
LP: O (2)
LP: O (2)
In DMS, the situation is fairly straightforward. The sulfur’s contribution to the S−C bonds is primarily its half-filled 3p orbitals; the 3s character is only 17%. Sulfur also has two unshared pairs of electrons, one in a pure 3p orbital and the other in what is mainly the 3s, although with 33% 3p. Thus the sulfur in DMS approximates its valence electron configuration in the free state, 3s23p23p13p1, with relatively little hybridization.
The above description of the S–C bonds also applies to DMSO, but in DMSO2, the sulfur is providing sp3 hybrids rather than essentially 3p orbitals. However, what is important in both DMSO and DMSO2, in light of the earlier discussion, is that they contain only single sulfur–oxygen bonds. These can be described roughly as composed of sp3 hybrid orbitals on the sulfurs and oxygens. Each oxygen also has three unshared pairs of electrons, two being in pure 2p orbitals and one in a primarily 2s. Table 3 shows no significant participation of d or f electrons in any bond or lone pair.
The fact that there is only one bond between sulfur and each oxygen, and that the oxygens have three unshared pairs of electrons, strongly indicates coordinate covalent single S−O bonds, in which both shared electrons come from the sulfur, S+→O−. In contrast, our NBO analysis of the SO molecule showed both a σ- and a π-bond between the sulfur and oxygen.
Electrostatic potential analyses
Most negative and most positive electrostatic potentials, VS,min and VS,max, on molecular surfaces of (CH3)2S, (CH3)2SO and (CH3)2SO2, computed at the B3PW91/6–31G** level. When the VS,min or VS,max occurs two or more times in the molecule, this is indicated in parentheses. Values are in kcal mol−1
S: −25.4 (2)
O: −38.5 (2)
S: negative region but no VS,min
H’s: 14.2 (4), 14.5 (2)
S: 30.2 (2)
H’s: 18.4 (2), 19.8 (2)
H’s: 23.2 (2)
Looking first at the potential on the surface of DMS (Fig. 2), the sulfur is seen to be entirely negative, with two VS,min of−25.4 kcal mol−1 (Table 4). These are located above and below the C–S–C plane and can be attributed to the overlapping electronic densities of the two unshared pairs of sulfur electrons (Table 3). The methyl hydrogens are just weakly positive, with VS,max between 14 and 15 kcal mol−1. The overall picture is consistent with relatively low polarity, and this is reflected in the dipole moment (Table 1).
When a half-filled p or hybridized p orbital interacts to form a covalent bond, or a filled one forms a coordinate covalent bond, some degree of electronic charge deficiency in its other, noninvolved lobe normally results. This “σ-hole” (the electron-deficient outer lobe of a bonding orbital) may result in a positive electrostatic potential, centered approximately along the extension of the covalent bond. The σ-hole becomes more positive as the atom becomes more polarizable, and as there is less mixing of s character into the p orbital. For these reasons, the σ-hole is typically enhanced in going from the lighter to the heavier elements in a given column of the periodic table. Thus the phosphorus atom in (CH3)3P is completely negative, whereas the arsenic in (CH3)3As has a VS,max of 7.7 kcal mol−1 along the extension of each C–As bond . The σ-holes also become more positive as the remainder of the molecule is more electron-withdrawing; in (CH3)2PF, there is a VS,max of 25.4 kcal mol−1 on the extension of the F–P bond, but still none along the C–P. Positive σ-holes have now been found computationally for covalently-bonded atoms of Groups V , VI  and VII [17, 18, 21], although only infrequently for the lightest members of these Groups (N, O and F).
In DMSO, the positive potential of the σ-hole merges with those of the neighboring hydrogens, but its presence is clearly evident in Fig. 3 by the single VS,max (red) between those hydrogens. In contrast, the other four methyl hydrogens have separate VS,max (Table 4). Overall, as Fig. 3 shows, there is an extended negative region on one side of the molecule, arising from the oxygen and sulfur, and an extended positive one on the other, due to the sulfur σ-hole and the adjoining two methyl hydrogens. The VS,min and VS,max, −46.0 and 26.2 kcal mol−1, are quite similar to those of the ammonia molecule, −46.3 and 25.5 kcal mol−1. The other methyl hydrogens also represent significant positive centers. The polarity seen in Fig. 3 readily explains the high dipole moment of DMSO.
The VS(r) analyses have shown that DMSO and DMSO2 offer a remarkable array of possible sites for intermolecular electrostatic interactions. Foremost are the strongly negative oxygens. The one in DMSO has a more negative VS,min, −46.0 kcal mol−1 (Table 4), than the oxygen in H2O, −39.6 kcal mol−1, while those in DMSO2 are about the same. In addition, the sulfur in DMSO has a significant negative region (Fig. 3). On the positive side must be considered the methyl hydrogens. DMSO and DMSO2 are often described as aprotic solvents, because methyl hydrogens are normally not viewed as having significant acidity. However those in DMSO and DMSO2 are more positive than is typical; for example, the hydrogens in n-butane all have VS,max ≤ 7 kcal mol−1, while in benzene the VS,max are 13.2 kcal mol−1. In fact, the hydrogens in DMSO2 are nearly as positive as those in NH3 (25.5 kcal mol−1), a prototypical hydrogen bond donor. Finally, and very importantly, there are the positive σ-holes on the sulfurs in DMSO and DMSO2.
It is well established, both experimentally [22–28] and computationally [18–21, 29–33], that sufficiently strongly positive σ-holes can interact electrostatically with negative regions on other molecules, e.g., lone pairs of Lewis bases. The resulting noncovalent bonding is highly directional, approximately along the extensions of the bonds that produced the σ-holes. These interactions, which are often called “halogen bonding” when the σ-hole is on a Group VII atom, are competitive with hydrogen bonding [22, 23, 34, 35].
With this variety of positive and negative sites, it follows that DMSO and DMSO2 can easily interact electrostatically with other molecules in several different ways, some of them simultaneously. We shall now look at some specific examples.
Computed properties of (CH3)2SO (DMSO) and (CH3)2SO2 (DMSO2) complexes. All geometry optimizations at B3PW91/6–311G(3df,2p) level. When the same separation occurs two or more times in the complex, this is indicated in parentheses
ΔE (kcal mol−1)
H---O: 2.39 (2)
S---O: 3.60 (2)
O−S---O: 173 (2)
H---O: 2.40 (4)
H---O: 2.77 (2)
H---O: 2.44 (2)
H---Oacetone: 2.56 (2)
H---Oacetone: 2.48 (2)
H---ODMSO: 2.85 (2)
Sums of van der Waals radii:a
H---O: 2.69; S---O: 3.35; H---S: 3.00.
The separations of the interacting atoms in complexes 1–11 are, for the most part, relatively large; many of them approach or even exceed the sum of the van der Waals radii (Table 5). Thus, it might be argued that there really is no significant σ-hole bonding in 2, 3, 4 and 11, because the S---O distances are greater than the sum of the sulfur and oxygen van der Waals radii. To test this, we reoptimized the geometry of 2 (Fig. 5), starting with the DMSO molecule on the right in such a position that its oxygen could still interact with the two methyl hydrogens but not with the σ-hole of the sulfur on the left. In the reoptimization process, however, the system reverted to the structure shown for 2, in which the oxygen in the molecule on the right is essentially on the extension of the O−S bond in the molecule on the left, as it would be in σ-hole bonding. Accordingly, the σ-hole interaction does play an important role. The large S---O separations in 2, 3, 4 and 11 may be due to steric factors and also because when there are several simultaneous interactions, the resulting structure is not likely to maximize any one of them.
Discussion and summary
The NBO analyses showed that the sulfur–oxygen linkages in DMSO and DMSO2 are coordinate covalent single S+→O− bonds, with both of the shared electrons coming from the sulfur. The molecular surface electrostatic potentials confirm the highly negative characters of the oxygens, and also reveal positive σ-holes on the sulfurs, on the extensions of the O−S bonds. The σ-hole potentials merge with the unusually strongly positive ones of the neighboring methyl hydrogens to create extended regions of positive potential, with one local maximum, VS,max, while the other hydrogens have their own VS,max. These features account for the large dipole moments and high dielectric constants of DMSO and DMSO2 (Table 1).
The resulting arrays of positive and negative sites in DMSO and DMSO2 (which includes the weak negative region on the sulfur in DMSO) make possible a variety of simultaneous intermolecular electrostatic interactions to which can be attributed the high boiling points and notable solvent powers [2, 4] of DMSO and DMSO2. The fact that they are effective solvents not only for polar solutes but for aromatic compounds as well can be explained in terms of interactions between the extended positive regions of DMSO and DMSO2 and the pi electrons of, for example, naphthalene and biphenyl.
This work was supported in part by the Deutsche Forschungsgemeinschaft as part of SFB583 Redox-active Metal Complexes: Control of Reactivity via Molecular Architecture.