Solute-solvent hydrogen bonds

Solvent effects on chemical equilibria and reactions have been an important issue in physical organic chemistry. Several empirical relationships have been proposed to characterize systematically the various types of properties in protic and aprotic solvents. One of the simplest models is the continuum reaction field characterized by the dielectric constant, e, of the solvent, which is still widely used. Taft and coworkers [30] presented more sophisticated solvent parameters that can take solute-solvent hydrogen bonding and polarity into account. Although this parameter has been successfully applied to rationalize experimentally observed solvent effects, it seems still far from satisfactory to interpret solvent effects on the basis of microscopic infomation of the solute-solvent interaction and solvation free energy. 

It is possible to distinguish axial and equatorial hydroxy groups in chroman-4-ols on the basis of their chemical shifts in DMSO (79BCJ2163). The pseudoaxial OH proton resonates at a significantly higher field the different behaviour is attributed to solute-solvent hydrogen bonding. 

The SM2/AM1 model was used to examine anomeric and reverse anomeric effects and allowed to state that aqueous solvation tends to reduce anomeric stabilization [58]. Moreover, SM2/AM1 and SM3/PM3 models were accounted for in calculations of the aqueous solvation effects on the anomeric and conformational equilibria of D-glucopy-ranose. The solvation models put the relative ordering of the hydroxymethyl conformers in line with the experimentally determined ordering of populations. The calculations indicated that the anomeric equilibrium is controlled primarily by effects that the gauche/trans 0-C6-C5-0 hydroxymethyl conformational equilibrium is dominated by favorable solute-solvent hydrogen bonding interactions, and that the rotameric equilibria were controlled mainly by dielectric polarization of the solvent [59]. On the other hand, Monte Carlo results for the effects of solvation on the anomeric equilibrium for 2-methoxy-tetrahydropyran indicated that the AM1/SM2 method tends to underestimate the hydration effects for this compound [60]. 

In general, a dilute solution in a nonpolar solvent furnishes the best (i.e., least distorted) spectrum. Nonpolar compounds give essentially the same spectra in the condensed phase (i.e., neat liquid, a mull, a KBr disk, or a thin film) as they give in nonpolar solvents. Polar compounds, however, often show hydrogen-bonding effects in the condensed phase. Unfortunately, polar compounds are frequently insoluble in nonpolar solvents, and the spectrum must be obtained either in a condensed phase or in a polar solvent the latter introduces the possibility of solute-solvent hydrogen bonding. 

Benigno AJ, Ahmed E, Berg M. The influence of solvent dynamics on the lifetime of solute-solvent hydrogen bonds. J Chem Phys 1996 104 7392-7394. 

This linear relationship demonstrates the similarity of the polarity effects on both homomorphic vibrators, trichloroacetic acid and its methyl ester. For HBA solvents B, however, the vc o data points are displaced below the reference line of Eq. (7-37). These deviations are caused by the formation of solute/solvent hydrogen bonds CCI3CO2H B, resulting in a decrease in the C=0 vibration wavenumber. The hydrogen-bond induced wavenumber shift Avc o for a HBA solvent is then calculated by Eq. (7-38) cf Eq. (7-33). 

Here and V refer to the volumes of stationary and mobile phases, respectively. This model of solute-solvent hydrogen bonding in LSC is pursued further in Section III,C. 

The use of isodensity to define fhs M) (Frisch et al., Gaussian94 Kolle and Jug, 1995) does not reflect the different strength of the interaction, and the effect this has on A(A) for different atoms. We suspect that, by using cavities based on isodensity surfaces, the model could no longer be able to properly treat solute-solvent hydrogen bonds. Analogous remarks apply for the use of the electronic component of the molecular potential to define fhs(M) (Rivail et al., 1985). 

Meanwhile, fluorinated alcohols cannot be good proton acceptors due to their electron-deficient lone pairs on the oxygen atoms. The (3 scale of both TFE and HFIP is 0.00, which is apparently smaller than those of ethanol (0.77), ether (0.47), water (0.18), or even toluene (0.11) [33]. Here, the (3 scale, i.e. the hydrogen-bond acceptor basicity of a solvent, describes the ability to accept a proton (donate an electron pair) in a solute-solvent hydrogen-bonding system [4]. These acidities (pFCa [31, 34]), hydrogen-bonding parameters (a and (3 scale... 

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