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Because of their specificity and directionality hydrogen bonds play an important role in determining the three dimensional structure of chemical and biological systems, especially in combination with other non-covalent forces such as ionic or hydrophobic interactions. The main draw back of hydrogen bonds is their limited strength. The more polar a solvent, the weaker are the hydrogen bonds. A single hydrogen bond thus has a substantial binding energy only in non-polar aprotic solvents, and not in water.

direct solvation by the solvent

formation of a new H-bond = exothermic desolvation = endothermic release of solvent molecules = entropy increases

Even for the most stable hydrogen bond in FHF- the binding energy decreases from 160 kJ mol-1 in the gas phase to only 3-4 kJ mol-1 in water! The effect of the polar solvent on the strength of H-bonds is twofold:

On a macroscopic level, the bulk solvent properties such as the dielectric constant are changed. According to the Coulomb law this weakens all electrostatic interactions and hence also hydrogen bonds.

On a molecular level, specific direct solvation of the donor and acceptor sites by individual solvent molecules occurs. In other words, the solvent functions as a competitive binding partner. Therefore, desolvation of both donor and acceptor sites must occur before a new hydrogen bond can be formed. The energetic price which must be paid for this desolvation, reduces the binding energy of the H-bond. In very polar solvents, the net energy change might then even be endothermic, but the association still can be favorable because of the release of ordered solvent molecules into the bulk solution upon desolvation, which increases the entropy of the system (Figure B.7.1).

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