Quantitative evaluation of hydration thermodynamics with a continuum model.

We attempt to analyze whether experimental entropies, enthalpies and free energies of hydration of small uncharged molecules can be quantitatively rationalized with a continuum model including a classical reaction field formalism. We find that a simple proportionality to accessible surface with five different atom types allows satisfactory (within 1-1.5 kcal/mol) reproduction of hydration entropies (T delta S) of over 40 solutes. The agreement with experiment can possibly be improved if proximity effects and configurational contributions to transfer entropies are taken into account. In calculations of hydration enthalpies a reasonable agreement with experimental data can be obtained only when solute polarizability is taken into account. Electrostatic contributions to calculated hydration enthalpies exhibit strong dependencies on both the magnitude and the direction of molecular dipole moments. We demonstrate that for 20 molecules with experimentally measured vacuum dipole moments density functional calculations with DZVPD basis set including diffuse functions on d-orbitals allows prediction of experimental dipole moments within 0.1 D. At a fixed direction of the molecular dipole moment, mu, the electrostatic component of hydration enthalpy varies as mu 2. Thus an uncertainty of 0.1 D corresponds to uncertainties of 0.5-0.7 kcal/mol in hydration enthalpies of most small dipolar solutes. A 30 degree change in the direction of the molecular dipole together with the corresponding change in the quadrupole moment can result in a change of hydration enthalpy of 3 kcal/mol. Changes in the quadrupole moment alone can result in hydration enthalpy changes of over 1 kcal/mol. Representations of multipole expansions by point charges on nuclei fitted to molecular electrostatic potentials cannot accurately reproduce all these factors. Use of such point charges in calculations of hydration enthalpies predictably leads to discrepancies with experiment of approximately 3 kcal/mol for some solutes. However, errors in hydration enthalpies and hydration entropies are usually compensating leading in most cases to agreement between calculated and experimental free energies of hydration within 1.5 kcal/mol.