ABSTRACT

We have seen in Chapters 2 and 3 how a “purist” would calculate the pKa of a chemical compound from first principles, employing a suitable thermodynamic cycle, a reasonably high level of computational theory, and an appropriate solvent model. The attraction of the purist, or absolute approach, is that it relies only on well-established theoretical methods (with the obvious exception of the value for the Gibbs energy of solvation for H+), and does not depend on the availability of experimental measurements, which may be difficult to perform, missing, or improperly measured. The drawback of the purist, absolute approach is that in order to achieve reasonable accuracy, a high level of theory is normally required, both in the calculation of the compound and in the solvent model. In fact, the absolute approach can be computationally demanding for even a modest level of accuracy, since one must determine the Gibbs energy change for the solution dissociation equilibrium to within ±1.36 kcal/mol (or ±5.69 kJ/mol) in order to obtain a pKa value accurate to ±1 pKa unit [8,9,11]. This presents a considerable challenge since although modern quantum chemical techniques can normally determine the aqueous Gibbs energies of solvation for neutral species to within ±1 kcal/mol, the corresponding accuracies for the accompanying ionic species are typically only about ±2-4 kcal/mol [37,55]. In addition, the high level of calculation needed for gas-phase calculations, as explained in Section 2.2, rules out absolute methods, at present, for larger molecules. In this chapter we describe an alternative approach that relies on experimental input, but that in turn often supplies reasonable accuracy with less computational demand [36].