An equilibrium constant is related to the standard Gibbs energy change for the reaction, so for an acid dissociation constant
R is the gas constant and T is the absolute temperature. Note that  and . At 25 °C, ΔG in kJ·mol ≈ 5.708 pK (1 kJ·mol = 1000 joules per mole). Free energy is made up of an enthalpy term and an entropy term.
The standard enthalpy change can be determined by calorimetry or by using the van t Hoff equation, though the calorimetric method is preferable. When both the standard enthalpy change and acid dissociation constant have been determined, the standard entropy change is easily calculated from the equation above. In the following table, the entropy terms are calculated from the experimental values of pK and ΔH'. The data were critically selected and refer to 25 °C and zero ionic strength, in water.
The first point to note is that, when pK is positive, the standard free energy change for the dissociation reaction is also positive. Second, some reactions are exothermic and some are endothermic, but, when ΔH is negative TΔS is the dominant factor, which determines that ΔG is positive. Last, the entropy contribution is always unfavourable () in these reactions. Ions in aqueous solution tend to orient the surrounding water molecules, which orders the solution and decreases the entropy. The contribution of an ion to the entropy is the partial molar entropy which is often negative, especially for small or highly charged ions. The ionization of a neutral acid involves formation of two ions so that the entropy decreases (). On the second ionization of the same acid, there are now three ions and the anion has a charge, so the entropy again decreases.
Note that the standard free energy change for the reaction is for the changes from the reactants in their standard states to the products in their standard states. The free energy change at equilibrium is zero since the chemical potentials of reactants and products are equal at equilibrium.