pKbC

pKbC is not a standard, widely recognized term in established chemical literature. In contexts where it appears, it is typically defined as the base dissociation constant of the conjugate base of a given compound C, in a specified solvent. The idea is that the species C− acts as a base, reacting with water according to a reaction such as C− + H2O ⇌ HC + OH−, where HC is the conjugate acid and Kb(C) is the corresponding base dissociation constant. Under this usage, pKbC is defined as −log10(Kb(C)), with the subscript C indicating the compound or its conjugate base.

Because pKbC depends on solvent and temperature, it is not universal. A common reference frame is aqueous solution at 25°C, in which pKbC may be related to the conjugate acid’s strength. The key thermodynamic relationship is Ka(HC) × Kb(C) = Kw, so that pKa(HC) + pKb(C) = pKw. In water at 25°C, pKw is 14, giving pKa(HC) + pKb(C) = 14.

Usage of pKbC would be as a comparative metric for base strength across a family of compounds C, or as a placeholder in databases where each compound is labeled with a subscript. Readers should exercise caution, as definitions can vary by author, and solvent and temperature must be specified.

Example: if a hypothetical HC has pKa = 9 in water at 25°C, then pKbC = 14 − 9 = 5. This illustrates the general relationship, not an actual measured value for a real compound.

See also: pKa, pKb, acid–base equilibrium, conjugate acid–base pairs.