pKbA

pKbA is the negative base-10 logarithm of the base dissociation constant Kb for a base A in aqueous solution. For a base A, the proton transfer to water is B + H2O ⇌ BH+ + OH−, and Kb(A) = [BH+][OH−]/[B]. The quantity pKbA = −log10(Kb(A)) provides a scale for base strength in water: smaller values indicate stronger bases because they produce higher concentrations of OH− at equilibrium, while larger values indicate weaker bases.

Relationship to pKa: The conjugate acid HA of base A has acid dissociation constant Ka(HA) = [H+][A−]/[HA]. For conjugate acid–base pairs in water, Ka(HA) × Kb(A) = Kw, where Kw ≈ 1.0×10−14 at 25°C. Therefore pKa(HA) + pKbA ≈ pKw ≈ 14, so pKa(HA) ≈ 14 − pKbA. This allows conversion between pKa and pKb for related species.

Measurement and conditions: Kb and pKb values are temperature- and solvent-dependent. They are typically determined from equilibrium concentrations, pH titration data, or by derivation from the pKa of the conjugate acid, using Kw for the given temperature. Statements about specific pKbA values assume standard conditions (aqueous solution, 25°C) unless otherwise noted.

Examples: Ammonia has pKb ≈ 4.75, methylamine about 3.3–3.4, and aniline about 9.4. Pyridine has pKb around 8.7–8.8. These values illustrate the general trend that more basic bases have lower pKb.

Notes: The notation pKbA may appear in literature where A denotes a specific base. pKbA is useful for comparing bases within a given solvent and temperature, but it is not universally identical across different conditions.