pKb

pKb is a measure of the base strength of a chemical species B in aqueous solution. It is defined as the negative base-10 logarithm of the base dissociation constant Kb for the equilibrium B + H2O ⇌ BH+ + OH−. The expression Kb = [BH+][OH−] / [B] describes how readily the base accepts a proton from water, producing hydroxide. A smaller pKb corresponds to a stronger base because it indicates a larger Kb.

In aqueous solution, pKb is related to the acidity of the conjugate acid BH+. At 25°C, pKa(BH+) + pKb(B) ≈ 14. This relationship reflects the interplay between conjugate acid–base pairs and the 14-log scale of pH and pOH. Thus pKb can be inferred from the pKa of the conjugate acid, and vice versa.

pKb is used to compare base strengths and to perform calculations involving pH and pOH. For example, the Henderson–Hasselbalch relation for basic systems can be written as pOH = pKb + log([BH+]/[B]). Strong bases have relatively low pKb values and may dissociate nearly completely, in which case the pKb concept becomes less informative; for such species, Kb is very large and pKb is correspondingly small or not practically defined.

Common bases show a range of pKb values: ammonia has pKb ≈ 4.75, corresponding to a conjugate acid pKa ≈ 9.25. Temperature and ionic strength affect the exact values, but pKb remains a key parameter for assessing and comparing the basicity of species in water.