pKa

pKa is the negative base-10 logarithm of the acid dissociation constant Ka for an acid in solution. For the acid HA that dissociates as HA ⇌ H+ + A-, Ka = [H+][A-]/[HA], and pKa = -log10(Ka). The pKa value indicates how readily an acid donates a proton under the chosen solvent and temperature.

In aqueous solution at 25°C, lower pKa corresponds to a stronger acid. The Henderson-Hasselbalch equation relates pH to pKa: pH = pKa + log([A-]/[HA]), showing how the protonation state of a molecule depends on pH relative to its pKa. For conjugate acid-base pairs in water, Ka × Kb = Kw, so pKa and pKb are linked by pKa + pKb ≈ 14 at 25°C.

Many acids are polyprotic and have several pKa values, each associated with successive deprotonation steps. Examples include acetic acid (pKa about 4.76) and carbonic acid (pKa1 about 6.35, pKa2 about 10.33). Solvent and temperature strongly influence pKa; values in nonaqueous solvents can differ markedly, and there is no single universal pKa across all environments.

Uses of pKa include estimating buffering ranges, predicting the predominant protonation state of a molecule at a given pH, guiding drug design, and interpreting reaction mechanisms. Measurement methods encompass titration with a strong base or acid, spectrophotometric techniques, and computational estimates.