Acid-base equilibrium

Under the Brønsted-Lowry definition, an acid donates a proton and a base accepts one. In water, an acid HA dissociates as HA ⇌ H⁺ + A⁻, and the position of that equilibrium is captured by the acid dissociation constant Ka = [H⁺][A⁻] / [HA]. A large Ka means dissociation strongly favors products, so the acid is strong; a small Ka means most of the acid remains intact. Because these values span many orders of magnitude, they are usually reported as pKa = −log Ka, where a lower pKa indicates a stronger acid.

Water itself participates in every aqueous acid-base equilibrium. Its autoionization gives Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C, which is why pH + pOH = 14 under those conditions. A conjugate acid-base pair is related by Ka × Kb = Kw, so the stronger an acid, the weaker its conjugate base.

The Henderson-Hasselbalch equation, pH = pKa + log([A⁻] / [HA]), links the equilibrium directly to pH and explains buffering. When the concentrations of acid and conjugate base are equal, the log term is zero and pH equals pKa — the point of maximum buffer capacity. This is why buffers are chosen with a pKa near the target pH, and it is the basis of the bicarbonate buffer system that stabilizes blood pH. On a titration curve, the pKa is read at the half-equivalence point, while the equivalence point is where stoichiometrically equivalent amounts of acid and base have reacted.

Acid-base equilibria are a core topic in the MCAT chemical and physical foundations section, and the same material underpins general chemistry questions on other science-heavy exams. Expect to estimate pH from a Ka value, identify conjugate pairs, apply Henderson-Hasselbalch to a buffer, and interpret the shape of a titration curve for a strong versus a weak acid.