acid-base equilibria

Meaning: In the Brønsted-Lowry model, an acid donates a proton and a base accepts one. A weak acid reacts reversibly: HA + H₂O ⇌ H₃O⁺ + A⁻. It ionizes only partly, so Ka sets how far the reaction goes. A larger Ka (smaller pKa) means a stronger acid. Its conjugate base is weaker, and Ka × Kb = Kw = 1.0 × 10⁻¹⁴ at 25 °C. Adding H₃O⁺, OH⁻, or a common ion shifts the equilibrium as Le Châtelier's principle predicts. A buffer is a weak acid plus its conjugate base. It resists pH change because added acid is consumed by A⁻ and added base is consumed by HA.

Equation: Ka = [H₃O⁺][A⁻]/[HA], where concentrations are in mol/L and Ka is treated as unitless. Buffers use pH = pKa + log([A⁻]/[HA]).

On the MCAT: Expect titration curves. The pH equals the pKa at the half-equivalence point of a weak acid, where [HA] = [A⁻]. Other questions ask you to rank acid strengths from Ka or pKa values, to predict the direction of a pH shift, or to choose a buffer whose pKa is close to the target pH.

Common confusion: Strong and weak describe the extent of ionization, not concentration. A dilute strong acid can have a higher pH than a concentrated weak acid. Also, a stronger acid has a weaker conjugate base, not a stronger one.