Acids and Bases
Acids, bases, and buffers
Brønsted–Lowry acids donate protons; bases accept them. Conjugate pairs differ by one H+. Strong acids (HCl, HBr, HI, HNO3, H2SO4 first proton, HClO4) ionize completely; weak acids establish Ka. For a weak acid, Ka = [H+][A−]/[HA]. Kw = [H+][OH−] = 1.0 × 10−14 at 25 °C, so pH + pOH = 14. pKa = −log Ka, and for a conjugate pair pKa + pKb = 14 at 25 °C.
pH calculations: for strong acid, [H+] equals the acid concentration times the number of acidic protons; for weak acid, set up an ICE table and solve, or use the approximation [H+] ≈ √(Ka·C) when C/Ka > 100. Percent ionization = [H+]/[HA]initial × 100.
Buffers contain a weak acid and its conjugate base (or a weak base and its conjugate acid). Henderson–Hasselbalch: pH = pKa + log([A−]/[HA]). Buffer capacity is greatest when pH ≈ pKa and concentrations are high. Titration: at half-equivalence, pH = pKa for a weak acid; at equivalence, the conjugate base hydrolyzes, giving pH > 7 for a weak acid–strong base titration.
Traps: dilution changes pH but not pKa; a strong acid has a very weak conjugate base; amphiprotic salts can act as either acid or base; and the equivalence point is not necessarily pH 7.
Worked example: 0.10 M acetic acid with Ka = 1.8 × 10−5. [H+] ≈ √(1.8 × 10−5 × 0.10) = 1.3 × 10−3 M, so pH ≈ 2.87.
Rule of thumb: identify strong versus weak first, then decide whether an ICE table is needed.