Acids, Bases and Salts

Three definitions of acids and bases

Arrhenius: acid releases H⁺ in water, base releases OH⁻. Brønsted-Lowry: acid is a proton donor, base is a proton acceptor — broader, includes NH₃ + HCl ⇌ NH₄⁺ + Cl⁻ in non-aqueous media. Lewis: acid accepts an electron pair, base donates — broadest, includes BF₃, AlCl₃ as Lewis acids despite having no protons. Atkins Chapter 5 and Clayden Chapter 8 both stress that the Lewis definition unifies acid-base behaviour with electrophile-nucleophile chemistry. A conjugate acid-base pair differs by one proton: HCl/Cl⁻, NH₄⁺/NH₃.

Water autoionisation, pH and pOH

Water self-ionises: 2 H₂O ⇌ H₃O⁺ + OH⁻ with K_w = [H⁺][OH⁻] = 1.0×10⁻¹⁴ at 25 °C (note K_w rises with T because the autoionisation is endothermic). pH = −log[H⁺]; pOH = −log[OH⁻]; pH + pOH = 14 at 25 °C. Pure water has [H⁺] = 10⁻⁷ M, pH = 7 — neutral. Lemon juice ~ pH 2; blood pH 7.35–7.45; ammonia ~ pH 11. A change of one pH unit corresponds to a tenfold change in [H⁺].

Strong vs weak acids and bases — K_a, K_b

Strong acids (HCl, HNO₃, H₂SO₄, HBr, HI, HClO₄) ionise completely; their conjugate bases are very weak. Weak acids partially ionise: K_a = [H⁺][A⁻]/[HA]. For CH₃COOH, K_a = 1.8×10⁻⁵, pK_a = 4.74. For weak base, K_b for ammonia = 1.8×10⁻⁵, pK_b = 4.74. Conjugate-pair relation: K_a·K_b = K_w; pK_a + pK_b = 14. For 0.1 M acetic acid, [H⁺] ≈ √(K_a·C) = √(1.8×10⁻⁶) ≈ 1.34×10⁻³ M, pH ≈ 2.87 — a percent ionisation of about 1.3%.

Buffer solutions and the Henderson-Hasselbalch equation

A buffer resists pH change on adding small amounts of acid or base; it is a weak acid + its conjugate base (e.g. CH₃COOH + CH₃COONa) or a weak base + its conjugate acid (NH₃ + NH₄Cl). Henderson-Hasselbalch: pH = pK_a + log([A⁻]/[HA]). When [A⁻] = [HA], pH = pK_a — the centre of buffer capacity. Blood is buffered around pH 7.4 by H₂CO₃/HCO₃⁻ (pK_a 6.1) plus phosphate and protein. The MDCAT loves "calculate pH of 0.10 M acid + 0.10 M conjugate base" — the answer is just pK_a.

Salt hydrolysis and indicators

Salts of strong acid + strong base (NaCl) give neutral solutions. Salts of strong base + weak acid (CH₃COONa) hydrolyse to give basic solutions because A⁻ + H₂O ⇌ HA + OH⁻. Salts of weak base + strong acid (NH₄Cl) give acidic solutions. Acid-base indicators are weak acids/bases whose protonated and deprotonated forms have different colours; phenolphthalein changes from colourless to pink between pH 8.3–10. Choose an indicator whose pK_In lies in the steep part of the titration curve — methyl orange for strong acid + weak base, phenolphthalein for weak acid + strong base, either for strong-strong.