Unit 8: Acids and Bases
AP Chemistry: 91 practice questions with detailed explanations.
Unit Study Guide
Executive Summary
Unit 8 quantifies proton transfer. Strong acids and bases dissociate completely; weak ones reach equilibrium described by Ka and Kb.
pH and pOH
pH = −log[H⁺]; pOH = −log[OH⁻]; pH + pOH = 14 at 25 °C. Strong acid/base pH comes straight from concentration (watch H₂SO₄'s two protons). Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C.
Weak acid and base equilibria
HA ⇌ H⁺ + A⁻ with Ka = [H⁺][A⁻]/[HA]. Solve with an ICE table; neglect x when Ka is small (5% rule). Ka × Kb = Kw for a conjugate pair. pKa = −log Ka; lower pKa means stronger acid. Stronger acid ↔ weaker conjugate base.
Buffers
A buffer is a weak acid plus its conjugate base (or weak base plus conjugate acid). It resists pH change: added acid is consumed by the base; added base by the acid. Henderson–Hasselbalch: pH = pKa + log([A⁻]/[HA]). Buffering is best within ±1 of pKa; buffer capacity grows with total concentration.
Titrations
Strong acid–strong base: equivalence point at pH 7. Weak acid–strong base: equivalence point above 7 (conjugate base hydrolyzes). Weak base–strong acid: below 7. At the half-equivalence point, pH = pKa. Indicators change color over pH ≈ pKa ± 1; pick one whose range brackets the equivalence point.
Structure and strength
Acid strength rises with bond polarity and conjugate-base stability: oxyacids strengthen with more oxygens and higher electronegativity; binary acids strengthen down a group (weaker H–X bond).
Quantitative skill-set
pH from strong and weak acids/bases; Ka/Kb ICE tables; percent ionization; buffer pH via Henderson–Hasselbalch; titration-curve landmarks; dilution of strong species.
Exam traps
[H⁺] from pH: 10^(−pH). Don't average pH values. At half-equivalence, [HA] = [A⁻] so pH = pKa. Polyprotic acids ionize stepwise (Ka1 >> Ka2). Adding water to a buffer changes concentration but barely changes pH (ratio preserved).