Topic 8 covers acid-base chemistry from definitions to quantitative pH calculations. SL focuses on the Bronsted-Lowry model and titrations. HL adds Ka/Kb calculations, pH curves, indicators, and buffer systems.
Acid: proton (H⁺) donor. Base: proton acceptor. Every acid-base reaction involves a transfer of H⁺. Conjugate pairs: an acid loses H⁺ to become its conjugate base, and vice versa. Amphoteric/amphiprotic substances (e.g., H₂O, HCO₃⁻) can act as both acid and base.
pH = −log₁₀[H⁺]. pOH = −log₁₀[OH⁻]. pH + pOH = 14 (at 25°C). Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C. Strong acids (HCl, HNO₃, H₂SO₄): fully dissociate → pH calculated directly from concentration. Strong bases (NaOH, KOH): fully dissociate → calculate [OH⁻] first, then pH.
Weak acids partially dissociate: HA ⇌ H⁺ + A⁻. Ka = [H⁺][A⁻]/[HA]. pKa = −log Ka. Stronger weak acid has larger Ka (smaller pKa). For weak acid calculations: use ICE table, assume [H⁺] = x, solve Ka = x²/(c−x) ≈ x²/c if x << c. pH of weak acid > pH of strong acid at same concentration.
Buffer: solution that resists pH change when small amounts of acid/base are added. Made from weak acid + its conjugate base (e.g., CH₃COOH/CH₃COONa) or weak base + its conjugate acid. Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA]). Buffer works because added H⁺ reacts with A⁻ and added OH⁻ reacts with HA.
Strong/weak refers to the degree of dissociation: a strong acid fully dissociates (HCl → H⁺ + Cl⁻); a weak acid partially dissociates (CH₃COOH ⇌ H⁺ + CH₃COO⁻). Concentrated/dilute refers to the amount of solute per volume. You can have a concentrated weak acid (lots of CH₃COOH in a small volume) or a dilute strong acid (small amount of HCl in lots of water). These are independent properties.
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