Revise: Acid-Base Reactions
Theories of acids and bases, pH and the ionisation of water, weak acid/base equilibria (Ka), salt hydrolysis, and how buffer solutions resist pH change.
Brønsted–Lowry acid = proton donor; base = proton acceptor.
NH₃ + H₂O ⇌ NH₄⁺ + OH⁻: water is the acid, ammonia is the base.
Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C, in ANY aqueous solution.
[H⁺]=1e-3 → [OH⁻]=1e-11.
Strong acids ionise completely; weak acids only partially.
0.1M HCl → pH≈1; 0.1M acetic acid → pH≈2.9.
A buffer needs a weak acid AND its conjugate base, both present.
Blood's H₂CO₃/HCO₃⁻ buffer keeps pH near 7.4.
Arrhenius acid increases [H⁺] in water; base increases [OH⁻] in water.
Vinegar (acid) turns litmus red; ammonia (base) turns it blue.
A broader theory than Arrhenius — works for NH₃, which has no OH⁻.
Water can act as either acid or base (amphoteric).
Kw links [H⁺] and [OH⁻] in every aqueous solution.
[OH⁻]=1e-2 → [H⁺]=1e-12.
Ka quantifies weak acid strength; use it to calculate weak-acid pH.
0.10M acetic acid (Ka=1.8e-5) → pH ≈ 2.87.
Strong bases (NaOH) dissociate completely; weak bases (NH₃) partially react with water.
NaOH is already Na⁺/OH⁻ ions in the solid — no equilibrium needed.
Weak-acid salts hydrolyse basic; weak-base salts hydrolyse acidic.
Sodium acetate solution is basic; ammonium chloride solution is acidic.
pH = pKa + log([A⁻]/[HA]) — when [A⁻]=[HA], pH = pKa.
pKa=4.74, ratio=2 → pH = 5.04.