Acids, Bases, and Salts
Core concepts
Acids donate protons (Brønsted–Lowry) and bases accept them. Strong acids and bases dissociate completely; weak ones partially. pH = −log[H⁺], and pOH = −log[OH⁻], with pH + pOH = 14 at 25 °C. Kw = [H⁺][OH⁻] = 1 × 10⁻¹⁴ at 25 °C.
Candidates must handle neutralisation, titration calculations (c₁V₁/n₁ = c₂V₂/n₂), buffer behaviour, and salt hydrolysis. Salts of strong acid and strong base are neutral; strong acid with weak base are acidic; weak acid with strong base are basic.
Indicators must be chosen to match the equivalence-point pH. Titration curves reveal the character of the acid and base involved.
Traps to watch
- Treating concentration as amount of substance - strong acids are not necessarily concentrated.
- Using pH + pOH = 14 at temperatures other than 25 °C.
- Forgetting stoichiometric ratios in titration calculations for diprotic acids.
- Assuming all salts are neutral.
- Confusing the equivalence point with the endpoint of an indicator.
Worked example
25.0 cm³ of 0.10 mol dm⁻³ NaOH neutralises 20.0 cm³ of HCl. c(HCl) = (0.10 × 25.0)/20.0 = 0.125 mol dm⁻³ (1:1 ratio). Rule of thumb: convert volumes to dm³ or keep both in cm³, since the ratio cancels.