Acids, Bases and Salts
Definitions and strength
Brønsted–Lowry acids donate protons; bases accept them. Strong acids ionise completely in water; weak acids ionise partially and have an equilibrium constant Ka. Strength is not concentration: a dilute strong acid can have a higher pH than a concentrated weak acid. Conjugate pairs differ by one proton.
pH and calculations
pH = −log₁₀[H⁺]; [H⁺] = 10^(−pH). For a strong monoprotic acid, [H⁺] equals the acid concentration. For a weak acid, use Ka = [H⁺]²/[HA] as an approximation. At 25 °C, pH + pOH = 14 and Kw = 1.0 × 10⁻¹⁴.
Neutralisation and salts
Acid + base → salt + water. Acid + carbonate → salt + water + carbon dioxide. Acid + metal → salt + hydrogen. Salt identity follows the acid used: hydrochloric → chlorides, sulfuric → sulfates, nitric → nitrates.
Buffers and titrations
A buffer resists pH change and consists of a weak acid with its conjugate base. During a titration, the equivalence point is where moles of acid and base match the equation; the endpoint is the indicator colour change. Strong acid–strong base titrations have equivalence pH 7; strong acid–weak base is acidic; weak acid–strong base is basic.
Traps
- Assuming pH 7 always marks equivalence.
- Treating a weak acid as fully ionised.
- Confusing the terms endpoint and equivalence point.
- Forgetting that each mole of H₂SO₄ supplies two moles of H⁺.
- Using an inappropriate indicator for a weak acid–weak base titration.
Rule of thumb
Always write the balanced equation first; then use n = cV to find the reacting moles, and only then apply pH relationships.