Acids, Bases, Buffers and Solubility
This domain tests acid–base equilibria, buffer calculations, titration curves and solubility products. STPM expects correct use of Ka, Kw and Ksp.
pH and Ka
pH = −log[H+]. At 25 °C, Kw = [H+][OH−] = 1.0×10−14. Strong acids dissociate completely, so pH = −log(c). Weak acids partially dissociate: for a monoprotic weak acid, [H+] ≈ √(Ka × c). A buffer resists pH change and is usually a weak acid with its conjugate base; pH = pKa + log([A−]/[HA]).
Titration and indicators
Titration curves show pH against volume of titrant. The equivalence point depends on the acid–base combination. Indicators must change colour within the steep part of the curve. Strong acid–strong base has an equivalence pH of 7; weak acid–strong base is basic at equivalence.
Solubility product
For a sparingly soluble salt, Ksp is the product of ion concentrations raised to their stoichiometric powers. The common ion effect decreases solubility. Precipitation occurs when the ionic product Q exceeds Ksp.
Worked example
A buffer contains 0.10 mol dm−3 CH3COOH and 0.20 mol dm−3 CH3COONa; Ka = 1.8×10−5. pKa = −log(1.8×10−5) = 4.74. pH = 4.74 + log(0.20/0.10) = 4.74 + 0.301 = 5.04.
Traps that cost marks
Do not use pH = −log(c) for a weak acid. Buffer calculations use the ratio of conjugate base to acid, not absolute concentrations alone. Ksp expressions exclude the solid. Common ions reduce solubility; they do not increase it. Salt hydrolysis can make equivalence pH acidic or basic.