Ionic Equilibrium
Unit 8 tests acid–base theories, pH, weak electrolyte equilibria, buffers, salt hydrolysis and solubility product.
Core rules
- Brønsted–Lowry acids donate protons and bases accept them; Lewis acids accept electron pairs.
- Kw = [H⁺][OH⁻] = 1 × 10⁻¹⁴ at 298 K, so pH + pOH = 14.
- Ostwald dilution law: Ka = Cα²/(1 − α); for weak acids with small α, [H⁺] = √(KaC) and pH = ½(pKa − log C).
- Buffer equation: pH = pKa + log([salt]/[acid]).
- Salts of weak acid and strong base hydrolyse to give basic solutions; strong acid and weak base give acidic solutions.
- Solubility product Ksp is the product of ionic concentrations in a saturated solution; precipitation occurs when the ionic product exceeds Ksp.
- Common ion effect suppresses ionization of a weak electrolyte.
Worked example
Calculate the pH of 0.01 M acetic acid with Ka = 1.8 × 10⁻⁵. Since α is small, [H⁺] = √(KaC) = √(1.8 × 10⁻⁵ × 0.01) = √(1.8 × 10⁻⁷) = 4.24 × 10⁻⁴ M. pH = −log(4.24 × 10⁻⁴) = 3.37.
Traps
- On diluting a weak acid, α increases but [H⁺] decreases.
- pH of 10⁻⁸ M HCl is about 6.98, not 8, because water contributes H⁺.
- Buffer capacity depends on the ratio of salt to acid, not on their absolute concentrations alone.
- Ksp expressions have no denominator; each ion is raised to its coefficient.
- pH + pOH = 14 holds only near 298 K.