Chemical Equilibrium and Acid-Base Chemistry
Scope
The exam tests equilibrium expressions and Le Chatelier shifts, ICE-table calculations, weak acid/base ionization, pH and buffers, titration curves, and solubility product equilibria.
Core rules
- Equilibrium constants include only species whose concentration changes; pure solids and liquids are omitted.
- Q compared with K predicts direction: Q < K shifts right, Q > K shifts left.
- Kp = Kc(RT)^Δn, where Δn counts gas moles of products minus reactants.
- For a weak acid, Ka = [H⁺][A⁻]/[HA]; with x = [H⁺], x ≈ √(Ka·C) when ionization is small. Validate the approximation with the 5% rule.
- pH = −log[H⁺]; pOH + pH = 14 at 25 °C; buffers follow pH = pKa + log([A⁻]/[HA]).
- Ksp expressions raise each ion concentration to its stoichiometric coefficient.
Worked example
Find the pH of 0.100 M acetic acid, Ka = 1.8 × 10⁻⁵. Set up: HA ⇌ H⁺ + A⁻, initial [HA] = 0.100, change −x. Then Ka = x²/(0.100 − x). Assume x is small: x² = 1.8 × 10⁻⁵ × 0.100 = 1.8 × 10⁻⁶, so x = 1.34 × 10⁻³ M. Check: 1.34 × 10⁻³/0.100 = 1.3%, below 5%, so the approximation holds. pH = −log(1.34 × 10⁻³) = 2.87.
Traps
- Thinking dilution changes Ka; it changes pH, not the constant.
- Applying the weak-acid approximation when ionization exceeds 5%, or when the acid is strong.
- Confusing titration endpoint (indicator color change) with equivalence point (stoichiometric neutralization).
- Omitting water autoionization for extremely dilute acids, where [H⁺] approaches 10⁻⁷ M.