Chemical reactivity: rates, equilibrium, acids and bases
This standard covers reaction rate, dynamic equilibrium and acid–base behaviour. It tests collision-theory explanations, predictions from Le Chatelier's principle with justification, and pH calculations for strong and weak acids and bases.
Rates
- Rate rises with concentration or pressure, temperature, surface area and a catalyst.
- Explain using collisions between particles with energy at or above the activation energy and correct orientation.
Equilibrium
- Dynamic: forward and reverse rates are equal and concentrations stay constant in a closed system.
- Le Chatelier: a change in concentration, pressure or temperature shifts the position to oppose the change; the value of Kc changes only with temperature.
- Kc = [products]^coefficients ÷ [reactants]^coefficients, excluding solids and pure liquids.
Acids and bases
- Brønsted–Lowry: an acid donates a proton, a base accepts one. Strong acids ionise completely; weak acids set up an equilibrium with Ka ≪ 1.
- pH = −log[H₃O⁺]; [H₃O⁺][OH⁻] = Kw = 1.0 × 10⁻¹⁴ at 25 °C; pH + pOH = 14.
- For a monoprotic strong acid, [H₃O⁺] equals the acid concentration; for a weak acid it is lower, so the pH is higher than a strong-acid calculation would predict.
Worked example
Find the pH of 0.0100 mol L⁻¹ NaOH. NaOH is a strong base, so [OH⁻] = 0.0100 mol L⁻¹; pOH = −log(0.0100) = 2.00; pH = 14.00 − 2.00 = 12.00. The same concentration of HCl would give pH 2.00, because [H₃O⁺] and [OH⁻] are linked through Kw.
Traps
- Assuming a weak acid is fully ionised and using its concentration directly as [H₃O⁺].
- Saying equilibrium 'shifts' without stating the direction and the resulting change in concentration.
- Claiming a catalyst increases yield - it changes only the rate.
- Forgetting that diluting a weak acid increases its degree of ionisation while lowering [H₃O⁺].