Rate stops being "how fast" and becomes a rate law

Rate is affected by temperature, concentration, surface area, catalyst, (pressure for gases). Explanation is collision theory: more frequent collisions, more energetic collisions. Rate measured as a gradient on a mass-loss or gas-volume graph.

Rate = k[A]^m[B]^n. The orders m and n can only be determined experimentally (Reactivity 2.2.9) — they are not the stoichiometric coefficients. The rate constant k has units that depend on overall order (2.2.11) and depends on temperature via the Arrhenius equation k = Ae^(−Ea/RT) (2.2.12), usually linearised as ln k = −Ea/R · (1/T) + ln A and handled by plotting ln k against 1/T. Mechanisms consist of elementary steps with molecularity (2.2.8); the rate-determining step is the slowest, and the rate law reflects the species present up to and including the RDS (2.2.6, 2.2.9). A proposed mechanism must (a) sum to the overall equation and (b) be consistent with the experimental rate law.

Two big things. (a) The rate law is an experimental fact about which the balanced equation tells you nothing — this is genuinely counter-intuitive and students fight it. (b) Kinetics becomes evidence about mechanism, i.e. it stops being descriptive and becomes inferential. Also the Maxwell–Boltzmann distribution is used quantitatively to explain why a modest temperature rise produces a large rate rise (it's the area beyond Ea, not the shift of the peak).

GCSE has no algebra for orders and no concept of mechanism.

Reading orders off the balanced equation (the classic error); saying "a catalyst lowers the activation energy of the reaction" without saying it provides an alternative pathway; saying "temperature increases rate because particles collide more often" as the main reason (the dominant factor is the exponential rise in the fraction of collisions exceeding Ea); forgetting that k changes with temperature but not with concentration; getting units of k wrong.