Reactions stop going to completion

Reactants → products. Reversible reactions exist as a special topic (the Haber process, hydrated copper sulfate) with the ⇌ symbol, and Le Châtelier's principle is taught as a set of rules for predicting yield. The default assumption everywhere else is that reactions run to completion.

Every reaction has an extent, quantified by the equilibrium constant K (Reactivity 2.3.2–2.3.3). Dynamic equilibrium means forward and reverse rates are equal and non-zero. K depends only on temperature — not on concentration, pressure, or catalyst. The reaction quotient Q (2.3.5) tells you which way a non-equilibrium mixture will move. At HL, K is connected to thermodynamics via ΔG⦵ = −RT ln K (Reactivity 2.3.7 / 1.4.4), and to acid–base chemistry via Ka, Kb, Kw (Reactivity 3.1).

"How far?" becomes a third independent question alongside "how much?" (stoichiometry) and "how fast?" (kinetics) — this is literally the title of Reactivity 2. Le Châtelier is demoted from an explanation to a qualitative prediction tool; the real explanation is Q vs K. And ICE-table calculations (initial / change / equilibrium) appear, which are algebraically much harder than anything at GCSE.

Equilibrium requires simultaneously holding a molecular-level dynamic picture and a macroscopic static observation (Johnstone's triangle problem). GCSE avoids the mathematics entirely.

Answering "the reaction stops at equilibrium"; asserting "concentrations of reactants and products are equal at equilibrium" (documented by RSC as stemming from confusing rate equality with concentration equality); believing a catalyst increases yield; believing adding more reactant changes K; treating weak acid dissociation as complete in a pH calculation. The RSC also documents students viewing the forward and reverse reactions as isolated — "if I add reactant the forward reaction runs until the extra is used up" — which makes Le Châtelier reasoning collapse.