"Exothermic = happens" becomes ΔG = ΔH − TΔS

Not addressed. Implicitly, reactions that release energy are the ones that happen; energetically "downhill" is treated as the reason a reaction occurs. Entropy does not exist at GCSE.

Entropy S is a measure of the dispersal/distribution of matter and energy (Reactivity 1.4.1), with absolute values (S⦵, in J K⁻¹ mol⁻¹ — note the units mismatch with ΔH) tabulated in the data booklet. Spontaneity at constant T and p is governed by Gibbs energy: ΔG = ΔH − TΔS, and a change is spontaneous iff ΔG < 0 (1.4.2–1.4.3). Four sign combinations give four behaviours (always spontaneous, never, spontaneous only at high T, only at low T). As a reaction proceeds toward equilibrium ΔG rises toward zero (1.4.4); at equilibrium ΔG = 0 and ΔG⦵ = −RT ln K.

The criterion for "will it happen?" is replaced. Endothermic reactions can be spontaneous (dissolving ammonium nitrate, ice melting above 0 °C) because TΔS beats ΔH. Also: "spontaneous" acquires a technical meaning that has nothing to do with speed — a spontaneous reaction can be infinitely slow (diamond → graphite).

Entropy requires a statistical/probabilistic idea of matter that GCSE has no room for.

Predicting that all endothermic reactions are impossible; equating "spontaneous" with "fast"; forgetting to convert ΔS from J to kJ before combining with ΔH in ΔG = ΔH − TΔS (the single most common HL arithmetic error in this topic); forgetting T must be in kelvin.