"Energy is given out / taken in" becomes enthalpy as a state function

Exothermic reactions give out heat (temperature goes up); endothermic take heat in. Energy level diagrams are drawn qualitatively. Bond breaking is endothermic, bond making exothermic; you can calculate "energy change = bonds broken − bonds made" using given bond energies.

ΔH is a defined, signed, per-mole quantity: the heat transferred at constant pressure under standard conditions and states (Reactivity 1.1.4), with the sign convention ΔH < 0 exothermic, ΔH > 0 endothermic. Because it is a state function, it is path-independent — this is Hess's law (1.2.2), and it enables ΔH⦵ = ΣΔH⦵f(products) − ΣΔH⦵f(reactants) (1.2.4) and, at HL, Born–Haber cycles for ionic lattices (1.2.5). Calorimetry becomes quantitative: q = mcΔT, then ΔH = −q/n, with attention to whose mass and whose moles. The reference state is elements in their standard states, ΔHf⦵ = 0 (using the plimsoll symbol ⦵ per D19).

Energy stops being a vague quantity and acquires: a sign convention, units (kJ mol⁻¹), a reference state (elements in standard states, ΔHf⦵ = 0), a "per mole of what?" specification, and an algebra (Hess cycles). Also, the system/surroundings distinction becomes load-bearing — a temperature rise in the water means heat left the system, so ΔH is negative. That sign flip is where most marks are lost.

GCSE needs students to predict "will the beaker feel hot?", not to compute a molar enthalpy with a sign.

Sign errors everywhere (the most common single arithmetic error in Reactivity 1); dividing by the wrong number of moles; using the mass of the solid reactant instead of the solution in q = mcΔT; treating average bond enthalpy calculations as exact; being unable to construct a Hess cycle because they don't believe path-independence is a theorem rather than a trick.