Chapter summary
Chemical Energetics: Thermochemistry, at a glance
Read reaction energy, measure heat and use supplied cycles without losing signs or stoichiometric factors.
A-Level 8873, revised syllabus (2026-2027)
Quick revision
Revisit the essentials, then return to an explanation when you need it.
| Term | Definition to write |
|---|---|
| Standard enthalpy change of formation, ΔHf° | The enthalpy change when one mole of a substance is formed from its constituent elements in their standard states under standard conditions (298 K, 100 kPa). |
| Standard enthalpy change of combustion, ΔHc° | The enthalpy change when one mole of a substance is completely burnt in excess oxygen under standard conditions (298 K, 100 kPa). |
| Standard enthalpy change of neutralisation | The enthalpy change when one mole of water is formed by the reaction of an acid with a base in dilute aqueous solution under standard conditions (298 K, 100 kPa). |
| Bond energy | The average energy required to break one mole of a particular covalent bond in gaseous molecules. |
| Lattice energy | The enthalpy change when one mole of a solid ionic compound is formed from its constituent gaseous ions under standard conditions (298 K, 100 kPa). It is always negative. |
| Hess's law | The enthalpy change of a reaction depends only on the initial and final states, and not on the route taken. |
| Step | Remember |
|---|---|
| Sign | Exothermic reaction negative; its surroundings gain heat. |
| One mole of what? | Formation: compound; combustion: fuel; neutralisation: water; lattice: solid. |
| Calorimetry | q = mcDeltaT; use all heated solution mass; reaction heat is -q. |
| Hess cycle | Reverse arrow → reverse sign; scaled equation → scaled enthalpy. |
| Bond energies | Broken minus formed, using gas-phase bonds. |
| Lattice comparison | Higher charges/smaller ions → greater magnitude, more negative formation energy. |
Quick recall
Define standard enthalpy change of formation.