Thermochemistry
What the exam tests
Thermochemistry covers enthalpy change, exothermic and endothermic reactions, calorimetry, Hess's law, bond energies and standard heats of combustion, formation and neutralisation.
Core rules
- Enthalpy change ΔH is negative for exothermic and positive for endothermic reactions.
- Calorimetry: q = mcΔT, with m the mass of solution being heated.
- Molar enthalpy change = q / n, where n is moles of the limiting reactant.
- Hess's law: the total enthalpy change is independent of the route taken.
- ΔH = Σ(bond energies broken) − Σ(bond energies formed).
- Standard conditions are 25 °C and 1 × 10⁵ Pa; standard enthalpy of formation refers to one mole of compound from its elements.
- Energy-level diagrams show reactants and products with the activation energy as a hump.
Worked example
50 g of water rises by 6.0 °C when 0.010 mol of a fuel burns. Find the molar enthalpy of combustion. q = mcΔT = 50 × 4.18 × 6.0 = 1254 J. ΔH = −q/n = −1254/0.010 = −1.25 × 10⁵ J/mol = −125 kJ/mol. The sign is negative because the water gained heat from an exothermic reaction.
Traps
- Dropping the negative sign for an exothermic reaction.
- Using the mass of the fuel instead of the water in q = mcΔT.
- Forgetting to divide by moles to get a molar value.
- Reversing a step in a Hess's law cycle without reversing its sign.
- Ignoring heat lost to the surroundings, which makes experimental values smaller.