Thermochemistry, Energetics and Chemical Thermodynamics
This domain tests enthalpy changes, Hess’s law and the direction of spontaneous change. STPM questions often combine calorimetry, bond energies and Gibbs free energy.
Enthalpy changes
ΔH is negative for exothermic reactions and positive for endothermic reactions. Standard conditions are 298 K, 1 atm and 1 mol dm−3. Standard enthalpy of formation is the enthalpy change when one mole of a compound forms from its elements in standard states. Bond breaking is endothermic; bond forming is exothermic.
Hess’s law and cycles
ΔH for a reaction is independent of route. Using formation enthalpies: ΔHr = ΣΔHf(products) − ΣΔHf(reactants). Using bond energies: ΔHr = Σbonds broken − Σbonds formed. Born–Haber cycles apply Hess’s law to ionic lattice formation.
Entropy and Gibbs free energy
Entropy S is a measure of disorder; ΔSuniverse must be positive for a spontaneous change. ΔG = ΔH − TΔS; a reaction is spontaneous when ΔG is negative. Temperature must be in kelvin.
Worked example
CH4 + 2O2 → CO2 + 2H2O(l). ΔHf: CH4 = −74.8, CO2 = −393.5, H2O = −285.8 kJ mol−1. ΔHr = [−393.5 + 2(−285.8)] − [−74.8 + 0] = −393.5 − 571.6 + 74.8 = −890.3 kJ mol−1.
Traps that cost marks
Watch signs carefully; subtracting a negative value adds. Include state symbols because ΔHf depends on state. Do not confuse bond energy with enthalpy of formation. Entropy units are J K−1 mol−1, while ΔH is usually kJ mol−1-convert before using ΔG.