Energetics and thermodynamics
Energetics and thermodynamics examines energy transfers, enthalpy cycles and the feasibility of reactions. The exam tests Hess's law, calorimetry, bond enthalpies, entropy and Gibbs free energy, always with attention to signs and states.
Core rules
- Enthalpy change ΔH is negative for exothermic reactions and positive for endothermic ones, at constant pressure.
- Standard enthalpy of formation ΔHf° is for one mole of compound from elements in their standard states; elements have ΔHf° = 0.
- Hess's law: the enthalpy change is independent of route. Construct cycles using formation, combustion or bond enthalpies.
- Calorimetry: q = mcΔT, where m is the mass of solution (assume density 1 g cm−3 for dilute aqueous), c = 4.18 J g−1 K−1, and ΔH = −q/n.
- Bond breaking is endothermic; bond making is exothermic. Mean bond enthalpies apply to gases and give approximate values.
- Entropy S is a measure of disorder; ΔS is positive when gas moles increase or temperature rises. Gibbs free energy ΔG = ΔH − TΔS; negative ΔG means the reaction is feasible, though it may be kinetically slow.
Worked example
Calculate ΔH for CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) using ΔHf°: CH4 = −74.8, CO2 = −393.5, H2O(l) = −285.8 kJ mol−1. ΔH = ΣΔHf(products) − ΣΔHf(reactants) = [−393.5 + 2(−285.8)] − [−74.8 + 2(0)] = −965.1 + 74.8 = −890.3 kJ mol−1. The negative sign shows heat is released.
Traps that cost marks
- Omitting state symbols: water vapour and liquid water have different formation enthalpies.
- Reversing the sign when reversing a reaction in a Hess cycle.
- Treating bond enthalpy as negative when breaking bonds; bond breaking always requires energy.
- Using mean bond enthalpies for liquids or solids, where they are not defined.
- Forgetting heat loss to the calorimeter, which makes experimental ΔH less exothermic than the true value.
- Assuming a negative ΔG means a fast reaction; thermodynamics says nothing about rate.