Chemical Energetics
Chemical energetics tests enthalpy changes, calorimetry, Hess's law and bond enthalpies, including the sign conventions that carry most of the marks.
Core rules
- Exothermic reactions release heat: ΔH is negative. Endothermic reactions absorb heat: ΔH is positive.
- Standard enthalpy changes are measured at 298 K and 100 kPa, with solutions at 1 mol dm−3.
- Calorimetry: q = mcΔT, where m is the mass of solution heated (assume 1 g cm−3 for dilute aqueous solutions), c = 4.18 J g−1 K−1, and ΔT is the temperature rise. Then ΔH = −q/n, with n the moles reacting.
- Hess's law: the enthalpy change is independent of route. Using formation data, ΔHr = ΣΔHf(products) − ΣΔHf(reactants).
- Using bond enthalpies, ΔHr = Σ bonds broken − Σ bonds formed. Mean bond enthalpies are averages, so answers are approximate.
Worked example
Estimate ΔH for CH4 + 2O2 → CO2 + 2H2O(g). Bonds broken: 4 C–H (4 × 413 = 1652) + 2 O=O (2 × 498 = 996) = 2648 kJ mol−1. Bonds formed: 2 C=O (2 × 805 = 1610) + 4 O–H (4 × 463 = 1852) = 3462 kJ mol−1. ΔH = 2648 − 3462 = −814 kJ mol−1. The negative sign shows heat is released; bond formation releases energy.
Traps
Reversing the sign convention in q = mcΔT; using the mass of the reactant instead of the solution; forgetting to divide by moles; using bond enthalpies for liquid water when the equation shows gas; treating mean bond enthalpies as exact; forgetting that ΔHf of an element in its standard state is zero.