Reactivity 1 - What drives chemical reactions?
Reactivity 1 explains why reactions release or absorb energy and whether they occur spontaneously. The exam tests calorimetry, Hess's law, bond enthalpies, entropy and Gibbs free energy.
Enthalpy and calorimetry
- ΔH is negative for exothermic reactions and positive for endothermic ones. Standard conditions are 100 kPa and a stated temperature, usually 298 K.
- q = mcΔT; for dilute aqueous solutions assume c = 4.18 J g⁻¹ K⁻¹ and density 1.00 g cm⁻³.
- ΔH = −q/n, with n the moles of the limiting reactant.
- Breaking bonds is endothermic; forming bonds is exothermic. Using average bond enthalpies, ΔH = Σ(bonds broken) − Σ(bonds formed).
Hess's law and entropy
- Hess's law: ΔH is independent of route.
- ΔH°rxn = ΣΔHf°(products) − ΣΔHf°(reactants), using stoichiometric coefficients.
- ΔHf° of an element in its standard state is zero.
- Entropy increases from solid to liquid to gas, with temperature, with more particles and on mixing.
- ΔG = ΔH − TΔS; negative ΔG means spontaneous. ΔG = 0 at equilibrium.
Worked example
Find ΔH for CH₄ + 2O₂ → CO₂ + 2H₂O using ΔHf°: CH₄ −74.8, CO₂ −393.5, H₂O −285.8 kJ mol⁻¹. Products: −393.5 + 2(−285.8) = −965.1. Reactants: −74.8 + 0 = −74.8. ΔH = −965.1 − (−74.8) = −890.3 kJ mol⁻¹.
Traps
- Reversing the sign of ΔH when reversing an equation.
- Using bond enthalpies for liquid water.
- Mixing J and kJ in ΔG calculations.