Thermochemistry, Kinetics, and Chemical Equilibrium
This domain tests energy flow, reaction rates, and reversible reactions. Expect items on calorimetry, enthalpy, Hess’s law, Gibbs free energy, rate laws, collision theory, equilibrium constants, reaction quotients, and Le Châtelier’s principle.
Core rules and concepts
- q = mcΔT for calorimetry; sign conventions describe heat gained or lost by the system.
- ΔH is negative for exothermic reactions and positive for endothermic reactions; Hess’s law adds enthalpy changes for steps.
- ΔG = ΔH − TΔS; negative ΔG indicates a spontaneous process under stated conditions.
- Rate depends on concentration, temperature, catalyst, and surface area; a catalyst lowers activation energy without changing ΔH or K.
- For aA + bB ⇌ cC + dD, Kc = [C]^c[D]^d / ([A]^a[B]^b); pure solids and liquids are omitted.
- Q compared with K predicts direction: Q < K shifts right, Q > K shifts left.
Worked example
For N₂(g) + 3H₂(g) ⇌ 2NH₃(g), Kc = [NH₃]² / ([N₂][H₂]³). If Q is less than K, the forward reaction is favored until Q = K. Increasing pressure shifts the system toward fewer gas moles, here toward NH₃. For the exothermic synthesis, increasing temperature shifts the equilibrium left and lowers K. Adding a catalyst speeds both directions equally and does not change the equilibrium position.
Traps that cost marks
- Reversing the sign of ΔH for an exothermic reaction.
- Thinking a catalyst increases equilibrium yield.
- Including solids or liquids in K expressions.
- Confusing reaction rate with equilibrium position.
- Assuming K changes with concentration or pressure.