Chemical Energetics, Kinetics and Equilibrium
Core concepts
Separate three ideas: energetics (enthalpy change, exothermic vs endothermic, bond breaking and forming, Hess's law), kinetics (rate, collision theory, activation energy, catalysts, factors affecting rate), and equilibrium (dynamic equilibrium, Le Chatelier's principle, equilibrium constant expression). Know that catalysts speed up both forward and reverse reactions equally and do not change the equilibrium position or the enthalpy change.
Traps that appear in questions
- Sign conventions: exothermic reactions have negative ΔH; bond breaking absorbs energy, bond forming releases it.
- Catalyst misconceptions: a catalyst lowers activation energy and speeds attainment of equilibrium but does not shift the position or change ΔH.
- Le Chatelier applications: increasing pressure favours the side with fewer gas moles; increasing temperature favours the endothermic direction.
- Rate vs yield: a catalyst or higher temperature may change rate without changing equilibrium yield, and temperature changes both.
- Equilibrium constant: only temperature changes K; concentration and pressure changes shift position but not K.
Worked example
For N₂ + 3H₂ ⇌ 2NH₃ (exothermic), which change increases ammonia yield? Lowering temperature favours the exothermic forward direction, and increasing pressure favours the side with fewer gas moles (the product side). Both increase yield, although low temperature slows the rate, which is why a catalyst is used industrially.
Rule of thumb
For any change, ask first: does it affect rate, equilibrium position, or both?