Chemical Kinetics and Equilibrium
Core concepts
Kinetics describes reaction rates; equilibrium describes the balance of forward and reverse reactions. Rate depends on concentration, temperature, surface area, pressure (for gases), and catalysts. Collision theory requires sufficient energy (activation energy) and correct orientation.
A catalyst lowers activation energy without changing ΔH or the position of equilibrium. Rate equations may be determined experimentally; order is not read from the balanced equation.
Le Chatelier's principle predicts how a system responds to changes in concentration, pressure, or temperature. For an exothermic forward reaction, raising temperature shifts equilibrium to the left. Equilibrium constants (Kc) depend only on temperature.
Traps to watch
- Assuming a catalyst increases yield - it only speeds attainment of equilibrium.
- Reading reaction order from stoichiometric coefficients.
- Forgetting that K changes with temperature but not with concentration or pressure.
- Confusing rate with extent: a fast reaction may still have a small equilibrium constant.
- Applying pressure changes to systems with equal moles of gas on both sides.
Worked example
For N₂ + 3H₂ ⇌ 2NH₃ (exothermic), increasing pressure shifts equilibrium to the right because fewer gas moles are on the product side. Rule of thumb: pressure changes favour the side with fewer gas molecules.