Le Chatelier's principle says that if you change a condition of a system at equilibrium, the position of equilibrium shifts to oppose the change. We use this to predict how to get more of a product.
- Concentration: increasing the concentration of a reactant, or removing a product, shifts the equilibrium towards the products.
- Temperature: raising the temperature favours the endothermic reaction. Lowering it favours the exothermic reaction.
- Pressure (gases only): increasing the pressure favours the side of the equation with fewer gas molecules. Decreasing it favours the side with more gas molecules.
A catalyst speeds up the forward and reverse reactions equally, so it does not change the position of equilibrium. It only helps equilibrium to be reached sooner.
In the Haber process, N2 + 3H2 ⇌ 2NH3, there are 4 molecules of gas on the left and 2 molecules of gas on the right. The forward reaction is exothermic. High pressure and a low temperature therefore give more ammonia, but a low temperature makes the reaction too slow. A compromise of about 450 °C is used, with a high pressure and an iron catalyst.