Hard

Effect of Pressure and Temperature on Equilibrium

Simple Explanation

For gas equilibria, increasing pressure shifts equilibrium toward the side with FEWER gas molecules. Increasing temperature shifts equilibrium in the endothermic direction (absorbing the extra heat); decreasing temperature shifts it in the exothermic direction.

Why Do We Need It?

These two factors are the main "dials" chemists and engineers use to push an industrial equilibrium toward higher yield, and understanding them requires connecting Le Chatelier's principle to genuinely different underlying reasons (molecule count for pressure, ΔH for temperature) rather than one single rule.

Worked Example

Predict the effect of pressure on the Haber process

For N₂(g) + 3H₂(g) ⇌ 2NH₃(g), predict the effect of increasing pressure on the position of equilibrium.

Why Does This Work?

Increasing pressure (by compressing the container) is equivalent to increasing the concentration of every gas — the system responds by shifting toward whichever side has fewer total gas molecules, since that reduces the total number of gas particles and partially counteracts the pressure increase. For temperature, since heat behaves like a 'reactant' in an endothermic reaction (and a 'product' in an exothermic one), raising temperature shifts equilibrium as if a reactant (for the endothermic direction) had been added.

Real-Life Example

Choosing conditions for the Haber process

The industrial Haber process operates at high pressure (~200 atm) and a compromise temperature (~450°C).

High pressure favours ammonia (fewer gas moles on the product side) and low temperature would too (the forward reaction is exothermic) — but low temperature makes the reaction too slow, so a moderate temperature is chosen as a practical compromise between equilibrium yield and reaction rate.

Practice

The forward reaction N₂(g) + 3H₂(g) ⇌ 2NH₃(g) is exothermic. What happens to the equilibrium if temperature is INCREASED?

Hard

Common mistake

Applying the pressure rule to reactions where the number of gas moles is EQUAL on both sides — when the mole counts match, changing pressure does not shift the equilibrium at all, since neither side is favoured by having 'fewer' gas particles.

Quick Review

Higher pressure shifts equilibrium toward the side with fewer gas moles.
Higher temperature shifts equilibrium in the endothermic direction.
Equal gas moles on both sides means pressure has no effect on the equilibrium position.