Le Chatelier's Principle – IGCSE Chemistry Definition

Le Chatelier’s principle is the rule used to predict how the position of equilibrium shifts when conditions change. It is a Supplement topic on the 0620 syllabus and is tested on Papers 2 and 4, particularly in the context of the Haber process and the Contact process.

The 0620 definition

Le Chatelier’s principle states that if a system at equilibrium is subjected to a change in conditions (temperature, pressure, or concentration), the position of equilibrium shifts to oppose (counteract) the change.

Applying the principle

1. Change in temperature

Change	Equilibrium shifts toward	Reason
Increase temperature	Endothermic direction	Absorbs the extra heat, opposing the rise
Decrease temperature	Exothermic direction	Releases heat, opposing the fall

Example: N₂ + 3H₂ ⇌ 2NH₃ (forward is exothermic)

• Increase temperature: shifts left (endothermic) — less NH₃, more N₂ and H₂
• Decrease temperature: shifts right (exothermic) — more NH₃

2. Change in pressure (gases only)

Change	Equilibrium shifts toward	Reason
Increase pressure	Side with fewer gas molecules	Reduces the number of gas molecules, opposing the pressure increase
Decrease pressure	Side with more gas molecules	Increases the number of gas molecules, opposing the pressure decrease

Example: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) — 4 moles of gas on the left, 2 on the right

• Increase pressure: shifts right (fewer gas molecules) — more NH₃
• Decrease pressure: shifts left (more gas molecules) — less NH₃

3. Change in concentration

Change	Equilibrium shifts toward	Reason
Increase concentration of reactant	Products	Uses up some of the added reactant
Increase concentration of product	Reactants	Uses up some of the added product
Remove a product	Products	Replaces the removed product

4. Catalyst

A catalyst has no effect on the position of equilibrium. It speeds up both the forward and reverse reactions equally, so equilibrium is reached faster but the proportions are unchanged.

Haber process conditions explained

N₂(g) + 3H₂(g) ⇌ 2NH₃(g) (forward is exothermic)

Condition used	Le Chatelier prediction	Industrial reason
450 degrees C (compromise)	Lower temp favours NH₃ but rate is too slow at low temp	Compromise between yield and rate
200 atm (high pressure)	Shifts right — more NH₃	Higher pressure favours fewer gas molecules
Iron catalyst	No effect on position	Reaches equilibrium faster

Worked exam question

In the Contact process: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) (forward is exothermic). (a) State the effect on the yield of SO₃ of increasing the pressure. Explain. (2) (b) Explain why a high temperature is not used even though it would increase the rate. (2)

Mark scheme

(a) Yield of SO₃ increases [1]; because increased pressure favours the side with fewer gas molecules (3 on left, 2 on right) / shifts equilibrium to the right [1]

(b) The forward reaction is exothermic, so increasing temperature shifts the equilibrium to the left / toward reactants [1]; this would decrease the yield of SO₃ [1]

Common exam mistakes

• Saying a catalyst shifts the equilibrium. It does not — it only makes equilibrium reached faster.
• Forgetting to count gas molecules correctly. Only count moles of gas, not solids or liquids. In N₂ + 3H₂ ⇌ 2NH₃, count 1 + 3 = 4 on the left and 2 on the right.
• Confusing yield with rate. Low temperature increases yield in an exothermic reaction but decreases the rate. Industrial processes use a compromise temperature.