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Le Chatelier's principle states that if conditions change, the equilibrium position shifts to counteract the change. Free notes and 5 questions answered.
Understanding Equilibrium Systems and Le Chatelier's Principle is part of Reversible reactions and dynamic equilibrium in GCSE AQA Higher Chemistry.
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Le Chatelier's principle states that if conditions change, the equilibrium position shifts to counteract the change. Free notes and 5 questions answered.
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Use this lesson to stay inside Reversible reactions and dynamic equilibrium while connecting the detail back to the wider Topic 6: The rate and extent of chemical change topic.
Read the core explanations from Understanding Equilibrium Systems and Le Chatelier's Principle before testing yourself from memory.
Picture a busy motorway junction where cars can take two routes to the same destination. Sometimes more cars go one way, sometimes the other. Chemical reactions can work the same way - they're not always one-way streets.
When you heat blue copper sulphate crystals, they turn white and release water vapour. Cool them down, add water back, and they turn blue again. The reaction goes both ways: $CuSO_4 \cdot 5H_2O_{(s)} \rightleftharpoons CuSO_4_{(s)} + 5H_2O_{(g)}$
Here's where it gets interesting. In a closed system, both reactions happen simultaneously. Forward reaction: reactants → products. Backward reaction: products → reactants. Eventually, they reach a balance.
Imagine two teams in a tug-of-war, pulling with exactly equal force. The rope doesn't move, but both teams are still pulling hard. That's dynamic equilibrium - everything looks still, but there's constant activity underneath.
The key word is 'dynamic' - molecules are constantly reacting, but the overall amounts don't change. Think of it like a busy shop where customers enter and leave at exactly the same rate. The number inside stays constant, but there's continuous movement.
**Trap**: Many students think equilibrium means the reactions stop. They don't - both forward and backward reactions continue at equal rates. The concentrations stay constant because the rates are balanced.
Henri Le Chatelier discovered something fascinating about equilibrium systems: they're stubborn. Disturb them, and they'll shift to undo your interference. It's like a thermostat - turn up the heat, and it works harder to cool things down.
Let's see this in action with the Haber process: $N_2_{(g)} + 3H_2_{(g)} \rightleftharpoons 2NH_3_{(g)}$ + heat
Add more nitrogen? The system shifts right to use it up. Remove ammonia? The system shifts right to make more. Increase temperature? The system shifts left to absorb the extra heat (since the forward reaction releases heat).
There are three main ways to upset an equilibrium system, and each one triggers a predictable response. Let's examine each disturbance and how the system fights back.
Temperature Changes: If the forward reaction is exothermic (releases heat), increasing temperature shifts equilibrium left (towards reactants). The system tries to absorb the extra heat by favouring the endothermic backward reaction.
Pressure Changes: Only affects reactions involving gases. Increasing pressure shifts equilibrium towards the side with fewer gas molecules. The system reduces pressure by making fewer gas particles.
Concentration Changes: Add more reactant, equilibrium shifts right. Add more product, equilibrium shifts left. Remove a substance, equilibrium shifts to replace it.
Worked example: Predicting Equilibrium Shifts
For the reaction $2SO_2_{(g)} + O_2_{(g)} \rightleftharpoons 2SO_3_{(g)}$ + heat, predict what happens when: a) temperature increases, b) pressure increases
Answer: a) Higher temperature shifts left (less $SO_3$ formed) b) Higher pressure shifts right (more $SO_3$ formed)
Chemical companies use Le Chatelier's Principle to maximise profits. The Haber process makes ammonia for fertilisers, but the equilibrium doesn't naturally favour products. Engineers have to be clever.
For $N_2_{(g)} + 3H_2_{(g)} \rightleftharpoons 2NH_3_{(g)}$ + heat, high pressure shifts right (good for yield), but high temperature shifts left (bad for yield). The compromise? Moderate temperature (450°C) with high pressure (200 atmospheres) plus an iron catalyst.
The catalyst doesn't change the equilibrium position - it just helps the system reach equilibrium faster. Think of it as a more efficient route to the same destination.
Equilibrium questions follow predictable patterns. Examiners love testing whether you can predict shifts and explain your reasoning. Here's how to nail these questions every time.
**Common Exam Mistakes:** Don't say equilibrium 'moves' - it 'shifts'. Don't confuse rate with position - catalysts affect rate, not equilibrium position. Always explain your answer using Le Chatelier's Principle.
For 6-mark questions, use this structure: State the change → Identify which direction equilibrium shifts → Explain using Le Chatelier's Principle → Give the effect on yield/concentration → Consider any industrial implications.
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Move to the related lessons or back to the topic page when you need the wider sequence before exam-style practice.
Le Chatelier's principle states that if a system at equilibrium is subjected to a change in conditions (temperature, pressure or concentration), the equilibrium position shifts to counteract that change. For example, adding more nitrogen to the Haber process shifts the equilibrium right to use it up.
For 2SO₂ + O₂ ⇌ 2SO₃ + heat, increasing the temperature shifts the equilibrium left, so less SO₃ forms, because the forward reaction is exothermic. Increasing the pressure shifts it right, so more SO₃ forms, because there are 3 gas molecules on the left and 2 on the right.
A 6-mark equilibrium answer should state the change, identify which direction the equilibrium shifts, explain it using Le Chatelier's principle, give the effect on yield or concentration, and consider any industrial implications. Say the equilibrium shifts rather than moves.
Adding more product shifts the equilibrium left, towards the reactants, to use up the extra product. In the same way, adding more reactant shifts it right, and removing a substance shifts the equilibrium to replace it.
Henri Le Chatelier discovered that equilibrium systems shift to undo a change. His principle states that if a system at equilibrium is subjected to a change in temperature, pressure or concentration, the equilibrium position will shift to counteract that change.