The rate and extent of chemical change: Le Chatelier's principle and effect of changing conditions on equilibrium

What you'll learn

This guide covers Le Chatelier's principle and how changing conditions affects systems at equilibrium. You'll learn how to predict the effect of altering concentration, temperature and pressure on reversible reactions, and apply these principles to industrial processes like the Haber process. This topic is essential for AQA GCSE Chemistry Paper 2 and links directly to industrial chemistry applications.

Key terms and definitions

Reversible reaction — a reaction where the products can react together to form the original reactants; represented by the symbol ⇌

Dynamic equilibrium — the state in a closed system where the rate of the forward reaction equals the rate of the backward reaction, and the concentrations of reactants and products remain constant

Closed system — a system where no substances can enter or leave, though energy can be transferred to or from the surroundings

Le Chatelier's principle — when a change is made to a system at equilibrium, the position of equilibrium shifts to counteract that change

Position of equilibrium — the relative amounts of reactants and products in an equilibrium mixture; shifts left (towards reactants) or right (towards products) when conditions change

Exothermic reaction — a reaction that transfers energy to the surroundings, shown by a temperature rise

Endothermic reaction — a reaction that takes in energy from the surroundings, shown by a temperature decrease

Core concepts

Reversible reactions and dynamic equilibrium

A reversible reaction can proceed in both directions. The forward reaction converts reactants to products, whilst the backward reaction converts products back to reactants.

Example: ammonium chloride ⇌ ammonia + hydrogen chloride

In a closed system, reversible reactions reach dynamic equilibrium when:

The forward and backward reactions continue to occur
The rates of both reactions are equal
The concentrations of reactants and products remain constant
Macroscopic properties (colour, pressure, concentration) appear unchanged

Key points about equilibrium:

Equilibrium can only be reached in a closed system
The equilibrium mixture may contain mostly products, mostly reactants, or significant amounts of both
Equilibrium does not mean equal amounts of reactants and products
At equilibrium, reactions have not stopped—both directions continue at equal rates

Le Chatelier's principle explained

Le Chatelier's principle states that if a system at equilibrium is subjected to a change, the system responds by shifting the position of equilibrium to counteract that change.

The principle applies to three main changes:

Changing concentration
Changing temperature
Changing pressure (for gaseous equilibria)

The equilibrium "opposes" the change but does not completely reverse it. Understanding this principle allows chemists to predict how industrial processes can be optimised for maximum yield.

Effect of changing concentration

When the concentration of a reactant or product changes, the equilibrium shifts to counteract this change.

If concentration of a reactant increases:

The equilibrium shifts right (towards products)
More products are formed
This reduces the concentration of the added reactant

If concentration of a product increases:

The equilibrium shifts left (towards reactants)
More reactants are formed
This reduces the concentration of the added product

If concentration of a reactant decreases:

The equilibrium shifts left (towards reactants)
This attempts to replace the removed reactant

If concentration of a product decreases:

The equilibrium shifts right (towards products)
This attempts to replace the removed product

Example: Consider the equilibrium: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

If more nitrogen is added, the equilibrium shifts right to use up the excess nitrogen, producing more ammonia.

If ammonia is removed (e.g., by liquefying it), the equilibrium shifts right to replace the ammonia, increasing the yield of product.

Effect of changing temperature

Temperature changes affect equilibrium position differently depending on whether the forward reaction is exothermic or endothermic.

For an exothermic forward reaction (releases energy):

Increasing temperature shifts equilibrium left (towards reactants) because the system counteracts the heat by favouring the endothermic backward reaction
Decreasing temperature shifts equilibrium right (towards products) because the system counteracts the cooling by favouring the exothermic forward reaction

For an endothermic forward reaction (absorbs energy):

Increasing temperature shifts equilibrium right (towards products) because the system counteracts the heat by favouring the endothermic forward reaction
Decreasing temperature shifts equilibrium left (towards reactants) because the system counteracts the cooling by favouring the exothermic backward reaction

Example: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) (forward reaction is exothermic, ΔH = -92 kJ/mol)

Increasing temperature shifts equilibrium left, reducing ammonia yield. However, higher temperatures increase the rate of reaction, so industrial processes use a compromise temperature (around 450°C in the Haber process).

Effect of changing pressure

Pressure changes only affect equilibria involving gases. The equilibrium shifts to reduce pressure increases or increase pressure decreases.

Count the moles of gas on each side of the equation:

If pressure increases:

Equilibrium shifts towards the side with fewer moles of gas
This reduces the total number of gas molecules, decreasing pressure

If pressure decreases:

Equilibrium shifts towards the side with more moles of gas
This increases the total number of gas molecules, increasing pressure

If equal moles of gas on both sides:

Changing pressure has no effect on the position of equilibrium

Example: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

Left side: 1 + 3 = 4 moles of gas Right side: 2 moles of gas

Increasing pressure shifts equilibrium right (towards ammonia), increasing yield. The Haber process uses high pressure (200 atmospheres) to maximise ammonia production, though very high pressures are expensive and require stronger equipment.

Effect of catalysts on equilibrium

Catalysts do not affect the position of equilibrium.

A catalyst:

Increases the rate of both forward and backward reactions equally
Allows equilibrium to be reached more quickly
Does not change the composition of the equilibrium mixture
Reduces energy costs in industrial processes by allowing lower temperatures

In the Haber process, an iron catalyst speeds up the reaction without affecting the equilibrium yield of ammonia.

Industrial applications: The Haber process

The Haber process produces ammonia from nitrogen and hydrogen:

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

Industrial conditions represent a compromise between rate and yield:

Temperature: 450°C

Lower temperature increases yield (equilibrium shifts right for exothermic reaction)
Higher temperature increases rate of reaction
450°C is a compromise providing reasonable yield at acceptable rate

Pressure: 200 atmospheres

Higher pressure increases yield (equilibrium shifts towards fewer moles)
Very high pressure requires expensive, reinforced equipment
200 atmospheres balances yield against cost

Catalyst: iron

Increases rate of reaction
Allows lower temperature to be used
Reduces energy costs
Does not affect equilibrium position

Continuous removal of ammonia:

Ammonia is liquefied and removed
Equilibrium shifts right to replace ammonia
Increases overall yield
Unreacted nitrogen and hydrogen are recycled

Worked examples

Example 1: Predicting equilibrium shifts

Question: Hydrogen and iodine react reversibly to form hydrogen iodide: H₂(g) + I₂(g) ⇌ 2HI(g)

The forward reaction is exothermic.

(a) Predict the effect on the equilibrium position of increasing the concentration of hydrogen. [2 marks]

(b) Predict the effect on the equilibrium position of increasing the temperature. [2 marks]

Mark scheme answers:

(a)

Equilibrium shifts to the right / towards products / towards HI [1 mark]
More hydrogen iodide is formed / concentration of hydrogen decreases [1 mark]

(b)

Equilibrium shifts to the left / towards reactants [1 mark]
Because the backward reaction is endothermic / the system opposes the temperature increase by favouring the endothermic reaction [1 mark]

(c)

There are equal numbers of moles of gas on both sides [1 mark]
1 + 1 = 2 (or similar calculation showing 2 moles on each side) [1 mark]

Example 2: The Contact process

Question: Sulfur dioxide reacts with oxygen to form sulfur trioxide in the Contact process: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g)

The forward reaction is exothermic.

(a) State the effect on the equilibrium yield of sulfur trioxide of increasing the pressure. Explain your answer. [3 marks]

(b) State and explain the effect of using a vanadium(V) oxide catalyst on the equilibrium yield. [2 marks]

Mark scheme answers:

(a)

Yield increases / equilibrium shifts right / more SO₃ formed [1 mark]
There are fewer moles of gas on the right / product side [1 mark]
(3 moles on left, 2 moles on right) Increasing pressure shifts equilibrium towards fewer moles [1 mark]

(b)

No effect on yield / position of equilibrium unchanged [1 mark]
Catalyst speeds up rate of forward and backward reactions equally / helps reach equilibrium faster [1 mark]

Example 3: Thermal decomposition equilibrium

Question: Calcium carbonate decomposes reversibly when heated: CaCO₃(s) ⇌ CaO(s) + CO₂(g)

The forward reaction is endothermic.

Predict and explain the effect of increasing temperature on the position of equilibrium. [3 marks]

Mark scheme answer:

Equilibrium shifts right / towards products / more calcium oxide and carbon dioxide formed [1 mark]
The forward reaction is endothermic / takes in heat [1 mark]
System opposes the temperature increase by favouring the endothermic reaction / the reaction that absorbs heat [1 mark]

Common mistakes and how to avoid them

Confusing "equilibrium shifts left" with "reaction stops" — At equilibrium, both reactions continue but at equal rates. Shifting left means the backward reaction temporarily goes faster until a new equilibrium is established.
Thinking catalysts increase yield — Catalysts only affect the rate of reaching equilibrium, not the position of equilibrium or the final amounts of products. Always state that yield is unchanged when discussing catalysts.
Not counting moles of gas correctly — Only count gaseous substances when predicting pressure effects. Solids and liquids don't contribute to pressure. In N₂(g) + 3H₂(g) ⇌ 2NH₃(g), there are 4 moles of gas on the left (1+3), not 2 (from coefficients).
Confusing "forward reaction is exothermic" with "increasing temperature increases yield" — If the forward reaction is exothermic, increasing temperature favours the backward (endothermic) reaction, shifting equilibrium left and decreasing product yield.
Writing "equilibrium moves to oppose the change" without being specific — Always state which direction (left/right) and explain why. For example: "equilibrium shifts right to reduce the concentration of the added reactant."
Forgetting that equilibrium requires a closed system — Reversible reactions in open systems cannot reach equilibrium because substances can escape. Always mention the closed system when defining dynamic equilibrium.

Exam technique for "The rate and extent of chemical change: Le Chatelier's principle and effect of changing conditions on equilibrium"

Command word "predict" requires you to state what happens (e.g., "equilibrium shifts right") but doesn't always need detailed explanation. For "explain," you must give reasons using Le Chatelier's principle.
Mark allocation guides detail needed — 1 mark questions need one clear point (e.g., "yield increases"). 2-3 mark questions require the prediction plus explanation, often referencing moles of gas, exothermic/endothermic, or concentration changes.
Always reference the specific equilibrium — Don't write generic answers. Use the equation provided: state which side has more moles, which substance's concentration increases, or which direction is exothermic based on information given.
In industrial process questions, mention compromise conditions — For 4-6 mark questions on the Haber or Contact processes, explain why moderate (not extreme) conditions are used, balancing yield against rate and economic factors.

Quick revision summary

Le Chatelier's principle states that systems at equilibrium respond to changes by shifting to oppose them. Increasing reactant concentration or decreasing product concentration shifts equilibrium right, increasing yield. For exothermic forward reactions, increasing temperature shifts equilibrium left (decreasing yield), whilst decreasing temperature shifts it right. Increasing pressure shifts equilibrium towards the side with fewer gas moles. Catalysts speed up reaching equilibrium but don't affect yield. Industrial processes like Haber use compromise conditions balancing yield, rate and cost.