Revision notes for AQA GCSE Chemistry The effect of pressure changes on equilibrium (HT only). Open the guide for explanations and worked examples. Written against the AQA GCSE Chemistry (8462) specification, so the content matches what's examinable rather than general Chemistry background.
The effect of pressure changes on equilibrium (HT only)
What you'll learn
- What dynamic equilibrium means in a reversible reaction.
- Why pressure changes only affect equilibria involving gases.
- How to count gas molecules from a balanced symbol equation.
- How to predict whether a pressure change increases or decreases the amount of product.
The starting point: reversible reactions
A reversible reaction is a reaction that can go in both directions. The forward reaction makes products; the reverse reaction turns products back into reactants.
We show this using the equilibrium symbol ⇌, not a one-way arrow.
For example:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
The reaction from left to right makes ammonia. The reaction from right to left breaks ammonia back into nitrogen and hydrogen.
Definition
Reversible reaction
A reversible reaction is a reaction where the products can react to reform the original reactants. It is shown using ⇌.
Dynamic equilibrium
In a closed system, no substances can enter or leave the reaction mixture. In a closed system, a reversible reaction can reach dynamic equilibrium.
At dynamic equilibrium:
- the forward reaction and reverse reaction are still happening
- the forward and reverse reactions happen at the same rate
- the amounts of reactants and products stay constant overall
Definition
Dynamic equilibrium
Dynamic equilibrium is reached in a closed system when the forward and reverse reactions occur at the same rate, so the amounts of reactants and products do not change overall.
Key Idea
Dynamic means still moving
At equilibrium, the reaction has not stopped. Both directions are still happening, but they balance each other exactly.
Equilibrium position
The equilibrium position tells you which side of the reversible reaction is favoured at equilibrium.
- If the equilibrium position shifts to the right, more products are formed.
- If the equilibrium position shifts to the left, more reactants are formed.
Definition
Equilibrium position
The equilibrium position describes the relative amounts of reactants and products present at equilibrium.
This Higher Tier topic is about predicting how the equilibrium position changes when the pressure is changed.
Why pressure matters for gases
Pressure is especially important in reactions involving gases because gas particles are far apart and can be squeezed into a smaller volume.
Definition
Pressure
Pressure is caused by gas particles colliding with the walls of their container. A higher pressure means the gas particles are more crowded.
If you compress a gas mixture into a smaller volume, the pressure increases. If you allow the gas mixture to expand into a larger volume, the pressure decreases.
For GCSE Chemistry, pressure changes are only considered for substances with the state symbol (g).
Le Chatelier’s principle
When a system at equilibrium is disturbed, it shifts to reduce the effect of the change. This is called Le Chatelier’s principle.
Definition
Le Chatelier’s principle
If the conditions of a system at equilibrium are changed, the equilibrium position shifts in the direction that opposes the change.
For pressure changes, the system “opposes” the change by shifting towards the side with a different number of gas molecules.
The pressure rule for gaseous equilibria
For gaseous reactions at equilibrium, compare the number of gas molecules on each side of the balanced equation.
Key Idea
The pressure rule
- Increasing pressure shifts the equilibrium position towards the side with fewer gas molecules.
- Decreasing pressure shifts the equilibrium position towards the side with more gas molecules.
The Haber reaction is a classic example:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
There are 4 gas molecules on the left and 2 gas molecules on the right, so increasing pressure favours ammonia.
Example
Predicting the shift in the Haber reaction
For the equilibrium:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
Predict the effect of increasing pressure.
-
Count the gas molecules on the left-hand side: one N₂ molecule plus three H₂ molecules gives 1+3=41 + 3 = 41+3=4 gas molecules.
-
Count the gas molecules on the right-hand side: the coefficient 2 in front of NH₃ means 2 gas molecules.
-
Increasing pressure favours the side with fewer gas molecules, so the equilibrium shifts to the right.
-
The amount of NH₃(g) increases, so the equilibrium yield of ammonia increases.
Counting gas molecules correctly
The “number of molecules” comes from the big numbers in the balanced equation. These big numbers are called coefficients.
For example:
2SO₂(g) + O₂(g) ⇌ 2SO₃(g)
On the left:
- 2SO₂(g) means 2 gas molecules
- O₂(g) has no big number, so it counts as 1 gas molecule
- total on the left is 2+1=32 + 1 = 32+1=3 gas molecules
On the right:
- 2SO₃(g) means 2 gas molecules
So increasing pressure shifts the equilibrium to the right, towards SO₃(g).
Common Mistake
Counting atoms instead of molecules
Do not count the atoms inside a formula. CO₂(g) counts as one gas molecule if there is no coefficient in front of it, even though it contains three atoms.
Only count gases
Pressure changes affect gases. If a substance has the state symbol (s), (l), or (aq), do not include it when counting gas molecules for this rule.
Common Mistake
Counting every substance
Only count substances with the state symbol (g). Ignore solids, liquids, and aqueous solutions when applying the pressure rule.
Example
Counting gases when solids are present
For the equilibrium:
CaCO₃(s) ⇌ CaO(s) + CO₂(g)
Predict the effect of increasing pressure.
-
Count the gas molecules on the left-hand side: CaCO₃ has state symbol (s), so there are 0 gas molecules on the left.
-
Count the gas molecules on the right-hand side: CO₂ has state symbol (g), so there is 1 gas molecule on the right.
-
Increasing pressure favours the side with fewer gas molecules, so the equilibrium shifts to the left.
-
The amount of CO₂(g) decreases because the equilibrium moves away from the side containing CO₂(g).
When pressure has no effect on equilibrium position
Sometimes both sides of the equation contain the same number of gas molecules. In that case, changing the pressure does not favour either side.
For example:
H₂(g) + I₂(g) ⇌ 2HI(g)
The left-hand side has 2 gas molecules. The right-hand side also has 2 gas molecules.
Common Mistake
Same number of gas molecules
If both sides have the same number of gas molecules, changing pressure does not shift the equilibrium position.
Example
Recognising no pressure shift
For the equilibrium:
N₂(g) + O₂(g) ⇌ 2NO(g)
Predict the effect of decreasing pressure.
-
Count the gas molecules on the left-hand side: one N₂ molecule plus one O₂ molecule gives 1+1=21 + 1 = 21+1=2 gas molecules.
-
Count the gas molecules on the right-hand side: 2NO(g) means 2 gas molecules.
-
Decreasing pressure would favour the side with more gas molecules, but both sides have the same number.
-
The equilibrium position does not shift, so the relative amounts of N₂(g), O₂(g), and NO(g) at equilibrium are unchanged.
Linking pressure changes to yield
The yield is the amount of product made. In equilibrium questions, you are often asked whether the yield of a named product increases or decreases.
Do not assume that increasing pressure always increases the yield. It depends on which side has fewer gas molecules.
For example:
CO(g) + 2H₂(g) ⇌ CH₃OH(g)
The product, methanol, is on the side with fewer gas molecules. So increasing pressure increases the equilibrium yield of methanol.
Tip
Quick memory check
High pressure favours the side with the smaller number of gas molecules. Low pressure favours the side with the larger number of gas molecules.
Common Mistake
Saying pressure increases the product automatically
A pressure increase only increases the product yield if the product side has fewer gas molecules. If the product side has more gas molecules, increasing pressure decreases the product yield.
A reliable method for any question
When you are given a pressure equilibrium question, use this method:
- Check the reaction is an equilibrium and look for the
⇌symbol. - Count only the substances with state symbol (g).
- Add up the coefficients on each side.
- Decide the direction of shift:
- increased pressure → fewer gas molecules
- decreased pressure → more gas molecules
- Link the direction of shift to the amount of the named substance.
Exam technique
In the exam
- Always count gas molecules, not atoms, masses, or all substances.
- State the direction of shift clearly: “to the left” or “to the right”.
- Finish by linking the shift to the named substance: “therefore the yield of ammonia increases” or “therefore less carbon dioxide is formed”.
Self review
Check yourself
- For 2SO₂(g) + O₂(g) ⇌ 2SO₃(g), which side is favoured by increasing pressure?
- Why does pressure have no effect on the equilibrium position for H₂(g) + I₂(g) ⇌ 2HI(g)?
- In a pressure equilibrium question, why should you ignore substances with state symbol (s) or (aq)?
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