Engineers designing the Haber process wanted to run it at thousands of atmospheres of pressure to maximise ammonia yield — Le Chatelier said it would work, but the steel vessels needed to contain that pressure would cost more to build than the ammonia was worth.
Wrong: Le Chatelier's Principle says the system opposes the disturbance by returning to the original concentrations.
Right: Le Chatelier's Principle states the system shifts to minimise the disturbance. After adding a reactant, the system consumes some of it — but the final concentration is still higher than before. The system minimises, not eliminates, the disturbance. Original concentrations are not restored.
Use the PDF for classwork, homework or revision. It includes key ideas, activities, questions, an extend task and success-criteria proof.
A student reads the following in a textbook: "Adding a platinum catalyst to the equilibrium 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) shifts the equilibrium to the right, producing more SO₃." The student underlines this as a key fact to remember. Before reading on — is the student correct? If not, what has the textbook apparently gotten wrong, and what is the correct statement? Write your analysis now.
Pressure/volume LCP rule:
Important: Count ONLY gaseous species — ignore solids and aqueous species
Catalyst rule: lowers Eₐ equally for both directions → no shift in equilibrium position; no change in Keq
⚠ Keq depends ONLY on temperature — not concentration, pressure, or catalyst
Pressure changes only affect equilibria that involve gases — and when they do, the system shifts toward whichever side has fewer gas molecules, because that reduces the pressure and partially counteracts the disturbance.
To apply this rule: count the moles of GAS on each side of the balanced equation (ignore solids, liquids, and aqueous species — only gases contribute to pressure).
Example 1: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
Example 2: H₂(g) + I₂(g) ⇌ 2HI(g)
Example 3: CaCO₃(s) ⇌ CaO(s) + CO₂(g)
Pressure change decision tree — always count gas moles on each side first
Le Chatelier gives you the direction of shift for pressure changes — collision theory explains the mechanism at the particle level.
When the volume of a gas-phase equilibrium system is decreased (pressure increased), the concentration of all gaseous species increases simultaneously. Both forward and reverse collision frequencies increase. However, the side with MORE moles of gas experiences a proportionally larger increase in collision frequency.
For the Haber process (4 moles gas left → 2 moles gas right): when pressure is doubled by halving the volume, the forward collision rate (involving 4 moles of gas per unit reaction) increases more than the reverse collision rate (involving 2 moles per unit reaction). Forward rate now > reverse rate → net forward reaction → equilibrium shifts right → more NH₃ produced.
The nitrogen dioxide/dinitrogen tetroxide equilibrium is a single experiment that simultaneously demonstrates pressure effects and temperature effects on equilibrium.
Equilibrium: 2NO₂(g) ⇌ N₂O₄(g)
NO₂: brown/reddish-brown | N₂O₄: colourless
ΔH = −57 kJ/mol (forward reaction exothermic)
Gas moles: Left = 2 mol NO₂; Right = 1 mol N₂O₄
Pressure effect — compressing the gas (decrease volume, increase pressure):
Temperature effect:
The catalyst misconception is the single most common error in IQ2, and it requires understanding both what catalysts do AND what they don't do.
What a catalyst does:
What a catalyst does NOT do:
For a system already at equilibrium: adding a catalyst increases both rates equally → they remain equal → no shift → concentrations unchanged.
For a system not yet at equilibrium: adding a catalyst allows the system to reach the same equilibrium position faster — but the destination is identical to without a catalyst.
The iron catalyst in the Haber process allows profitable ammonia production at 400–500°C — without it, the rate at this temperature would be too slow. The catalyst does not improve yield — only temperature and pressure can do that.
One of the most reliably tested HSC questions in Module 5 is distinguishing between factors that change Keq and factors that only change the equilibrium position.
| Factor Changed | Effect on Equilibrium Position | Effect on Keq |
|---|---|---|
| Add reactant | Shifts right (toward products) | No change |
| Remove reactant | Shifts left | No change |
| Add product | Shifts left | No change |
| Remove product | Shifts right | No change |
| Increase pressure (gas, unequal moles) | Shifts toward fewer gas moles | No change |
| Decrease pressure (gas, unequal moles) | Shifts toward more gas moles | No change |
| Add catalyst | No shift | No change |
| Increase temperature (exothermic forward) | Shifts left | Decreases |
| Decrease temperature (exothermic forward) | Shifts right | Increases |
| Increase temperature (endothermic forward) | Shifts right | Increases |
| Decrease temperature (endothermic forward) | Shifts left | Decreases |
The most commonly tested Module 5 fact: temperature is the ONLY factor that changes the value of Keq
You diagnosed the misconception that a catalyst shifts equilibrium. A catalyst does not shift the position of equilibrium — it speeds up both the forward and reverse reactions equally, so the system reaches equilibrium faster but the final concentrations are unchanged. For pressure/volume: increasing pressure (or decreasing volume) shifts equilibrium toward the side with fewer moles of gas. Adding an inert gas at constant volume has no effect on equilibrium position.
Look back at what you wrote in the Think First section. What has changed? What did you get right? What surprised you?
Problem: For each equilibrium, predict the effect of increasing pressure by halving the container volume. State direction of shift (or no shift) and explain using mole counts. (a) N₂(g) + 3H₂(g) ⇌ 2NH₃(g). (b) H₂(g) + Cl₂(g) ⇌ 2HCl(g). (c) Fe₃O₄(s) + 4H₂(g) ⇌ 3Fe(s) + 4H₂O(g).
Problem: The Haber process N₂(g) + 3H₂(g) ⇌ 2NH₃(g), ΔH = −92 kJ/mol, is at equilibrium at 450°C and 200 atm. An engineer simultaneously increases the temperature to 550°C and increases the pressure to 300 atm. (a) Predict the direction of shift from temperature change alone. (b) Predict the direction of shift from pressure change alone. (c) Predict the overall direction when both are applied simultaneously — do the effects reinforce or oppose?
1 mark
Q1: The equilibrium 2NO(g) + O₂(g) ⇌ 2NO₂(g) has 3 moles of gas on the left and 2 moles of gas on the right. Which correctly predicts the effect of decreasing pressure by increasing the volume?
1 mark
Q2: Which statement about the effect of a catalyst on a chemical equilibrium is correct?
1 mark
Q3: For PCl₃(g) + Cl₂(g) ⇌ PCl₅(g), increasing pressure shifts equilibrium right. Which correctly explains why Keq is unchanged by this pressure increase?
3 marks
Q4: A syringe containing a mixture of NO₂(g) and N₂O₄(g) at equilibrium is suddenly compressed to half its volume. Describe and explain the two-stage colour change observed, including both the immediate observation and the observation after the system re-establishes equilibrium.
3 marks
Q5: A student claims: "If I add a catalyst to a flask containing SO₂, O₂, and SO₃ at equilibrium, the equilibrium will shift right and more SO₃ will be produced." Identify what is incorrect in this statement and write the correct version.
4 marks
Q6: The reaction N₂(g) + O₂(g) ⇌ 2NO(g), ΔH = +180 kJ/mol, has equal moles of gas on both sides. An industrial chemist proposes: (a) increasing temperature; (b) increasing pressure. For each proposal, predict whether it would shift the equilibrium, which direction, and whether Keq would change. Then evaluate which variable is more useful for maximising NO yield.
Climb platforms, hit checkpoints, and answer questions on Le Chatelier's Principle — Concentration & Pressure. Quick recall from lessons 1–6.