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Inquiry Question 2: What factors affect equilibrium and how?

Investigate the effects of temperature, concentration, volume and/or pressure on a system at equilibrium and explain how Le Chatelier's principle can be used to predict such effects

A focused answer to the HSC Chemistry Module 5 dot point on Le Chatelier's principle. How concentration, pressure, volume and temperature shift equilibrium position, the role of catalysts, the Haber process worked example, and the past HSC questions markers reward.

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What this dot point is asking

NESA wants you to apply Le Chatelier's principle to predict how an equilibrium responds to changes in concentration, pressure, volume, and temperature, and to explain industrial applications (especially the Haber process). This builds directly on dynamic equilibrium and is one of the highest-frequency Module 5 questions, appearing every HSC paper in either short answer or extended response form.

The answer

Le Chatelier's principle

If a system at equilibrium is disturbed by a change in conditions, the system shifts in the direction that partially opposes the disturbance.

The principle predicts the direction of the shift but not its magnitude. The magnitude depends on Kc and the size of the disturbance.

Concentration changes

Adding a reactant (or removing a product) shifts the equilibrium to the right (toward products), because the system responds to "use up" the extra reactant.

Removing a reactant (or adding a product) shifts the equilibrium to the left (toward reactants).

Kc does not change. Concentration disturbances shift the position, not the constant.

Pressure changes (gas reactions)

Increasing pressure (by decreasing volume) shifts the equilibrium toward the side with fewer moles of gas.

Decreasing pressure (by increasing volume) shifts the equilibrium toward the side with more moles of gas.

If both sides have equal moles of gas (for example, H2+I2⇌2HIH_2 + I_2 \rightleftharpoons 2HI), pressure changes have no effect on the position.

Adding an inert gas at constant volume does not shift the equilibrium because partial pressures of reactants and products are unchanged.

Temperature changes

This is the only disturbance that changes Kc.

Increasing temperature shifts the equilibrium in the endothermic direction (the system absorbs the added heat).

Decreasing temperature shifts the equilibrium in the exothermic direction.

For an exothermic forward reaction (ΔH<0\Delta H < 0), the reverse is endothermic. Heating shifts left, cooling shifts right.

Catalysts

A catalyst increases the rate of both forward and reverse reactions equally. It does not shift the equilibrium position and does not change Kc. It only reduces the time taken to reach equilibrium.

Worked example

The Haber process is the canonical HSC application.

N2(g)+3H2(g)⇌2NH3(g),ΔH=−92 kJ/molN_{2(g)} + 3H_{2(g)} \rightleftharpoons 2NH_{3(g)}, \quad \Delta H = -92 \text{ kJ/mol}

Industrial conditions are chosen to maximise the rate of ammonia production while keeping yield economically viable.

Pressure: 200 atm. Forward reaction goes from 4 moles of gas to 2. High pressure shifts equilibrium right (toward ammonia), increasing yield. Pressure is limited by reactor cost.

Temperature: 400°C. Forward reaction is exothermic. Le Chatelier favours low temperature for yield. But low temperature reduces the rate. 400°C is the kinetics-thermodynamics compromise.

Iron catalyst. Speeds up the approach to equilibrium without changing position.

Continuous removal of ammonia. NH3NH_3 is liquefied and removed. Le Chatelier predicts the system shifts further right to replace the removed product.

Result: a sustained, economically viable rate of production at modest yield (around 15-20% per pass), with unreacted gases recycled.

Common traps

Confusing rate and position. A catalyst changes the rate, not the position. Temperature changes both.

Forgetting that Kc only changes with temperature. Concentration and pressure shifts change the position but not the constant.

Saying high pressure favours fewer molecules (without "of gas"). Pure solids and liquids do not count. State symbols matter.

Mixing up endothermic and exothermic. If the forward reaction is exothermic, heat is a product. Adding heat (raising T) shifts equilibrium left. Treat heat as a chemical species when in doubt.

Adding inert gas at constant pressure. This expands the volume and decreases partial pressures, so the equilibrium does shift. At constant volume, it does not. Read the question carefully.

In one sentence

Le Chatelier's principle predicts the direction in which a disturbed equilibrium will shift to partially oppose the change: add a reactant or remove a product to shift right, increase pressure to shift to the side with fewer moles of gas, heat to shift in the endothermic direction, and a catalyst changes rate but not position.

Past exam questions, worked

Real questions from past NESA papers on this dot point, with our answer explainer.

2022 HSC5 marksThe industrial production of ammonia by the Haber process is represented by the equation N₂(g) + 3H₂(g) ⇌ 2NH₃(g), ΔH = -92 kJ/mol. Apply Le Chatelier's principle to explain the choice of industrial conditions (200 atm, 400°C, iron catalyst) used in this process.
Show worked answer →

A 5 mark answer needs to apply Le Chatelier to each condition and resolve the temperature compromise.

Pressure (200 atm). The forward reaction converts 4 moles of gas (1 + 3) to 2 moles of gas. Increasing pressure shifts the equilibrium toward the side with fewer moles of gas, that is, toward ammonia. High pressure increases yield. The pressure is limited to ~200 atm because higher pressures require expensive, thick-walled reactors.

Temperature (400°C). The forward reaction is exothermic (ΔH<0\Delta H < 0). Le Chatelier predicts that decreasing temperature shifts the equilibrium right (toward ammonia), increasing yield. However, lower temperature slows the rate of reaction. 400°C is a compromise. High enough for an acceptable rate, low enough to maintain a workable equilibrium position.

Iron catalyst. A catalyst does not shift the equilibrium position. It increases the rate of both forward and reverse reactions equally, allowing equilibrium to be reached faster at the chosen temperature.

Removing ammonia. Liquid ammonia is condensed and removed continuously, decreasing product concentration and shifting the equilibrium further right (Le Chatelier's response to product removal).

Markers reward (1) correct prediction of direction for each factor, (2) explicit reference to "fewer moles of gas" and "endothermic/exothermic", (3) recognising the temperature compromise (yield vs rate), (4) noting that a catalyst affects rate not position.

2019 HSC3 marksA sealed flask contains the equilibrium 2NO₂(g) ⇌ N₂O₄(g), ΔH = -57 kJ/mol. The flask is heated. Predict and explain the colour change observed.
Show worked answer →

NO2NO_2 is brown; N2O4N_2O_4 is colourless. The forward reaction is exothermic (ΔH=−57\Delta H = -57 kJ/mol), so the reverse reaction is endothermic.

Le Chatelier's principle states that when temperature is increased, the equilibrium shifts in the endothermic direction to absorb the added heat. Here, the endothermic direction is the reverse (right to left), forming more NO2NO_2.

Therefore the equilibrium shifts left, [NO2][NO_2] increases, and the gas mixture becomes a darker brown.

Markers reward (1) identifying that the reverse reaction is endothermic, (2) stating the equilibrium shifts in the endothermic direction, (3) connecting the shift to a darker brown colour.

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