NSWChemistrySyllabus dot point

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.

Generated by Claude Opus 4.89 min answerUpdated 2026-06-22

Reviewed by: AI editorial process; not yet individually human-reviewed

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What this dot point is asking
The answer
Examples in context
Try this

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, $H_2 + I_2 \rightleftharpoons 2HI$ ), 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 ( $\Delta H < 0$ ), 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.

N_{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: $NH_3$ 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.

Examples in context

Example 1. BHP Newcastle blast furnace carbon monoxide equilibrium. Iron oxide reduction in the blast furnaces formerly operated by BHP at Newcastle relied on $Fe_2O_{3(s)} + 3CO_{(g)} \rightleftharpoons 2Fe_{(l)} + 3CO_{2(g)}$ , $\Delta H < 0$ . Plant operators kept CO partial pressures high by burning coke against a hot air blast, pushing the equilibrium to the right and harvesting iron at the hearth. Excess heat from the bottom of the furnace would have shifted equilibrium left, lowering yield, so coolant water jackets stabilised the temperature. The HSC Le Chatelier framework explains why historical operators tuned blast-air ratios, coke addition rates and slag composition simultaneously rather than treating any one variable in isolation.

Example 2. Sulfuric acid contact process at Port Kembla. The contact-process step $2SO_{2(g)} + O_{2(g)} \rightleftharpoons 2SO_{3(g)}$ , $\Delta H = -198 \text{ kJ mol}^{-1}$ , runs at the Incitec Pivot sulfuric acid plant near Port Kembla. Plant chemists are caught between two competing pressures: low temperatures shift equilibrium right (more $SO_3$ ) but slow the reaction unbearably. They compromise at 450 degrees C with a vanadium oxide catalyst that speeds the kinetics without changing $K_{eq}$ . They also feed slight excess oxygen to push the equilibrium further right under Le Chatelier, and remove $SO_3$ by absorption into 98 percent sulfuric acid, again driving net forward reaction.

Try this

Q1. State Le Chatelier's principle and explain why adding a catalyst does not change the position of equilibrium. [3 marks]

Cue. Statement of the principle (system partially opposes imposed change); catalyst speeds forward and reverse rates equally, so the equilibrium position is unchanged.

Q2. For the reaction $N_2O_{4(g)} \rightleftharpoons 2NO_{2(g)}$ , $\Delta H = +57 \text{ kJ mol}^{-1}$ , predict and justify the shift in equilibrium position when (a) the temperature is increased, (b) the volume is halved. [2+2 marks]

Cue. (a) Endothermic forward, so heating shifts right (more $NO_2$ ). (b) Halving volume doubles pressure; equilibrium shifts to the side with fewer moles of gas (left, towards $N_2O_4$ ).

Q3. Consider $2SO_{2(g)} + O_{2(g)} \rightleftharpoons 2SO_{3(g)}$ , $\Delta H = -198 \text{ kJ mol}^{-1}$ . (a) State the effect on yield of $SO_3$ of increasing temperature. (b) State the effect of increasing pressure. (c) Explain why industry operates at a moderate 450 degrees C rather than the lowest possible temperature. [1+1+3 marks]

Cue. (a) Lower yield. (b) Higher yield. (c) Lower temperature would maximise yield but the rate becomes too slow to be economic; the chosen temperature is a kinetic-thermodynamic compromise.

Exam-style practice questions

Practice questions written in the style of NESA exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.

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 ( $\Delta H < 0$ ). 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 →

$NO_2$ is brown; $N_2O_4$ is colourless. The forward reaction is exothermic ( $\Delta H = -57$ 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 $NO_2$ .

Therefore the equilibrium shifts left, $[NO_2]$ 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.