The Haber process produces ammonia using the reversible reaction: N₂(g) + 3H₂(g) ⇌ 2NH₃(g). The forward reaction is exothermic. Industrial conditions use a temperature of around 450 °C and a pressure of around 200 atmospheres, along with an iron catalyst. Explain how each of these three conditions affects the position of equilibrium and/or the rate of reaction, and why these conditions are chosen rather than more extreme alternatives.
Written & reviewed by James Millett — Biology (Imperial College London), PGCE Science (University of Cambridge).
In the Haber process, nitrogen and hydrogen gases react to form ammonia. The reaction is reversible and is carried out in a closed system, meaning the reaction eventually reaches dynamic equilibrium. Industrial chemists must balance the rate of reaction against the yield of ammonia when choosing operating conditions.
Model answer (5 marks)
A high pressure (≈200 atm) reduces the volume of the system, so the equilibrium shifts towards the side with fewer gas moles – the product side (2 NH₃ vs 4 moles of reactants). This gives a higher yield of ammonia.
Higher pressure also increases the collision frequency of N₂, H₂ and NH₃ molecules, so the rate of reaction is faster.
The forward reaction is exothermic. Lowering the temperature would favour the forward reaction and increase the equilibrium yield, but the rate would drop sharply. 450 °C is chosen as a compromise: it is low enough to give a reasonable equilibrium yield yet high enough to keep the reaction rate acceptable.
The iron catalyst provides an alternative reaction pathway with a lower activation energy, so the rate of reaction is increased and equilibrium is reached more quickly. The catalyst does not alter the equilibrium position or the final yield of ammonia.
Higher pressure also increases the collision frequency of N₂, H₂ and NH₃ molecules, so the rate of reaction is faster.
The forward reaction is exothermic. Lowering the temperature would favour the forward reaction and increase the equilibrium yield, but the rate would drop sharply. 450 °C is chosen as a compromise: it is low enough to give a reasonable equilibrium yield yet high enough to keep the reaction rate acceptable.
The iron catalyst provides an alternative reaction pathway with a lower activation energy, so the rate of reaction is increased and equilibrium is reached more quickly. The catalyst does not alter the equilibrium position or the final yield of ammonia.
Examiner tips
- Mention Le Chatelier’s principle for pressure and temperature effects; state that the catalyst only affects the rate, not the equilibrium. Use correct stoichiometry (4 → 2 moles).
- Show the temperature–rate trade‑off explicitly. Include the catalyst’s role in lowering activation energy.
Mark scheme (5 marks)
- A high pressure shifts the equilibrium position towards the side with fewer moles of gas (the products side, 2 moles vs 4 moles), so a higher yield of ammonia is produced.
- Higher pressure also increases the rate of reaction because gaseous particles are closer together, leading to more frequent successful collisions.
- A lower temperature would shift the equilibrium towards the exothermic (forward) reaction, giving a greater yield of ammonia, but the rate of reaction would be too slow; 450 °C is a compromise between yield and rate.
- The iron catalyst increases the rate of reaction by providing a different reaction pathway with a lower activation energy, so equilibrium is reached more quickly.
- The catalyst does not change the position of equilibrium or the yield of ammonia; it only increases how quickly equilibrium is reached.
Key terms in this question
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