# 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.

> OCR A-Level Chemistry B: Salters (H433) — 5.3 Transition elements · Explain · 5 marks

> 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.

## Mark scheme (5 marks)

1. 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.
2. Higher pressure also increases the rate of reaction because gaseous particles are closer together, leading to more frequent successful collisions.
3. 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.
4. 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.
5. The catalyst does not change the position of equilibrium or the yield of ammonia; it only increases how quickly equilibrium is reached.

## Key terms

- [catalyst](https://www.gradenine.co.uk/glossary/catalyst)

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Source: [GradeNine](https://www.gradenine.co.uk/q/the-haber-process-produces-ammonia-using-8079c580) · Published by Druglandscape Ltd.