# The industrial production of ammonia uses the reversible reaction: N₂(g) + 3H₂(g) ⇌ 2NH₃(g). The forward reaction is exothermic. Explain how increasing the temperature and increasing the pressure each affect the position of equilibrium in this reaction, and explain why a catalyst is used in this process.

> OCR A-Level Chemistry B: Salters (H433) — 5.1 Rates, equilibrium and pH · Explain · 5 marks

> Ammonia is manufactured on a large scale using the Haber process. The reaction vessel is a closed system, and the conditions used must be carefully chosen to produce ammonia efficiently.

## Mark scheme (5 marks)

1. Increasing temperature shifts the equilibrium position in the direction of the endothermic reaction
2. So the yield / amount of ammonia decreases when temperature is increased
3. Increasing pressure shifts the equilibrium position towards the side with the smaller number of moles of gas
4. So increasing pressure increases the yield / amount of ammonia produced
5. A catalyst increases the rate of reaction by providing a different reaction pathway with a lower activation energy, so equilibrium is reached more quickly (but it does not change the position of equilibrium)

## Key terms

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

## Related

- [Revision notes for OCR A-Level Chemistry B: Salters (H433)](https://www.gradenine.co.uk/learn)
- [How to answer "Explain" questions](https://www.gradenine.co.uk/tools/command-word-cheatsheet)
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Source: [GradeNine](https://www.gradenine.co.uk/q/the-industrial-production-of-ammonia-uses-3480b389) · Published by Druglandscape Ltd.