In the Haber process, nitrogen and hydrogen react to form ammonia.
N2(g)+3H2(g)⇌2NH3(g)\mathrm{N_{2}(g)} + \mathrm{3H_{2}(g)} \rightleftharpoons \mathrm{2NH_{3}(g)}N2(g)+3H2(g)⇌2NH3(g)
The table shows the percentage of ammonia in the equilibrium mixture at 200 atm and different temperatures.
| Temperature / °C | Percentage of ammonia at equilibrium (%) |
|---|---|
| 300 | 63 |
| 350 | 50 |
| 400 | 39 |
| 450 | 28 |
| 500 | 18 |
Use the table to deduce whether the forward reaction is exothermic or endothermic. Explain your answer.
The industrial conditions are a temperature of 450 °C, a pressure of 200 atm and an iron catalyst. Explain why these conditions are used. You should refer to rate, yield and cost in your answer.
A plant uses 5.6 tonnes of nitrogen. Calculate the maximum mass of ammonia, in tonnes, that could be made. Relative formula masses: N2=28,NH3=17\mathrm{N_{2}} = 28, \mathrm{NH_{3}} = 17N2=28,NH3=17
After recycling the unreacted gases, the plant obtains 6.12 tonnes of ammonia from the 5.6 tonnes of nitrogen. Calculate the percentage yield.
45 exam-style questions on Edexcel GCSE Chemistry 5.2 Reversible reactions and equilibria, covering 5.2.1 Reversible reactions and dynamic equilibrium, 5.2.2 The Haber process, and 5.2.3 Changing the position of a dynamic equilibrium. Each one has a worked solution and a mark scheme showing where the marks go.