Reversible Reactions and Equilibrium
Reversible reactions and equilibrium, in the separate GCSE Chemistry course, are treated in more depth than in Combined Science: a precise definition of dynamic equilibrium, Le Chatelier's principle for predicting how a change in concentration, pressure or temperature shifts the position of equilibrium, and the Haber process as the worked industrial example of why manufacturers use compromise conditions rather than the conditions that would give the highest possible yield.
Method
- Define dynamic equilibrium precisely: in a closed system, a reversible reaction reaches dynamic equilibrium when the forward and reverse reactions are happening at exactly the same rate, so the concentrations of reactants and products remain constant, even though both reactions are still taking place.
- Apply Le Chatelier's principle in words: if a system at equilibrium is subjected to a change in concentration, pressure or temperature, the position of equilibrium shifts to counteract that change.
- Apply the concentration rule: increasing the concentration of a substance on one side of the equation shifts the equilibrium position towards the other side, to use up some of the extra substance; decreasing a concentration shifts the equilibrium position towards that side, to replace what was removed.
- Apply the pressure rule for reactions involving gases: increasing the pressure shifts the equilibrium position towards the side of the equation with fewer moles of gas molecules; decreasing the pressure shifts it towards the side with more moles of gas molecules; if both sides have equal moles of gas, changing pressure has no effect on the position of equilibrium.
- Apply the temperature rule: increasing the temperature shifts the equilibrium position in the endothermic direction, absorbing some of the extra energy; decreasing the temperature shifts it in the exothermic direction, releasing energy to compensate.
- Learn the Haber process as the named example: nitrogen and hydrogen react in a reversible reaction to form ammonia, N2(g) + 3H2(g) -> 2NH3(g), the forward reaction is exothermic, and the industrial conditions are a compromise temperature of about 450 degrees C, a compromise pressure of about 200 atmospheres, and an iron catalyst.
- For an extended question on industrial equilibrium conditions, structure the answer around each condition in turn: state its effect on the position of equilibrium (and so the yield), state its separate effect on the rate of reaction, and then explain why a compromise is chosen between yield, rate and the cost or safety of the equipment.
Worked example
In the Haber process, nitrogen and hydrogen react to form ammonia in a reversible reaction: N2(g) + 3H2(g) -> 2NH3(g), and the forward reaction is exothermic. Predict and explain the effect of increasing the pressure on the position of equilibrium and on the yield of ammonia.
- Count the total moles of gas on the reactants side of the equation: 1 mol N2 + 3 mol H2 = 4 moles of gas.
- Count the moles of gas on the products side: 2 mol NH3 = 2 moles of gas.
- Compare the two sides: there are fewer moles of gas on the product (ammonia) side than on the reactant side.
- Apply the pressure rule: increasing pressure shifts the position of equilibrium towards the side with fewer gas molecules.
- Apply this to the Haber process: the equilibrium shifts towards the ammonia (product) side.
- Final answer: increasing the pressure increases the yield of ammonia, because the equilibrium position shifts towards the product side.
Practice questions
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Q1State what is meant by dynamic equilibrium.Show answer
Answer: The point in a reversible reaction, in a closed system, at which the forward and reverse reactions are happening at exactly the same rate, so the concentrations of reactants and products stay constant.
Q2State Le Chatelier's principle in your own words.Show answer
Answer: If a system at equilibrium is subjected to a change in concentration, pressure or temperature, the position of equilibrium shifts so as to counteract (oppose) that change.
Q3In the Haber process, N2(g) + 3H2(g) -> 2NH3(g), state the effect of decreasing the pressure on the position of equilibrium.Show answer
Answer: The position of equilibrium shifts towards the reactants (the side with more moles of gas), so the yield of ammonia decreases.
Q4In the Haber process, state the effect of removing some ammonia from the equilibrium mixture as it forms on the position of equilibrium.Show answer
Answer: The position of equilibrium shifts towards the products, to replace some of the ammonia that was removed, so more nitrogen and hydrogen react to form ammonia.
Q5State the two named reactants used in the Haber process.Show answer
Answer: Nitrogen and hydrogen.
Q6State the catalyst commonly used in the Haber process.Show answer
Answer: An iron catalyst.
Q7The forward reaction of a particular reversible reaction is exothermic. State whether the reverse reaction is exothermic or endothermic.Show answer
Answer: Endothermic, and by exactly the same amount of energy as the forward (exothermic) reaction.
Q8State the effect of increasing the temperature on the position of equilibrium of a reaction whose forward reaction is exothermic.Show answer
Answer: The position of equilibrium shifts towards the reactants (in the endothermic direction), so the yield of product decreases.
Exam-style questions
Written in the style of a GCSE Science exam paper, with a full mark scheme.
The Haber process is carried out at a pressure of around 200 atmospheres. Explain, in terms of the number of gas molecules on each side of the equation N2(g) + 3H2(g) -> 2NH3(g), why using a high pressure increases the yield of ammonia.
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Nitrogen dioxide, a brown gas, exists in equilibrium with dinitrogen tetroxide, a colourless gas, in the reaction 2NO2(g) -> N2O4(g). The forward reaction (forming N2O4) is exothermic. A sealed tube containing this equilibrium mixture is placed in a beaker of hot water. (a) State and explain the colour change that would be observed. (b) State the effect of increasing the pressure on this equilibrium, referring to the number of gas molecules on each side.
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The Haber process uses a temperature of about 450 degrees C and a pressure of about 200 atmospheres, with an iron catalyst. Both a lower temperature and a higher pressure would increase the percentage yield of ammonia at equilibrium. Explain why the Haber process instead uses a compromise temperature and pressure, rather than the conditions that would give the highest possible yield.
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This topic is chapter 17 of GCSE Chemistry Workbook, the whole course as one free printable PDF.
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