AQA A-Level Chemistry: Reversible Reactions and Le Chatelier's Principle — mark scheme explained

The question

Consider the reaction: N 2 (g) + 3H 2 (g) ⇌ 2NH 3 (g). Predict how increasing the temperature will affect the position of equilibrium. [Paraphrased for study — not reproduced from any exam paper.]

Mark scheme, decoded

What each mark is really for — in plain English — and the wording trap that loses it.

• S1

  Step 1: Identify the type of reaction (exothermic or endothermic).

• S2

  The Haber process is exothermic (ΔH < 0).

• S3

  Step 2: Apply Le Chatelier’s principle to predict the effect of increasing temperature.

• S4

  For an exothermic reaction, increasing the temperature shifts the equilibrium to the left, favoring the reactants.

• S5

  Step 3: State the final answer.

• S6

  Increasing the temperature will shift the equilibrium to the left, favoring N 2 (g) and H 2 (g).

Model answer

Worked through, with each step tagged to the mark it earns.

1. S1

   Step 1: Identify the type of reaction (exothermic or endothermic).

2. S2

   The Haber process is exothermic (ΔH < 0).

3. S3

   Step 2: Apply Le Chatelier’s principle to predict the effect of increasing temperature.

4. S4

   For an exothermic reaction, increasing the temperature shifts the equilibrium to the left, favoring the reactants.

5. S5

   Step 3: State the final answer.

6. S6

   Increasing the temperature will shift the equilibrium to the left, favoring N 2 (g) and H 2 (g).

7. Final answer: The equilibrium will shift to the left, favoring N 2 (g) and H 2 (g).

Common mistakes

• Misunderstanding the direction of equilibrium shifts for exothermic reactions when temperature increases. — Remember that increasing the temperature for an exothermic reaction (ΔH < 0) shifts the equilibrium to the left, favoring the reactants.
• Forgetting that catalysts do not affect the position of equilibrium. — Always remember that a catalyst only speeds up both the forward and reverse reactions equally, allowing the system to reach equilibrium faster but not changing the position of equilibrium.
• Misinterpreting the effect of pressure on reactions involving gases. — Practice calculating Δn for different reactions and applying Le Chatelier’s principle to predict the effect of changes in pressure. For Δn 0, increasing pressure shifts the equilibrium to the left.
• Incorrectly predicting the effect of concentration changes on equilibrium. — Practice applying Le Chatelier’s principle to predict the direction of equilibrium shifts when concentrations change. Increasing the concentration of a reactant shifts the equilibrium to the right, and decreasing it shifts the equilibrium to the left.
• Failing to explain why compromise conditions are used in industrial processes. — Practice explaining how compromise conditions (temperature and pressure) optimize both yield and cost-effectiveness in industrial processes. For example, in the Haber process, a high temperature increases the rate but favors the reverse reaction, while a low temperature favors the forward reaction but decreases the rate.
• Confusing the effects of changes in conditions on equilibrium with other chemical concepts. — Review and practice applying Le Chatelier’s principle specifically to reversible reactions. Focus on how changes in temperature, pressure, and concentration affect the position of equilibrium.

Where the marks go

• Full worked solution (all marking points)3 marks

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