The reaction between hydrogen and iodine vapour is reversible. In a closed container, the following reaction reaches dynamic equilibrium: H₂(g) + I₂(g) ⇌ 2HI(g). The forward reaction is exothermic. Explain what happens to the position of equilibrium and to the concentration of HI when the temperature of the system is increased.

Pearson Edexcel International GCSE Chemistry (4CH1) — 3.3 Reversible reactions and equilibria · Explain · 4 marks · View as Markdown

Written & reviewed by James Millett — Biology (Imperial College London), PGCE Science (University of Cambridge).

Model answer (4 marks)

Increasing the temperature supplies energy to the system. Because the forward reaction is exothermic, the reverse reaction is endothermic and is favoured by the added heat. Consequently the equilibrium shifts to the left, towards the reactants H₂ and I₂. The concentration of the product HI therefore decreases.

Examiner tips

  • State that the forward reaction is exothermic and the reverse is endothermic.
  • Explain that increasing temperature favours the endothermic direction.
  • Show that the equilibrium shifts left and HI concentration falls.
  • Use the phrase "shifts to the left" to demonstrate understanding of Le Chatelier’s principle.

Common mistakes

  • Confusing the direction of shift (saying it moves right instead of left).
  • Failing to mention that the reverse reaction is endothermic.
  • Not linking the temperature change to the shift in equilibrium position.

Mark scheme (4 marks)

  1. Increasing temperature shifts the equilibrium position in the direction of the endothermic reaction
  2. The backward/reverse reaction is endothermic (because the forward reaction is exothermic)
  3. Therefore the equilibrium position shifts to the left (towards the reactants)
  4. The concentration of HI decreases

Key terms in this question

dynamic equilibrium

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