CCEA GCSE Chemistry 2.4 Equilibrium Exam Questions

About CCEA GCSE Chemistry 2.4 Equilibrium

Equilibrium is a Unit 2 topic in CCEA GCSE Chemistry that explains reactions that can go in both directions. You need to describe reversible reactions, define dynamic equilibrium and use Le Chatelier's Principle to predict how changes in temperature, pressure and concentration move the position of equilibrium. The topic is applied to the Haber process for making ammonia, where the conditions chosen are a compromise between yield, rate and cost. It links to rates of reaction, energy changes and the properties of ammonia in gas chemistry. Examiners commonly ask for explanations of why a particular temperature or pressure is used and for predictions of equilibrium shifts.

What you need to know

• Reversible reactions Some reactions can go both forwards and backwards, shown by ⇌. If the forward reaction is exothermic, the reverse is endothermic by the same amount. An example is heating hydrated copper(II) sulfate (blue) to form anhydrous copper(II) sulfate (white); adding water reverses it.
• Dynamic equilibrium In a closed system, equilibrium is reached when the forward and reverse reactions happen at the same rate, so the concentrations of reactants and products stay constant. Both reactions are still happening.
• Le Chatelier's Principle If a condition of a system at equilibrium is changed, the position of equilibrium moves to oppose the change. Raising temperature favours the endothermic direction; raising pressure favours the side with fewer gas molecules; adding a reactant moves equilibrium to the right.
• The Haber process N2(g) + 3H2(g) ⇌ 2NH3(g); the forward reaction is exothermic. Nitrogen comes from the air and hydrogen from natural gas. Typical conditions are about 450 °C, 200 atmospheres and an iron catalyst. Unreacted gases are recycled and ammonia is removed by cooling and condensing it.
• Compromise conditions A lower temperature would give a higher yield but a slower rate, so a moderate temperature is used. Higher pressure increases yield and rate but is expensive and dangerous. The catalyst increases rate but does not change the position of equilibrium.

Common mistakes in 2.4 Equilibrium

• Saying that at equilibrium the reaction has stopped. Both reactions continue at equal rates.
• Stating that at equilibrium the concentrations of reactants and products are equal. They are constant, not necessarily equal.
• Claiming a catalyst increases the yield of ammonia. It only increases the rate at which equilibrium is reached.
• Explaining high pressure without counting gas molecules. There are 4 molecules on the left and 2 on the right, so high pressure favours ammonia.
• Forgetting that equilibrium can only be reached in a closed system.

Try three 2.4 Equilibrium exam-style questions