Consider the following reaction equation: X(g) + Y(g) ⇌ XY(g); ∆H = + 220KJ $mol^{-1}$ If the temperature of the system is increased, the

Analyze the question: The reaction X(g) + Y(g) ⇌ XY(g) has ΔH = +220 kJ/mol. We are asked what happens when the temperature is increased.

Concept: Le Chatelier's Principle When a system at equilibrium is disturbed, it shifts to oppose the change.

• Increasing temperature favours the endothermic direction (the one that absorbs heat).
• Decreasing temperature favours the exothermic direction (the one that releases heat).

Solution:

1. ΔH is positive (+220 kJ/mol), so the forward reaction (X + Y → XY) is endothermic. It absorbs heat.
2. The backward reaction (XY → X + Y) is therefore exothermic.
3. When the temperature is raised, the system tries to absorb the extra heat by favouring the endothermic direction, which is the forward reaction.
4. The result is more XY formed, so the equilibrium shifts to the right and the equilibrium constant increases.

Checking the other options:

• A: The backward reaction is favoured only if the temperature is lowered, so this is wrong.
• C: A reaction does not stop when the temperature rises. It only shifts position (and the rate actually increases).
• D: The system is already at equilibrium. The question asks about the shift, and the shift is toward the forward reaction, so D does not answer it.

Answer: The forward reaction would be favoured, which is option B.

Exam tip: Remember "Heat is a reactant in endothermic reactions." Adding heat is like adding a reactant, so the equilibrium shifts forward. For exothermic reactions (ΔH negative), heat is a product, so raising the temperature shifts it backward.