Consider the following reaction equation N$_{2}$(g) + 3H$_{2(g)}$  ⇌ 2NH3(g) ∆H = -92kJ Increasing the temperature of the reaction would

Analyze the question: The Haber process reaction is exothermic (ΔH = -92 kJ). We are asked what happens to the equilibrium when the temperature is increased.

Concept: Le Chatelier's Principle When a system at equilibrium is disturbed, it shifts in the direction that opposes the change.

• Increasing temperature means adding heat. The system responds by absorbing that heat, so it favors the endothermic direction.
• Decreasing temperature means removing heat. The system favors the exothermic direction.

Solution:

1. Forward reaction: N₂ + 3H₂ → 2NH₃, ΔH = -92 kJ, which is exothermic (releases heat).
2. Backward reaction: 2NH₃ → N₂ + 3H₂, which is endothermic (absorbs heat).
3. Raising the temperature makes the system absorb heat, so the equilibrium shifts to the left (backward, endothermic direction).
4. Results of shifting left: • NH₃ is decomposed, so the yield of ammonia decreases. • N₂ and H₂ increase, so the amount of hydrogen increases.

Checking the options:

• A. Shift to the right: wrong, it shifts left.
• B. Increase the yield of ammonia: wrong.
• C. Decrease the amount of hydrogen: wrong, hydrogen increases.
• D. Decrease the yield of ammonia: correct.

Tip: For exothermic reactions, think of heat as a product: "Heat + reactants ⇌ products + heat". Adding heat is like adding a product, so it pushes the reaction backward. In industry, the Haber process uses a compromise temperature of about 450°C, which is high enough for a reasonable rate but not so high that it ruins the yield.