Consider the reaction represented by the following equation: N2(g) + 3H2(g) ⇌ 2NH3(g); ∆Hϴ = -xkJ mol^ -1…

Analyze the question: This is a Haber process equilibrium. We need the condition that shifts the equilibrium to the right and so increases the yield of NH₃.

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

Information from the equation:

• Reactants: 1 mol N₂ + 3 mol H₂ = 4 moles of gas
• Products: 2 moles of gas of NH₃
• ΔH is negative (-x kJ/mol), so the forward reaction is exothermic.

Step-by-step evaluation of each option:

• A. Iron catalyst: A catalyst speeds up both the forward and backward reactions equally. It helps equilibrium be reached faster but does not change the yield. ✗
• B. Increase in temperature: The forward reaction is exothermic, so higher temperature shifts the equilibrium to the left (the endothermic direction). This reduces NH₃ yield. ✗
• C. Increase in volume: This lowers the pressure, so the equilibrium shifts toward the side with more gas moles (the reactants). This reduces NH₃ yield. ✗
• D. Increase in total pressure: Higher pressure favors the side with fewer gas moles. The product side has 2 moles versus 4 on the reactant side, so the equilibrium shifts right and more NH₃ forms. ✓

Answer: D

Tip: For gaseous equilibria, compare the moles of gas on each side. High pressure favors the side with fewer moles. For exothermic reactions, low temperature favors products. (In industry, a compromise temperature of about 450°C is used with an iron catalyst to get a reasonable rate.)