Reviewed CSCA Chemistry question · Hard

In the industrial synthesis of ammonia, N2(g) + 3H2(g) ⇌ 2NH3(g), the reaction is exothermic. The process uses an iron catalyst and a temperature of about 450°C. Which of the following best explains why a temperature higher than the optimum (e.g., 600°C) would be undesirable?

The catalyst becomes inactive at higher temperatures.
The equilibrium yield of ammonia would decrease because the reaction is exothermic.
The rate of the forward reaction would become too slow.
The pressure inside the reactor would become too high.

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Correct answer

B. The equilibrium yield of ammonia would decrease because the reaction is exothermic.

Principle or equation

For an exothermic reaction, increasing temperature shifts the equilibrium to the left (towards reactants), decreasing the yield of product, according to Le Chatelier's principle.

Why this answer is correct

The forward reaction is exothermic, so heat is a product. Increasing temperature adds heat, shifting equilibrium to the left, reducing the equilibrium concentration of ammonia. Although higher temperature increases reaction rate, the loss in yield makes it undesirable. The catalyst remains active, the rate would increase (not slow), and pressure is controlled separately.

Example

In the Haber process, raising the temperature from 450°C to 600°C would decrease the equilibrium conversion of N2 and H2 to NH3 because the equilibrium constant for this exothermic reaction decreases with temperature.

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