When ammonia is made industrially, the mixture of ${}^{{\mathbf{N}}_{\mathbf{2}}}$, ${}^{{\mathbf{H}}_{\mathbf{2}}}$ and${}^{{\mathbf{NH}}_{\mathbf{3}}}$ that emerges from the reaction chamber is far from equilibrium. Why does the plant supervisor use reaction conditions that produce less than the maximum yield of ammonia?

Haber's process is used in theindustrial production of ammonia. Nitrogen from gas is combined with hydrogen from natural gas (methane) in a 1:3 ratio to produce ammonia. The reaction is exothermic and reversible.

The best conditions for producing ammonia are a pressure of ${}^{\mathbf{200}\mathbf{\times }{\mathbf{10}}^{\mathbf{5}}}$Pa, a temperature of 4700 K, and an iron oxide catalyst containing details of ${}^{{\mathbf{AI}}_{\mathbf{2}}{\mathbf{O}}_{\mathbf{3}}}$and${}^{{\mathbf{K}}_{\mathbf{2}}\mathbf{O}}$ .

The production of ammonia from ${}^{{\mathbf{H}}_{\mathbf{2}}}$and ${}^{{\mathbf{N}}_{\mathbf{2}}}$is given by:

${\mathbf{N}}_{\mathbf{2}\left(g\right)}\mathbf{+}\mathbf{3}{\mathbf{H}}_{\mathbf{2}\left(g\right)}\mathbf{}\mathbf{\rightleftharpoons }\mathbf{2}{\mathbf{NH}}_{\mathbf{3}\left(g\right)}$

The reaction is exothermic. Thus, low temperature is required to produce more amount of ${}^{{\mathbf{NH}}_{\mathbf{3}}}$by Le-Chatelier’s principle. But at low temperatures, the reaction rate is very slow. Hence, the optimum temperature is used to make the reaction rate fast. Hence, the amount or yield of ammonia is less than predicted.