How can Le Chatelier's principle help maximize the yield of reactions?

May 25, 2014

Le Chatelier's Principle can be used in certain reactions to help maximize the yield of a given product. I'm going to use the Haber Process as an example of this.

The Haber Process is a process for making ammonia from elemental nitrogen and hydrogen using an iron catalyst. The reaction is shown below:

${N}_{2} \left(g\right) + 3 {H}_{2} \left(g\right) \to 2 N {H}_{3} \left(g\right) + 92.4 k J$

Now if my goal is to make as much $N {H}_{3}$ as possible, I can use Le Chatelier's principle to shift my equilibrium towards the product side as much as possible. To do this let's look at two stresses we can use temperature and pressure.

Since the forward reaction is exothermic, I can shift my equilibrium to the product side by cooling the reaction environment down,

In addition, because there are 4 molecules of gas reactants on the forward reaction, and only 2 molecules of gas reactants in the reverse reaction; increasing the pressure will cause the forward reaction to proceed faster and shift the equilibrium to the right, again causing more $N {H}_{3}$ to be produced.

We could also look at devising a way to remove the $N {H}_{3}$ from the reaction chamber, which will again cause a equilibrium shift to the right.

All these examples are ways that Le Chatelier's Principle can be used to maximize the amount of product obtained in a chemical reaction. This process also works on multi-step reactions as well.