Industrial Applications

Industrial Applications

In industrial processes, it is important to get the product as quickly and as efficiently as possible. The less expensive the process the better.

The Haber process is a good example of an industrial process which uses the equilibrium principles that have been discussed. The equation for the process is as follows:

\(\text{N}_{2}(\text{g}) + 3\text{H}_{2}(\text{g})\) \(\leftrightharpoons\) \(2\text{NH}_{3}(\text{g})\) + energy

Since the forward reaction is exothermic, to produce a lot of product and favour the forward reaction the system needs to be colder. However, cooling a system slows down all chemical reactions and so the system can't be too cold. This process is carried out at a much higher temperature to ensure the speed of production.

Because high temperature favours the reverse reaction, the \(\text{NH}_{3}\) product is actually removed as it is made (product concentration decreased) to prevent ammonia being used in the reverse reaction. The decrease of product concentration favours the forward reaction.

High pressure is also used to ensure faster reaction time and to favour the production of \(\text{NH}_{3}\). The forward reaction is favoured by higher pressures because there are \(\text{2}\) gas molecules of product for every \(\text{4}\) gas molecules of reactant.

We will look at the Haber process and other industrial applications some more in another tutorial.