One of the most famous chemical reactions and today’s accepted method for mass industrial production of ammonia, we have the Haber process. It is an important industrial application as it produces large amount of NH3 that is used in further production of many products. For this application, we shall study the temperature as a key factor in this process.
The equation for the Haber process is
N2(g) + 3H2(g) ⇌ 2NH3(g) ΔH < 0
To increase the yield of NH3, the equilibrium position needs to shift to the right. Since the reaction is exothermic, conventionally we should decrease the system temperature to achieve this. But in reality, a whooping 450°C is used, which contradicts the concepts that we have studied!
The reason for this is because when the temperature is too low, the rate of reaction will be severely impaired, which greatly decreasing the yield. Thus, there was no choice but to increase the temperature. To compensate, the pressure of the system was also increased to force the equilibrium position to the right.
So in short, Le Chatelier’s Principle holds, but in real life, we also have to consider industrial limitations.