Chemical equilibrium is very important in industrial processes because it helps maximize the yield of desired products. The Haber process (for making ammonia) is a classic example.
Here’s how equilibrium applies, in simple terms:
- The reaction is reversible:
- Nitrogen and hydrogen react to form ammonia, but ammonia can also break back into nitrogen and hydrogen.
- At equilibrium, the forward and backward reactions happen at the same rate.
- Controlling conditions:
- Pressure: High pressure favors ammonia formation because there are fewer gas molecules on the product side.
- Temperature: Moderate temperature is used because the reaction is exothermic — too high temperature reduces ammonia yield.
- Use of catalysts:
- Catalysts speed up the reaction, helping the system reach equilibrium faster, without changing the maximum yield.
- Continuous removal of product:
- Ammonia is removed as it forms, which shifts equilibrium forward, increasing overall production.
In short: Understanding and controlling equilibrium allows industries to optimize product yield efficiently.