Equilibria In Industrial Processes

Haber & Contact Processes

Equilibrium reactions are involved in some stages of large-scale production of certain chemicals
An understanding of equilibrium and Le Chatelier’s principle is therefore very important in the chemical industry

N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH_r = -92 kJ mol^-1

An increase in pressure will result in the equilibrium shifting in the direction of the fewest molecules of gas formed to reduce the pressure
In this case, the equilibrium shifts towards the right so the yield of ammonia increases
An increase in pressure will cause the particles to be closer together and therefore increasing the number of successful collisions leading to an increased reaction rate
Very high pressures are expensive to produce therefore a compromise pressure of 200 atm is chosen

To get the maximum yield of ammonia the position of equilibrium should be shifted as far as possible to the right as possible
Since the Haber process is an exothermic reaction, according to Le Chatelier’s principle the equilibrium will shift to the right if the temperature is lowered
A decrease in temperature will decrease the energy of the surroundings so the reaction will go in the direction in which energy is released to counteract this
Since the reaction is exothermic, the equilibrium shifts to the right
However, at a low temperature the gases won’t have enough kinetic energy to collide and react and therefore equilibrium would not be reached therefore compromise temperature of 400-450 ^oC is used in the Haber process
A heat exchanger warms the incoming gas mixture to give molecules more kinetic energy such that the gas molecules collide more frequently increasing the likelihood of a reaction

Removing ammonia by condensing it to a liquid causes the equilibrium position to shift to the right to replace the ammonia causing more ammonia to be formed from hydrogen and nitrogen
The removed ammonia is stored at very low temperatures and there is no catalyst present with the stored ammonia so the decomposition reaction of ammonia to decompose back into hydrogen and nitrogen will be too slow

In the absence of a catalyst the reaction is so slow that hardly anything happens in a reasonable time!
Adding an iron catalyst speeds up the rate of reaction

2SO₂(g) + O₂(g) ⇌ 2SO₃(g) ΔH_r = -197 kJ mol^-1

An increase in pressure will result in the equilibrium shifting in the direction of the fewest molecules of gas formed to reduce the pressure
In this case, the equilibrium shifts towards the right so the yield of sulfuric acid increases
In practice, the reaction is carried out at only 1 atm
This is because K_p for this reaction is already very high meaning that the position of the equilibrium is already far over to the right
Higher pressures than 1 atm will be unnecessary and expensive

The same principle applies to increasing the temperature in the Contact process as in the Haber process
A compromise temperature of 450 ^oC is used

The Contact process uses vanadium(V) oxide as a catalyst to increase the rate of reaction