Industrial Processes (CIE AS Chemistry)

• 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

Haber process

• The Haber process involves the synthesis of ammonia according to:

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

• Le Chatelier’s principle is used to get the best yield of ammonia

Maximising the ammonia yield

Pressure

• 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

Temperature

• 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

• 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

Catalysts

• 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

Contact process

• The Contact process involves the synthesis of sulfuric acid according to:

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

• Le Chatelier’s principle is used to get the best yield of sulfuric acid

Maximising the sulfuric acid yield

Pressure

• 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 Kp 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

Temperature

• 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

Removing sulfuric acid

• SO₃ is removed by absorbing it in 98% sulfuric acid
• The SO₃ reacts with the solution and more H₂SO₄ is formed

Catalysts

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