Applications of equilibrium principles in the chemical industry

1. Importance of Equilibrium in Industry

Many industrial processes involve reversible chemical reactions that reach equilibrium. The economic viability of these processes depends on:

✔ The yield (amount of product formed compared to reactants).

✔ The reaction rate (how quickly products form).

✔ The cost-effectiveness of production.

2. Key Industrial Applications of Equilibrium

Two major industrial processes that rely on equilibrium principles are:

2.1 The Haber Process (Ammonia Production)

• Reaction: $N_2(g) + 3H_2(g) ⇌ 2NH_3(g) \quad \Delta H < 0$

• Reactants: Nitrogen $(N_2)$ and Hydrogen $(H_2)$

• Products: Ammonia $(NH_3)$

• Reaction Conditions:

  • Temperature: ~450°C
  • Pressure: ~200 atm
  • Catalyst: Iron $(Fe)$

• Equilibrium Considerations:

  • Forward reaction is exothermic → Lower temperatures increase yield but slow the reaction.
  • High pressure favours product formation (fewer gas molecules).
  • Catalyst speeds up equilibrium attainment without shifting equilibrium position.

2.2 The Contact Process (Sulphuric Acid Production)

• Step 2 Reaction: $2SO_2(g) + O_2(g) ⇌ 2SO_3(g) \quad \Delta H < 0$

• Reactants: Sulphur dioxide $(SO_2)$ and Oxygen $(O_2)$

• Product: Sulphur trioxide $(SO_3)$

• Reaction Conditions:

  • Temperature: ~450°C
  • Pressure: ~1 atm
  • Catalyst: Vanadium pentoxide $(V_2O_5)$

• Equilibrium Considerations:

  • Forward reaction is exothermic → Lower temperatures increase yield but slow reaction rate.
  • Increasing pressure slightly favours product formation but is unnecessary due to high conversion at $1$ $atm.$

3. Optimising Industrial Processes Using Equilibrium

To maximise yield and efficiency, industries use:

✔ Lower temperatures (for exothermic reactions) while balancing reaction rate.

✔ Higher pressure (for reactions producing fewer gas molecules).

✔ Catalysts to speed up reaction rates without shifting equilibrium.

✔ Continuous removal of products to drive equilibrium forward.