Le Chatelier's Principle

What is Le Chatelier's Principle?

Le Chatelier's Principle states that when a system at equilibrium experiences a change (or "stress") in concentration, temperature, or pressure, the system will adjust to counteract that change and restore equilibrium.

This principle is particularly useful in industrial processes to maximize the production of desired products.

Types of Stress Affecting Equilibrium

Le Chatelier's principle helps predict how a change in conditions will affect the position of equilibrium.

Concentration

• Increasing the concentration of a reactant will shift the equilibrium towards the products, as the system tries to use up the added reactant.
• Increasing the concentration of a product will shift the equilibrium towards the reactants, as the system tries to reduce the amount of product.

Pressure (only affects gases)

• Increasing the pressure favours the side of the reaction with fewer gas molecules.
• Decreasing the pressure favours the side with more gas molecules. Note: If there are equal numbers of gas molecules on both sides of the reaction, pressure changes will have no effect.

Temperature

• Increasing the temperature favours the endothermic direction (absorbing heat).
• Decreasing the temperature favours the exothermic direction (releasing heat).

Catalysts

Catalysts do not affect the position of equilibrium. They only increase the rate at which equilibrium is reached by speeding up both the forward and reverse reactions.

Industrial Applications of Le Chatelier's Principle

Le Chatelier's principle is applied in several key industrial processes to optimize yield while considering practical limitations.

Haber Process (Ammonia Production)

The Haber process synthesizes ammonia from nitrogen and hydrogen gases:

$$N_2(g) + 3H_2(g) ⇌ 2NH_3(g)$$ $$\quad (\Delta H = -92 \, \text{kJ/mol})$$

Conditions:

• Pressure: High pressure favours ammonia production (since fewer gas molecules are produced). Industrial pressures of around 200 atm are used, but extremely high pressures are avoided for safety and cost reasons.
• Temperature: A lower temperature would favour the production of ammonia, but this would slow down the reaction rate. Industrial temperatures of 400-500°C are used as a compromise.
• Catalyst: An iron catalyst is used to speed up the reaction without affecting the equilibrium position.

Contact Process (Sulfuric Acid Production)

The Contact process produces sulfur trioxide ($SO_3$) from sulfur dioxide and oxygen:

$$2SO_2(g) + O_2(g) ⇌ 2SO_3(g)$$ $$\quad (\Delta H = -196 \, \text{kJ/mol})$$

Conditions:

• Pressure: Increasing pressure shifts the equilibrium to the right (producing more sulfur trioxide). However, since there are only slight differences in gas molecules, a pressure just above atmospheric pressure is sufficient.
• Temperature: A lower temperature favours $SO_3$ production, but, as with the Haber process, a moderate temperature (450°C) is used to balance the yield and reaction rate.
• Catalyst: Vanadium($V$) oxide ($V_2O_5$) is used to speed up the reaction.