What Le Chatelier’s principle is

Le Chatelier’s principle is used to predict the effect of changes in conditions on the position of equilibrium of a reversible reaction.

The position of equilibrium refers to the relative amounts of products and reactants present at equilibrium.

Le Chatelier's principle states that if a chemical system at equilibrium is subjected to a change in concentration, pressure or temperature, the position of equilibrium will shift to counteract the change.

In other words, the equilibrium tries to reduce the stress and resist the changes.

• Shifting 'to the right' means more products form.
• Shifting 'to the left' means more reactants form.

Using Le Chatelier’s principle

We can use the following rules to predict how the position of equilibrium will change:

Concentration changes

• Increasing the concentration of a reactant - Equilibrium shifts towards the products (to the right) to use up the extra reactant.
• Increasing the concentration of a product -Equilibrium shifts towards the reactants (to the left) to use up the extra product.
• Decreasing the concentrations has the opposite effects.

For example, in the reaction:

N_2(g)_ + 3H_2(g)_ ⇌ 2NH_3(g)_

• If [H_2_] is increased, the equilibrium position will shift right towards NH_3_ to use up the extra H_2_.
• If [NH_3_] is decreased, the equilibrium position will shift right to produce more NH_3_.

Pressure changes

Pressure changes only apply to gaseous equilibria.

• Increasing the pressure* -* Equilibrium shifts towards the side with fewer gas molecules to lower the pressure.
• Decreasing the pressure* -* Equilibrium shifts towards the side with more gas molecules to raise the pressure.

For example, in the reaction:

2SO_2(g)_ + O_2(g)_ ⇌ 2SO_3(g)_

The left side has 3 gas molecules, and the right side has 2. So, raising the pressure moves the equilibrium to the right, where there are fewer gas molecules.

Temperature changes

• Increasing the temperature - Equilibrium shifts in the endothermic direction to absorb the excess heat.
• Decreasing the temperature* -* Equilibrium shifts in the exothermic direction to release heat.

For example, in the reaction:

N_2(g)_ + O_2(g)_ ⇌ 2NO_(g)_, ΔH = +180 kJ mol^-1^

• Heating this endothermic reaction shifts the equilibrium to the right to absorb the additional heat.
• Cooling this reaction shifts the equilibrium to the left, releasing more heat.

Effect of catalysts

• Adding a catalyst does not affect the position of equilibrium.
• A catalyst speeds up both the forward and reverse reactions equally, so the equilibrium constant remains unchanged.
• Catalysts lower the activation energy for the reaction, allowing equilibrium to be reached faster, but they do not alter the final equilibrium composition.

Industrial production of ammonia: balancing yield, rate and cost

Ammonia is industrially produced via the Haber process:

N_2(g)_ + 3H_2(g)_ ⇌ 2NH_3(g)_, ΔH = $-$92 kJ mol^-1^

The conditions used are:

• Temperature: 450°C
• Pressure: 200 atm
• Iron catalyst

Effects of pressure:

• A high pressure (>200 atm) increases the yield of NH₃ but requires expensive high-pressure equipment.
• A low pressure (<200 atm) is more economical but reduces the ammonia yield.

Effects of temperature:

• A high temperature (>450°C) speeds up the reaction rate but decreases the yield of NH₃ for this exothermic reaction.
• A low temperature (<450°C) favours a higher yield but results in a slower reaction rate.

Thus, the industrial conditions used in the Haber process balance:

• Maximising the yield of ammonia.
• Achieving a sufficiently fast reaction rate.
• Minimising the production costs. A temperature of 450°C and a pressure of 200 atm are chosen as the optimal compromise considering these factors. The iron catalyst increases the rate of reaction, allowing the use of a lower temperature.

Industrial production of sulfuric acid: balancing yield, rate and cost

Sulfur trioxide is industrially produced via the Contact process:

2SO_2(g)_ + O_2(g)_ ⇌ 2SO_3(g)_, ΔH = $-$196 kJ mol^-1^

The conditions used are:

• Temperature: 450°C
• Pressure: 2 atm
• Vanadium(V) oxide catalyst

Effects of pressure:

• A high pressure (>2 atm) increases the yield of SO₃ but is not used due to the costs of high-pressure equipment and safety concerns, as SO₃ is a highly acidic gas.
• A low pressure (<2 atm) is more economical and maintains an acceptable yield.

Effects of temperature:

• A high temperature (>450°C) increases the reaction rate but reduces the yield of SO₃ for this exothermic reaction.
• A low temperature (<450°C) favours a higher yield but results in a slower reaction rate.

Thus, the industrial conditions used in the Contact process balance:

• Maximising the yield of SO₃.
• Achieving a sufficiently fast reaction rate.
• Minimising the production costs. A temperature of 450°C and a pressure of 2 atm are chosen as the optimal compromise considering these factors. The vanadium(V) oxide catalyst increases the rate of reaction, allowing for a lower temperature to be used.