Changing the Conditions of an Equilibrium Reaction

This lesson covers: 

1. What Le Chatelier’s principle is
2. How to use Le Chatelier’s principle
3. Examples of using Le Chatelier’s principle in industry

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 [H₂] is increased, the equilibrium position will shift right towards NH₃ to use up the extra H₂.
• If [NH₃] is decreased, the equilibrium position will shift right to produce more NH₃.

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 ethanol: balancing yield, rate and cost

Ethanol is industrially produced from the exothermic reaction between ethene gas and steam:

C₂H_4(g) + H₂O_(g) ⇌ C₂H₅OH_(g),  ΔH = -46 kJ mol^-1

The conditions used are:

• Temperature: 300°C
• Pressure: 60-70 atm
• Phosphoric(V) acid catalyst

Effects of pressure:

• A high pressure (>70 atm) promotes more product formation but requires costly equipment.
• A low pressure (<60 atm) is more economical but decreases the yield of ethanol.

Effects of temperature:

• A high temperature (>300°C) increases reaction rate but lowers the yield for exothermic reactions.
• A low temperature (<300°C) increases yield but slows down the reaction rate.

Thus, the industrial conditions represent a balance between:

• Maximising the yield of ethanol
• Ensuring a rapid reaction rate
• Keeping production costs low

A temperature of 300°C and a pressure of 60-70 atm are selected as the best compromise among these considerations.