Changing the Position of Equilibrium

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:

N2(g) + 3H2(g) ⇌ 2NH3(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:

2SO2(g) + O2(g) ⇌ 2SO3(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:

N2(g) + O2(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:

N2(g) + 3H2(g) ⇌ 2NH3(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 increasing the rate of reaction, allowing a lower temperature to be used.

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

Sulfur trioxide is industrially produced via the Contact process:

2SO2(g) + O2(g) ⇌ 2SO3(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 still provides 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.