Predicting Equilibrium Shifts

This lesson covers Higher Tier content.

A reversible reaction can make products and turn those products back into reactants. In a closed system, it can reach dynamic equilibrium: the forward and reverse reactions continue at equal rates, so the concentrations remain constant. They do not have to be equal.

The equilibrium position describes the relative amounts of reactants and products present. Changing the conditions can change this balance. A shift to the right favours the forward reaction and produces more product; a shift to the left favours the reverse reaction and produces more reactant.

Le Chatelier’s principle predicts these changes: when conditions change in a system at equilibrium, the equilibrium shifts in the direction that opposes the change. The response reduces the effect of the change; it does not necessarily undo it completely. After the shift, a new equilibrium is established, with the forward and reverse rates equal again.

Concentration: using up what is added

Concentration describes how much of a substance is present in a given volume. Adding a reactant increases its concentration. The equilibrium responds by favouring the reaction that uses up that reactant, so it shifts towards the products.

Removing a product also favours the forward reaction: more product forms to replace some of what was removed. This is useful when the aim is to produce as much product as possible.

The opposite changes favour the reverse reaction:

For example, consider this gaseous equilibrium:

Adding nitrogen dioxide at constant temperature and volume favours the forward reaction, which consumes nitrogen dioxide and forms more dinitrogen tetroxide. Removing some dinitrogen tetroxide also shifts the equilibrium right, forming more of it.

Temperature: favouring energy absorption or release

In a reversible reaction, one direction is exothermic, transferring energy to the surroundings, and the other is endothermic, taking in energy from the surroundings.

Increasing the temperature favours the endothermic direction, which absorbs energy and opposes the heating. Decreasing the temperature favours the exothermic direction, which releases energy and opposes the cooling.

This means heating does not always produce more product. The answer depends on the energy change of the forward reaction.

For example, iodine monochloride reacts reversibly with chlorine to form iodine trichloride:

The forward reaction is exothermic, so the reverse reaction is endothermic. Heating favours the reverse reaction and shifts the equilibrium left. More dark-brown iodine monochloride forms, so the mixture becomes browner. Cooling favours the forward reaction and increases the amount of yellow iodine trichloride.

Pressure: counting gaseous particles

For a gaseous equilibrium, pressure can be increased by compressing the mixture into a smaller volume at constant temperature. The equilibrium opposes this by shifting towards the side with fewer gas molecules. Decreasing pressure favours the side with more gas molecules.

Use the coefficients in the balanced equation to compare the amounts of gas. State symbols matter: solids and liquids are not included in this count.

For the nitrogen dioxide equilibrium,

the left side has two molecules of gas for every one on the right. Increasing pressure therefore shifts the equilibrium right, giving a greater proportion of dinitrogen tetroxide. Decreasing pressure shifts it left, favouring nitrogen dioxide.

If both sides have equal numbers of gas molecules, changing pressure does not shift the equilibrium position. Pressure changes also do not shift an equilibrium involving only solids and liquids.

Choosing conditions for the greatest product yield

To maximise the amount of a particular product, choose conditions that shift the equilibrium towards it. First identify the product side, then check the energy change and count the gaseous particles.

For the iodine trichloride reaction above, a lower temperature favours the exothermic forward reaction. A higher pressure also favours the product side: there is one gaseous reactant, chlorine, but no gaseous product. Both changes therefore increase the equilibrium yield of iodine trichloride.

Keep yield separate from speed. Conditions that favour a large amount of product at equilibrium are not necessarily the conditions that produce it most quickly.