How does pressure influence the equilibrium in gaseous reactions?

Pressure influences the equilibrium in gaseous reactions by shifting it towards the side with fewer gas molecules.

In more detail, the principle that explains this behaviour is known as Le Chatelier's Principle. This principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to counteract the change. In the context of gaseous reactions, an increase in pressure will shift the equilibrium towards the side of the reaction with fewer gas molecules. This is because fewer gas molecules occupy a smaller volume, thus reducing the pressure.

For example, consider the equilibrium reaction between nitrogen and hydrogen to form ammonia:

N2(g) + 3H2(g) ⇌ 2NH3(g)

In this reaction, there are four moles of gas on the left-hand side (one mole of nitrogen and three moles of hydrogen) and two moles of gas on the right-hand side (two moles of ammonia). If the pressure is increased, the equilibrium will shift to the right, towards the side with fewer gas molecules, to reduce the pressure.

Conversely, if the pressure is decreased, the equilibrium will shift towards the side with more gas molecules, to increase the pressure. In the above example, a decrease in pressure would shift the equilibrium to the left.

It's important to note that this only applies to reactions where the number of moles of gas is different on each side of the equation. If the number of moles of gas is the same on both sides, a change in pressure will not affect the position of equilibrium.

Understanding how pressure affects equilibrium in gaseous reactions is crucial in many industrial processes. For instance, the Haber process for the production of ammonia is carried out at high pressure to shift the equilibrium towards the right, favouring the production of ammonia.