How does changing pressure affect gaseous equilibria?

Changing pressure affects gaseous equilibria by shifting the equilibrium position towards the side with fewer gas molecules.

In more detail, the principle that explains this behaviour is 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 equilibria, if the pressure is increased, the equilibrium will shift to the side with fewer gas molecules to reduce the pressure. Conversely, if the pressure is decreased, the equilibrium will shift to the side with more gas molecules to increase the pressure.

For example, consider the equilibrium between nitrogen gas and hydrogen gas producing ammonia:

N2(g) + 3H2(g) ⇌ 2NH3(g)

In this reaction, there are four moles of gas on the left-hand side and two moles of gas on the right-hand side. If the pressure is increased, the equilibrium will shift to the right, towards the side with fewer gas molecules, to decrease the pressure. This results in more ammonia being produced. If the pressure is decreased, the equilibrium will shift to the left, towards the side with more gas molecules, to increase the pressure. This results in less ammonia being produced.

It's important to note that not all changes in pressure will cause a shift in the equilibrium position. If the number of gas molecules is the same on both sides of the equation, a change in pressure will not affect the position of equilibrium. For example, in the reaction:

2H2(g) + O2(g) ⇌ 2H2O(g)

There are three moles of gas on both sides of the equation, so a change in pressure will not shift the equilibrium position.

Remember, understanding how changing conditions affect equilibria is key to controlling chemical reactions in industrial processes.