In Gaseous equilibria we saw that the pressure of a gas is proportional to its concentration. So we can use Le Chatelier's Principle to predict the effect on the equilibrium position of changing the total pressure of a gaseous equilibrum.

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Think about the equilibrium system: 2A(g) + B(g) _∏ C(g) + D(g)

If we increase the total pressure, Le Chatelier's Principle predicts that the equilibrium position will move in a way that counteracts this change, ie moves to reduce the pressure.

The pressure is caused by molecules colliding with the walls of the container. The fewer molecules, the lower the pressure. Because the right-hand side of the system has fewer molecules than the left-hand side a movement from left to right will result in fewer molecules and therefore a lower pressure.

What will happen to the equilibrium position if we decrease the total pressure?

The equilibrium position will move in a way that increases the pressure. C and D will react to give A and B because this produces more molecules, which in turn increases the pressure.

Think about the equilibrium system: A(g) + B(g) _∏ C(g) + D(g). How will the system react if the total pressure is increased?

There will be no change because the numbers of molecules on each side is the same. The forward reaction or the reverse reaction will not change the number of molecules so there will be no change in equilibrium position.