Increasing pressure can enhance the spontaneity of certain reactions by shifting the equilibrium toward the side with fewer gas molecules.

This phenomenon is explained by Le Chatelier’s Principle, which states that when a dynamic equilibrium is disturbed by changes in conditions, the equilibrium position shifts to counteract the disturbance. Specifically, in the case of pressure, an increase in pressure will prompt the equilibrium to shift in a direction that reduces pressure. This occurs by favoring the side of the reaction that contains fewer gas molecules, as fewer gas molecules result in lower pressure.

For instance, consider the reaction where nitrogen and hydrogen combine to form ammonia:

In this reaction, there are four moles of gas on the reactant side (one mole of nitrogen and three moles of hydrogen) and only two moles of gas on the product side (two moles of ammonia). When the pressure is increased, the equilibrium will shift to the right, favoring the formation of ammonia, which has fewer gas molecules. This shift makes the formation of ammonia more spontaneous under higher pressure conditions.

It is crucial to recognize that this principle applies only to reactions where there is a change in the number of moles of gas. If the number of gas moles is the same on both sides of the reaction, an increase in pressure will not influence the position of equilibrium.

Moreover, the impact of pressure on the spontaneity of reactions is particularly pronounced for gaseous reactions, as gases are significantly more compressible than solids or liquids. Therefore, changes in pressure can substantially affect the volume of a gas and consequently the position of equilibrium.

In summary, increasing pressure can promote the spontaneity of reactions by shifting the equilibrium toward the side with fewer gas molecules, in accordance with Le Chatelier’s Principle. However, this effect is most notable in reactions involving gases where there is a change in the number of moles of gas.