What occurs to equilibrium when the pressure is increased in a gaseous reaction?

In a gaseous reaction, when the pressure is increased, the equilibrium will shift to the side with fewer moles of gas. This behavior is a consequence of Le Chatelier's principle, which states that if a system at equilibrium is subjected to a change in pressure, temperature, or concentration, the system will adjust in a way that counteracts the effect of that change.

Specifically, when pressure is increased, the system will strive to decrease the pressure by favoring the direction that produces fewer gas molecules, thereby reducing the overall volume of the gas present. If one side of the equilibrium reaction has more moles of gas than the other, the shift will be towards the side with fewer moles to relieve the pressure.

This can be illustrated with a typical reaction like:

[ aA(g) + bB(g) \rightleftharpoons cC(g) + dD(g) ]

If the left side (reactants) contains more moles of gas than the right side (products), an increase in pressure will shift the equilibrium to the right, favoring the formation of products which have fewer moles of gas.

Understanding this principle helps predict how changes in pressure affect equilibria in gaseous systems, ensuring a correct interpretation