Final answer:
Changing the volume or pressure of a system at equilibrium affects the position of the equilibrium depending on the number of moles of gases present. An increase in pressure favors the side with fewer moles of gas, while a decrease in pressure favors the side with more moles.
Step-by-step explanation:
When a system at equilibrium experiences a change in volume or pressure, its response is explained by Le Chatelier's principle. If you increase the pressure of the system (perhaps by reducing its volume), the system will adjust to reduce this pressure. This is achieved by shifting the equilibrium toward the side with the fewest moles of gas. For instance, if the reactants' side of an equilibrium system has 1 mole of gas and the products' side has 2 moles, an increase in pressure will shift the equilibrium to the left, favoring the reverse reaction. Conversely, decreasing pressure (by increasing the volume) causes the equilibrium to shift to the side with the most moles of gas.
In a reaction where there are different numbers of moles on each side, such as three moles of reactants and two moles of products, decreasing the volume will result in an increased pressure, and the equilibrium will shift toward the products side, which has fewer moles.
These shifts apply only if there are gases involved and the number of moles of gases on each side is not equal. If the number of moles of reactant and product gases is the same, then a change in volume or pressure will not result in a shift in equilibrium. Also, adding a catalyst or altering the pressure or volume in a system with no gases will not affect the position of the equilibrium.