If there is no change in volume in a gaseous reaction, the pressure will

Le Chatelier's principle states that if a system at equilibrium is subjected to a change in conditions, the equilibrium shifts in the direction that tends to counteract that change. For pressure changes, the key factor is the difference in the total number of moles of gas on each side of the equation.

If there is no change in volume during a gaseous reaction, this means the total number of moles of gaseous products equals the total number of moles of gaseous reactants. In such a case, increasing or decreasing the pressure gives the system no direction in which to shift, because neither the forward nor the backward reaction would reduce the total number of gas molecules.

Therefore, a change in pressure will have no effect on the equilibrium position when the reaction involves equal moles of gas on both sides.

For example, in the reaction:

\[ \text{H}_2(g) + \text{I}_2(g) \rightleftharpoons 2\text{HI}(g) \]

there are 2 moles of gas on each side, so pressure changes do not shift the equilibrium.

Pressure only affects equilibrium when there is an unequal number of moles of gas on either side. In those cases, increasing pressure favours the side with fewer moles of gas, and decreasing pressure favours the side with more moles.