Increasing the pressure in a chemical reaction involving gases affects the equilibrium position according to Le Chatelier's principle, which states that if a system at equilibrium is subjected to a change in pressure, temperature, or concentration of reactants or products, the system will adjust its equilibrium position to counteract the change.In the case of a gas-phase reaction, increasing the pressure will shift the equilibrium position towards the side with fewer moles of gas, as this will help to reduce the overall pressure in the system. Conversely, decreasing the pressure will shift the equilibrium position towards the side with more moles of gas.An example of a gas-phase reaction is the Haber process for the synthesis of ammonia:N2 g + 3H2 g 2NH3 g In this reaction, there are 4 moles of gas on the reactants side 1 mole of N2 and 3 moles of H2 and 2 moles of gas on the products side 2 moles of NH3 . If the pressure is increased, the equilibrium position will shift towards the side with fewer moles of gas, which is the products side formation of ammonia . This will result in an increase in the concentration of ammonia and a decrease in the concentrations of nitrogen and hydrogen.It is important to note that the equilibrium constant K itself is not affected by changes in pressure, as it is only dependent on temperature. However, the equilibrium concentrations of the reactants and products will change as the system adjusts to the new pressure, which will affect the reaction quotient Q . If the pressure is increased, the reaction quotient Q will initially be less than the equilibrium constant K , and the reaction will proceed in the forward direction towards the formation of ammonia until a new equilibrium is established, and Q = K again.