The effect of pressure on the equilibrium position of a chemical reaction can be explained using 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.When the pressure of a system at equilibrium is increased, the equilibrium will shift in the direction that reduces the pressure. This typically occurs by favoring the side of the reaction with fewer moles of gas. Conversely, when the pressure is decreased, the equilibrium will shift in the direction that increases the pressure, favoring the side with more moles of gas.It is important to note that pressure changes only affect reactions involving gases, and the effect is more significant when there is a difference in the number of moles of gas between the reactants and products.Example reaction: N2 g + 3H2 g 2NH3 g In this reaction, 1 mole of nitrogen gas reacts with 3 moles of hydrogen gas to form 2 moles of ammonia gas. The total number of moles of gas on the reactant side is 4, while on the product side, it is 2.1. When the pressure is increased:The equilibrium will shift towards the side with fewer moles of gas to counteract the increase in pressure. In this case, it will shift towards the formation of ammonia NH3 , favoring the product side.2. When the pressure is decreased:The equilibrium will shift towards the side with more moles of gas to counteract the decrease in pressure. In this case, it will shift towards the formation of nitrogen N2 and hydrogen H2 , favoring the reactant side.In summary, increasing the pressure will favor the side of the reaction with fewer moles of gas, while decreasing the pressure will favor the side with more moles of gas.