The reaction between nitrogen and hydrogen gases to form ammonia gas is represented by the following balanced equation:N2 g + 3H2 g 2NH3 g This reaction is an example of a reversible reaction, meaning it can proceed in both the forward and reverse directions. The position of equilibrium in a reversible reaction is determined by the relative rates of the forward and reverse reactions, and it can be affected by changes in pressure, temperature, and concentration.According to Le Chatelier's principle, if a system at equilibrium is subjected to a change in pressure, temperature, or concentration, the system will adjust its position of equilibrium to counteract the change. In this case, we are interested in the effect of increasing pressure on the position of equilibrium.To analyze the effect of pressure on the equilibrium, we can use the reaction's stoichiometry. On the left side of the equation, there are 1 mole of N2 and 3 moles of H2, totaling 4 moles of gas. On the right side, there are 2 moles of NH3. When the pressure is increased, the system will try to counteract the change by shifting the position of equilibrium to the side with fewer moles of gas, which in this case is the side with ammonia NH3 .To further support this conclusion, we can use the equilibrium constant Kp expression for this reaction:Kp = [NH3]^2 / [N2] * [H2]^3 Where [NH3], [N2], and [H2] represent the equilibrium concentrations of ammonia, nitrogen, and hydrogen gases, respectively.When the pressure is increased, the concentrations of all gases in the reaction will increase. However, since there are fewer moles of gas on the right side of the equation, the increase in the concentration of ammonia will have a greater effect on the equilibrium constant than the increase in the concentrations of nitrogen and hydrogen. As a result, the value of Kp will increase, indicating that the position of equilibrium has shifted to the right, favoring the formation of ammonia.In conclusion, increasing the pressure in the reaction between nitrogen and hydrogen gases to form ammonia gas will shift the position of equilibrium to the right, favoring the formation of ammonia. This is consistent with Le Chatelier's principle and can be supported by analyzing the stoichiometry of the reaction and the equilibrium constant expression.