Final answer:
Decreasing the temperature and increasing the pressure in the Haber process for producing ammonia will favor the formation of ammonia according to Le Chatelier's principle.
Step-by-step explanation:
According to Le Chatelier's principle, when a system at equilibrium is disturbed, it will respond to minimize that disturbance. For the reaction N2(g) + 3H2(g) ⇌ 2NH3(g) with ΔH = -92.38 kJ/mol (the Haber process for synthesizing ammonia), we can adjust temperature and pressure to increase the yield of ammonia.
Since the reaction is exothermic (releases heat), decreasing the temperature will favor the production of ammonia. However, it is important to keep a balance, because if the temperature is too low, the rate at which equilibrium is reached is significantly slowed down. Conversely, increasing the pressure favors the formation of ammonia because the forward reaction results in a decrease in the number of gas molecules, thus reducing the pressure according to Le Chatelier's principle.
The effect of adding an inert gas like argon to the system at constant volume does not shift the equilibrium position because argon does not react with the other components. It may increase the total pressure but it doesn't change the partial pressures of the reacting gases, therefore, it has no effect on the equilibrium concentrations of the reactants and products.
Overall, for an industrial process like the Haber process aimed at producing ammonia, the goal is to find an optimal combination of increased pressure and moderately decreased temperature to maximize yield while maintaining an acceptable reaction rate.