The Haber process is used to manufacture ammonia. Explain the optimum conditions for this reaction and why these conditions may not be used in industry

N₂(g)+3H₂(g) ⇋ 2NH₃(g) ΔH=-92kJmol^-1 (equation for question above)
According to Le Chatelier's principle, a system in dynamic equilibrium will shift the position of equilibrium to minimise the change made to it. In this case, optimum conditions would be high pressure and low temperature. High pressure would mean that the system would favour the side that produces the least amount of moles, thereby decreasing the pressure, which in this case is the right hand side.Also, due to the fact the forward reaction is exothermic, the system would favour this side with a lower temperature, because the system wants to negate the change by favouring the exothermic reaction.
However, in industry, these conditions may not be the best options. A high pressure is not only hazardous, but it also requires a large amount of energy to create and maintain. Whilst, a low temperature will cause a lower rate of reaction compared to a higher temperature, due to less collisions occurring. This is why, in industry, they would most likely favour a lower pressure and a higher temperature compared to the optimum conditions.