Le Chatelier's principle says that if a change is made to the conditions of a system at equilibrium, the position of equilibrium shifts to oppose the change. Shifting to the right means more products and a higher yield. Shifting to the left means more reactants.
Concentration: increasing the concentration of a reactant shifts the position to the right, to use up the extra reactant. Removing a product has the same effect, because the system makes more product to replace it. Temperature: increasing the temperature favours the endothermic direction, which takes in the extra energy. Decreasing the temperature favours the exothermic direction, which gives out energy.
Pressure only matters when gases are involved. Increasing the pressure favours the side of the equation with fewer gas molecules. Decreasing the pressure favours the side with more gas molecules. If both sides have the same number of gas molecules, pressure has no effect on the position. A catalyst does not change the position of equilibrium. It speeds up the forward and reverse reactions equally, so equilibrium is reached faster.
In the Haber process, nitrogen + hydrogen ⇌ ammonia, the balanced equation has four gas molecules on the left and two on the right. The forward reaction is exothermic. A high pressure and a lower temperature both shift the equilibrium to the right and give a higher yield of ammonia. However, a low temperature makes the reaction too slow, so industry uses a compromise temperature.