Final answer:
In the Haber process, when pressure is increased, the equilibrium shifts to the right, favoring the production of ammonia because the reaction side with fewer moles of gas (2 moles of ammonia) is favored over the side with more moles of gas (4 moles of nitrogen and hydrogen).
Step-by-step explanation:
When considering the Haber process, which is used to synthesize ammonia (NH3), it is important to understand how the equilibrium can be affected by changes in conditions, such as pressure. According to Le Chatelier's principle, if the pressure is increased, the equilibrium shifts to the side with fewer moles of gas. In the case of the Haber process, the balanced chemical equation is:
N2(g) + 3H2(g) → 2NH3(g)
On the left, we have 1 mole of nitrogen gas and 3 moles of hydrogen gas, making a total of 4 moles of gaseous reactants. On the right, we have 2 moles of ammonia gas. When the pressure is increased, the equilibrium shifts to the right (towards the products) where there are fewer moles of gas, so the forward reaction is favored.
This shift towards the formation of ammonia enables the Haber process to be more efficient by increasing the yield of ammonia under high-pressure conditions.