Final answer:
The equilibrium amounts in a reaction can change when external conditions such as pressure and volume are altered, according to Le Chatelier's Principle. If the number of moles of gas differs between reactants and products, equilibrium will shift in response to these changes to re-establish balance.
Step-by-step explanation:
Whether the equilibrium amounts change depending on the direction of the reaction in a system starting with the same amounts of nitrogen and oxygen atoms depends on the reaction's stoichiometry and the external conditions imposed on the reaction. In a scenario where the reaction is at equilibrium and the system experiences a pressure or concentration change, Le Chatelier's Principle can be used to predict the direction in which the reaction will shift to re-establish equilibrium.
For gas-phase reactions, if the number of moles of gases on the reactant side is the same as on the product side, a change in volume or pressure does not affect the position of equilibrium. But if the number of moles of reactant and product gases is different, an increase or decrease in volume or pressure will shift the equilibrium. Decreasing the volume (or increasing pressure) will shift the equilibrium toward the side with fewer moles of gas, and increasing volume (or decreasing pressure) will favor the side with more moles of gas.
In the case of the nitrogen and hydrogen reacting to form ammonia (N2 + 3H2 → 2NH3), the reaction would shift toward products (NH3) if the pressure were increased (or volume decreased), as there are 4 moles of gas on the reactants side and only 2 moles of gas as product. Conversely, a decrease in pressure (or increase in volume) would shift the reaction toward reactants (N2 and H2).