For the reaction of nitrogen gas and hydrogen gas to make methane, describe FOUR stresses that would shift the equilibrium to the left toward the reactants? N2(g) + 3 H2(g) ⇌ 2 NH3(g) + Energy

- Addition of NH₃(g)

- Removal of N₂(g)

- Increase of temperature

- Pressure decrease

Explanation:

According to Le Chatelier's principle, if we apply an stress to a reaction at equilibrium, the system will try to shift the equilibrium in order to decrease the stress. If we add reactants, the equilibrium will shift toward the formation of more products (to the consumption of reactants) and vice versa.

The stresses we can apply to this equilibrium are the following:

- Addition of NH₃(g) : it is a product, thus its addition will result in a shift toward reactants.

- Removal of N₂(g): it is a reactant, thus its removal from the reaction mixture will result in a shift toward reactants.

- Increase of temperature: the reaction is exothermic, so it releases energy. Energy is a product. If we add energy (increase the temperature), we are adding a product, so the equilibrium will shift toward the reactants.

- Pressure decrease: because both reactants and products are in the gas phase. A decrease in pressure shifts an equilibrium to the side of the reaction with greater number of moles of gas. In this case, the reactants side has greater number of moles of gas (1 mol + 3 moles= 4 moles) than the products side (2 moles). Thus, the equilibrum will shift toward reactants.

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