When an inert gas, such as helium, is added to a system at equilibrium

It would effect equilibrium only if it produces changes in concentration (or partial pressures) of the reactants or products.

The correct answer is option 3. It would effect equilibrium only if it produces changes in concentration (or partial pressures) of the reactants or products.

When considering the addition of an inert gas to a system at equilibrium, it's important to analyze the circumstances under which the gas is added. An inert gas, such as helium, is chemically non-reactive and does not participate in the chemical reaction. Therefore, its addition primarily affects the physical conditions (like total pressure or volume) of the system rather than the chemical nature of the equilibrium itself. Let's break down the detailed effects:

Constant Volume Addition

When an inert gas is added to a system at constant volume:

Total Pressure: The total pressure of the system increases because the inert gas adds to the overall number of gas molecules in the container.

Partial Pressures: The partial pressures of the reactants and products remain unchanged because the volume of the container does not change, and the inert gas does not alter the number of moles of the reactants or products.

Equilibrium Position: According to Le Chatelier's principle, the equilibrium position of a reaction is influenced by changes in concentration (or partial pressure) of the reactants and products. Since the partial pressures of the reactants and products remain unchanged in this scenario, there is no shift in the equilibrium position.

Therefore, at constant volume, the addition of an inert gas has no effect on the equilibrium position. The correct interpretation in this case is aligned with option (1): "No effect on the equilibrium would occur."

Constant Pressure Addition

When an inert gas is added to a system while maintaining constant pressure:

Volume: The volume of the container must increase to keep the total pressure constant. This expansion decreases the partial pressures of all gaseous species in the system because the same number of moles of gas are now distributed over a larger volume.

Partial Pressures: The partial pressures of the reactants and products decrease as a result of the increased volume.

Equilibrium Position: According to Le Chatelier's principle, if the partial pressures of the reactants and products change, the equilibrium will shift to counteract this change. The direction of the shift depends on the stoichiometry of the reaction:

If the reaction produces more moles of gas on the product side, the equilibrium will shift towards the products to increase the pressure.

If the reaction produces fewer moles of gas on the product side, the equilibrium will shift towards the reactants to increase the pressure.

For example, consider the reaction: \(\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g) \)

In this reaction, 4 moles of gas on the reactant side produce 2 moles of gas on the product side. If an inert gas is added at constant pressure, the equilibrium will shift to the left (towards the reactants) to counteract the decrease in partial pressures by producing more gas molecules.

Analysis of the Given Options:

(1) No effect on the equilibrium would occur: This is true for constant volume addition but not necessarily true for constant pressure addition.

(2) The product would be favored: This is not universally true; it depends on the specific reaction and the conditions of inert gas addition.

(3) It would affect equilibrium only if it produces changes in concentration (or partial pressures) of the reactants or products: This is correct. The equilibrium is affected if the addition of the inert gas changes the partial pressures of the reactants or products, which can occur under constant pressure conditions.

(4) It would affect equilibrium only if it reacts with one of the substances involved in the equilibrium: While this statement is true in general, it is not relevant to the specific scenario involving an inert gas, which by definition does not react with the equilibrium components.

The most accurate statement is (3) It would affect equilibrium only if it produces changes in concentration (or partial pressures) of the reactants or products. This aligns with the detailed understanding that the addition of an inert gas at constant volume has no effect, while at constant pressure, it can affect the equilibrium by changing the partial pressures of the gases involved.