The equilibrium: \(\mathrm{SOCl}_{2}(\mathrm{~g}) \rightleftharpoons \mathrm{SO}_{2}(\mathrm{~g})\) \(+\mathrm{Cl}_{2}(\mathrm{~g})\) is attained at \(25^{\circ} \mathrm{C}\) in a closed container and helium gas is introduced. Which of the following statements is correct? (a) concentration of \(\mathrm{SO}_{2}\) is increased (b) more \(\mathrm{Cl}_{2}\) is formed (c) concentrations of all change (d) concentrations will not change

When studying chemical equilibrium in closed systems, understanding the inert gas effect is crucial. An inert gas, such as helium, neon, or argon, is known for not reacting with other substances under certain conditions. Adding an inert gas to a system in equilibrium affects the total pressure but, significantly, it does not alter the individual partial pressures of the reactants or the products if the volume remains constant.

This aspect is essential because chemical equilibrium is dictated by the partial pressures of the reactants and products rather than the total pressure of the system. Therefore, in the example of SOCl₂ equilibrium, adding helium to the reaction vessel at constant volume will increase the total pressure, but will have no effect on the composition of the equilibrium mixture. The concentrations of SOCl₂, SO₂, and Cl₂ will remain unchanged, which is also supported by Le Chatelier’s principle.

Exploring the concept of chemical equilibrium, specifically in the context of a closed container, is foundational in understanding many chemical reactions. Equilibrium represents a state where the rates of the forward and reverse reactions are equal, leading to constant concentrations of reactants and products over time, albeit not necessarily equal concentrations.

In a closed container, any change to the system, such as temperature, pressure, concentration, or volume, can potentially disturb the equilibrium. However, if these changes do not affect the factors on which the position of equilibrium depends, the system will remain at equilibrium. In our exercise example, despite adding helium gas, which increases the total pressure, the closed nature of the container ensures that the initial conditions of equilibrium (partial pressures of reacting gases and the volume) are maintained. Consequently, the system does not shift in favor of the formation of reactants or products, thus answering the exercise.

Partial pressures are a concept in chemistry that describe the individual pressure contributions of each gas in a mixture of gases. According to Dalton's Law of Partial Pressures, the total pressure exerted by a mixture of non-reacting gases is equal to the sum of the partial pressures of individual gases.

This concept becomes particularly important when dealing with reactions in a gaseous state, as in the given equilibrium reaction with SOCl₂. For a reaction at equilibrium in a closed container, the partial pressures of the gases involved are used to determine the reaction quotient and assess the position of equilibrium. Interestingly, when an inert gas is added to such a system at constant volume, its presence does not alter the partial pressures of the other gases, provided it is indeed inert and does not participate in the reaction. Therefore, the equilibrium position remains unchanged, reinforcing the understanding that chemical reactions in closed systems are sensitive to changes in partial pressures of the reactive gases rather than the total pressure.