Chemical equilibrium describes a state where the forward and reverse reaction rates are equal, meaning the concentrations of all substances appear constant over time, even though reactions continue at the molecular level. An inert gas, such as argon or helium, is a gas that does not chemically react with any of the species in the system. The impact of adding such a non-reactive gas depends entirely on the specific physical conditions under which the reaction is contained.
Understanding Chemical Equilibrium and Inert Gases
The analysis of how a system at equilibrium responds to changes relies on Le Chatelier’s Principle, which states that a system will adjust to partially offset any imposed stress. These stresses typically involve changes in temperature, concentration, or pressure. For gaseous reactions, the crucial factor determining the reaction’s direction is not the overall pressure of the container but the pressure exerted by each reacting gas, known as its partial pressure.
Partial pressure is the theoretical pressure a specific gas would exert if it were the only gas occupying the container’s entire volume. Chemical reactions are fundamentally driven by the concentration or partial pressure of the species actively participating in the reaction.
The equilibrium constant expression, which dictates the ratio of products to reactants at equilibrium, is solely a function of these reacting species. Because an inert gas does not participate in the reaction, its partial pressure never appears in the equilibrium constant expression.
Its addition only affects the total pressure of the system, according to Dalton’s Law of Partial Pressures, which states that the total pressure is the sum of the partial pressures of all gases present. Therefore, any effect on the equilibrium must be an indirect consequence of the inert gas’s presence altering the partial pressures of the reactants and products.
The Effect When Volume is Held Constant
When a gaseous reaction is confined within a rigid container, the system’s volume remains fixed. Introducing an inert gas into this sealed container causes the total number of moles of gas to increase. Consequently, the total pressure within the container rises because the inert gas contributes its own partial pressure to the total.
Despite the increase in total pressure, the partial pressures of the reacting gases remain completely unchanged. This is because the volume is constant, and the number of moles of the reactants and products has not changed.
Since the concentration of a gas is defined as moles divided by volume, the concentrations of the reacting species are preserved. Le Chatelier’s Principle only dictates a shift in equilibrium if the concentration or partial pressure of a reacting species is disturbed. Because the partial pressures of the reactants and products are unaffected by the inert gas addition under these constant volume conditions, the equilibrium position does not change.
The Effect When Pressure is Held Constant
The scenario changes significantly when the system is maintained at a constant pressure, such as in a container with a movable piston that can adjust the volume. When an inert gas is added to this system, the total number of gas moles increases, which would normally lead to an increase in total pressure. To counteract this stress and keep the total pressure constant, the piston must move outward, which expands the system’s volume.
This increase in the system’s volume has a crucial consequence: it decreases the concentration and partial pressure of every gas present in the mixture, including the reactants and products. The system experiences a uniform dilution, which acts as the stress that Le Chatelier’s Principle responds to.
To relieve the stress of this overall dilution, the equilibrium will shift in the direction that produces a greater number of moles of gas. By producing more gas molecules, the system attempts to partially restore the lost concentration that resulted from the volume increase.
For example, if the reaction \(A(g) + 2B(g) \rightleftharpoons C(g)\) is diluted, the system has three moles of gas on the reactant side and one mole on the product side. The equilibrium will shift toward the reactants (left) to increase the total number of moles of gas from one to three, counteracting the dilution.
Only reactions where the number of moles of gaseous reactants equals the number of moles of gaseous products will remain unaffected, as there is no side to shift toward to increase the total number of moles.
Why the Distinction is Essential
The contrasting outcomes in these two experimental setups underscore the necessity of defining the reaction conditions precisely. The distinction between a constant volume system and a constant pressure system is the difference between no effect on the equilibrium and a potential shift in the composition of the reaction mixture.
In chemical analysis and industrial processes, predicting the behavior of a reaction when a non-reactive gas is present is paramount for controlling the final yield. The core principle remains that the reaction equilibrium is sensitive to the partial pressures or concentrations of the reacting species, not simply the total pressure of the container.
While adding an inert gas always increases the total system pressure, this increase is only translated into a shift in equilibrium when it forces a volume change. This volume change then alters the concentration of the reactants and products, providing the necessary stress for Le Chatelier’s Principle to take effect.