Question

Consider the equilibrium $$ 2 \mathrm{I}(g) \rightleftharpoons \mathrm{I}_{2}(g) $$ What would be the effect on the position of equilibrium of (a) increasing the total pressure on the system by decreasing its volume, (b) adding \(I_{2}\) to the reaction mixture, (c) decreasing the temperature?

Short answer

According to Le Chatelier's Principle, (a) increasing pressure will cause the equilibrium to shift towards the side with fewer gaseous molecules—in this case, the formation of \(I_{2}\), (b) adding \(I_{2}\) will make the equilibrium shift to the left, favoring dissociation into two \(I\) atoms, and (c) lowering the temperature will shift the equilibrium to the right, producing more \(I_{2}\) since the reaction is endothermic.

Step-by-step solution

Effect of Increasing Pressure

According to Le Chatelier's Principle, if pressure is increased, the equilibrium shifts in the direction where fewer gaseous molecules exist, to reduce the pressure. In the equation \(2 \mathrm{I}(g) \rightleftharpoons \mathrm{I}_{2}(g)\), the right-hand side (1 mol of \(I_{2}\)) has fewer molecules than the left-hand side (2 mol of \(I\)). Therefore, when the pressure is increased, the equilibrium will shift to the right to favor the formation of more \(I_{2}\).

Effect of Additional Iodine \(I_{2}\)

When more \(I_{2}\) is added to the mixture, the increased concentration of \(I_{2}\) would cause a shift to the side with less \(I_{2}\) to establish a new equilibrium according to Le Chatelier's principle. As a result, the equilibrium will shift to the left, favoring the dissociation into two \(I\) molecules.

Effect of Decreasing Temperature

This reaction is endothermic (absorbs heat) because the dissociation of iodine into two separate atoms requires energy. So according to Le Chatelier's principle, if the temperature decreases (heat is removed), the equilibrium will shift in the direction that produces heat. That would be the left hand side, as 2I is converted to \(I_{2}\). So the equilibrium will shift to the right, yielding a significant increase in \(I_{2}\).

Key concepts

Le Chatelier's Principle

Le Chatelier's Principle is a fundamental concept in chemistry that helps us predict how a change in conditions affects a chemical equilibrium. Simply put, it states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifts to counteract the change and re-establish equilibrium.
Le Chatelier's Principle can be applied to changes in pressure, temperature, and concentration:

• **Pressure:** For gaseous reactions, changing pressure by altering volume affects the equilibrium in the direction where fewer or more gas molecules are present.
• **Temperature:** A decrease in temperature will shift the equilibrium in the direction of the exothermic reaction, while an increase will favor the endothermic reaction.
• **Concentration:** Adding or removing reactants or products will influence the equilibrium to shift to use up or produce more of the added or removed substances.

By understanding these principles, we can predict the behavior of gases and liquids under various chemical settings, making this principle essential for controlling reactions in industrial and laboratory contexts.

Reaction Shifts

Reaction shifts refer to the change in the position of equilibrium in response to a disturbance in the system. According to Le Chatelier's Principle, equilibrium shifts aim to neutralize changes imposed on the system, striving to reach a new state of balance.For example, in the reaction \[2 \mathrm{I}(g) \rightleftharpoons \mathrm{I}_{2}(g)\]various external changes can cause shifts:

• **Increased Pressure:** The equilibrium shifts towards fewer moles of gas. In this case, to the right, forming \(I_{2}\).
• **Added \(I_{2}\):** An increase in \(I_{2}\) concentration causes the equilibrium to shift towards the left, favoring the formation of more \(I\) atoms.
• **Decreased Temperature:** Since the reaction is endothermic, a temperature drop shifts the equilibrium to the side that releases heat, shifting to the right.

Understanding reaction shifts allows for better manipulation and optimization of chemical processes, crucial for industrial applications.

Gas Phase Reactions

Gas phase reactions are chemical reactions that occur within gases. Due to the fluid nature of gases, reactions in the gas phase can be significantly influenced by changes in pressure and temperature, making them a perfect subject for Le Chatelier's Principle. The key aspects of gas phase reactions include:

• **Molecular Collisions:** Gas molecules move quickly and collide with each other frequently, which influences the rate of reaction. Higher pressure increases collision frequency.
• **Volume and Pressure:** According to Boyle's Law, for a fixed amount of gas at constant temperature, the volume of a gas is inversely proportional to its pressure. Therefore, changes in volume greatly impact the behavior of gas phase reactions.
• **Temperature Dependence:** The reactions are sensitive to temperature changes, with endothermic and exothermic characteristics playing a role in shifts in equilibrium depending on whether the temperature rises or falls.

These characteristics make gas phase reactions particularly dynamic and susceptible to external changes, challenging but also providing opportunities for control in chemical processes.