Question

For the reaction: \(\mathrm{PCl}_{5}(\mathrm{~g}) \rightleftharpoons \mathrm{PCl}_{3}(\mathrm{~g})+\mathrm{Cl}_{2}(\mathrm{~g})\) the forward reaction at constant temperature is favoured by 1\. Introducing an inert gas at constant volume 2\. Introducing chlorine gas at constant volume 3\. Introducing an inert gas at constant pressure 4\. Increasing the volume of the container 5\. Introducing \(\mathrm{pc} 1_{5}\) at constant volume (a) \(1,2,3\) (b) 4,5 (c) \(2,3,5\) (d) \(3,4,5\)

Short answer

Option (d) 3, 4, 5. Both steps favoring product formation.

Step-by-step solution

Understanding the Reaction

The reaction is an equilibrium between \(\mathrm{PCl}_{5}(\mathrm{~g})\) and \(\mathrm{PCl}_{3}(\mathrm{~g}) + \mathrm{Cl}_{2}(\mathrm{~g})\). Increasing the concentration of products or removing products can shift the equilibrium towards the reactants (\(\mathrm{PCl}_{5}\)), while increasing the concentration of reactants or removing reactants can shift it towards the products (\(\mathrm{PCl}_{3}\) and \(\mathrm{Cl}_{2}\)).

Key concepts

Le Chatelier's Principle

Le Chatelier's Principle is a fundamental concept in chemistry that explains how a system at equilibrium responds to changes in conditions. Imagine a balance that is perfectly level. If you push one end down, it will try to balance itself ever again by adjusting in opposition to your push. This principle applies to chemical reactions in equilibrium in much the same way.When a change is introduced to a system at equilibrium—such as a change in concentration, temperature, or pressure—the system will attempt to counteract that change to restore a new equilibrium. For instance, if you increase the concentration of the products in a reaction, the system will try to counteract this by shifting to form more reactants in an effort to rebalance.In our example reaction of \[\mathrm{PCl}_{5}(\mathrm{~g}) \rightleftharpoons \mathrm{PCl}_{3}(\mathrm{~g})+\mathrm{Cl}_{2}(\mathrm{~g}),\]if we add chlorine gas (a product of the reaction), the equilibrium will shift towards the formation of more \(\mathrm{PCl}_{5}\) to minimize the change. Alternatively, if \(\mathrm{PCl}_{5}\) is added to the system, the equilibrium would shift to produce more \(\mathrm{PCl}_{3}\) and \(\mathrm{Cl}_{2}\). Understanding and applying Le Chatelier's Principle helps predict the behavior of gaseous reactions under different conditions.

Gaseous Reactions

In gaseous reactions, like the one involving \[\mathrm{PCl}_{5}(\mathrm{~g}) \rightleftharpoons \mathrm{PCl}_{3}(\mathrm{~g}) + \mathrm{Cl}_{2}(\mathrm{~g}),\]both reactants and products are in the gaseous state. The behavior of gaseous reactions can be significantly influenced by changes in pressure and volume, primarily because gases can be compressed or expanded.

• At constant temperature, increasing the volume of the container decreases the overall pressure.
• This change favors the formation of more moles of gas as the equilibrium shifts to restore pressure, in accordance with Le Chatelier's Principle.

This means that if you increase the size of the container in which the reaction occurs, the equilibrium for our example reaction shifts towards the products, as more moles of gase are found on the product side (total of two moles) than the reactant side (one mole).Because these reactions involve gases, they are especially sensitive to these physical changes, making understanding the dynamics of gaseous reactions critical to predicting the outcome of such changes.

Equilibrium Shift

In the context of chemical reactions, an equilibrium shift refers to the movement of a reaction equilibrium position in response to a change in conditions. This shift is an attempt by the system to achieve a new balance between reactants and products when disturbed.For the given reaction example\[\mathrm{PCl}_{5}(\mathrm{~g}) \rightleftharpoons \mathrm{PCl}_{3}(\mathrm{~g})+\mathrm{Cl}_{2}(\mathrm{~g}),\]equilibrium can be shifted by altering factors like pressure, volume, or concentrations of the reacting species.

• Introducing more \(\mathrm{PCl}_{5}\) to the system shifts equilibrium towards the products \(\mathrm{PCl}_{3}\) and \(\mathrm{Cl}_{2}\).
• A rise in container volume lowers pressure and prompts a shift towards products, which produce more overall moles of gas.

Understanding equilibrium shifts allows us to predict how a reaction will proceed under various conditions. The goal in many chemical processes is to either maximize product formation or reactant recovery, making this concept crucial for practical applications in chemistry.