Problem 15 - A Complete Resource Book in Chemistry for JEE Main

Question

Which of the following equilibrium, in gaseous phase, would be unaffected by an increase in pressure? (a) \(\mathrm{N}_{2}+3 \mathrm{H}_{2} \rightleftharpoons 2 \mathrm{NH}_{3}\) (b) \(\mathrm{N}_{2}+\mathrm{O}_{2} \rightleftharpoons 2 \mathrm{NO}\) (c) \(\mathrm{N}_{2} \mathrm{O}_{4} \rightleftharpoons 2 \mathrm{NO}_{2}\) (d) \(\mathrm{CO}_{2}+1 / 2 \mathrm{O}_{2} \rightleftharpoons \mathrm{CO}_{2}\)

Step-by-Step Solution

Answer

Reaction (b) is unaffected by an increase in pressure because moles are equal on both sides.

Step 1: Understand How Pressure Affects Equilibrium

When pressure on a gaseous equilibrium system is increased, the equilibrium position shifts towards the side with fewer moles of gas. This behavior can help determine which reactions would be unaffected by a change in pressure.

Step 2: Write Down the Number of Moles on Each Side

Identify the number of moles of gases on each side of the reactions: - (a) Reaction: Left (N2 + 3H2) = 4 moles; Right (2NH3) = 2 moles. - (b) Reaction: Left (N2 + O2) = 2 moles; Right (2NO) = 2 moles. - (c) Reaction: Left (N2O4) = 1 mole; Right (2NO2) = 2 moles. - (d) Reaction: Left (CO2 + 1/2 O2) = 1.5 moles; Right (CO2) = 1 mole.

Step 3: Determine Which Equilibrium is Unaffected by Pressure

Compare the total moles of gas on both sides of each reaction: - (a) Total moles decrease when moving to the right. - (b) Total moles are equal on both sides. - (c) Total moles increase when moving to the right. - (d) Total moles decrease when moving to the right. Since (b) has the same number of moles on both sides, it is unaffected by changes in pressure.

Key Concepts

Effect of Pressure on Equilibrium

In a gaseous chemical reaction that reaches equilibrium, the pressure can have a significant impact on the position of equilibrium. When pressure is increased in a closed system, the system will try to counteract this by favoring the side of the reaction that has fewer moles of gas. This is because fewer molecules occupy less space, thereby reducing the pressure.

To visualize this, consider a scenario where you have a reaction with a different number of moles on each side. By increasing the pressure, you effectively 'squeeze' the gas particles, which encourages the system to shift towards the side with fewer particles.

For instance:

• In reaction (a), where you start with 4 moles of reactants and end up with 2 moles of product, increasing pressure shifts the equilibrium towards the products.
• Conversely, for reaction (b), where there is no net change in the number of moles (2 moles on both sides), the position of the equilibrium is not affected by pressure changes.

Understanding this concept is crucial in predicting how changes in conditions can affect chemical reactions and is a fundamental principle of chemical equilibrium.

Le Chatelier's Principle

Le Chatelier's Principle provides a simple way to predict the effect of changes in conditions on a chemical equilibrium. According to this principle, if an external change is applied to a system at equilibrium, the system will adjust itself to partially counteract the effect of that change, thereby establishing a new equilibrium position.

Consider an example where pressure is increased for a gaseous reaction. The equilibrium will shift toward the side with fewer moles of gas, mitigating the increase in pressure by reducing volume.

Le Chatelier's Principle is not just limited to pressure changes. It applies to:

• Concentration changes: Adding more reactants shifts equilibrium to the right, and adding more products shifts it to the left.
• Temperature changes: For exothermic reactions, increasing temperature shifts equilibrium to the left, favoring the reactants. For endothermic reactions, it shifts to the right, favoring products.

This principle serves as a guiding rule in many chemical processes, offering valuable insights into how systems respond to various stresses.

Gaseous Reactions

Gaseous reactions involve substances whose states are gaseous, and they are a central part of discussions about chemical equilibrium. In these reactions, both the reactants and products are often in the gas phase, making them particularly sensitive to changes in pressure and volume.

One key aspect of gaseous reactions is how they respond to external conditions set by their container, like volume adjustments or pressure changes. Unlike solids and liquids, gases are compressible, meaning their molecules can be squeezed closer together with changes in pressure.

When analyzing gaseous reactions, it is essential to:

• Count the number of moles on each side of the reaction to understand potential shifts in equilibrium.
• Recognize that reactions with equal moles of gas on both sides, like reaction (b) in the example, typically remain unaffected by changes in pressure.

These properties make gaseous reactions a fascinating study for understanding chemical equilibria and the dynamic nature of chemical reactions.