Question

If an inert gas is added in the reaction: \(\mathrm{N}_{2}+3 \mathrm{H}_{2} \rightleftharpoons 2 \mathrm{NH}_{3}\) at constant volume, then its equi- librium (a) Remains unaffected (b) Favours the backward reaction (c) Favours the forward reaction (d) Increases the dissociation of reactants

Short answer

(a) Remains unaffected.

Step-by-step solution

Understand the Reaction

The reaction given is the formation of ammonia (NH₃) from nitrogen (N₂) and hydrogen (H₂). It is represented as:\[ \mathrm{N}_{2}+3 \mathrm{H}_{2} \rightleftharpoons 2 \mathrm{NH}_{3} \] This is an equilibrium reaction where both the forward and backward reactions are happening simultaneously.

Examine the Effect of Inert Gas Addition

An inert gas is one that does not react with the substances in the system. The addition of an inert gas at constant volume increases the total pressure, but it does not change the partial pressures of the reactants or products.

Apply Le Chatelier's Principle

Le Chatelier's Principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifts to counteract the change. However, since the partial pressures of the involved gases are unchanged, the equilibrium position remains unaffected by the addition of an inert gas at constant volume.

Determine the Reaction Outcome

Since neither the forward nor the backward reactions are being favoured due to the presence of the inert gas under constant volume, the equilibrium of the reaction remains unchanged.

Key concepts

Le Chatelier's Principle

Le Chatelier's Principle is a fundamental concept in chemistry that helps us understand how a chemical equilibrium responds to changes in conditions. According to this principle, if a system at equilibrium experiences a change in concentration, temperature, or pressure, the system will adjust itself to counteract the effect of the disturbance and restore a new equilibrium. This adjustment might involve shifting the equilibrium position to favor either the forward or the backward reaction.

This principle is particularly useful in predicting how a reaction might shift under different circumstances. However, it's important to note that when adding an inert gas at constant volume—as in the reaction where nitrogen and hydrogen form ammonia—the total pressure may change, but since the partial pressures remain constant, Le Chatelier's Principle indicates that the equilibrium position won't be affected.

Effect of Inert Gas

When an inert gas is introduced into a reaction vessel at constant volume, many students wonder how this affects the equilibrium of the reaction. Inert gases are gases that do not participate in the chemical reaction itself. Common examples include noble gases like helium or neon.

The key point to understand is that when these gases are added to a system at constant volume, they increase the total pressure inside the system. However, they do not change the partial pressures of the individual reactants or products. Since the position of equilibrium is dependent on the partial pressures and concentrations, the equilibrium itself remains unaffected by the presence of these gases.

• Inert gases do not react with the components of the reaction.
• They do not alter the concentrations or partial pressures of the reacting species.
• Thus, the addition of such gases doesn't disrupt the chemical equilibrium.

Equilibrium Constant

The equilibrium constant, denoted as \( K \), is a crucial concept when dealing with reactions at equilibrium. It is a number that signifies the ratio of the concentration of products to the reactants at equilibrium at a certain temperature. For a general reaction \[ aA + bB \rightleftharpoons cC + dD \]The equilibrium constant expression is given by:\[ K = \frac{[C]^c [D]^d}{[A]^a [B]^b} \]

The strength of \( K \) indicates whether the forward or backward reaction is favored:

• If \( K \) is much greater than 1, products are favored at equilibrium.
• If \( K \) is much less than 1, reactants are favored.

Importantly, the value of the equilibrium constant remains unchanged when an inert gas is added at constant volume. This is because the ratio of products to reactants does not change, and thus \( K \) stays constant.

Gaseous Reactions

Gaseous reactions involve reactants and products in the gaseous state and are often influenced by pressure and volume changes. One of the classic examples is the formation of ammonia from nitrogen and hydrogen gases, represented as:\[ \mathrm{N}_{2} + 3\mathrm{H}_{2} \rightleftharpoons 2\mathrm{NH}_{3} \]

These reactions are particularly sensitive to changes in conditions due to the compressibility of gases. Unlike liquids and solids, gases can occupy more or less volume quite easily, which means their concentration can vary significantly with changes in pressure or temperature.

• Gaseous reactions often depend on both the total pressure and the partial pressures of the components involved.
• The balance of these pressures directs the progression and direction of equilibrium.

An understanding of gaseous reactions is essential in industries and laboratories, especially when designing processes that rely on gas-phase chemistry. Overall, recognizing how gaseous conditions alter equilibrium is crucial for both predicting and controlling the outcome of these reactions.