Problem 30 - A Complete Resource Book in Chemistry for JEE Main

Question

In the reaction \(\mathrm{H}_{2}+\mathrm{I}_{2}\) 2HI at equilibrium, some \(I_{2}\) is added. What happens to the equilibrium? (a) it gets shifted to the right (b) it remains unchanged (c) it gets shifted to the left (d) first (b) then (c)

Step-by-Step Solution

Answer

(a) it gets shifted to the right

Step 1: Identify the Initial Reaction

The given reaction is the formation of hydrogen iodide (HI) from hydrogen gas (\( \mathrm{H}_2 \)) and iodine gas (\( \mathrm{I}_2 \)). At equilibrium, the equation can be represented as: \( \mathrm{H}_2 + \mathrm{I}_2 \rightleftharpoons 2\mathrm{HI} \).

Step 2: Apply Le Chatelier's Principle

Le Chatelier's Principle states that if a system at equilibrium is disturbed, the system will adjust itself to counteract the disturbance and restore a new equilibrium. Adding more \( \mathrm{I}_2 \) will increase the concentration of iodine.

Step 3: Predict Direction of Shift

To counteract the increase in \( \mathrm{I}_2 \) concentration, the equilibrium will shift to the right, towards the formation of more \( \mathrm{HI} \), to reduce the excess \( \mathrm{I}_2 \).

Step 4: Evaluate Result of the Shift

As the equilibrium shifts to the right, more hydrogen iodide is produced. This increases the concentration of \( \mathrm{HI} \), while decreasing the concentration of additional \( \mathrm{I}_2 \).

Key Concepts

Chemical Equilibrium

In chemistry, many reactions do not go to completion. Instead, they reach a state called chemical equilibrium. This occurs when the forward and reverse rates of a reaction are equal. At this point, the concentrations of reactants and products remain constant over time. It is important to note that equilibrium does not mean the reactants and products are equal in concentration, but rather that their rates of formation are balanced.

For example, consider the reaction between hydrogen gas (\( \mathrm{H}_2 \) ) and iodine gas (\( \mathrm{I}_2 \) ) to form hydrogen iodide (\( \mathrm{HI} \) ). At equilibrium, the equation is written as:

• \( \mathrm{H}_2 + \mathrm{I}_2 \rightleftharpoons 2\mathrm{HI} \)

This double arrow indicates that the reaction is reversible and can proceed in both directions. Understanding chemical equilibrium is key in predicting how a reaction mixture will behave when subjected to changes in conditions.

Reaction Shifts

According to Le Chatelier's Principle, any change in the conditions of a chemical equilibrium system will cause the system to adjust to minimize that change. When we say that an equilibrium "shifts," we're talking about the direction the reaction moves to counteract the disturbance.

Suppose you disturb the equilibrium by adding more iodine (\( \mathrm{I}_2 \) ) to the HI system:

• The equilibrium would shift to the right.
• This means the system produces more hydrogen iodide (\( \mathrm{HI} \) ).

By shifting to the right, the system uses up the added \( \mathrm{I}_2 \) , reducing its concentration and reestablishing equilibrium. This concept of reaction shifts is crucial for controlling product yields in industrial and laboratory settings.

Concentration Changes

In a chemical equilibrium, concentration changes can have significant effects on the position of equilibrium. If you increase the concentration of one of the reactants or products, the system will shift in a direction that reduces this change.

Consider the equation:

• \( \mathrm{H}_2 + \mathrm{I}_2 \rightleftharpoons 2\mathrm{HI} \)

Adding more \( \mathrm{I}_2 \) increases its concentration in the mixture:

• The system reacts by trying to "use up" the excess \( \mathrm{I}_2 \) .
• To do this, it shifts right, forming more \( \mathrm{HI} \) .

This change continues until a new equilibrium is reached, balancing the concentrations once again. By understanding how concentration changes affect equilibrium, chemists can predict and control the outcomes of reactions.