Problem 11 - A Complete Resource Book in Chemistry for JEE Main

Question

Which of the following change will shift the reaction in forward direction? \(\mathrm{I}_{2}(\mathrm{~g}) \rightleftharpoons 21(\mathrm{~g})\) Take \(\Delta \mathrm{H}^{\circ}=+150 \mathrm{~kJ}\) (a) increase in concentration of I (b) increase in total pressure (c) decrease in concentration of \(\mathrm{I}_{2}\) (d) increase in temperature

Step-by-Step Solution

Answer

Increasing temperature shifts the reaction forward by favoring the endothermic direction.

Step 1: Understanding the Reaction Equation

This is a reversible reaction: \( \mathrm{I}_{2}(\mathrm{g}) \rightleftharpoons 2 \mathrm{I}(\mathrm{g}) \). It involves the dissociation of iodine gas \( \mathrm{I}_2 \) into iodine atoms \( \mathrm{I} \), which is an endothermic process as indicated by the positive \( \Delta \mathrm{H}^{\circ} = +150 \) \( \mathrm{kJ} \).

Step 2: Determining Effect of Concentration Changes

According to Le Chatelier's Principle, increasing the concentration of a product will shift the reaction towards the reactants to counteract the change. Thus, increasing \( \mathrm{I} \) would cause the reaction to shift left. Conversely, decreasing \( \mathrm{I}_2 \) shifts equilibrium to the right to produce more \( \mathrm{I}_2 \).

Step 3: Analyzing Pressure Changes

In this reaction, the forward reaction increases the number of moles of gas as it goes from 1 mole of \( \mathrm{I}_2 \) to 2 moles of \( \mathrm{I} \). Increasing total pressure generally shifts the equilibrium towards the side with fewer moles of gas. Here, increased pressure would favor the backward reaction (more moles are on the left side).

Step 4: Considering Temperature Effects

The reaction is endothermic; thus, increasing the temperature favors the endothermic direction, which is the forward reaction in this case. This happens because the system will absorb the extra heat to form more products \( 2\mathrm{I} \).

Step 5: Conclusion

Of the given options, increasing temperature will shift the reaction in the forward direction as it favors the endothermic process by absorbing more heat and promoting the dissociation of \( \mathrm{I}_2 \) into \( 2\mathrm{I} \).

Key Concepts

Chemical Equilibrium

Chemical equilibrium is a fundamental concept in chemistry where the forward and reverse reactions occur at the same rate, meaning that the concentrations of reactants and products remain constant over time. In the chemical equation \( \mathrm{I}_{2}(\mathrm{g}) \rightleftharpoons 2 \mathrm{I}(\mathrm{g}) \), the equilibrium can be influenced by various factors such as concentration, pressure, and temperature.
Applying Le Chatelier's Principle, if a system at equilibrium is disturbed, it will shift in a direction that counteracts the disturbance. For instance, increasing the concentration of iodine \( \mathrm{I} \) will shift the equilibrium to the left, favoring the formation of \( \mathrm{I}_{2} \).
Similarly, decreasing the concentration of \( \mathrm{I}_{2} \) will shift the equilibrium to the right, favoring the production of more \( \mathrm{I} \). Understanding such shifts is crucial for controlling chemical reactions in industrial and laboratory settings.

Endothermic Reactions

Endothermic reactions are processes that absorb heat from their surroundings. In the reaction \( \mathrm{I}_{2}(\mathrm{g}) \rightleftharpoons 2 \mathrm{I}(\mathrm{g}) \), the positive \( \Delta \mathrm{H}^{\circ} = +150 \mathrm{~kJ} \) indicates it is an endothermic reaction. This means energy is required to break the iodine molecules into individual iodine atoms.
When the temperature increases, the system absorbs additional heat, which shifts the equilibrium toward the formation of products—in this case, more \( \mathrm{I} \) atoms. Consequently, endothermic reactions such as this one favor the forward reaction when the temperature rises, as they utilize the excess heat to proceed.

• Endothermic reactions require heat input.
• Higher temperatures favor the forward direction.
• Common in processes like photosynthesis and melting.

Pressure Effects

Pressure can greatly affect reactions involving gases, as is the case with \( \mathrm{I}_{2}(\mathrm{g}) \rightleftharpoons 2 \mathrm{I}(\mathrm{g}) \). According to Le Chatelier's Principle, an increase in pressure will shift the equilibrium towards the side with fewer gas moles.
In this reaction, the forward direction results in an increase from 1 mole of \( \mathrm{I}_{2} \) to 2 moles of \( \mathrm{I} \). Therefore, increasing pressure will shift the equilibrium to the left, favoring the formation of \( \mathrm{I}_{2} \) rather than \( \mathrm{I} \).

• Increased pressure shifts equilibrium towards fewer moles of gas.
• Useful in optimizing industrial chemical processes.
• Pressure effects are significant in gaseous reactions.

Understanding these effects helps chemists and engineers to optimize reactions by manipulating conditions to favor the desired product.