Question

Assertion A: An increase of temperature of a chemical reaction at constant pressure, the equilibrium value of the reaction shifts in the direction in which the heat of reaction absorbed.
Reason R: To minimize the Gibbs free energy change due to increase in temperature, the heat of reaction needs to be absorbed by shifting the equilibrium value.
In the light of the above statements, choose the most appropriate answer:

• A. Both A and R are correct and R is the correct explanation of A
• B. Both A and R are correct but R is NOT the correct explanation of A
• C. A is correct but R is not correct
• D. A is not correct but R is correct

Hint:

Remember: Increasing temperature favors the Endothermic direction ($\Delta H > 0$), while decreasing temperature favors the Exothermic direction ($\Delta H < 0$).

Answer 1

The Correct Option is A

Concept: This is a direct application of Le Chatelier's Principle regarding temperature changes and chemical equilibrium, fundamentally rooted in the van't Hoff equation.

Step 1: Analyze Assertion A based on Le Chatelier's Principle.
Le Chatelier's Principle states that if a system at equilibrium is disturbed, it will shift to counteract the disturbance. Adding heat (increasing temperature) causes the system to favor the endothermic direction (where heat is absorbed) to "consume" the added energy. Thus, A is correct.

Step 2: Analyze the thermodynamic basis (Reason R).
Chemical equilibrium always moves toward the state of minimum Gibbs free energy ($G$). The temperature dependence of the equilibrium constant $K$ is governed by $\frac{d(\ln K)}{dT} = \frac{\Delta H^\circ}{RT^2}$. Absorbing heat allows the system to reach a new equilibrium state that minimizes the overall free energy at the higher temperature. Thus, R is correct.

Step 3: Conclusion.
The thermodynamic drive to minimize free energy at a new temperature is the underlying reason why the equilibrium shifts to absorb heat. Therefore, R explains A.