Question

A reaction having equal energies of activation for forward and reverse reactions has :-

• A. $\Delta H = \Delta G = \Delta S = 0 $
• B. $\Delta S = 0$
• C. $\Delta G = 0 $
• D. $ \Delta H = 0 $

Answer 1

The Correct Option is D

To solve the question regarding the reaction with equal energies of activation for both forward and reverse reactions, we must explore the thermodynamic implications of this scenario.

In a chemical reaction, the energy of activation (EA) is the minimum energy that reacting species must have in order to form the activated complex during the reaction. When the energies of activation for the forward (\( E_{A, \text{forward}} \)) and reverse (\( E_{A, \text{reverse}} \)) reactions are equal, it indicates a particular relation in terms of the thermodynamic properties of the reaction.

The difference in energies of activation for the forward and reverse reactions is related to the change in enthalpy (\( \Delta H \)) of the reaction. This can be expressed using the following equation:

E_{A, \text{forward}} - E_{A, \text{reverse}} = \Delta H

If \( E_{A, \text{forward}} = E_{A, \text{reverse}} \), then:

\Delta H = 0

This implies that the enthalpy change for the reaction is zero, meaning the reaction is thermoneutral or isothermal. That is, there is no net heat change in the course of the reaction.

Let us evaluate the given options:

• \Delta H = \Delta G = \Delta S = 0: All these properties being zero is a stronger condition not necessary in this context. Only \( \Delta H = 0 \) is required.
• \Delta S = 0: The entropy change (\( \Delta S \)) being zero is not implied by equal activation energies.
• \Delta G = 0: This is the condition for equilibrium, not necessarily related to equal activation energies.
• \Delta H = 0: Correct as the enthalpy change is zero when activation energies are equal.

Hence, the correct answer is that the enthalpy change (\( \Delta H \)) is zero when the energies of activation for forward and reverse reactions are equal.