For an endothermic reaction, energy of activation is $E_a$, and enthaply of reaction is ΔH (both of these are in kJ mol^-1), minimum value of $E_a$, will be:

More than ΔH

The correct answer is Option (3) → More than ΔH

Endothermic Reaction:

In an endothermic reaction, the enthalpy change (\( \Delta H \)) is positive, meaning that the products have higher energy than the reactants. Therefore, energy is absorbed during the reaction.

Activation Energy (\( E_a \)):

The activation energy is the minimum energy required for the reactants to transform into products. It represents the energy barrier that must be overcome for the reaction to proceed.

Relationship Between Activation Energy and Enthalpy Change:

For an endothermic reaction:

The activation energy (\( E_a \)) must be greater than the change in enthalpy (\( \Delta H \)). This is because, in order to reach the transition state (the highest energy point during the reaction), the reactants must not only provide enough energy to overcome the enthalpy difference but also account for the energy needed to reach that transition state.

Mathematically, this can be expressed as:

\(E_a = \Delta H + E_{\text{required to reach the transition state}}\)

Since there is always some additional energy required to reach the transition state, \( E_a \) must be greater than \( \Delta H \).

Conclusion:

Thus, for an endothermic reaction, the minimum value of activation energy \( E_a \) will be more than \( \Delta H \).