The relationship between the activation energy of the forward reaction ($E_{a(forward)}$), the activation energy of the backward reaction ($E_{a(backward)}$), and the enthalpy of the reaction ($\Delta H$) is given by: $\Delta H = E_{a(forward)} - E_{a(backward)}$. Rearranging this, we get: $E_{a(backward)} = E_{a(forward)} - \Delta H$.

• If the reaction is exothermic ($\Delta H$ is negative), then $E_{a(backward)} > E_{a(forward)}$.
• If the reaction is endothermic ($\Delta H$ is positive), then $E_{a(backward)} < E_{a(forward)}$. Therefore, the activation energy for the reverse reaction can be less than or more than $E_a$ depending on whether the reaction is endothermic or exothermic.