In a common interpretation of the Arrhenius rate equation
$$k = A\exp\left(-\frac{E_\mathrm a}{RT}\right),$$
the activation energy $E_\mathrm a$ is understood to represent the difference in the energy of the reactants and the minimum energy required to form the activated complex or the transition state, as shown in the following diagram
But this must be wrong! As $E_\mathrm a$ is the minimum energy itself which the molecules must possess to reach the activated complex. That is, it is not the difference but the minimum energy level itself. This is the only way in which it will remain constant i.e. independent of temperature (otherwise with increase in $T$, $E_\mathrm a$ will go down) and can be used in the Arrhenius Equation where the exponential term denotes the fraction of molecules having energy higher than or equal to $E_\mathrm a$ according to Maxwell-Boltzmann distribution. Sometimes, threshold energy ($E_\mathrm t$) is defined as the minimum energy required for reaction to proceed and $E_\mathrm a = E_\mathrm t-\mathrm{KE_{average}}$. If this definition is true, then threshold energy and not activation energy should be in the Arrhenius equation. I therefore need to clarify the actual definition of activation energy $E_\mathrm a$.