The correct statements for the activation energy of a reaction are:

(A) Activation energy is the energy difference between the activated complex and the total energy of the reactants.
(B) Higher is the activation energy, the faster is the rate of reaction.
(C) Catalyst increases the rate of the reaction by following the path of lower activation energy.
(D) In the Arrhenius equation, $E_a$ corresponds to the activation energy

Choose the correct answer from the options given below:

(A), (C) and (D) only

The correct answer is Option (3) → (A), (C) and (D) only

Based on the principles of chemical kinetics and the Arrhenius theory, let's evaluate each statement:

(A) Activation energy is the energy difference between the activated complex and the total energy of the reactants. Correct. Activation energy ($E_a$) is the minimum energy required to reach the high-energy "transition state" or "activated complex" starting from the energy level of the reactants.

(B) Higher is the activation energy, the faster is the rate of reaction. Incorrect. A higher activation energy means a taller "energy barrier" for the reactants to overcome. Consequently, fewer molecules will have enough kinetic energy to react, resulting in a slower reaction rate.

(C) Catalyst increases the rate of the reaction by following the path of lower activation energy. Correct. A catalyst provides an alternative reaction mechanism or pathway that has a lower activation energy. This allows a larger fraction of reactant molecules to successfully transition into products at a given temperature.

(D) In the Arrhenius equation, $E_a$ corresponds to the activation energy. Correct. In the Arrhenius equation, $k = Ae^{-E_a/RT}$, the term $E_a$ explicitly represents the activation energy of the reaction.