According to the Arrhenius equation ($k = Ae^{-E_a/RT}$), if the activation energy ($E_a$) of a reaction were zero, the rate constant $k$ would become equal to the Arrhenius constant $A$ ($k = A e^0 = A$). This implies that the rate of the reaction would be independent of temperature, which is practically incorrect for standard chemical reactions. Therefore, a chemical reaction cannot have zero activation energy, making the assertion (A) false. The activation energy is indeed the extra amount of energy that must be supplied to the reactant molecules so that their total energy becomes equal to the threshold energy required to form the activated complex. Hence, the reason (R) is true.