Assertion (A): If the activation energy of a reaction is zero, temperature will have no effect on rate constant. 

Reason (R):The lower the activation energy, the faster is the reaction.

Arrhenius Equation is $\text{k}={\text{Ae}}^{-\text{Ea}/\text{RT}}$

If Activation energy is zero

Their $k={\text{A.e}}^{-0/\text{RT}}$

$k=\text{A}({\text{e}}^0=1)$

So, temperature has no effect on rate constant.

At two different temperatures ${\text{T}}_{\text{1}}{\text{ and T}}_2$

Arrhenius Equation may be written as

$\log \frac{k_2}{k_1}=\frac{\text{Ea}}{2.303\text{R}}\left[\frac{{\text{T}}_{\text{2}}-{\text{T}}_{\text{1}}}{{\text{T}}_{\text{1}}{\text{T}}_{\text{2}}}\right]$

Therefore, if the activation energy is low or zero and then the rate of reaction is faster.