Which statement is incorrect according to the Arrhenius equation?

Correct statements regarding Arrhenius equation among the following are:
$(A)$ Factor $e^{-Ea/RT}$ corresponds to fraction of molecules having kinetic energy less than $Ea$.
$(B)$ At a given temperature,lower the $Ea$,faster is the reaction.
$(C)$ Increase in temperature by about $10^{\circ}C$ doubles the rate of reaction.
$(D)$ Plot of $\log k$ vs $\frac{1}{T}$ gives a straight line with $slope = -\frac{Ea}{2.303R}$.
Choose the correct answer from the options given below:

Given below is an expression for the rate constant of a first order reaction occurring at a certain temperature,$T (\text{K})$.
$\ln k = 14.34 - \frac{1.25 \times 10^4}{T}$
The energy of activation in $\text{kcal mol}^{-1}$ for the reaction is :
(Given : $k$ is $\text{s}^{-1}$,$R = 1.987 \text{ cal mol}^{-1} \text{ K}^{-1}$)

The activation energy for a reaction which doubles the rate when the temperature is raised from $298 \ K$ to $308 \ K$ is ........... $kJ \ mol^{-1}$

Regarding the velocity in a reversible reaction,which is the correct explanation of the effect of a catalyst?

For a certain gaseous reaction,the temperature is increased by $10\,^{\circ}C$ from $25\,^{\circ}C$ to $35\,^{\circ}C$. If the rate of the reaction doubles,what will be the value of the activation energy $(E_a)$?