Which of the following statements is in accordance with the Arrhenius equation?

Consider $A \xrightarrow{k_1} B$ and $C \xrightarrow{k_2} D$ are two reactions. If the rate constant $(k_1)$ of the $A \rightarrow B$ reaction can be expressed by the following equation $\log_{10} k = 14.34 - \frac{1.5 \times 10^4}{T/K}$ and activation energy of $C \rightarrow D$ reaction $(Ea_2)$ is $\frac{1}{5}$th of the $A \rightarrow B$ reaction $(Ea_1)$, then the value of $(Ea_2)$ is . . . . . . $kJ \ mol^{-1}$. (Nearest Integer)

For a reaction,activation energy $E_a = 0$ and rate constant $K = 3.2 \times 10^6 \ s^{-1}$ at $300 \ K$. What is the value of the rate constant at $300 \ K$?

Consider the following energy profile for the reaction $X + Y \rightarrow R + S$. Which of the following deductions about the reaction is not correct?

For $A + B \longrightarrow C + D$; $\Delta H = -20 \ kJ \ mol^{-1}$,the activation energy of the forward reaction is $85 \ kJ \ mol^{-1}$. The activation energy for the backward reaction is.....$kJ \ mol^{-1}$

For a gaseous reaction,a large increase in the rate of reaction with a small increase in temperature indicates: