$A \rightarrow B$
The rate constants of the above reaction at $200 \, K$ and $300 \, K$ are $0.03 \, min^{-1}$ and $0.05 \, min^{-1}$ respectively. The activation energy for the reaction is $.... \, J$ (Nearest integer).
(Given: $\ln 10 = 2.3$,$R = 8.3 \, J \, K^{-1} \, mol^{-1}$,$\log 5 = 0.70$,$\log 3 = 0.48$,$\log 2 = 0.30$)

If for two reactions $E_{a_1} > E_{a_2}$ and $\mu_1$ and $\mu_2$ are their respective temperature coefficients,then which of the following is correct?

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)

Which of the following graphs between the rate constant $(k)$ and temperature $(T)$ represents the Arrhenius equation?

The activation energy of a reaction is zero. The rate constant of this reaction

The rate of a chemical reaction doubles for every $10\,^{\circ}C$ rise in temperature. If the temperature is increased to $50\,^{\circ}C$ from $0\,^{\circ}C$,by how many times will the rate of reaction increase?