How are the values of $E_a$ and $A$ obtained from the Arrhenius equation?

For a first order decomposition of a certain reaction,rate constant is given by the equation $\log k \left( s^{-1} \right) = 7.14 - \frac{1 \times 10^4 \ K}{T}$. The activation energy of the reaction (in $kJ \ mol^{-1}$) is $(R = 8.3 \ J \ K^{-1} \ mol^{-1})$ (in $.1$)

In respect of the equation $k = A e^{-E_a/RT}$ in chemical kinetics,which one of the following statements is correct?

The first order rate constant $k$ is related to temperature $T$ as $\log \, k = 15.0 - (10^{6} / T)$. Which of the following pairs of values for the Arrhenius factor $A$ and activation energy $E_a$ is correct?

$A \rightarrow B$
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: $\ln 10 = 2.3$,$R = 8.3 \, J \, K^{-1} \, mol^{-1}$,$\log 5 = 0.70$,$\log 3 = 0.48$,$\log 2 = 0.30$)

The activation energy of one of the reactions in a biochemical process is $532611 \, J \, mol^{-1}$. When the temperature falls from $310 \, K$ to $300 \, K$,the change in rate constant observed is $k_{300} = x \times 10^{-3} \, k_{310}$. The value of $x$ is $.....$ [Given: $\ln 10 = 2.3$,$R = 8.3 \, J \, K^{-1} \, mol^{-1}$]