If for a hypothetical reaction,$E_a = 0$ at $273 \ K$,then find the ratio of the rate constants at $383 \ K$ and $273 \ K$.

The rate constant of a first order reaction was doubled when the temperature was increased from $300 \ K$ to $310 \ K$. What is its approximate activation energy (in $kJ \cdot mol^{-1}$)? ($R = 8.3 \ J \cdot mol^{-1} \cdot K^{-1}$; $\log 2 = 0.3$)

For a reversible chemical reaction where the forward process is exothermic, which of the following statements is correct?

The rate coefficient $(k)$ for a particular reaction is $1.3 \times 10^{-4} \ M^{-1} \ s^{-1}$ at $100 \ ^oC$ and $1.3 \times 10^{-3} \ M^{-1} \ s^{-1}$ at $150 \ ^oC$. What is the energy of activation $(E_a)$ (in $kJ \ mol^{-1}$) for this reaction? $(R = 8.314 \ J \ K^{-1} \ mol^{-1})$

The rate constants $k_1$ and $k_2$ for two different reactions are $10^{16} \times e^{-2000/T}$ and $10^{15} \times e^{-1000/T}$ respectively. The temperature at which $k_1 = k_2$ is

What is the value of the slope of the plot $\ln K$ versus $\frac{1}{T}$ for a reaction having $E_{a} = 33.256 \ J \ mol^{-1}$?