When the temperature of a reaction increases from 300 K to 310 K, the rate of the reaction doubles. What is the activation energy for this reaction?

 \((R=8.314 \mathrm{~J} \mathrm{~K}^{-1} \mathrm{~mol}^{-1} \text { and } \log 2=0.301)\)

1.	\(53.6 \mathrm{~kJ} \mathrm{~mol}^{-1} \)	2.	\(68.6 \mathrm{~kJ} \mathrm{~mol}^{-1} \)
3.	\(59.5 \mathrm{~kJ} \mathrm{~mol}^{-1} \)	4.	\(70.5 \mathrm{~kJ} \mathrm{~mol}^{-1}\)

For a reaction of the type 2A + B $\to$ 2C, the rate of the reaction is given by $k{\left[A\right]}^2\left[B\right]$. When the volume of the reaction vessel is reduced to $\frac{1}{4}$ th of the original volume, the rate of reaction changes by a factor of:

1. 0.25                                                          

2. 16

3. 64                                                             

4. 4

If a reaction A + B → C is exothermic to the extent of 30 kJ mol⁻¹ and the forward reaction has an activation energy of 249 kJ mol⁻¹, the activation energy for the reverse reaction in kJ mol^-1 will be:

Select the correct option based on statements below:

For a reaction A → B, the Arrhenius equation is given as  \(log_{e}k \ = \ 4 \ - \ \frac{1000}{T}\) the activation energy in J/mol for the given reaction will be:

1. 8314

2. 2000

3. 2814

4. 3412

The rate constant for a first-order reaction is $4.606\times {10}^{-3}$ $s^{-1}$. The time required to reduce 2.0 g of the reactant to 0.2 g will be:

For the reaction, \(2 A+B \rightarrow 3 C+D\)

Which of the following is an incorrect expression for the rate of reaction?

For a general reaction A $\to$ B, the plot of the concentration of A vs. time is given in the figure.
 

The slope of the curve will be:

The decomposition of hydrocarbons follows the equation: k = (4.5 × 10¹¹s^–1) e^–28000K/T

The activation energy (E_a) for the reaction would be:
1. 232.79 kJ mol^–1
2. 245.86 kJ mol^–1
3. 126.12 kJ mol^–1
4. 242.51 kJ mol^–1

The role of a catalyst is to change: