The plot of ln k vs \({1 \over T}\) for the following reaction 
\(2N_2O_5(g) \rightarrow 4NO_2 (g) + O_2(g) \) gives a straight line with the slope of the line equal to \(-1.0 \times 10^4 K \).
The activation energy for the reaction in J mol^-1 is:
(Given R = 8.3 J K^-1 mol^-1)

1. \(4.0 \times 10^2 \)
2. \(4.0 \times 10^{-2} \)
3. \(8.3 \times 10^{-4} \)
4. \(8.3 \times 10^4 \)

For a reaction A$\to$B, enthalpy of reaction is $-4.2$ $\mathrm{kJ}\hspace{0.17em}{\mathrm{mol}}^{-1}$ and enthalpy of activation is $9.6$ $\mathrm{kJ}\hspace{0.17em}{\mathrm{mol}}^{-1}$. The correct potential energy profile for the reaction is:

1.		2.
3.		4.

The slope of Arrhenius Plot (ln k v/s $\frac{1}{\mathrm{T}}$) of the first-order reaction is $-5\times {10}^3$ $\mathrm{K}$. The value of E_a of the reaction is:

[Given R = 8.314 JK^-1 mol^-1]

An increase in the concentration of the reactants of a reaction leads to a change in:

1. heat of reaction

2. threshold energy

3. collision frequency

4. activation energy

For a reaction, activation energy $E_a=0$ and the rate constant at 200 K is    1.6 $X$ ${10}^6{s}^{-1}$. The rate constant at 400K will be [Given that gas constant, R=8.314 J $K^{-1}$ $mo{l}^{-1}$]

1.  3. 2 x 10⁴ s^-1
2. 1.6 x 10⁶s^-1

3. 1.6 x10³ s^-1

4. 3.2 x 10⁶ s^-1

A reaction having equal energies of activation for forward and reverse reaction has:

1. ΔG = 0

2. ΔH = 0

3. ΔH = ΔG = ΔS = 0

4. ΔS = 0

For an endothermic reaction, the energy of activation is E_a, and the enthalpy of reaction is ΔH (both of these in kJ/mol). The minimum value of E_a will be:

1. Less than $∆$H

2. Equal to $∆$H

3. More than $∆$H

4. Equal to zero