For the reaction, $\mathrm{aA}+\mathrm{bB}\to \mathrm{cC}+\mathrm{dD}$ the plot of log k vs $\frac{1}{T}$ is given below :

Find the temperature(K) at which the rate constant of the reaction is 10^–4s^–1 ?

(Rounded-off to the nearest integer) 

[Given: The rate constant of the reaction is ${10}^{-5}{\mathrm{s}}^{-1}\text{\hspace{0.17em}at\hspace{0.17em}}500\mathrm{K}\text{.]\hspace{0.17em}}$

The rate constant of a reaction increases by five times on increase in temperature from 27°C to 52°C. The value of activation energy in kJ mol^–1 is-

(Rounded-off to the nearest integer) [R = 8.314 J K^–1 mol^–1]

1. 50

2. 56

3. 52

4. 60

The results given in the below table were obtained during kinetic studies of the following reaction:
2A + B $\to$ C + D

X and Y in the given table are respectively :

1. 0.3, 0.4

2. 0.4, 0.3

3. 0.4, 0.4

4. 0.3, 0.3

For the reaction $2A<mspace\; width="1em"></mspace>+<mspace\; width="1em"></mspace>3B<mspace\; width="1em"></mspace>+\frac{3}{2}C<mspace\; width="1em"></mspace>\to <mspace\; width="1em"></mspace>3P$, the correct statement is:

If 75 % of a first-order reaction was completed in 90 minutes, 60 % of the same reaction would be completed in approximately (in minutes):

(Take : log 2 = 0.30 ; log 2.5 = 0.40)

1. 50 min

2. 60 min

3. 70 min

4. 65 min

The rate constant (k) of a reaction is measured at different temperatures (T), and the data are plotted in the given figure. The activation energy of the reaction in kJ mol^–1 is:
(R is gas constant)

1.  2R
2.  R
3.  1/R
4.  2/R

The rate of a reaction is decreased by 3.555 times when the temperature was changed from 40°C to 30°C. The activation energy (in kJ ${\mathrm{mol}}^{-1}$) of the reaction is:
(Take R=8.314 J ${\mathrm{mol}}^{-1}<mspace\; width="1em"></mspace>{\mathrm{K}}^{-1}$ In 3.555=1.268)

1. 100 kJ/mol
2. 120 kJ/mol
3. 95 kJ/mol
4. 108 kJ/mol

For the reaction $2A<mspace\; width="1em"></mspace>+<mspace\; width="1em"></mspace>B<mspace\; width="1em"></mspace>\to C$, the values of initial rate at different reactant concentrations are given in the table below. The rate law for the reaction is:

1. Rate = k[A] [B]
2. Rate = k [A] ${\left[B\right]}^2$
3. Rate =$K<mspace\; width="1em"></mspace>{\left[A\right]}^2<mspace\; width="1em"></mspace>{\left[B\right]}^2$
4. Rate = $K<mspace\; width="1em"></mspace>{\left[A\right]}^2<mspace\; width="1em"></mspace>\left[B\right]$