In a first-order reaction A $\to$ products, the concentration of the reactant decreases to 6.25 % of its initial value in 80 minutes. The value of the rate constant, if the initial concentration is 0.2 mole/litre, will be:

1. $2.17$ $\times$ ${10}^{-2}$ $mi{n}^{-1}$

2. $3.46$ $\times$ ${10}^{-2}$ $mi{n}^{-1}$

3. $3.46$ $\times$ ${10}^{-3}mi{n}^{-1}$

4. $2.16$ $\times$ ${10}^{-3}$ $mi{n}^{-1}$

For a given reaction, the presence of a catalyst reduces the energy of activation by 2 kcal at 27 ^oC. The rate of reaction will be increased by:

1. 20 times

2. 14 times

3. 28 times

4. 2 times

The rate constant for a chemical reaction that takes place at 500 K is expressed as K = A e^-1000. The activation energy of the reaction will be:

1. 100 cal/mol

2. 1000 kcal/mol

3. 10⁴ kcal/mol

4. 10⁶ kcal/mol

The relationship between temperature and the variance in reaction rate is:

1.		2.
3.		4.

A catalyst lowers the activation energy of a reaction from 20 kJ mol^–1 to 10 kJ mol^-1. The temperature at which the uncatalysed reaction will have the same rate as that of the catalysed at 27 ^oC will be:
1. \(-123\ ^{\circ}C\)
2. \(-327\ ^{\circ}C\)
3. \(327\ ^{\circ}C\)
4. \(23\ ^{\circ}C\)

The rate of reaction triples when the temperature changes from \(20{ }^{\circ} \mathrm{C} \text { to } 50^{\circ} \mathrm{C}\). The energy of activation for the reaction will be:

Given the following reaction:
N₂O₅   as   N₂O₅   ⇌ 2NO₂ + (1/2)O₂
The values of rate constants for the above reaction are 3.45 × 10^-5 and 6.9 × 10^-3 at 27 ^oC and 67 ^oC respectively. The activation energy for the above reaction is : 

1. $102$ $\times {10}^2$ $J$ $$ $$                                    

2. $488.5$ $kJ$

3. $112$ $J$ $$                                             

4. $112.5$ $kJ$

For a chemical reaction $\mathrm{A}\to$ product, the postulated mechanism of the reaction is as follows.

$\mathrm{A}\underset{{\mathrm{k}}_2}{\overset{{\mathrm{k}}_1}{\rightleftharpoons }}3\mathrm{B}\underset{\mathrm{R}.\mathrm{D}.\mathrm{S}}{\overset{{\mathrm{k}}_3}{\to }}{\mathrm{C}}_{}$
If the rate constants for individual reactions are  k₁, k₂ and k₃, and activation energies are 
\(E_{a_{1}} = 180 \ kJ \ mol^{-1}, \)
\( E_{a_{2}} = 90 \ kJ \ mol^{-1}, \)
\( E_{a_{3}} = 40 \ kJ \ mol^{-1}\)
then overall activation energy for the reaction given above is

1. 70 kJ mol^-1

2. -10 kJ mol^-1

3. 310 kJ mol^-1

4. 130 kJ mol^-1

For A + B $\to \hspace{0.17em}$C + D, $∆$H = -20 kJ mol^-1 , the activation energy of the forward reaction is 85 kJ mol^-1. The activation energy for the backward reaction is…. kJ mol^-1.

If a reaction A + B $\to$ C is exothermic to the extent of 30 kJ/mol and the forward reaction has an activation energy of 70 kJ/mol, the activation energy for the reverse reaction will be:

1. 30 kJ/mol                                       

2. 40kJ/mol

3. 70 kJ/mol                                       

4. 100 kJ/mol