The rate constant for a reaction 2×10^–2 s^–1 at 300 K and 8×10^–2 s^–1 at 340 K. The energy of activation of the reaction is:

1. 14.69 kJ mol^–1

2. 29.39 kJ mol^–1

3. 44.34 kJ mol^–1

4. 22.05 kJ mol^–1

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

1.  3.2 × 10⁴ s^-1
2. 1.6 × 10⁶s^-1
3. 1.6 × 10³ s^-1
4. 3.2 × 10⁶ s^-1

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?

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

If the rate constant for a first order reaction is k, the time (t) required for the completion of 99% of the reaction is given by:

1. t = 2.303/k

2. t = 0.693/k

3. t = 6.909/k

4. t = 4.606/k

If 60% of a first-order reaction was completed in 60 min, 50% of the same reaction would be completed in approximately: 
(log 4 = 0.60, log 5 = 0.69)

In a reaction, A + B → Product, the rate is doubled when the concentration of B is doubled, and the rate increases by a factor of 8, when the concentrations of both the reactants (A and B) are doubled. The rate law for the reaction can be written as:

1. Rate = k[A][B]²

2. Rate = k[A]²[B]²

3. Rate = k[A][B]

4. Rate = k[A]²[B]

The rate constant of the given reaction-

R(g) \(\xrightarrow[]{\Delta }\) 2P(g) is 2.48 × 10^–4 s^–1.

A 1 : 1 molar ratio of R to P in the reaction mixture is attained after: 

1. 32 min

2. 27.3 min

3. 20 min

4. 0 min

For a zero-order reaction, the initial amount of reaction is 20 g and half-life is 30 minutes. The amount of reactant left after 60 minutes would be:

1. 5 g

2.  10 g

3.  2.5 g

4.  Zero

The rate constant of a reaction is
3 x 10^-3 mol^-2 L²s^-1. Hence, the order is:

1. 1

2. Zero

3. 2

4. 3

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