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]

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)

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

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

When the initial concentration of a reactant is doubled in a reaction, its half-life period is not affected. The order of the reaction will be:
1. 0
2. 1
3. 1.5
4. 2 

The rate Constant of reaction A → B is 0.6 × 10^–3 \(\mathrm{molL}^{-1} \mathrm{~S}^{-1}\). If the Concentration of A is 5M, then the concentration of B after 20 min is:

1. 1.08M

2. 3.60M

3. 0.36M

4. 0.72M

When the initial concentration of the reactant is doubled,
the half-life period of a zero-order reaction:

The incorrect statement among the following is:

A first-order reaction has a specific reaction rate of
10^–2 sec^–1. How much time will it take for 20 g of the reactant to reduce to 5 g?