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

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]

In a zero-order reaction for every 10 ° rise of temperature, the rate is doubled.
If the temperature is increased from 10 °C to 100 °C,
the rate of the reaction will become:

1. 256 times

2. 512 times

3. 64 times

4. 128 times

The incorrect statement regarding the order of reaction is:

For the reaction, 
${\mathrm{N}}_2{\mathrm{O}}_5\left(\mathrm{g}\right)\to 2{\mathrm{NO}}_2\left(\mathrm{g}\right)+\frac{1}{2}{\mathrm{O}}_2\left(\mathrm{g}\right)$
the value of rate of disappearance of N₂O₅ is given as 6.25 x 10^-3mol L^-1s^-1.The rate of formation of NO₂ and O₂ is given respectively as:
 

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

During the kinetic study of the reaction, 2A + B\( \rightarrow\)C + D, following results were obtained:
 

Based on the above data which one of the following is correct?

1. rate= k[A]²[B]

2. rate= k[A][B]

3. rate= k[A]²[B]²

4. rate= k[A][B]²

The half-life period of a first-order reaction is 1386 s. The specific rate constant of the reaction is:

1. $5.0\times {10}^{-3}{s}^{-1}$

2. $0.5\times {10}^{-2}{s}^{-1}$

3. $0.5\times {10}^{-3}{s}^{-1}$

4. $5.0\times {10}^{-2}{s}^{-1}$

For the reaction, A + B → products, it is observed that-
(1) On doubling the initial concentration of A only, the rate of reaction is also doubled and 
(2) On doubling the initial concentrations of both A and B, there is a change by a factor of 8 in the rate of the reaction. 
The rate of this reaction is given by:

1. $rate=k$ $\left[A{]}^2\right[B]$

2. $rate=k$ $\left[A\right][B{]}^2$

3. $rate=k$ $\left[A{]}^2\right[B{]}^2$

4. $rate=k$ $\left[A\right]\left[B\right]$

For the reaction, \(\mathrm{N}_2+3 \mathrm{H}_2 \rightarrow 2 \mathrm{NH}_3,\) if, \(\frac{d[NH_{3}]}{dt} \ = \ 2\times 10^{-4} \ mol \ L^{-1} \ s^{-1}\), the value of  \(\frac{-d[H_{2}]}{dt}\) would be: