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 incorrect statement regarding the order of reaction is:

The unit of rate constant for a zero-order reaction is:

1. ${\mathrm{s}}^{-1}$

2. $\mathrm{mol}$ ${\mathrm{L}}^{-1}{\mathrm{s}}^{-1}$

3. $\mathrm{L}$ ${\mathrm{mol}}^{-1}{\mathrm{s}}^{-1}$

4. ${\mathrm{L}}^2{\mathrm{mol}}^{-2}{\mathrm{s}}^{-1}$

The rate of the reaction $2{\mathrm{N}}_2{\mathrm{O}}_5\to 4{\mathrm{NO}}_2+{\mathrm{O}}_2$  can be written in three ways:

$\frac{-d\left[N_2{O}_5\right]}{dt}=k\left[N_2{O}_5\right]$
$\frac{d\left[N{O}_2\right]}{dt}=k\text{'}\left[N_2{O}_5\right]$
$\frac{d\left[O_2\right]}{dt}=k\text{'}\text{'}\left[N_2{O}_5\right]$

The relationship between k and k′ and between
k and k′′ are-

1. k′ = k, k′′= k 

2. k′= 2k; k′′= k

3. k′= 2k, k′′= k/2 

4. k′ = 2k; k′′= 2k

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 rate of the reaction

2NO + Cl₂ → 2NOCl is given by the rate equation
rate = k[NO]²[Cl₂]

The value of the rate constant can be increased by: 

1. Increasing the concentration of NO
2. Increasing the concentration of Cl₂
3. Increasing the temperature
4. All of the above

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:

In the reaction, 
${\mathrm{BrO}}_3^{-}\left(\mathrm{aq}\right)+5{\mathrm{Br}}^{-}\left(\mathrm{aq}\right)+6{\mathrm{H}}^{+}\to$ 3Br₂(l)+3H₂O(l) 
The rate of appearance of bromine (Br₂) is related to the rate of disappearance of bromide ions:

1. $\frac{d\left[B{r}_2\right]}{dt}=-\frac{3}{5}\frac{d\left[B{r}^{-}\right]}{dt}$

2. $\frac{d\left[B{r}_2\right]}{dt}=-\frac{5}{3}\frac{d\left[Br\right]}{dt}$

3. $\frac{d\left[B{r}_2\right]}{dt}=\frac{5}{3}\frac{d\left[B{r}^{-}\right]}{dt}$

4. $\frac{d\left[B{r}_2\right]}{dt}=\frac{3}{5}\frac{d\left[B{r}^{-}\right]}{dt}$