Which of the following statements about the order of reaction is incorrect?

1.	Order is not influenced by the stoichiometric coefficient of the reactants.
2.	Order of reaction is the sum of power to the concentration terms of reactants to express the rate of reaction.
3.	The order of reaction is always a whole number.
4.	Order can be determined by experiments only.

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. $5.0\times {10}^{-2}{s}^{-1}$

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

The reaction of hydrogen and iodine monochloride is given as: 
H₂(g) + 2ICl(g) → 2HCl(g) + I₂(g) 
This reaction is of first order with respect to H₂(g) and ICl(g), for which of the following proposed mechanisms:
Mechanism A: 
H₂(g) + 2ICl(g) → 2HCl(g) + I₂(g) 
Mechanism B: 
H₂(g) + ICl(g) →HCl(g) + HI(g); slow 
HI(g) + ICl(g) →HCl(g) + I₂(g); fast

1. B Only

2. A and B both

3. Neither A nor B

4. A only

For the reaction, \(2 A+B \rightarrow 3 C+D\)

Which of the following is an incorrect expression for the rate of reaction?

Consider the reaction 
N₂(g) + 3H₂(g) → 2NH₃(g) 
The equality relationship between \( \frac{{d}\left[{{NH}_{3}}\right]}{dt}\) and \( {-}\frac{{d}\left[{{H}_{2}}\right]}{dt}\) is :

1. $\frac{d\left[N{H}_3\right]}{dt}=-\frac{1}{3}\frac{d\left[H_2\right]}{dt}$

2. $+\frac{d\left[N{H}_3\right]}{dt}=-\frac{2}{3}\frac{d\left[H_2\right]}{dt}$

3. $+\frac{d\left[N{H}_3\right]}{dt}=-\frac{3}{2}\frac{d\left[H_2\right]}{dt}$

4. $+\frac{d\left[N{H}_3\right]}{dt}=-\frac{d\left[H_2\right]}{dt}$

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

For the chemical reaction ${\mathrm{N}}_2\left(\mathrm{(g)}\right)+3{\mathrm{H}}_2\left(\mathrm{(g}\right)$ )$\rightleftharpoons$ $2{\mathrm{NH}}_3\left(\mathrm{(g)}\right)$ the correct option is:

1. \(3\frac {d[H_2]} {dt} = 2 \frac {d[NH_3]} {dt} \)

2. \(-\frac {1} {3}\frac {d[H_2]} {dt} = -\frac {1} {2} \frac {d[NH_3]} {dt} \)

3. \(- \frac{d \left[N_{2}\right]}{dt} = 2 \frac{d [NH_{3}]}{dt}\)

4. \(- \frac{d \left[N_{2}\right]}{dt} = \frac{1}{2} \frac{d [NH_{3}]}{dt}\)