Following mechanism has been proposed for a reaction,

2A+B $\to$D+E

A+B$\to$ C+D    ...(Slow)

A+ C$\to$ E         ...(Fast)

The rate law expression for the reaction is:

1. r = K[A]²[B]

2. r=K[A][B]

3. r= K[A]²

4. r= K[A][C]

For the reaction N₂ + 3H₂ $\to$ 2NH₃, the rate $\frac{\mathrm{d}\left[{\mathrm{NH}}_3\right]}{\mathrm{dt}}$= 2 x 10^-4 M s^-1 .Therefore, the rate $\frac{-\mathrm{d}\left[{\mathrm{N}}_2\right]}{\mathrm{dt}}$ is given as:

1. 10^-4 Ms^-1

2. 10⁴ Ms^-1

3. 10^-2 sM^-1

4. 10^-4 sM^-1

If 'I' is the intensity of absorbed light and 'c' is the concentration of AB for the photochemical process AB + hv→  AB *, the rate of formation of AB * is directly proportional to:

1. c

2. I

3. I²

4. cI

When ethyl acetate was hydrolysed in pressure of 0.1 N HCl, the rate constant was found to be 5.40 x 10^-5 sec^-1 . But when 0.1 N H₂SO₄ was used for hydrolysis, the rate constant was found to be 6.25 X10^-5sec^-1. Thus, it may be concluded that:

1. H₂SO₄ is stronger than HCI

2. H₂SO₄ is weaker than HCl

3. H₂SO₄ and HCl both have the same strength

4. The data are not sufficient to compare the strength of H₂SO₄ and HCI

The chemical reaction, 2O₃ $\to$3O₂ proceeds as follows;

O₃ $\rightleftharpoons$O₂ + O .....(Fast)

O+O₃ $\to$ 2O₂ ....(Slow)

The rate law expression should be:

1. r = K[O₃]²

2. r = K[O₃]²[O₂]^-1

3. r = K[O₃][O₂]

4. unpredictable

The rate of reaction becomes 2 times for every 10°C rise in temperature. How the rate of reaction will increase when temperature is increased from 30°C to 80°C?

1. 16

2. 32

3. 64

4. 128

For the reaction 2NO₂ + F₂ → 2NO₂F, following

mechanism has been provided,

 NO₂ + F₂   $\stackrel{slow}{\to }$  NO₂F+F

NO₂ + F   $\stackrel{fast}{\to }$ NO₂F

Thus, rate expression of the above

reaction can be written as:

1. r = K[NO₂]²[F₂]

2. r = K[NO₂ ][F₂]

3. r = K[NO₂]

4. r = K[F₂]

For the reaction:

[Cu(NH₃)₄]²⁺ + H₂O_{$\rightleftharpoons$}[Cu(NH₃)₃H₂O]²⁺ + NH₃

the net rate of reaction at any time is given by, net rate =

2.0x10^-4 [Cu(NH₃)₄]²⁺[H₂O] - 3.0x10⁵ [Cu(NH₃ )₃ H₂0]²⁺[NH₃]

Then correct statement is/are :

1. rate constant for forward reaction = 2 x 10^-4

2. rate constant for backward reaction = 3 x 10⁵

3. equilibrium constant for the reaction = 6.6 x 10^-10

4. all of the above

Consider the chemical reaction,

${\mathrm{N}}_2\left(\mathrm{g}\right)+3{\mathrm{H}}_2\left(\mathrm{g}\right)\to 2{\mathrm{NH}}_3\left(\mathrm{g}\right)$

The rate of this reaction can be expressed in terms of time derivative of concentration of  ${\mathrm{N}}_2 \left(\mathrm{g}\right), {\mathrm{H}}_2\left(\mathrm{g}\right)$ and ${\mathrm{NH}}_3\left(\mathrm{g}\right)$.

The correct relationship amongest the rate expressions is: 

(1) Rate $=\frac{-\mathrm{d}\left[{\mathrm{N}}_2\right]}{dt}=-\frac{1}{3}\frac{ \mathrm{d}\left[{\mathrm{H}}_2\right]}{dt}=\frac{1}{2}\frac{ \mathrm{d}\left[{\mathrm{NH}}_3\right]}{dt}$

(2) Rate $=\frac{-\mathrm{d}\left[{\mathrm{N}}_2\right]}{dt}=-\frac{3 \mathrm{d}\left[{\mathrm{H}}_2\right]}{dt}=\frac{2 \mathrm{d}\left[{\mathrm{NH}}_3\right]}{dt}$

(3) Rate $=\frac{\mathrm{d}\left[{\mathrm{N}}_2\right]}{dt}=\frac{1}{3}\frac{ \mathrm{d}\left[{\mathrm{H}}_2\right]}{dt}=\frac{1}{2}\frac{ \mathrm{d}\left[{\mathrm{NH}}_3\right]}{dt}$

(4) Rate $=\frac{-\mathrm{d}\left[{\mathrm{N}}_2\right]}{dt}=-\frac{ \mathrm{d}\left[{\mathrm{H}}_2\right]}{dt}=\frac{ \mathrm{d}\left[{\mathrm{NH}}_3\right]}{dt}$

Select the intermediate in the following reaction mechanism:

O₃(g) $\rightleftharpoons$ O₂(g) +O(g)

O(g) +O₃(g) $\to$ 2O₂(g)

1. O₃(g)

2. O(g)

3. O₂(g)

4. none of these