The decomposition of NH₃ on a platinum surface is a zero-order reaction. The rates of production of N₂ and H₂
 will be respectively:
(given ; k = 2.5 × 10^–4 mol^–1 L s^–1 ) 

1. 2.5 x 10^-4 mol L^-1 s^-1 and 5.5 x 10^-4 mol L^-1 s^-1

2. 2.5 x 10^-4 mol L^-1 s^-1 and 7.5 x 10^-4 mol L^-1 s^-1

3. 1.5 x 10^-4 mol L^-1 s^-1 and 4.5 x 10^-4 mol L^-1 s^-1

4. 0.5 x 10^-4 mol L^-1 s^-1 and 3.5 x 10^-4 mol L^-1 s^-1

The rate equation of a reaction is expressed as, Rate = \(k(P_{CH_{3}OCH_{3}})^{\frac{3}{2}}\)

(Unit of rate = bar min^-1)

The units of the rate constant will be:

1. bar^1/2 min    
2. bar² min^-1    
3. bar^-1min^-2  
4. bar^-1/2min^-1

The factor(s) that affect the rate of a chemical reaction is/are: 

1. Concentration/Pressure of reactants.

2. Temperature.

3. Presence of a catalyst. 

4. All of the above.

$1.$ $6.67$ $\times {10}^{-2}$ $\mathrm{mol}$ ${\mathrm{L}}^{-1}$ ${\mathrm{s}}^{-1}$
$2.$ $2.67$ $\times$ ${10}^{-4}$ $\mathrm{mol}$ ${\mathrm{L}}^{-1}$ ${\mathrm{s}}^{-1}$ $$
$3.$ $4.67$ $\times$ ${10}^{-3}$ $\mathrm{mol}$ ${\mathrm{L}}^{-1}$ ${\mathrm{s}}^{-1}$
$4.$ $4.27$ $\times$ ${10}^3$ $\mathrm{mol}$ ${\mathrm{L}}^{-1}$ ${\mathrm{s}}^{-1}$

A reaction is first-order with respect to A and second-order with respect to B. The concentration of B is increased three times. The new rate of the reaction would:

In a reaction between A and B, the initial rate of reaction (r₀) was measured for different initial concentrations of A and B as given below:

The order of the reaction with respect to A and B would be:

For a reaction, 2A + B → C + D, the following observations were recorded:

The rate law applicable to the above mentioned reaction would be:

1. Rate = k[A]²[B]³

2. Rate = k[A][B]²

3. Rate = k[A]²[B]  

4. Rate = k[A][B]  

For the  reaction $2\mathrm{A} + \mathrm{B}\to {\mathrm{A}}_2\mathrm{B}$, rate = $\mathrm{k}\left[\mathrm{A}\right]{\left[\mathrm{B}\right]}^2$ with k = $2.0\times {10}^{-6} {\mathrm{mol}}^{-2}{\mathrm{L}}^2{\mathrm{s}}^{-1}$ and $\left[\mathrm{A}\right]=0.1 \mathrm{M}; \left[\mathrm{B}\right]=0.2 \mathrm{M}$. The initial rate of the reaction will be :

1. 0.04$\times {10}^{-9}$ mol ${\mathrm{L}}^{-1} {\mathrm{s}}^{-1}$

2. 8$\times {10}^{-8} {\mathrm{mol}}^{-1}\mathrm{L} {\mathrm{s}}^{-1}$

3. 8$\times {10}^{-9} \mathrm{mol} {\mathrm{L}}^{-1} {\mathrm{s}}^{-1}$

4. 8$\times {10}^{-7}$ mol ${\mathrm{L}}^{-1} {\mathrm{s}}^{-1}$

Given that $\mathrm{Rate}=\mathrm{k}\left[{\mathrm{H}}_2{\mathrm{O}}_2\right]\left[{\mathrm{I}}^{-}\right]$. The dimension of the rate constant in the given rate law is -

1.  ${\mathrm{mol}}^{-1}{\mathrm{Ls}}^{-1}$

2.  ${\mathrm{mol}}^{-1}{\mathrm{Ls}}^{-2}$

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

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

Consider the following rate expression.

$\mathrm{Rate}=\mathrm{k}{\left[C{H}_{3}CHO\right]}^{\frac{3}{2}}$

The order of reaction and dimension of the rate constant are, respectively-

1. $\raisebox{1ex}{$3$}\!\left/ \!\raisebox{-1ex}{$2$}\right.$; k = ${\mathrm{mol}}^{-1}{\mathrm{Ls}}^{-1}$

2. 3; k = ${\mathrm{mol}}^{-1}{\mathrm{Ls}}^{-1}$

3. $\raisebox{1ex}{$3$}\!\left/ \!\raisebox{-1ex}{$2$}\right.$;  k = ${\mathrm{mol}}^{\frac{1}{2}}{\mathrm{L}}^{-\frac{1}{2}}{\mathrm{s}}^{-1}$

4. $\raisebox{1ex}{$3$}\!\left/ \!\raisebox{-1ex}{$2$}\right.$; k = ${\mathrm{mol}}^{-\frac{1}{2}}{\mathrm{L}}^{\frac{1}{2}}{\mathrm{s}}^{-1}$