Rate of Reaction

Q. The rate of reaction between two reactants A and B decreases by a factor of 4, if the concentration of reactant B is doubled. The order of this reaction with respect to reactant B is [CBSE AIPMT 2005]

1. - 1
2. 2
3. - 2
4. 1

Q. For the gaseous elementary reaction:
$2A\left(g\right)+B\left(g\right)⟶\text{product}$ , if the pressure is increased three times keeping temperature and volume constant then the rate of reaction will be increased by how many times?

1. $27$
2. $9$
3. $3$
4. $12$

Q.

Select the rate law that conforms to the following data for the reaction : $A+B\to Products$
$\begin{array}{cccc}ExpNo.& \left[A\right]\left(mol{L}^{-1}\right)& \left[B\right]\left(mol{L}^{-1}\right)& Initialrate\left(mol{L}^{-1}{S}^{-1}\right)\\ 1& 0.0012& 0.0035& 1\times {10}^{-4}\\ 2& 0.024& 0.070& 8\times {10}^{-3}\\ 3& 0.024& 0.035& 1\times {10}^{-3}\\ 4& 0.012& 0.070& 8\times {10}^{-3}\end{array}$

1. $-\frac{d\left[A\right]}{dt}=k\left[A\right][B{]}^2$

2. $-\frac{d\left[A\right]}{dt}=k\left[A{]}^2\right[B]$
   

3. $-\frac{d\left[A\right]}{dt}=k\left[A\right][B{]}^3$

4. $-\frac{d\left[A\right]}{dt}=k[B{]}^3$

Q. For the reaction 2A + B + C $\to A_{2}B$ + C the rate law is found to be Rate =k[A][B]${}^2$ with $k=2\times {10}^{-6}mo{l}^{-2}{L}^2{s}^{-1}$
The initial rate of reaction with $\left[A\right]=0.1mol{L}^{-1},\left[B\right]=0.2mol{L}^{-1}and\left[C\right]=0.8mol{L}^{-1}$ is

1. $6.4\times {10}^{-8}mol{L}^{-1}{s}^{-1}$
2. $4\times {10}^{-3}mol{L}^{-1}{s}^{-1}$
3. $8\times {10}^{-9}mol{L}^{-1}{s}^{-1}$
4. $4\times {10}^{-7}mol{L}^{-1}{s}^{-1}$

Q. For a reaction, $N_2{O}_5⟶2N{O}_2+\frac{1}{2}{O}_2$
Given:
$\frac{-d\left[N_2{O}_5\right]}{dt}=k_1\left[N_2{O}_5\right]$
$\frac{d\left[N{O}_2\right]}{dt}=k_2\left[N_2{O}_5\right]$
$\frac{d\left[O_2\right]}{dt}=k_3\left[N_2{O}_5\right]$
The relation between $k_1,k_2$ and $k_3$ are:

1. $2k_1=k_2=4k_3$
2. $k_1=k_2=k_3$
3. $2k_1=4k_2=k_3$
4. $\frac{1}{2}{k}_1=k_2=\frac{1}{4}{k}_3$

Q. Ammonia and oxygen reacts at higher temperatures as
$4N{H}_3\left(g\right)+5O_2\left(g\right)\to 4NO\left(g\right)+6H_{2}O\left(g\right)$
In an experiment, the concentration of NO increases by $2.16\times {10}^{–2}mollitr{e}^{–1}$ in 3 seconds.

Which of the following is correct?

1. Rate of the reaction is $7.2\times {10}^{-3}mol{L}^{-1}{s}^{-1}$
2. Rate of disappearance of ammonia is $7.2\times {10}^{-3}mol{L}^{-1}{s}^{-1}$
   
3. Rate of formation of water is $4.32\times {10}^{-2}mol{L}^{-1}{s}^{-1}$
4. Rate of disappearance of oxygen is $3.6\times {10}^{-2}mol{L}^{-1}{s}^{-1}$

Q. For the elementary reaction, M $\to$ N the rate of disappearance of M increases by a factor of 8 upon doubling the concentration of M. The order of the reaction with respect to M is

1. 4
2. 3
3. 2
4. 1

Q. $xA+yB\to zC.If-\frac{d\left[A\right]}{dt}=-\frac{d\left[B\right]}{dt}=1.5\frac{d\left[C\right]}{dt}$ then x.y and z are

1. 1, 1, 1
2. 3, 2, 3
3. 3, 3, 2
4. 2, 2, 3

Q.

For the reaction, $N_2{O}_5\left(g\right)\to N{O}_2\left(g\right)+\frac{1}{2}{O}_2\left(g\right)$ The value of rate of disappearance of $N_2{O}_5$ is given as $6.25\times {10}^{-3}mol{L}^{-1}{s}^{-1}.$ The rate of formation of $N{O}_2$ and $O_2$ is given respectively as:

1. $6.25\times {10}^{-3}mol{L}^{-1}{s}^{-1}and6.25\times {10}^{-3}mol{L}^{-1}{s}^{-1}$

2. $1.25\times {10}^{-2}mol{L}^{-1}{s}^{-1}and3.125\times {10}^{-3}mol{L}^{-1}{s}^{-1}$

3. $6.25\times {10}^{-3}mol{L}^{-1}{s}^{-1}and3.125\times {10}^{-3}mol{L}^{-1}{s}^{-1}$

4. $1.25\times {10}^{-}2mol{L}^{-1}{s}^{-1}and6.25\times {10}^{-3}mol{L}^{-1}{s}^{-1}$

Q. The rate constant for the reaction
$2N_2{O}_5\to 4N{O}_2+O_2$
is $3.0\times {10}^{-5}{\text{sec}}^{-1}$. If the rate is $2.4\times {10}^{-5}mol{L}^{-1}{\text{sec}}^{-1}$, then the concentration of $N_2{O}_5(inmole{L}^{-1}$) is

1. $1.2$
2. $0.8$
3. $0.004$
4. $1.4$

Q.

For the reaction, $N_2{O}_5\left(g\right)\to N{O}_2\left(g\right)+\frac{1}{2}{O}_2\left(g\right)$ The value of rate of disappearance of $N_2{O}_5$ is given as $6.25\times {10}^{-3}mol{L}^{-1}{s}^{-1}.$ The rate of formation of $N{O}_2$ and $O_2$ is given respectively as:

1. $6.25\times {10}^{-3}mol{L}^{-1}{s}^{-1}and6.25\times {10}^{-3}mol{L}^{-1}{s}^{-1}$

2. $1.25\times {10}^{-2}mol{L}^{-1}{s}^{-1}and3.125\times {10}^{-3}mol{L}^{-1}{s}^{-1}$

3. $6.25\times {10}^{-3}mol{L}^{-1}{s}^{-1}and3.125\times {10}^{-3}mol{L}^{-1}{s}^{-1}$

4. $1.25\times {10}^{-}2mol{L}^{-1}{s}^{-1}and6.25\times {10}^{-3}mol{L}^{-1}{s}^{-1}$

Q. The rate of a reaction is expressed in different ways as follows:
$\frac{1}{2}\left(\frac{d\left[C\right]}{dt}\right)=-\frac{1}{3}\left(\frac{d\left[D\right]}{dt}\right)=\frac{1}{4}\left(\frac{d\left[A\right]}{dt}\right)=-\left(\frac{d\left[B\right]}{dt}\right)$
The reaction is

1. $B+3D\to 4A+2C$
2. $4A+B\to 2C+3D$
3. $4A+2B\to 2C+3D$
4. $B+\left(1/2\right)D\to 4A+3C$

Q.

When dragons burn their victims, the following reaction takes place.

If the dragon burned 1000 moles of O₂ in 10 seconds, what would be the rate of production of CO₂?

1. 100 $molse{c}^{-1}$
2. 200 $molse{c}^{-1}$
3. 2000 $molse{c}^{-1}$
4. 1000 $molse{c}^{-1}$

Q. The rate of a chemical reaction doubles for every $10{}^{\circ }C$ rise of temperature. If the temperature is raised by $50{}^{\circ }C$, the rate of the reaction increase by about

Q. For the reaction 2A + B + C $\to A_{2}B$ + C the rate law is found to be Rate =k[A][B]${}^2$ with $k=2\times {10}^{-6}mo{l}^{-2}{L}^2{s}^{-1}$
The initial rate of reaction with $\left[A\right]=0.1mol{L}^{-1},\left[B\right]=0.2mol{L}^{-1}and\left[C\right]=0.8mol{L}^{-1}$ is

1. $6.4\times {10}^{-8}mol{L}^{-1}{s}^{-1}$
2. $4\times {10}^{-3}mol{L}^{-1}{s}^{-1}$
3. $8\times {10}^{-9}mol{L}^{-1}{s}^{-1}$
4. $4\times {10}^{-7}mol{L}^{-1}{s}^{-1}$

Q. In the following reaction, how is the rate of appearance of product related to the rate of disappearance of the reactant?
$Br{O}_3^{-}\left(aq\right)+5B{r}^{-1}\left(aq\right)+6H^{+}\to 3B{r}_2\left(l\right)+2H_{2}O\left(l\right)$

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

Q. A reaction follows the given concentration (M)-time graph. The rate for this reaction at 20 seconds will be

1. $4\times {10}^{-3}M/s$
2. $8\times {10}^{-2}M/s$
3. $2\times {10}^{-2}M/s$
4. $7\times {10}^{-3}M/s$

Q. The rate constant for the reaction
$2N_2{O}_5\to 4N{O}_2+O_2$
is $3.0\times {10}^{-5}{\text{sec}}^{-1}$. If the rate is $2.4\times {10}^{-5}mol{L}^{-1}{\text{sec}}^{-1}$, then the concentration of $N_2{O}_5(inmole{L}^{-1}$) is

1. $1.4$
2. $1.2$
3. $0.004$
4. $0.8$

Q. Nitric oxide (NO) reacts with oxygen to produce nitrogen dioxide:
$2NO\left(g\right)+O_2\left(g\right)\to 2N{O}_2\left(g\right)$
If the mechanism of reaction is;
$NO+O_2\stackrel{k}{\rightleftharpoons }N{O}_3\left(fast\right)$
$N{O}_3+NO\stackrel{k_1}{⟶}N{O}_2+N{O}_2\left(slow\right)$
The rate law is:

1. Rate=$k^{\prime }\left[NO\right]\left[O_2\right]$
2. Rate=$k^{\prime }\left[NO\right][O_2{]}^2$
3. Rate=$k^{\prime }\left[NO{]}^2\right[O_2]$
4. Rate=$k^{\prime }\left[NO{]}^3\right[O_2]$

Q. For the reaction $2A+B\to C$, the values of initial rate at different reactant concentrations are given in the table below: The rate law for the reaction is:

$\left[A\right]\left(mol{L}^{-1}\right)$	$\left[B\right]\left(mol{L}^{-1}\right)$	$\text{Initial Rate}\left(mol{L}^{-1}{s}^{-1}\right)$
0.05	0.05	0.045
0.10	0.05	0.090
0.20	0.10	0.72

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

Q.

To study the rate of a reaction, it is necessary to

1. identify the reactants

2. know relative amounts of reactants used

3. know the chemical equation for the reaction

4. all of the above

Q. The decomposition of a compound A, at temperature T according to the equation
$2P\left(g\right)\to 4Q\left(g\right)+R\left(g\right)+S\left(l\right)$
is the first order reaction. After 30 minutes from the start of decomposition in a closed vessel, the total pressure developed is found to be 160.25 mm of Hg and after a long period of time the total pressure observed to be 220.25 mm of Hg. Calculate the total pressure of the vessel just after 2 hours in mm of Hg, if volume of liquid S is supposed to be
negligible.
Given : Vapour pressure of S (l) at temperature T = 20.25 mm of Hg.

Q. The data for the reaction is A + B $\to$C is
$\begin{array}{cccc}Exp.& [A{]}_0& [B{]}_0& Initialrate\\ 1& 0.012& 0.035& 0.10\\ 2& 0.024& 0.035& 0.80\\ 3& 0.012& 0.070& 0.10\\ 4& 0.024& 0.070& 0.80\end{array}$

1. $r=k[B{]}^3$
2. $r=k[A{]}^3$
3. $r=k\left[A\right][B{]}^4$
4. $r=k\left[A{]}^2\right[B{]}^2$

Q. Nitrogen tetra oxide $\left(N_2{O}_4\right)$ decomposes as $N_2{O}_4\left(g\right)\to 2N{O}_2\left(g\right)$. If the pressure of $\left(N_2{O}_4\right)$ falls from 0.50 atm to 0.32 atm is 30 minutes, the rate of appearance of $N{O}_2\left(g\right)$ is

1. 0.006 atm/min
2. 0.003 atm/min
3. 0.012 atm/min
4. 0.024 atm/min

Q. When dragons burn their victims, the following reaction takes place:

$2C+O_2⟶2C{O}_2$

The average rate of reaction could be expressed as?

1. 
2. 
3. 
4. All of the above

Q. The rate at which a substance reacts depends upon:

1. Atomic weight
2. Atomic number
3. Molecular weight
4. Active mass

Q. In the following reaction, how is the rate of appearance of product related to the rate of disappearance of the reactant?
$Br{O}_3^{-}\left(aq\right)+5B{r}^{-1}\left(aq\right)+6H^{+}\to 3B{r}_2\left(l\right)+2H_{2}O\left(l\right)$

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

Q.

When dragons burn their victims, the following reaction takes place.

If the dragon burned 1000 moles of O₂ in 10 seconds, what would be the rate of production of CO₂?

1. 100 $molse{c}^{-1}$
2. 200 $molse{c}^{-1}$
3. 1000 $molse{c}^{-1}$
4. 2000 $molse{c}^{-1}$

Q. The following results have been obtained during the kinetic study of the reaction
$A+B⟶C+D$
$\begin{array}{cccc}\text{Experiment}& \text{[A]}& \text{[B]}& \text{Initial rate of}\\ & & & \text{formation of}\\ & & & \left(Mse{c}^{-1}\right)\\ \text{I}& 0.1& 0.1& 6.0\times {10}^{-3}\\ \text{II}& 0.3& 0.2& 7.2\times {10}^{-2}\\ \text{III}& 0.3& 0.4& 2.88\times {10}^{-1}\\ \text{IV}& 0.4& 0.1& 2.40\times {10}^{-2}\end{array}$
Determine the rate law of the above reaction.

1. Rate law $=k\left[A{]}^2\right[B]$
2. Rate law $=k\left[A\right]\left[B\right]$
3. Rate law $=k\left[A{]}^2\right[B{]}^2$
4. Rate law $=k\left[A\right][B{]}^2$

Q. In the following reaction, how is the rate of appearance of product related to the rate of disappearance of the reactant?
$Br{O}_3^{-}\left(aq\right)+5B{r}^{-1}\left(aq\right)+6H^{+}\to 3B{r}_2\left(l\right)+2H_{2}O\left(l\right)$

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