Chemistry - Chemical Kinetics-PYQ

For the chemical reaction: \[ \text{N}_2(g) + 3\text{H}_2(g) \rightarrow 2\text{NH}_3(g) \], the correct option is:

The rate of the reaction: \[ 2\text{N}_2\text{O}_5 \rightarrow 4\text{NO}_2 + \text{O}_2 \] can be written in three ways:
\[ -\frac{d[\text{N}_2\text{O}_5]}{dt} = k[\text{N}_2\text{O}_5] \]
\[ \frac{d[\text{NO}_2]}{dt} = k'[\text{N}_2\text{O}_5] \]
\[ \frac{d[\text{O}_2]}{dt} = k''[\text{N}_2\text{O}_5] \]
The relationship between k and k' and between k and k'' are:

For the reaction \[ \text{N}_2\text{O}_5(g) \rightarrow 2\text{NO}_2(g) + \frac{1}{2}\text{O}_2(g) \], the value of rate of disappearance of N₂O₅ is given as \[ 6.25 \times 10^{-3} \, \text{mol L}^{-1} \text{s}^{-1} \]. The rate of formation of NO₂ and O₂ is given respectively as:

For the reaction, \[ \text{N}_2 + 3\text{H}_2 \rightarrow 2\text{NH}_3 \], if \[ \frac{d[\text{NH}_3]}{dt} = 2 \times 10^{-4} \, \text{mol L}^{-1} \text{s}^{-1} \], the value of \[ -\frac{d[\text{H}_2]}{dt} \] would be:

In the reaction: \[ \text{BrO}_3^-(aq) + 5\text{Br}^-(aq) + 6\text{H}^+(aq) \rightarrow 3\text{Br}_2(l) + 3\text{H}_2\text{O}(l) \]
The rate of appearance of bromine (Br₂) is related to rate of disappearance of bromide ions as:

Consider the reaction: \[ \text{N}_2(g) + 3\text{H}_2(g) \rightarrow 2\text{NH}_3(g) \]
The equality relationship between \[ \frac{d[\text{NH}_3]}{dt} \] and \[ -\frac{d[\text{H}_2]}{dt} \] is:

For the reaction, \[ 2A + B \rightarrow 3C + D \], which of the following does not express the reaction rate?

For the reaction \[ 3A \rightarrow 2B \], rate of reaction \[ +\frac{d[B]}{dt} \] is equal to:

For the reaction: \[ \text{H}^+ + \text{BrO}_3^- + 3\text{Br}^- \rightarrow 5\text{Br}_2 + \text{H}_2\text{O} \]
Which of the following relations correctly represents the consumption and formation of products?

For the reaction \[ \text{H}_2(g) + \text{I}_2(g) \rightleftharpoons 2\text{HI}(g) \], the rate of reaction is expressed as:

Mechanism of a hypothetical reaction, \[ \text{X}_2 + \text{Y}_2 \rightarrow 2\text{XY} \], is given below:
(i) \[ \text{X}_2 \rightleftharpoons \text{X} + \text{X} \] (fast)
(ii) \[ \text{X} + \text{Y}_2 \rightarrow \text{XY} + \text{Y} \] (slow)
(iii) \[ \text{X} + \text{Y} \rightarrow \text{XY} \] (fast)
The overall order of the reaction will be:

The decomposition of phosphine (PH₃) on tungsten at low pressure is a first-order reaction. It is because the:

The rate constant of the reaction A → B is \[ 0.6 \times 10^{-3} \, \text{mol L}^{-1} \text{s}^{-1} \]. If the concentration of A is 5 M, then concentration of B after 20 minutes is:

For a reaction between A and B the order with respect to A is 2 and the order with respect to B is 3. The concentrations of both A and B are doubled, the rate will increase by a factor of:

In a reaction, A + B → product, rate is doubled when the concentration of B is doubled, and rate increases by a factor of 8 when the concentration of both the reactants (A and B) are doubled, rate law for the reaction can be written as:

Which one of the following statements for the order of a reaction is incorrect?

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

During the kinetic study of the reaction, 2A + B → C + D, following results were obtained:
Run [A]/mol L⁻¹ [B]/mol L⁻¹ Initial rate of formation of D/mol L⁻¹ min⁻¹
I. 0.1 0.1 6.0 × 10⁻³
II. 0.3 0.2 7.2 × 10⁻²
III. 0.3 0.4 2.88 × 10⁻¹
IV. 0.4 0.1 2.40 × 10⁻²
Based on the above data which one of the following is correct?

For the reaction, A + B → products, it is observed that:
(i) on doubling the initial concentration of A only, the rate of reaction is also doubled and
(ii) on doubling the initial concentration 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:

The bromination of acetone that occurs in acid solution is represented by:
\[ \text{CH}_3\text{COCH}_3(aq) + \text{Br}_2(aq) \rightarrow \text{CH}_3\text{COCH}_2\text{Br}(aq) + \text{H}^+(aq) + \text{Br}^-(aq) \]
These kinetic data were obtained:
Initial concentrations (M): [CH₃COCH₃] [Br₂] [H⁺] Initial rate (M s⁻¹)
0.30 0.05 0.05 5.7 × 10⁻⁵
0.30 0.10 0.05 5.7 × 10⁻⁵
0.30 0.10 0.10 1.2 × 10⁻⁴
0.40 0.05 0.20 3.1 × 10⁻⁴
Based on these data, the rate equation is:

The reaction of hydrogen and iodine monochloride is given as:
\[ \text{H}_2(g) + 2\text{ICl}(g) \rightarrow 2\text{HCl}(g) + \text{I}_2(g) \]
This reaction is of first order with respect to H₂(g) and ICl(g), following mechanisms were proposed.
Mechanism A: \[ \text{H}_2(g) + 2\text{ICl}(g) \rightarrow 2\text{HCl}(g) + \text{I}_2(g) \]
Mechanism B:
\[ \text{H}_2(g) + \text{ICl}(g) \rightarrow \text{HCl}(g) + \text{HI}(g) \] (slow)
\[ \text{HI}(g) + \text{ICl}(g) \rightarrow \text{HCl}(g) + \text{I}_2(g) \] (fast)
Which of the above mechanism(s) can be consistent with the given information about the reaction?

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:

If the rate of the reaction is equal to the rate constant, the order of the reaction is:

For reaction 2A → B + C, It would be a zero order reaction when:

For the reaction; \[ 2\text{N}_2\text{O}_5 \rightarrow 4\text{NO}_2 + \text{O}_2 \] rate and rate constant are \[ 1.02 \times 10^{-4} \] and \[ 3.4 \times 10^{-5} \, \text{sec}^{-1} \] respectively, then concentration of N₂O₅ at that time will be:

The experimental data for the reaction, \[ 2A + B_2 \rightarrow 2AB \] is: (refer image)

The rate equation for the above data is?

The given reaction, \[ 2\text{FeCl}_3 + \text{SnCl}_2 \rightarrow 2\text{FeCl}_2 + \text{SnCl}_4 \] is an example of:

The data for the reaction A + B → C, is:
Exp. [A]₀ [B]₀ Initial rate
1 0.012 0.035 0.10
2 0.024 0.070 0.80
3 0.024 0.035 0.10
4 0.012 0.070 0.80
The rate law corresponds to the above data is:

The rate constant for a first order reaction is \[ 4.606 \times 10^{-3} \, \text{s}^{-1} \]. The time required to reduce 2.0 g of the reactant to 0.2 g is:

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:

A first order reaction has a rate constant of \[ 2.303 \times 10^{-3} \, \text{s}^{-1} \]. The time required for 40 g of this reactant to reduce to 10 g will be [Given that \[ \log_{10} 2 = 0.3010 \]]:

The correct difference between first and second order reactions is that:

When initial concentration of the reactant is doubled, the half-life period of a zero order reaction:

A first order reaction has a specific reaction rate of \[ 10^{-2} \, \text{sec}^{-1} \]. How much time will it take for 20 g of the reactant to reduce to 5 g?

The rate of first-order reaction is 0.04 mol L⁻¹ s⁻¹ at 10 seconds and 0.03 mol L⁻¹ s⁻¹ at 20 seconds after initiation of the reaction. The half-life period of the reaction is:

When initial concentration of a reactant is doubled in a reaction, its half-life period is not affected. The order of the reaction is:

A reaction is 50% complete in 2 hours and 75% complete in 4 hours. The order of reaction is:

The half-life of a substance in a certain enzyme-catalysed reaction is 138 s. The time required for the concentration of the substance to fall from 1.28 mg L⁻¹ to 0.04 mg L⁻¹ is:

Half-life period of a first order reaction is 1386 seconds. The specific rate constant of the reaction is:

If 60% of a first order reaction was completed in 60 minutes, 50% of the same reaction would be completed in approximately (log 4 = 0.60, log 5 = 0.69):

In a first-order reaction, A → B, if k is rate constant and initial concentration of the reactant A is 0.5 M, then the half-life is:

For a first order reaction A → B the reaction rate at reactant concentration of 0.01 M is found to be \[ 2.0 \times 10^{-5} \, \text{mol L}^{-1} \text{s}^{-1} \]. The half-life period of the reaction is:

The rate of a first order reaction is \[ 1.5 \times 10^{-2} \, \text{mol L}^{-1} \text{min}^{-1} \] at 0.5 M concentration of the reactant. The half-life of the reaction is:

The reaction A → B follows first order kinetics. The time taken for 0.8 mole of A to produce 0.6 mole of B is 1 hour. What is the time taken for conversion of 0.9 mole of A to produce 0.675 mole of B?

For a first-order reaction, the half-life period is independent of:

For a reaction, activation energy E_a = 0 and the rate constant at 200 K is \[ 1.6 \times 10^{6} \, \text{s}^{-1} \]. The rate constant at 400 K will be [Given that gas constant R = 8.314 J K⁻¹ mol⁻¹]:

The addition of a catalyst during a chemical reaction alters which of the following quantities?

The activation energy of a reaction can be determined from the slope of which of the following graphs?

What is the activation energy for a reaction if its rate doubles when the temperature is raised from 20°C to 35°C? (R = 8.314 J mol⁻¹ K⁻¹)

In a zero-order reaction, for every 10°C rise of temperature, the rate is doubled. If the temperature is increased from 10°C to 100°C, the rate of the reaction will become:

Activation energy (E_a) and rate constants (k₁ and k₂) of a chemical reaction at two different temperatures (T₁ and T₂) are related by:

The rate of the reaction, \[ 2\text{NO} + \text{Cl}_2 \rightarrow 2\text{NOCl} \] is given by the rate equation, rate = k[NO]²[Cl₂]. The value of the rate constant can be increased by:

The rate constants k₁ and k₂ for two different reactions are \[ 10^{16} \cdot e^{-2000/T} \] and \[ 10^{15} \cdot e^{-1000/T} \], respectively. The temperature at which k₁ = k₂ is:

The temperature dependence of rate constant (k) of a chemical reaction is written in terms of Arrhenius equation, \[ k = A \times e^{-E_a/RT} \]. Activation energy (E_a) of the reaction can be calculated by plotting:

The activation energy for a simple chemical reaction A ⇌ B is E_a in forward direction. The activation energy for reverse reaction:

When a biochemical reaction is carried out in laboratory, outside the human body in absence of enzyme, then rate of reaction obtained is 10⁻⁶ times, the activation energy of reaction in the presence of enzyme is:

How enzymes increases the rate of reactions?

Activation energy of a chemical reaction can be determined by:

By the action of enzymes, the rate of biochemical reaction:

An increase in the concentration of the reactants of a reaction leads to change in:

If the rate constant of a reaction is \[ 0.03 \, \text{s}^{-1} \], how much time does it take for a \[ 7.2 \, \text{mol L}^{-1} \] concentration of the reactant to get reduced to \[ 0.9 \, \text{mol L}^{-1} \] (Given: \[ \log 2 = 0.301 \])

If the half-life (t_1/2) for a first-order reaction is 1 minute, then the time required for 99.9% completion of the reaction is closest to:

Activation energy of any chemical reaction can be calculated if one knows the value of:

The rate of a reaction quadruples when temperature changes from 27°C to 57°C. Calculate the energy of activation. Given R = 8.314 J K⁻¹ mol⁻¹, log 4 = 0.6021

For the reaction: \[ 2\text{HI}(g) \rightarrow \text{H}_2(g) + \text{I}_2(g) \] Which of the following rate expressions is correct?

Effective collisions are known to possess:
A: Energy greater than threshold energy.
B: Breaking of old bonds in reactants.
C: Formation of new bonds in products.
D: High activation energy.
E: Proper orientation.
Choose the correct answer from the options given below:

The time taken by the first-order decomposition of SO₂Cl₂ to decompose to 40% is 560 seconds. The rate constant for the reaction is: (Given: log 2.5 = 0.3979)

Rate constants of a reaction at 500 K and 700 K are 0.04 s⁻¹ and 0.14 s⁻¹, respectively. The activation energy of the reaction is: (Given: log 3.5 = 0.5441, R = 8.31 J K⁻¹ mol⁻¹)

For a certain reaction, the rate = k[A]²[B], when the initial concentration of A is tripled keeping concentration of B constant, the initial rate would:

Given below are two statements: one is labelled as Assertion A and the other is labelled as Reason R:
Assertion (A): A reaction can have zero activation energy.
Reason (R): The minimum extra amount of energy absorbed by reactant molecules so that their energy becomes equal to threshold value, is called activation energy.
In the light of the above statements, choose the correct answer from the options given below:

The correct options for the rate law that corresponds to overall first order reaction is:

For a reaction: \[ 3A \rightarrow 2B \] The average rate of appearance of B is given by \[ \frac{\Delta[B]}{\Delta t} \]. The correct relation between the average rate of appearance of B and the average rate of disappearance of A is:

For a first order reaction A → Products, initial concentration of A is 0.1 M, which becomes 0.001 M after 5 minutes. Rate constant for the reaction in min⁻¹ is:

For a chemical reaction \[ 4A + 3B \rightarrow 6C + 9D \] Rate of formation of C is \[ 6 \times 10^{-2} \, \text{mol L}^{-1} \text{s}^{-1} \] and rate of disappearance of A is \[ 4 \times 10^{-2} \, \text{mol L}^{-1} \text{s}^{-1} \]. The rate of reaction and amount of B consumed in interval of 10 seconds, respectively will be:

For a reaction A → B, enthalpy of reaction is –4.2 kJ mol⁻¹ and enthalpy of activation is 9.6 kJ mol⁻¹. The correct potential energy profile for the reaction is shown in option.

The slope of Arrhenius Plot (ln k vs 1/T) of the first-order reaction is –5 × 10³ K. The value of E_a of the reaction is. [Given R = 8.314 JK^–1 mol^–1]