Class 12 Chemistry Chapter 4 Notes Chemical Kinetics

Chemical kinetics is considered a branch of chemistry that handles the study of speeds or rate of chemical reactions, the factor affecting the rates of the reactions and the mechanism by which the reaction proceeds. 

Class 12 Chemistry Chapter 4 Chemical Kinetics is a topic that defines chemical reactions, variables, and mechanisms. It includes both the chemical reaction and the physical process. This chapter describes all the aspects students need to understand the concept and farewell in their examination. The Class 12 Chemistry Chapter 4 notes compiled by Extramarks will help students with that. 

Chemical kinetics Class 12 Chemistry Chapter 4 notes are categorized into swift, prolonged, and moderate reactions based on their varying rate. Students learn the key elements related to this topic, such as dependence on the rate of concentration, integrated equations, collision theory, catalyst and Arrhenius equation under class 12 Chemistry chapter 4 notes.

Chemical kinetics relates to multiple aspects of cosmology, geology, and even in some cases, psychology. As a result, in class 12 chemistry chapter 4 notes we have described a wide variety of chemical kinetics usage and how it helps humanity create a better world of tomorrow. 

Key Topics Covered In Class 12 Chemistry Chapter 4 Notes

Some of the important topics covered under Class 12 Chemistry Chapter 4 Notes by Extramarks include the following.

Introduction

Chemical kinetics, also known as reaction kinetics, helps us understand the rates of reactions and how certain conditions influence them. It further helps to define the characteristics of a chemical reaction by gathering and analyzing information about the reaction mechanism.

A chemical reaction involves breaking bonds in reactant molecules and making bonds in product molecules. Different reactions differ in the strength of broken bonds, occurring at different rates.

Rate of Chemical reaction

It is explained as the rate of change in either reactant or product per unit mole concentration. “Rate of change in concentration of reactant or product”, i.e. the disappearance of the amount of reactant or appearance of the product in a unit interval of time. The rate of the reaction unit is mole L-1 sec-1. Under this exercise of class 12 Chemistry chapter 4 notes, students will learn about

The Rate Law: The rate law (also called the rate equation) for a chemical reaction is an expression that shows a relationship between the rate of the reaction and the concentrations of the reactants participating in it. 

Expression: For a  chemical reaction given by aA + bB → cC + dD

a, b, c, & d are the stoichiometric coefficients of the reactants and products. 

The rate equation for the reaction is obtained as below.

Rate ∝ [A][B]y ⇒ Rate = k[A][B]y

Where [A] & [B] represents the concentrations of the reactants A and B. x & y shows the partial reaction orders for reactants A & B (that may or may not be the same as their stoichiometric coefficients a & b). The proportionality constant is the rate constant of the reaction.

It’s key to understanding that the expression of the rate law for a specific reaction can only be determined experimentally. The rate law expression cannot be achieved from the balanced chemical equation (since the partial orders of the reactants are not necessarily equal to the stoichiometric coefficients). Students can refer to Class 12 Chemistry Chapter 4 notes by Extramarks for more information on the ‘rate of chemical reaction’.

Reaction Orders: The combination of the partial orders of the reactants in the rate law expression equals the total order of the reaction.

If Rate = k[A]x[B]y, the total order of the reaction (n) = x+y

The order of a reaction highlights the changes in the reaction rate that can be anticipated by an increase in the reactants’ concentration. 

For example,

• If this reaction is a zero-order reaction, doubling the reactant concentration will not affect the reaction rate.
• If it is a first-order reaction, doubling the reactant concentration will double the reaction rate.
• In the case of second-order reactions, doubling the concentration of the reactants will quadruple the overall rate of reaction.
• For third-order reactions, the total rate increases by eight times when the reactant concentration is doubled. 

Students may refer to NCERT and CBSE Solutions in addition to class 12 Chemistry chapter 4 notes for a more detailed explanation.

Rate Constants: Rearranging the rate of the equation, the value of the rate constant ‘k’ is 

shown as below k = Rate/[A]x[B]y

Hence, the units of rate constant k (considering the representation of concentration as mol.L-1 or M and time is shown as seconds) can be calculated via the following equation.

k = (M.s-1)*(M-n) = M(1-n).s-1

While class 12 Chemistry chapter 4 notes provide a detailed explanation, students may also refer to various other study materials.

Factors influencing the rate of reaction

The reaction rate depends upon the experimental conditions such as reactant concentration, temperature and catalyst. Under this exercise of class 12 Chemistry chapter 4 notes, students will learn about the following.

a)Nature of reactant: The reactants’ bonding nature determines the reaction rate. The ionic compounds react faster than covalent compounds due to the energy requirement in covalent compounds to cleave the existing binds. 

The reaction between ionic compounds:

AgNO3 + NaCl –> AgCl + NaNO3

Precipitation of AgCl

The reactions between covalent compounds in the presence of sulphuric acid:

CH3COOH + C2H5OH CH3COOC2H5 + H2O

Acetic acid ethyl acetate

b)Temperature: The rate of reaction increases with the increase in temperature due to an increase in average kinetic energy, increasing the number of molecules having more incredible energy than the threshold energy and consequently increasing the number of effective collisions. The rate of a reaction is increased by 100% or doubled with a 10oC increase in temperature. This rate of increase factor denotes the temperature coefficient.

c)Pressure: An increase in partial pressure increases the number of collisions. Hence the rate of reactions involving gaseous reactants increases with the increase in partial pressures.

d)Presence of catalyst: As per CBSE solutions class 12 Chemistry chapter 4 notes, a catalyst increases the rate of reaction by giving an alternative path with decreasing activation energy (Ea‘) for the reaction to proceed, but a negative catalyst retards a chemical reaction.

e)The concentration of reactants: Increasing concentration increases the number of collisions and the activated collisions between the reactant molecules. The Frequency of collision is directly proportional to the rate as per the collision theory. So, the reaction rate increases with the increase in the concentration of reactant.

f)Exposure to radiation: Rates of specific reactions increase by absorption of photons of a particular radiation. These reactions are known as photochemical reactions, e.g. the photosynthesis of plants forming ozone in the stratosphere. 

g)Surface area of solid reactant: According to Class 12 Chemistry chapter 4 notes, the rate of reaction increases with an increase in the surface area of a solid reactant, that is,

Rate of a reaction surface area of solid reactant

Wood shavings ignite more rapidly than a log of wood of the same molecular mass. Coal dust combusts at a faster rate than a large piece of coal. 

Apart from class 12 Chemistry Chapter 4 notes, students may refer to other study materials such as CBSE previous year question papers, important questions and CBSE revision notes while preparing for the competitive examination.

The average rate of reaction: It is the ratio of change in concentration of reactants to the change in time. It is explained by the change in concentration of reactants or products and the time taken for the change. As the reaction proceeds forward, the collisions between the molecules of the participating reactants are reduced, decreasing the reaction’s average rate.

Generally,

 Average reaction rate = change in concentration / time taken = (mol/litre)/time

Example:   For the reaction R → P, the concentration of a reactant converts from 0.03 M to 0.02 M in 25 min. Calculates the average rate of reaction using units of time, both in minutes and seconds.

Numerical solution:  R2= 0.02 M

R1= 0.03 M

t2 – t1= 25 Min

∆[R]/∆t =∆[R2 – R1]/ t2 – t1

=- (0.02 – 0.03) / 25 = 6.67 X 10-6 Ms-1

=0.005 ML-1 min-1

Chapter 4 notes that students may refer to other study materials such as CBSE previous year question papers, important questions and CBSE revision notes while preparing for the examination.

Order of a reaction and molecularity

a)Order of a reaction: The sum of powers of the concentration of a reactant in the rate law expression is known as the order of the chemical reaction. It can be a fraction, zero, or any whole number.

 A + 2B –> C + D is a chemical reaction.

By the equation of rate law,  R = k [A]x [B]y

Now the order of the reaction is defined as the addition of the order of all the reactants participating in a chemical reaction.

Overall order of  reaction = (x + y)

The order of reaction can be a whole number(0, 1,2, 3) and even a fraction. A zero-order reaction depicts that the reaction rate is independent of the concentration of reactants. 

1. i) Zero-order reaction: 

Those reactions in which the rate of reaction does not charge with a reactant’s concentration.

For a general reaction, aA bB

The rate law for such a reaction is denoted as

k = rate / [A]0 

Unit = (mol.L-1s-1/mol L-1) 0

       = mol L-1 sec-1

1. ii) For First order reaction: 

A reaction is said to be first order if its reaction rate is identified by the variation of only one concentration term. The rate law for such a reaction is denoted as.

k = rate / [A]1

Unit of rate constant for first order reaction is = (mol.L-1s-1/molL-1) 1= sec-1

iii)For the second-order reaction: 

The reaction in which the sum of powers of concentration term in rate law equation is two. The rate law for such a reaction is expressed as.

k = rate / [A]2

Unit of rate constant for first order reaction is = (mol.L-1s-1/molL-1) 2=mol.L-1s-1

There are various study materials available in addition to class 12 Chemistry Chapter 4 notes for a more detailed explanation of the order of reactions.

b)Molecularity: Molecularity of reaction is defined as the number of reacting species(atom molecules or ions)participating in an elementary chemical reaction that collides simultaneously to bring about the chemical reaction. The value of the molecularity of a reaction is always positive. It is a theoretical concept. It is never more than three. It cannot be zero.

Let us consider a following reaction:

A + 2B –> C + D

Here reactants are: 1 molecule of A, two molecules of B

Products are one molecule of C and one molecule of D.

Hence we can conclude that it is a trimolecular reaction.

A reaction with molecularity equal to one is called unimolecular.

Example, PCl5 –> PCl3 + Cl2

A reaction with molecularity equal to 2 is called bimolecular.

Example, Cl + CH4 –> HCl + CH3

A reaction with molecularity =3 is called trimolecular.

2FeCl3 + SnCl2 –> 2FeCl2 + SnCl4

It is a theoretical value and does not determine the reaction rate. It does not depend upon external factors like temperature or pressure, etc.

For example, the order of reaction and molecularity of electrolysis of water of ethyl acetate is given by

CH3COOC2H5 + H20   CH3COOH + C2H5OH

The order of a reaction is 1( because the concentration of water does not change during the reaction)

And Molecularity is 2.

Mechanism: The reactions which take place in one step are called elementary reactions. If a sequence of elementary reactions (called mechanism) gives us the products, the reactions are known as complex reactions. The various steps in which the complex reaction takes place are called the mechanism of the reaction. 

Students may refer to various study materials such as NCERT Solutions and CBSE revision notes and Class 12 Chemistry Chapter 4 notes for a more detailed explanation.

Integrated Rate of Equations

The integrated rate of equations expresses the concentration of the reactants in a chemical reaction as a function of time. Thus, such a rate of equations can be employed to check how long it would take for a given percentage of the reactants to be consumed in a chemical reaction. It is essential to observe that different ordered reactions have different integrated rates of equations.

Integrated Rate of Equation for Zero-Order Reactions: With reference to Class 12 Chemistry Chapter 4 notes, Integrated Rate of Equation for Zero-Order Reactions is given by:

Consider the reaction

aA + bB –> cC + dD

Rate = k [A]x[B]y

-dR/dt = k[A]x[B]y

dR/dt is an instantaneous rate.

Integrated rate equation for zero order reaction

-dR/dt = k[R]0=k

dR/dt = – k

∫ dR = – k ∫ dt

[R] = – kt +I

 Where I is the constant of integration.

At t = 0, concentration of reactant R =[R]0, where [R]0 is initial concentration of the reactant. Substituting in equation 

R0= – k .0 + I

I = R0

So the equation becomes R = -kt + R0

kt = [R]0 – [R] (or) k = ([R]0 – [R])/t

[R]0 is the initial reactant concentration ( t = 0)

[R] is the reactant concentration at a time ‘t.’

and the rate constant is k

The Integrated Rate Equation for First-Order Reactions: The integrated rate law for first-order reactions is:

Integrated rate equation for first order reaction

Rate = -dR/dt = k[R]

 ∫ dR/R= – ∫ kt

ln R = – kt + I

At t = 0, R = R0, where [ R0] is the initial concentration of the reactant.

Therefore, equation can be written as 

ln [R]0= – k X 0 +I

I = ln[R]0

So equation becomes

ln R = -kt + ln[R]0

ln [R0/ R]= – kt

kt = 2.303log([R]0/[R]) (or) k = (2.303/t)log([R]0/[R])

Integrated Rate Equation for Second-Order Reactions: The integrated rate of the equation is:

kt = (1/[R]) – (1/[R]0)

The Integrated Rate of reaction details is described in Class 12 Chemistry Chapter 4 notes.

Half-Life of a Reaction: As per class 12 Chemistry chapter 4 notes, the half-life of a chemical reaction can give as the time taken for the concentration of a given reactant to reach 50% of its initial concentration (i.e. the time taken for the reactant concentration to get half of its initial). It is denoted by the  ‘t1/2‘ symbol and is usually represented in seconds.

Half-life for first order reactions:

t = 2.303/k log a/(a-x)

For Half-life x = a/2; and  t = t 1/2

t ½ = 2.303 / k log a/(a-a/2) –> t1/2 = 2.303/ k log 2

t ½ = 2.303 / k X 0.301

t ½ = 0.693/ k

Half-life Formula: It is essential to express that the formula for the half-life of a reaction varies with the order of a reaction.

• For a zero-order reaction, the mathematical expression that can be employed to determine the half-life is: t1/2 = [R]/2k
• For a first-order reaction, the half-life is given by: t1/2 = 0.693/k
• For a second-order reaction, the formula for the half-life of the reaction is 1/k[R]0

Where

• t1/2 is the half-life of the reaction (unit: seconds)
• [R]0 is the initial reactant concentration (unit: mol.L-1 or M)
• k denotes the rate constant of the reaction (unit: M(1-n)s-1 where ‘n’ is the reaction order)

Students may refer to various other study materials in addition to Class 12 Chemistry Chapter 4 notes for more information on this exercise. 

Pseudo order reaction

The reaction that appears to be an nth order reaction yet belongs to some different order is called Pseudo order reaction. Referring to class 12 Chemistry chapter 4 notes, a pseudo-first-order reaction is a chemical reaction between two reactants participating in a chemical reaction and, therefore, should be a second-order reaction. But it resembles a first-order response due to the presence of reactants in negligible quantities. 

Let R` + R“ –> P

Rate = k[A]1[B]1

Order of reaction = 2.

Let us consider another equation,

CH3Br + OH−→ CH3OH + Br−

The rate law reaction is shown below

Rate = k [OH−] [CH3Br]

Rate = k [OH−][CH3Br] 

= k(constant)[CH3Br]=k′[CH3Br]

so only the concentration of methyl bromide would change during the reaction. The rate would purely depend upon the changes in the CH3Br reaction. Students may refer to NCERT Solutions parallel to Class 12 Chemistry Chapter 4 notes for a more detailed explanation.

Rate determining step

The slowest step during a chemical reaction determines the overall speed of a reaction towards completion and is called the rate-determining step.

Consider the reaction,

NO2(g)+CO(g)→NO(g)+CO2(g)

The elementary way of the reaction are as: 

1st step:   NO2+NO2→NO+NO3 (Rate constant = k1, slow)

Second step: NO3+CO→NO2+CO2 (Rate constant = k2, fast)

When the first step is the slowest in the reaction, it will determine the overall reaction rate. Hence the first step is the rate-determining step of the given reaction, and thus the rate expression for the given reaction is the product of the rate constant and the reactants of this step.

Rate = k1[NO2][NO2]=k1[NO2]2

Students can access the class 12 chemistry chapter 4 notes on Extramarks.

Activation energy: The minimum amount of external energy needed to convert reactant into a product or produce an unstable intermediate is called activation energy. It is denoted as E . Rate of reaction is inversely proportional to the activation energy. Hence, the greater value of activation energy leads to a lower rate of reaction and increased influence of temperature change on the rate constant. Unless particles collide with enough energy to supply the activation energy, they simply do not react. Students may refer to various study materials in addition to Class 12 Chemistry Chapter 4 notes.

Arrhenius equation

The Arrhenius equation represents the actual dependence of the rate constants on temperature.

Students will learn about the Arrhenius equation in the class 12 Chemistry chapter 4 notes. The formula used to describe the energy of activation and justify the effect of temperature on the rate of reaction is called Arrhenius Equation.

The formula is,

              K = A e-Ea/RT                  

Where

k = Rate constant

A=Arrhenius equation for Frequency factor

e = mathematical quantity

Ea= activation energy

R = gas constant

T = kelvin temperature

The above relation was created by Swedish chemist Svante Arrhenius and hence named after him.

ln K = ln A – Ea/(2.303RT)

Equation of a straight line with slope = –Ea /R.

When Ea = 0 , Temperature = Infinity

e-Ea/RT =Boltzmann factor.

For a chemical reaction, the rate constant gets doubled for a rise of 10° temperature due to the Arrhenius Equation. 

The class 12 Chemistry chapter 4 notes offer more examples. 

K = Ae-Ea/RT

Taking log on both sides of the equation

Ln k = ln A – Ea/RT

Compared with the straight-line equation

y= mx+c,  

[ where m= slope of the line

c= y-intercept]

So we have

y = ln k

x = 1/T

m = -Ea / R

c = ln A

Collision theory

According to the collision theory,” the molecules of a reactant are assumed to be hard spheres, and the reactions are believed to occur only when these spheres collide with each other”.

As per class 12 Chemistry Chapter 4 collision theory is when the molecules collide with significant activation energy to create a chemical reaction. The molecules of the reacting particles collide through space in a rectilinear motion. The chemical rate of reaction is proportional to the number of collisions between the molecules of the reacting species. In collision theory, activation energy and ideal orientation of the molecules altogether determine the criteria for an effective collision and hence the rate of reaction.

The molecules with sufficient energy and proper orientation collide to give a product. Such collisions are called effective collisions—the rate of successful collisions Rate ∝ Fraction of successful collisions X Overall collision frequency. The no. of collisions per second per unit volume of the chemical reaction mixture is called collision frequency (Z).

Let A+B –> C + D

Rate = ZABe-Ea/RT

Where ZAB = collision frequency of A and B.

In many reactions Rate = P ZABe-Ea/RT

Where p= steric factor, which considers the proper orientation of the molecules participating in a chemical reaction. 

Radiocarbon dating: carbon C14 dating is based on the fundamental assumption that the intensity of cosmic rays and C14  has been constant over thousands of years in the atmosphere. Radiocarbon dating creates the initial activity of C14 corresponding to when a plant or an animal died, and further assimilation of radiocarbon ceased.

           7N14 + 0n1 6C14 + 1H1

           6C14 7N14 + 1e0

Time period can be identified as

 t = 2.303/ log N0/Nt

t ½  of carbon C 14 = 5770 years

Students may refer to CBSE revision notes, CBSE sample papers, essential questions and more apart from Class 12 Chemistry Chapter 4 notes for more details on carbon dating and collision theory.

A brief summary of the key topics included in chapter 4

CBSE  Class 12 Chemistry short  revision notes Chapter 4 Chemical kinetics:

• Chemical Kinetics is the field of chemistry that deals with the study of reaction rates of chemical reactions, the effect of various factors affecting it and the mechanism by which the reactions proceed.
• The reaction rate changes the concentration of reactants or products per unit of time.
• The negative sign denotes that the concentration is getting decreased with time. Unit for a reaction rate is mol L-1s-1.
• The rate of a reaction is not a constant quantity (except for zero-order reactions). It decreases as the reaction move in the forward direction.
• A rate law shows the mathematical relationship between the reaction rate and the molar concentration of one or more reactants. It can be deduced experimentally and cannot be predicted.
• The rate constant is the reaction rate in which the concentration of each reacting species is unity. It is represented by ‘k’. It is also called a specific reaction rate or velocity constant of the reaction.

The ratio of change in concentration in a chemical reaction to the time period is called the instantaneous reaction rate. 

-d[R]/dt = change in chemical concentration over a short period of time/ the short time elapsed = (mol/litre) / time.

• Order of reaction is described as the total of the exponents to which the concentration terms are raised in the reaction’s rate equation (or rate law). It can be a fraction, zero or any whole number.

The order of reaction concerning reactant is the power of its concentration term, which appears in the rate law. It is an experimentally determined quantity.

• The Molecularity of reaction is the number of reacting species (atoms or molecules or any other species) taking part in elementary reactions, which must collide simultaneously to bring about a chemical reaction.

It is a theoretical concept. Its value is always a whole number, and it  is never more than three. It cannot be zero.

• The rate constant is the proportionality constant in the rate law.
• First-order reaction: A reaction is said to be first order if its reaction rate is determined by the variation of one concentration term only.
• The integrated rate equation expresses the concentration of reactants as a function of time.
• The integrated rate equation can be determined by integrating the differential rate equations.
• Second-order reaction: The reaction in which the sum of powers of concentration terms in the rate law equation is two.
• Zero-order reaction: Those reactions in which the rate of reaction does not change with the concentration of the reactants. The rate law for such reaction is expressed – Rate = k [A]°[B]°
• Half-life period: It is the time required for the initial concentration of the reactant to be reduced to half its value. It is the time when half of the initial amount of reactant is converted into products(time used for 50% completion of reaction). 

Several factors such as the nature of reactant and product, exposure to radiation, temperature, the concentration of reactants, catalyst, and surface area of solid reactant affect the reaction rate.

• It has been found that for a chemical reaction with a rise in temperature by 10 °C, the rate constant gets nearly doubled.
• The temperature coefficient of a reaction is the ratio of the reaction’s rate constants at two temperatures that differ by 10°C. The two temperatures usually are 35 °C and 25 °C.
• The Arrhenius equation can represent the variation of rate constants with temperature, K=Ae-Ea/Rt, where A is a constant called the frequency factor, and Ea is called the activation energy.
• There are two critical theories of reaction rates:

(i) Collision theory and,

(ii) Transition state theory.

• In collision theory, activation energy and proper orientation of the molecules together determine the criteria for an effective collision and, hence, the reaction rate.
• Nuclear chemistry studies chemical reactions involving changes in nuclei of atoms. It provides information about the kinetics of radioactive decay and its period, as mentioned in Chemistry Class 12 Chapter 4 Notes.

All students must go through chapter 4 class 12 chemistry notes because the notes provide them with evidence of the rate of reaction, chemical mechanism, and how it takes place in the reaction. The chemical reactions are of genuine scientific interest. It will also help students, teachers, research students, and even scientists determine which catalyst works best for the reaction to complete in a minimum time.

Class 12 Chemistry Chapter 4: Exercises & Answer Solutions

Students may refer to class 12 Chemistry chapter 4 notes on Chemical kinetics on Extramarks. The exercise and answer solutions are explained in detail to help students understand the various concepts mentioned in the chapter. Every minute detail that a student may need to understand chemical kinetics is clearly described in the notes.

Furthermore, students can click on the links provided below to access the study material they may need to prepare for the examination. In addition, to exercise and answer solutions, at Extramarks, we provide various other study materials that can be accessed below. 

CBSE Syllabus

CBSE Revision Notes

CBSE Sample Papers

CBSE Extra Questions

CBSE Previous Year Question Papers

The Class 12 Chemistry Chapter 4 notes on Chemical kinetics add to students’ learning skills and tests their information recall, comprehension, analytical thinking, and problem-solving ability. It is the most comprehensive study material students can rely on to study, practice various questions, and prepare for board and various other entrance examinations. Students may find various information under one channel, making studying easier, especially during board exams. 

Key Features of NCERT Solutions Class 12 Chemistry Chapter 4 Notes

Chemical kinetics is an important topic for exams like IIT, JEE and NEET. Hence while studying Chemistry, the best quality notes and solutions provided by Extramarks have proved to be helpful for students. Class 12 chemistry chapter 4 notes can help students understand the topics covered under the chapter. 

• Educated and experienced subject matter experts prepare them.
• These important notes include a detailed explanation of every topic present in the chapter. 
• They can be used to study right after school or revise notes during examinations. 
• The class 12 chemistry chapter 4 notes are helpful for all types of school examinations board or various competitive examinations.