CBSE Class 12 Chemistry Revision Notes Chapter 3 Chemical Kinetics

Chemical kinetics studies the rate of chemical reactions and the factors that control them. In CBSE Class 12 Chemistry, this chapter explains rate law, order, molecularity, integrated rate equations and collision theory.

Chemical Kinetics explains how fast chemical reactions take place. Thermodynamics tells whether a reaction is feasible, but chemical kinetics tells the speed of the reaction. For example, diamond can convert into graphite, but the rate is so slow that the change is not seen easily.

Use these CBSE Class 12 Chemistry Revision Notes Chapter 3 to revise the 2026–27 chapter in a quick and exam-ready way. Start with rate of reaction, then revise rate law, order, molecularity, integrated rate equations, half-life, Arrhenius equation, catalyst and collision theory.

Key Takeaways

• Rate of Reaction: It is the change in concentration of reactant or product per unit time.
• Rate Law: It is determined experimentally and cannot be predicted only from the balanced equation.
• Order of Reaction: It is the sum of powers of reactant concentrations in the rate law.
• First Order Half-Life: It is independent of the initial concentration of the reactant.

Need help revising Chemical Kinetics formulas and numericals?
Access interactive practice, chapter-wise notes and doubt-solving support on the Extramarks Learning App. Sign Up Free

Access 30 Minutes Class 12 Chemistry Chapter 3 Chemical Kinetics Notes

Chemical Kinetics becomes easier when you revise it through formulas and graphs. First revise average rate, instantaneous rate and rate law.

Then move to order, molecularity, zero order reaction, first order reaction and half-life. Keep Arrhenius equation, catalyst and collision theory for the final round.

Chemical Kinetics Class 12 Chemistry Chapter 3 Overview

Chemical kinetics is the branch of chemistry that studies reaction rates and mechanisms.

It helps explain why some reactions are very fast, some are very slow and some occur at moderate speed.

The rate of a reaction depends on concentration, temperature, pressure and catalyst.

Rate of a Chemical Reaction

The rate of reaction is the change in concentration of a reactant or product in unit time.

For a reaction:

R → P

Rate of disappearance of R:

Rate = -Δ[R] / Δt

Rate of appearance of P:

Rate = +Δ[P] / Δt

The negative sign is used for reactants because their concentration decreases with time.

Average Rate of Reaction

Average rate is the rate measured over a definite time interval.

For reactant R:

Average rate = -Δ[R] / Δt

For product P:

Average rate = +Δ[P] / Δt

Average rate depends on the change in concentration and the time taken for that change.

Instantaneous Rate of Reaction

Instantaneous rate is the rate of reaction at a particular moment.

It is found when the time interval becomes very small.

Instantaneous rate = -d[R] / dt

or

Instantaneous rate = +d[P] / dt

Graphically, it is obtained from the slope of the tangent drawn on the concentration-time curve.

Average Rate and Instantaneous Rate Difference

Units of rate are usually mol L⁻¹ s⁻¹.

For gaseous reactions, rate can also be expressed in atm s⁻¹.

Rate Expression and Rate Constant

Rate expression shows the relation between reaction rate and concentration of reactants.

For a general reaction:

aA + bB → cC + dD

Rate ∝ [A]x[B]y

Rate = k[A]x[B]y

Here, k is the rate constant.

The rate law is found experimentally. It cannot always be predicted from the balanced chemical equation.

Example:

2NO(g) + O₂(g) → 2NO₂(g)

Rate = k[NO]²[O₂]

Order of Reaction in Chemical Kinetics

Order of reaction is the sum of powers of the concentration terms in the rate law.

For:

Rate = k[A]x[B]y

Order = x + y

Order can be 0, 1, 2, 3 or even a fraction.

A zero order reaction means the rate is independent of reactant concentration.

Units of Rate Constant

The unit of rate constant depends on the order of reaction.

A first order rate constant has the unit s⁻¹.

A second order rate constant has the unit L mol⁻¹ s⁻¹.

Molecularity of a Reaction

Molecularity is the number of reacting species taking part in an elementary reaction.

These species must collide simultaneously for the reaction to occur.

Molecularity is used only for elementary reactions.

Reactions with molecularity more than three are rare because simultaneous collision of many particles is unlikely.

Order and Molecularity Difference

For a complex reaction, the slowest step controls the overall rate. This slowest step is called the rate determining step.

Rate Determining Step

The rate determining step is the slowest step in a reaction mechanism.

It controls the overall rate of a complex reaction.

Example: Decomposition of hydrogen peroxide in alkaline medium is catalysed by iodide ion.

Overall reaction:

2H₂O₂ → 2H₂O + O₂

Steps:

H₂O₂ + I⁻ → H₂O + IO⁻

H₂O₂ + IO⁻ → H₂O + I⁻ + O₂

The first step is slow. So, it determines the rate of the reaction.

Integrated Rate Equations

Integrated rate equations connect concentration, time and rate constant.

They help calculate rate constant without using the tangent method.

The NCERT chapter mainly covers zero order and first order reactions.

Zero Order Reactions

A zero order reaction has a rate independent of reactant concentration.

For:

R → P

Rate = k[R]⁰

Rate = k

Integrated rate equation:

[R] = [R]₀ - kt

or

kt = [R]₀ - [R]

Rate constant:

k = ([R]₀ - [R]) / t

For a zero order reaction, a plot of [R] against time gives a straight line.

Slope = -k

Example: Decomposition of ammonia on hot platinum surface at high pressure.

2NH₃(g) → N₂(g) + 3H₂(g)

First Order Reactions

A first order reaction has rate proportional to the first power of reactant concentration.

For:

R → P

Rate = k[R]

Integrated rate equation:

ln[R] = -kt + ln[R]₀

or

k = 1/t ln([R]₀/[R])

Using common logarithm:

k = 2.303/t log([R]₀/[R])

For a first order reaction, a plot of ln[R] against time gives a straight line.

Slope = -k

Examples:

• Hydrogenation of ethene
• Radioactive decay
• Decomposition of N₂O₅

Zero Order and First Order Reaction Formula Table

Half-Life of a Reaction

Half-life is the time in which the concentration of a reactant becomes half of its initial concentration.

It is written as t₁/₂.

Half-Life of Zero Order Reaction

For zero order reaction:

t₁/₂ = [R]₀ / 2k

The half-life of a zero order reaction depends on initial concentration.

Half-Life of First Order Reaction

For first order reaction:

t₁/₂ = 0.693 / k

The half-life of a first order reaction is independent of initial concentration.

This is a very important point for numericals.

Pseudo First Order Reactions

Some reactions are actually higher order but behave like first order reactions.

This happens when one reactant is present in large excess. Its concentration remains almost constant during the reaction.

Example: Hydrolysis of ethyl acetate.

CH₃COOC₂H₅ + H₂O → CH₃COOH + C₂H₅OH

Water is taken in large excess. So, the rate depends mainly on ethyl acetate.

Inversion of cane sugar is another pseudo first order reaction.

C₁₂H₂₂O₁₁ + H₂O → C₆H₁₂O₆ + C₆H₁₂O₆

Rate = k[C₁₂H₂₂O₁₁]

Temperature Dependence and Arrhenius Equation

Most chemical reactions become faster when temperature increases.

For many reactions, a rise of 10° nearly doubles the rate constant.

The temperature dependence of rate constant is explained by the Arrhenius equation.

k = Ae-Ea/RT

Here:

Taking natural log:

ln k = -Ea/RT + ln A

A plot of ln k against 1/T gives a straight line.

Slope = -Ea/R

Intercept = ln A

Activation energy is the minimum energy needed to form the activated complex.

Effect of Catalyst on Reaction Rate

A catalyst increases the rate of reaction without undergoing permanent chemical change.

It provides an alternative pathway with lower activation energy.

Lower activation energy means more reactant molecules can cross the energy barrier.

Important points:

Example:

2KClO₃ → 2KCl + 3O₂

MnO₂ acts as a catalyst in this reaction.

Collision Theory of Chemical Reactions

Collision theory explains how reactions occur at the molecular level.

According to this theory, reactant molecules must collide to form products.

For a bimolecular elementary reaction:

A + B → Products

Rate = ZAB e-Ea/RT

Here, ZAB is the collision frequency.

But every collision does not form products. A collision becomes effective only when molecules have:

• sufficient energy
• proper orientation

So, the modified form is:

Rate = PZAB e-Ea/RT

Here, P is the steric factor.

Collision theory explains why temperature, concentration and catalyst affect reaction rate.

Quick Formula Table for Chemical Kinetics

Important Terms in Chemical Kinetics

Useful Links for Class 12 Chemistry