CBSE Class 11 Chemistry Revision Notes for All Chapters

CBSE Class 11 Chemistry Revision Notes explain the core concepts, formulas, reactions and numerical methods included in the current syllabus. The notes cover all nine units across NCERT Chemistry Parts I and II for the CBSE 2026–27 academic year.

Class 11 Chemistry builds the foundation required to understand physical, inorganic and organic chemistry. The course begins with matter, measurements and atomic structure before moving to periodic properties, chemical bonding, thermodynamics, equilibrium, redox reactions and carbon compounds.

Use these CBSE Class 11 Chemistry Revision Notes to review essential definitions, formulas, trends and reaction principles. The current NCERT textbooks contain nine units, with six in Chemistry Part I and three in Chemistry Part II.

Key Takeaways

• Nine units: The current course begins with Some Basic Concepts of Chemistry and ends with Hydrocarbons.
• 70 theory marks: Organic Chemistry carries 11 marks, while Hydrocarbons carries 10 marks.
• 30 practical marks: Practical assessment includes volumetric analysis, salt analysis, experiments, project work and records.
• Connected concepts: Atomic structure supports periodicity and bonding, while organic principles support Hydrocarbons.

Access CBSE Class 11 Chemistry Revision Notes in 30 Minutes

Revise the course in three parts:

• First 10 minutes: Mole concept, atomic structure, periodic trends and chemical bonding
• Next 10 minutes: Thermodynamics, chemical equilibrium, ionic equilibrium and redox reactions
• Final 10 minutes: Organic nomenclature, electronic effects, reaction intermediates and hydrocarbons

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Chapter-Wise CBSE Class 11 Chemistry Notes

The Chapter-wise Class 11 Chemistry Notes follow the current NCERT sequence for 2026–27.

Some Basic Concepts of Chemistry Notes

Chemistry studies the composition, structure, properties and transformations of matter.

The first unit introduces the numerical language used throughout Chemistry.

Nature and Classification of Matter

Matter occupies space and has mass.

It may be classified as:

• Elements: Contain one type of atom
• Compounds: Contain two or more elements chemically combined
• Mixtures: Contain substances physically combined

Mixtures may be homogeneous or heterogeneous.

Physical Quantities and SI Units

Every measurement contains:

• A numerical value
• A unit

Common SI units include:

Accuracy shows how close a measurement is to the true value. Precision shows how closely repeated measurements agree with one another.

Significant Figures

Significant figures indicate the precision of a measurement.

Important rules include:

• All non-zero digits are significant.
• Zeros between non-zero digits are significant.
• Leading zeros are not significant.
• Trailing zeros after a decimal point are significant.
• Exact numbers have unlimited significant figures.

For multiplication and division, the answer should contain the same number of significant figures as the least precise value.

Laws of Chemical Combination

Law of Conservation of Mass: Mass is neither created nor destroyed during a chemical reaction.

Law of Definite Proportions: A compound always contains the same elements in the same proportion by mass.

Law of Multiple Proportions: When two elements form more than one compound, the masses of one element combining with a fixed mass of the other occur in simple whole-number ratios.

Gay-Lussac’s Law of Gaseous Volumes: Gases react in simple whole-number ratios by volume under similar conditions.

Avogadro’s Law: Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules.

Mole Concept

One mole contains (6.022 \times 10^{23}) representative particles.

This value is called Avogadro’s constant.

Important relationships include:

Number of moles = Given mass ÷ Molar mass

Number of particles = Number of moles × Avogadro’s constant

Molar mass = Mass of one mole of a substance

Percentage Composition

The percentage of an element in a compound is:

Mass percentage = Mass of element in one mole of compound ÷ Molar mass of compound × 100

Percentage composition helps determine empirical and molecular formulas.

Empirical and Molecular Formula

The empirical formula gives the simplest whole-number ratio of atoms.

The molecular formula gives the actual number of each type of atom.

Molecular formula = Empirical formula × n

Here:

n = Molar mass ÷ Empirical formula mass

Stoichiometry

Stoichiometry uses a balanced chemical equation to calculate quantities of reactants and products.

The usual steps are:

1. Write and balance the equation.
2. Convert the given quantity into moles.
3. Apply the mole ratio.
4. Convert the result into the required unit.

The reactant consumed first is called the limiting reagent. It determines the maximum amount of product formed.

Structure of Atom Notes

Atoms contain electrons, protons and neutrons.

The atomic number represents the number of protons. The mass number equals the total number of protons and neutrons.

Important Atomic Terms

Atomic Models

Thomson’s model: Electrons were embedded in a positively charged sphere.

Rutherford’s model: The atom contains a small, dense and positively charged nucleus.

Bohr’s model: Electrons move in fixed energy levels around the nucleus.

Bohr’s model explained the hydrogen spectrum but could not fully describe multi-electron atoms.

Electromagnetic Radiation

The relationship between wavelength and frequency is:

c = νλ

where:

• (c) is the speed of light
• (ν) is frequency
• (λ) is wavelength

Planck proposed that energy is emitted or absorbed in packets called quanta.

E = hν

Dual Nature of Matter

De Broglie proposed that moving particles have wave-like properties.

λ = h ÷ mv

Heisenberg’s uncertainty principle states that the exact position and momentum of an electron cannot be known simultaneously.

Quantum Numbers

For a given value of (n), the possible values of (l) range from 0 to (n-1).

The subshells are represented as:

• (l = 0): s
• (l = 1): p
• (l = 2): d
• (l = 3): f

Rules for Electronic Configuration

Aufbau principle: Electrons fill lower-energy orbitals first.

Pauli exclusion principle: No two electrons in an atom can have the same set of four quantum numbers.

Hund’s rule: Electrons occupy degenerate orbitals singly before pairing.

Classification of Elements and Periodicity in Properties Notes

The modern periodic law states that the properties of elements are periodic functions of their atomic numbers.

The periodic table has:

• Seven periods
• Eighteen groups
• s, p, d and f blocks

Effective Nuclear Charge

Effective nuclear charge is the net positive charge experienced by an electron after accounting for shielding by other electrons.

It generally increases across a period.

Atomic Radius

Atomic radius generally:

• Decreases across a period
• Increases down a group

Across a period, nuclear charge increases while electrons enter the same shell.

Down a group, new electron shells are added.

Ionic Radius

Cations are smaller than their parent atoms because they lose electrons.

Anions are larger than their parent atoms because added electrons increase repulsion.

In an isoelectronic series, size decreases as nuclear charge increases.

Ionisation Enthalpy

Ionisation enthalpy is the energy required to remove an electron from an isolated gaseous atom.

It generally:

• Increases across a period
• Decreases down a group

Exceptions occur because of stable half-filled and fully filled subshells.

Electron Gain Enthalpy

Electron gain enthalpy is the enthalpy change when an electron is added to an isolated gaseous atom.

More negative values indicate a greater tendency to accept an electron.

Electronegativity

Electronegativity is the ability of an atom in a molecule to attract the shared electron pair.

It generally increases across a period and decreases down a group.

Valency and Oxidation State

Valency depends on the number of electrons lost, gained or shared during bond formation.

Across a period, valency usually increases from one to four and then decreases towards zero.

Chemical Bonding and Molecular Structure Notes

Atoms form chemical bonds to reach a more stable energy state.

Octet Rule

According to the octet rule, atoms tend to gain, lose or share electrons to obtain eight electrons in their valence shell.

The rule has exceptions, including:

• Incomplete octets
• Odd-electron molecules
• Expanded octets

Ionic Bond

An ionic bond forms through electron transfer and electrostatic attraction between oppositely charged ions.

Ionic compounds generally have:

• High melting points
• Crystalline structures
• Electrical conductivity in molten or aqueous form
• Solubility in polar solvents

Lattice enthalpy affects the stability of ionic compounds.

Covalent Bond

A covalent bond forms when atoms share electron pairs.

It may be:

• Single
• Double
• Triple
• Polar
• Non-polar

Lewis Structures and Formal Charge

Lewis structures represent valence electrons through dots and bonds through lines.

Formal charge is:

Formal charge = Valence electrons − Non-bonding electrons − ½(Bonding electrons)

The preferred structure generally has minimum formal charges and places negative charge on the more electronegative atom.

VSEPR Theory

VSEPR theory predicts molecular shape by assuming that electron pairs around the central atom repel one another.

The repulsion order is:

Lone pair–lone pair > Lone pair–bond pair > Bond pair–bond pair

Hybridisation

Hybridisation is the mixing of atomic orbitals of similar energy to form equivalent hybrid orbitals.

Molecular Orbital Theory

Atomic orbitals combine to form bonding and antibonding molecular orbitals.

Bond order is:

Bond order = ½(Number of bonding electrons − Number of antibonding electrons)

A positive bond order indicates a stable molecule.

Paramagnetic substances contain unpaired electrons. Diamagnetic substances contain only paired electrons.

Hydrogen Bonding

Hydrogen bonding occurs when hydrogen bonded to a highly electronegative atom is attracted to another electronegative atom.

It may be intermolecular or intramolecular.

Hydrogen bonding affects boiling point, solubility and molecular shape.

Thermodynamics Notes

Thermodynamics studies energy changes accompanying physical and chemical processes.

System and Surroundings

The part of the universe under study is the system. Everything outside it forms the surroundings.

Systems may be:

• Open
• Closed
• Isolated

State Functions

State functions depend only on the initial and final states.

Examples include:

• Internal energy
• Enthalpy
• Entropy
• Pressure
• Volume
• Temperature

Heat and work are path functions.

First Law of Thermodynamics

The first law expresses conservation of energy.

ΔU = q + w

Here:

• (ΔU) is change in internal energy
• (q) is heat supplied to the system
• (w) is work done on the system

For expansion work:

w = −PₑₓΔV

Enthalpy

Enthalpy is:

H = U + PV

At constant pressure:

ΔH = qₚ

For gaseous reactions:

ΔH = ΔU + Δn₉RT

Types of Enthalpy Change

Important enthalpy changes include:

• Enthalpy of formation
• Enthalpy of combustion
• Enthalpy of neutralisation
• Enthalpy of atomisation
• Bond enthalpy
• Enthalpy of solution
• Enthalpy of hydration

Hess’s Law

Hess’s law states that the total enthalpy change of a reaction is independent of the path followed.

It allows unknown enthalpy changes to be calculated using known reactions.

Spontaneity and Entropy

A spontaneous process can occur without continuous external assistance.

Entropy measures the degree of energy dispersal or randomness.

Spontaneity depends on both enthalpy and entropy.

Gibbs Energy

ΔG = ΔH − TΔS

• (ΔG < 0): Spontaneous process
• (ΔG > 0): Non-spontaneous process
• (ΔG = 0): Equilibrium

Equilibrium Notes

Equilibrium is established when opposing processes occur at equal rates.

It is dynamic because both forward and reverse processes continue.

Physical Equilibrium

Examples include:

• Solid–liquid equilibrium
• Liquid–vapour equilibrium
• Dissolution equilibrium
• Gas–solution equilibrium

Chemical Equilibrium

For a reaction:

aA + bB ⇌ cC + dD

The equilibrium constant is:

Kc = [C]ᶜ[D]ᵈ ÷ [A]ᵃ[B]ᵇ

Pure solids and pure liquids are omitted from equilibrium-constant expressions.

For gaseous reactions:

Kp = Kc(RT)Δn

Reaction Quotient

The reaction quotient has the same form as the equilibrium constant but may be calculated before equilibrium.

• (Q < K): Forward reaction is favoured
• (Q > K): Reverse reaction is favoured
• (Q = K): System is at equilibrium

Le Chatelier’s Principle

When a system at equilibrium is disturbed, it shifts in a direction that reduces the effect of the disturbance.

Changes may involve:

• Concentration
• Pressure
• Volume
• Temperature

A catalyst speeds up both directions equally. It does not change the equilibrium constant or equilibrium composition.

Acids and Bases

Arrhenius acid: Produces (H^+) in water.

Arrhenius base: Produces (OH^-) in water.

Brønsted acid: Proton donor.

Brønsted base: Proton acceptor.

Lewis acid: Electron-pair acceptor.

Lewis base: Electron-pair donor.

Ionic Product of Water

At 298 K:

Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴

pH and pOH

pH = −log[H⁺]

pOH = −log[OH⁻]

At 298 K:

pH + pOH = 14

Buffer Solutions

A buffer resists changes in pH when small amounts of acid or base are added.

An acidic buffer contains a weak acid and its salt. A basic buffer contains a weak base and its salt.

Solubility Product

For a sparingly soluble salt, the solubility product is the product of ionic concentrations raised to their stoichiometric powers.

Precipitation occurs when the ionic product exceeds the solubility product.

Redox Reactions Notes

Redox reactions involve oxidation and reduction occurring together.

Oxidation and Reduction

Oxidation may involve:

• Addition of oxygen
• Removal of hydrogen
• Loss of electrons
• Increase in oxidation number

Reduction may involve:

• Removal of oxygen
• Addition of hydrogen
• Gain of electrons
• Decrease in oxidation number

Oxidising and Reducing Agents

An oxidising agent causes oxidation and is itself reduced.

A reducing agent causes reduction and is itself oxidised.

Oxidation Number Rules

• The oxidation number of a free element is zero.
• The oxidation number of a monatomic ion equals its charge.
• Oxygen is generally −2.
• Hydrogen is generally +1 with non-metals and −1 with metals.
• Fluorine is always −1.
• The sum of oxidation numbers equals the charge on the species.

Balancing Redox Reactions

Redox equations may be balanced by:

• Oxidation-number method
• Ion-electron method

The final equation must balance atoms and total charge.

Disproportionation

In a disproportionation reaction, the same substance is oxidised and reduced.

Organic Chemistry: Some Basic Principles and Techniques Notes

Organic chemistry studies carbon compounds and their reactions.

Carbon forms a large number of compounds because of tetravalency and catenation.

Classification of Organic Compounds

Organic compounds may be:

• Acyclic
• Cyclic
• Homocyclic
• Heterocyclic
• Alicyclic
• Aromatic
• Saturated
• Unsaturated

Functional Groups

A functional group determines the characteristic reactions of an organic compound.

Common functional groups include:

IUPAC Nomenclature

The general process is:

1. Select the longest carbon chain.
2. Identify the principal functional group.
3. Number the chain to give the lowest locants.
4. Name substituents.
5. Arrange prefixes alphabetically.
6. Add the correct suffix.

Isomerism

Isomers have the same molecular formula but different arrangements.

Structural isomerism includes:

• Chain isomerism
• Position isomerism
• Functional isomerism
• Metamerism

Stereoisomerism arises from different spatial arrangements.

Bond Fission

Homolytic fission: Each atom takes one electron, forming free radicals.

Heterolytic fission: One atom takes both bonding electrons, forming ions.

Reaction Intermediates

Important intermediates include:

• Carbocations
• Carbanions
• Free radicals

Their stability depends on structure and electronic effects.

Electronic Effects

Inductive effect: Permanent displacement of sigma electrons due to electronegativity differences.

Resonance effect: Delocalisation of electrons across conjugated structures.

Hyperconjugation: Delocalisation involving a sigma bond adjacent to a multiple bond or positive centre.

Electromeric effect: Temporary complete transfer of pi electrons in the presence of a reagent.

Types of Organic Reactions

• Substitution
• Addition
• Elimination
• Rearrangement
• Oxidation
• Reduction

Purification Methods

Organic compounds may be purified through:

• Sublimation
• Crystallisation
• Distillation
• Differential extraction
• Chromatography

Hydrocarbons Notes

Hydrocarbons contain only carbon and hydrogen.

They may be aliphatic or aromatic.

Alkanes

Alkanes are saturated hydrocarbons with the general formula:

CₙH₂ₙ₊₂

Carbon atoms are (sp^3)-hybridised.

Preparation of Alkanes

Methods include:

• Hydrogenation of alkenes and alkynes
• Wurtz reaction
• Decarboxylation
• Kolbe’s electrolytic method
• Reduction of alkyl halides

Reactions of Alkanes

Alkanes undergo:

• Combustion
• Halogenation
• Isomerisation
• Aromatisation
• Pyrolysis

Halogenation proceeds through a free-radical chain mechanism involving initiation, propagation and termination.

Alkenes

Alkenes contain at least one carbon-carbon double bond.

Their general formula is:

CₙH₂ₙ

The double-bonded carbon atoms are (sp^2)-hybridised.

Preparation of Alkenes

Common methods include:

• Dehydration of alcohols
• Dehydrohalogenation of alkyl halides
• Dehalogenation of vicinal dihalides

Reactions of Alkenes

Alkenes mainly undergo electrophilic addition.

Important reactions include:

• Hydrogenation
• Halogenation
• Addition of hydrogen halides
• Hydration
• Oxidation
• Ozonolysis
• Polymerisation

Markovnikov’s Rule

During the addition of an unsymmetrical reagent to an unsymmetrical alkene, hydrogen generally attaches to the carbon already carrying more hydrogen atoms.

In the presence of peroxide, HBr may show anti-Markovnikov addition.

Alkynes

Alkynes contain at least one carbon-carbon triple bond.

Their general formula is:

CₙH₂ₙ₋₂

The triple-bonded carbon atoms are (sp)-hybridised.

Terminal alkynes show acidic behaviour because the (sp)-hybridised carbon has greater s-character.

Alkynes undergo addition reactions involving hydrogen, halogens, hydrogen halides and water.

Aromatic Hydrocarbons

Benzene is a planar cyclic molecule with delocalised pi electrons.

It shows unusual stability due to resonance.

Electrophilic Substitution in Benzene

Important reactions include:

• Nitration
• Sulphonation
• Halogenation
• Friedel-Crafts alkylation
• Friedel-Crafts acylation

Substituents already present on the benzene ring influence the position of further substitution.

Some direct incoming groups towards the ortho and para positions, while others direct towards the meta position.

Important Class 11 Chemistry Formulas

CBSE Class 11 Chemistry Unit-Wise Marks 2026–27

The remaining 30 marks are assigned to practical work according to the prescribed school assessment structure.