CBSE Class 9 Science Atoms and Molecules Revision Notes 2026-27

The concept of atoms and molecules forms the bedrock of modern chemistry. Long before laboratory instruments could detect individual particles, ancient thinkers reasoned about the nature of matter.

Maharishi Kanad, the Indian philosopher, proposed that if matter is divided repeatedly, we eventually reach the smallest indivisible particle which he called paramanu. Centuries later, John Dalton formalised this idea scientifically, naming the fundamental particle the atom (from the Greek word meaning "indivisible").

Dalton also recognised that atoms typically exist in combined forms called molecules, and that different combinations of atoms give rise to the diverse materials we observe in nature.

All matter is composed of tiny particles called atoms and molecules. Their differing properties at the atomic level explain why different substances behave differently in the physical world.

Laws of Chemical Combination

Before Dalton built his atomic theory, chemists established key experimental laws that described how substances react. These Laws of Chemical Combination laid the quantitative foundation for understanding atoms.

(a) Law of Conservation of Mass

Proposed by: Antoine Lavoisier, 1774

Statement: Matter can neither be created nor destroyed during a chemical reaction. The total mass of the reactants always equals the total mass of the products.

Mathematical Expression:

If reactants A (mass = a g) and B (mass = b g) produce products C (mass = c g) and D (mass = d g):

a + b = c + d

Verified Example:

When calcium carbonate (CaCO₃) is heated, it decomposes into calcium oxide (CaO) and carbon dioxide (CO₂):

CaCO₃ + Heat → CaO + CO₂ 100 g 56 g 44 g Mass of reactants = 100 g Mass of products = 56 + 44 = 100 g ✓

The mass is conserved no matter is lost or gained in the reaction.

(b) Law of Constant Proportions (Law of Definite Proportions)

Proposed by: Joseph Louis Proust, 1779

Statement: A chemical compound always contains the same elements combined in the same fixed ratio by mass, regardless of its source or method of preparation.

Example:

Water (H₂O), whether sourced from a river, a well, or synthesised in a laboratory, always contains:

• 1 part hydrogen by mass
• 8 parts oxygen by mass

This fixed 1:8 ratio of hydrogen to oxygen never changes in pure water.

Important Note: The converse is not always true when the same elements combine in the same proportion, it does not guarantee the same compound will always be formed.

Dalton's Atomic Theory

In 1808, John Dalton proposed his atomic theory of matter the first scientific attempt to explain chemical behaviour in terms of atoms. The key postulates are:

How Dalton's Theory Connects to the Laws:

• Postulate (iii) "atoms cannot be created or destroyed" directly explains the Law of Conservation of Mass.
• Postulates (iv) and (viii) "fixed number and kind of atoms in a compound" explain the Law of Constant Proportions.

Dalton's atomic theory was the first modern framework to describe matter in terms of atoms and remains one of the most significant intellectual achievements in the history of science.

Atoms

An atom is the smallest particle of an element that can take part in a chemical reaction.

Facts about atoms:

• Atoms of most elements are highly reactive and do not exist in a free (isolated) state.
• They exist in combination with atoms of the same or different elements.
• Atomic size is measured in nanometres (nm), where 1 nm = 10⁻⁹ m.
• Atoms are too small to be seen even under the most powerful optical microscope.
• The hydrogen atom is the smallest of all atoms, with an atomic radius of 0.037 nm.

Symbols of Elements

A chemical symbol is a short, universally accepted notation for an element.

System Proposed by: J.J. Berzelius

Rules for Writing Symbols:

• The symbol is either the first letter (capitalised) of the element's English or Latin name, or the first letter plus one other letter.
• In a two-letter symbol, the first letter is always capitalised and the second letter is always lowercase.

Examples:

Significance of a Chemical Symbol:

A symbol is far more than a shorthand label. It simultaneously represents:

1. The name of the element.
2. One atom of the element.
3. One mole (6.023 × 10²³ atoms) of the element.
4. The atomic mass (in grams) of the element.

Example: The symbol H represents the element hydrogen, one atom of hydrogen, one mole of hydrogen atoms, and one gram of hydrogen.

Atomic Mass

The actual masses of atoms are inconveniently small. For example, one hydrogen atom has a mass of just 1.673 × 10⁻²⁴ grams. To avoid such unwieldy numbers, atomic mass is expressed relative to a standard.

Standard Reference: The Carbon-12 atom (⁶¹²C) an atom with 6 protons and 6 neutrons is assigned an atomic mass of exactly 12 atomic mass units (12 u).

Definition of Atomic Mass:

Atomic mass expresses how many times heavier an atom of an element is compared to 1/12th the mass of a Carbon-12 atom.

1 Atomic Mass Unit (u):

1 u = 1/12 × mass of a Carbon-12 atom = 1.6605 × 10⁻²⁴ g

How Atoms Exist in Nature

Most atoms do not exist freely. They exist either as:

1. Molecules groups of two or more atoms chemically bonded together.
2. Ions atoms or groups of atoms carrying an electrical charge.

The only exceptions are the noble gases (helium, neon, argon, krypton, xenon, radon), which exist as single atoms because they are chemically inert.

Molecules and Atomicity

A molecule is an electrically neutral group of two or more atoms held together by chemical bonds. It is the smallest particle of a substance that can exist independently and retains all the chemical properties of that substance.

Rule: Every compound is a molecule, but every molecule is not necessarily a compound.

Types of Molecules

1. Molecules of Elements

Formed by combining atoms of the same element.

2. Molecules of Compounds

Formed by combining atoms of different elements.

Atomicity

Atomicity is the number of atoms present in one molecule of an element or compound.

• Noble gases: Atomicity = 1 (monatomic)
• H₂, O₂, N₂: Atomicity = 2 (diatomic)
• O₃: Atomicity = 3 (triatomic)
• P₄: Atomicity = 4
• S₈: Atomicity = 8

Molecular Mass

The molecular mass of a substance is the relative mass of one molecule of the substance compared to 1/12th the mass of a Carbon-12 atom.

How to Calculate: Molecular mass = Sum of atomic masses of all atoms in one molecule.

Example Sulphuric Acid (H₂SO₄):

Molecular mass of H₂SO₄ = (2 × 1) + (1 × 32) + (4 × 16) = 2 + 32 + 64 = 98 u

Example Carbon Dioxide (CO₂):

Molecular mass of CO₂ = (1 × 12) + (2 × 16) = 12 + 32 = 44 u

Ions

An ion is an atom or group of atoms that carries an electrical charge due to loss or gain of electrons. Neutral atoms have equal numbers of protons (+) and electrons (−), so they carry no net charge.

(a) Cations Positively Charged Ions

Formed when an atom loses electrons. The positive charge equals the number of electrons lost.

All metal elements form cations.
A cation is always smaller in size than its parent neutral atom.

(b) Anions Negatively Charged Ions

Formed when an atom gains electrons. The negative charge equals the number of electrons gained.

All non-metal elements form anions (except hydrogen).
An anion is always larger in size than its parent neutral atom.

Monoatomic vs Polyatomic Ions

Ionic Compounds and Formula Mass

Ionic compounds consist of cations and anions held together by strong electrostatic (ionic) bonds. The overall charge on any ionic compound is always zero.

Molecular formula cannot be determined for ionic compounds (since they exist as vast lattice structures of ions, not discrete molecules). Instead, they are represented by formula units the simplest electrically neutral combination of ions.

Example Sodium Chloride (NaCl):

• Na⁺ and Cl⁻ in equal numbers → simplest formula = NaCl

Formula Mass

The formula mass is the sum of the atomic masses of all atoms in the formula unit of the ionic compound.

Example Potassium Carbonate (K₂CO₃):

Formula mass of K₂CO₃ = (2 × 39) + (1 × 12) + (3 × 16) = 78 + 12 + 48 = 138 u

Chemical Formulas

Steps for Writing the Formula of a Molecular Compound

1. Write the symbols of the elements in the compound.
2. Write the valency below each symbol.
3. Exchange the valencies (valency of first element becomes the subscript of the second, and vice versa).
4. If the subscripts share a common factor, simplify by dividing through.

Example Carbon Dioxide (CO₂):

Symbols: C O Valencies: 4 2 Exchange → C₂O₄ → divide by common factor 2 → CO₂

Steps for Writing the Formula of an Ionic Compound

1. Write the cation on the left, anion on the right.
2. Write each ion's valency (charge) below.
3. Exchange the charges as subscripts.
4. Ensure the compound is electrically neutral.

Example Aluminium Sulphate:

Symbols: Al SO₄ Charge: +3 −2 Exchange → Al₂(SO₄)₃

Naming Rules for Compounds

Mole Concept

Atoms and molecules are too small to count individually. Chemists use a special counting unit called the mole, similar to how a "dozen" means 12 items.

Definition of a Mole

1 mole is the amount of a substance that contains as many entities (atoms, molecules, or ions) as there are atoms in exactly 12 grams of Carbon-12.

This number is called Avogadro's Number (Nₐ):

Nₐ = 6.023 × 10²³

1 mole of any substance contains 6.023 × 10²³ particles.

The Mole as a Bridge

Gram Atomic Mass and Gram Molecular Mass

• Gram Atomic Mass: Atomic mass of an element expressed in grams.
  Example: Atomic mass of N = 14 u → Gram atomic mass = 14 g
• Gram Molecular Mass: Molecular mass expressed in grams.
  Example: Molecular mass of O₂ = 32 u → Gram molecular mass = 32 g
• Molar Mass: Mass of 1 mole of a substance, expressed in g/mol.

Atoms and Molecules Solved Examples

Example 1 Applying the Law of Conservation of Mass

Problem: 4 g of magnesium reacts completely with 2.67 g of oxygen. What mass of magnesium oxide is formed?

Solution:

By the Law of Conservation of Mass:

Mass of reactants = Mass of products 4 g (Mg) + 2.67 g (O) = Mass of MgO Mass of MgO = 6.67 g ✓

Example 2 Calculating Molecular Mass

Problem: Find the molecular mass of glucose (C₆H₁₂O₆).

Solution:

Molecular mass = (6 × 12) + (12 × 1) + (6 × 16) = 72 + 12 + 96 = 180 u

Example 3 Mole Calculation (Mass to Moles)

Problem: Calculate the number of moles in 46 g of sodium (Na). (Atomic mass of Na = 23 u)

Solution:

Number of moles = Mass ÷ Gram atomic mass = 46 ÷ 23 = 2 moles

Example 4 Moles to Number of Particles

Problem: How many atoms are present in 2 moles of oxygen atoms?

Solution:

Number of atoms = Moles × Avogadro's number = 2 × 6.023 × 10²³ = 1.2046 × 10²⁴ atoms

Example 5 Mass of a Given Number of Molecules

Problem: Calculate the mass of 3.011 × 10²³ molecules of water (H₂O). (Molecular mass of H₂O = 18 u)

Solution:

Step 1: Find number of moles.

Moles = N ÷ Nₐ = (3.011 × 10²³) ÷ (6.023 × 10²³) = 0.5 moles

Step 2: Calculate mass.

Mass = Moles × Gram molecular mass = 0.5 × 18 = 9 g