Chapter 3 — Atomic Foundation of Matter | Class 9 CBSE Chemistry 2026-27

Section 3.2 – 3.3

Laws of Chemical Combination

Careful experiments showed that substances combine in fixed, consistent amounts. These regularities — the laws of chemical combination — were strong evidence that matter is made of tiny particles (atoms).

❖ Law of Conservation of Mass (Lavoisier, 1789)

In a chemical reaction, matter can neither be created nor destroyed — the total mass of the reactants equals the total mass of the products. Atoms are only rearranged.

Heating 50 g of calcium carbonate gives 28 g of calcium oxide + 22 g of carbon dioxide — total mass unchanged

❖ Law of Constant (Definite) Proportions (Proust)

A pure chemical compound always contains the same elements combined in the same fixed proportion by mass, no matter how it is prepared. Magnesium oxide always has Mg : O = 3 : 2; water always has H : O = 1 : 8.

Worked ExampleConservation of mass check

5.3 g sodium carbonate + 6.0 g ethanoic acid → 8.2 g sodium ethanoate + 2.2 g CO₂ + 0.9 g water. Reactants = 5.3 + 6.0 = 11.3 g; products = 8.2 + 2.2 + 0.9 = 11.3 g. Equal → law obeyed.

⚠️

Open vessel trap: if a gas escapes (e.g. burning ethanol or decomposing CaCO₃ in an open dish), the leftover mass seems to fall — but mass is still conserved; the gas simply left. Always reason as if in a closed system.

Section 3.4

Dalton's Atomic Theory (1808)

Dalton explained the laws of combination by proposing that matter is made of indivisible atoms.

Matter is made of small indivisible particles called atoms.
Atoms of an element are identical in mass and chemical properties.
Atoms of different elements are different in all respects.
Atoms combine in simple whole-number ratios to form compounds.
Atoms can neither be created nor destroyed in a chemical reaction.

Explains conservation of mass

Since atoms are neither created nor destroyed, the kinds and numbers of atoms in the products equal those in the reactants → mass is conserved.

Explains constant proportions

Because the number and kind of atoms in a compound are fixed, the elements always combine in the same fixed mass ratio.

⚠️

Drawbacks: atoms are not indivisible (they contain electrons, protons, neutrons); isotopes show atoms of one element can differ in mass; and isobars show atoms of different elements can have the same mass.

💡 Did You Know?

The red pigment cinnabar (called hingula in the Indian subcontinent) is mercury sulphide, HgS. Heating it splits it into mercury and sulphur — and the two recombine to re-form cinnabar, an early hint of the laws of chemical combination.

Section — Formation of Molecules

Why and How Atoms Combine

All atoms except the noble gases are reactive, so they don't exist free — they join into molecules or ions. They do this to complete their outermost shell (octet, or duplet for the K shell).

❖ Molecule

The smallest unit of a pure substance (element or compound) that can exist independently and retains all its physical and chemical properties. Atoms are held together in a molecule by a chemical bond.

Sharing electrons → Covalent bond

Two atoms share one or more pairs of valence electrons. Common between non-metals.

Transferring electrons → Ionic bond

One atom transfers electrons to another, forming oppositely charged ions that attract. Common between a metal and a non-metal.

Section 3.5

Covalent Bonding

When atoms share electron pairs, the number of shared pairs sets the bond type: single, double or triple.

Number of shared electron pairs decides the bond: H₂ (single), O₂ (double), N₂ (triple)

Molecule	Each atom needs	Pairs shared	Bond
H₂ / Cl₂	1 electron	1	Single
O₂	2 electrons	2	Double
N₂	3 electrons	3	Triple

Covalent compoundsWater, ammonia, carbon dioxide

H₂O

Oxygen (2,6) shares 1 electron with each of two H atoms → two single bonds.

NH₃

Nitrogen (2,5) shares 1 electron with each of three H atoms → three single bonds.

CO₂

Carbon (2,4) shares 2 pairs with each oxygen → two double bonds.

❖ Naming covalent compounds

Use prefixes mono(1), di(2), tri(3), tetra(4), penta(5), hexa(6)... and end the second element in -ide; mono- is dropped for the first element. CO = carbon monoxide, CO₂ = carbon dioxide, PCl₅ = phosphorus pentachloride, SF₆ = sulphur hexafluoride.

Atomicity

Atomicity of Elements

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

Type	Atoms per molecule	Examples
Monoatomic	1	All noble gases (He, Ne, Ar...); all metals
Diatomic	2	H₂, O₂, N₂, Cl₂, F₂, Br₂, I₂
Triatomic	3	Ozone (O₃)
Polyatomic	4 or more	Phosphorus (P₄, tetra-atomic); Sulphur (S₈, octa-atomic)

✓

Notation: O₂ = one molecule of oxygen (2 atoms bonded); 4O₂ = four oxygen molecules; 2O = two separate oxygen atoms.

Section 3.6

Ionic Bonding

Instead of sharing, one atom transfers electrons to another. The atom that loses electrons becomes a positive ion (cation); the one that gains becomes a negative ion (anion). Their electrostatic attraction is the ionic (electrovalent) bond.

Sodium transfers its 1 valence electron to chlorine; the resulting Na⁺ and Cl⁻ ions attract to form NaCl

Cation (positive)

Formed by losing electrons (more protons than electrons). Na⁺ (loses 1), Mg²⁺ (loses 2), Al³⁺ (loses 3). Metals form cations.

Anion (negative)

Formed by gaining electrons (more electrons than protons). Cl⁻ (gains 1), O²⁻ (gains 2), N³⁻ (gains 3). Non-metals form anions.

❖ Ionic compounds

Made of metal cations and non-metal anions held by strong electrostatic forces (ionic bonds). The compound is overall neutral. Examples: NaCl, MgCl₂, CaO, CuSO₄. A radical is a charged group of atoms acting as a single ion (e.g. SO₄²⁻, NO₃⁻, NH₄⁺).

Section 3.7

Writing Chemical Formulae

A formula must be electrically neutral — positive and negative valencies must balance. The criss-cross method makes this easy.

Criss-cross method: Al³⁺ with SO₄²⁻ gives Al₂(SO₄)₃ — a radical taking a subscript > 1 is bracketed

Compound	Ions	Formula
Calcium chloride	Ca²⁺, Cl⁻	CaCl₂
Magnesium sulphate	Mg²⁺, SO₄²⁻	MgSO₄
Aluminium sulphate	Al³⁺, SO₄²⁻	Al₂(SO₄)₃
Potassium dichromate	K⁺, Cr₂O₇²⁻	K₂Cr₂O₇
Calcium bicarbonate	Ca²⁺, HCO₃⁻	Ca(HCO₃)₂

Section 3.8

Ionic vs Covalent Compounds

Their very different properties come from the difference in bonding — strong ionic attractions vs weaker forces between molecules.

Property	Ionic compounds	Covalent compounds
Made of	Ions (in a crystal lattice)	Molecules
Physical state	Hard, brittle crystalline solids	Gases, liquids or soft solids
Melting / boiling point	High	Low
Solubility	Usually soluble in water	Usually insoluble in water, soluble in organic solvents
Electrical conductivity	Conduct when molten or dissolved	Do not conduct
Examples	NaCl, MgCl₂, CaO	H₂O, CO₂, NH₂

💡 Did You Know?

In the conductivity test, salt and copper-sulphate solutions light the bulb (free ions carry charge), but a sugar solution does not — sugar dissolves as neutral molecules, not ions. Solid ionic compounds don't conduct either, because their ions are locked in place.

Section 3.9 – 3.10

Molecular Mass & Formula Unit Mass

Both are simply the sum of the atomic masses of all atoms in the formula — "molecular mass" for covalent compounds, "formula unit mass" for ionic compounds.

How to calculate

Molecular / Formula mass = sum of (atomic mass × number of atoms) for every element

Numerically the same idea for both; we just use "formula unit mass" for ionic compounds since they have ions, not molecules.

Example ACO₂ (C=12, O=16)

= 12 + (2 × 16) = 12 + 32 = 44 u

Example BGlucose C₆H₁₂O₆ (C=12, H=1, O=16)

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

Example CCa(HCO₃)₂ (Ca=40, H=1, C=12, O=16)

= 40 + (2×1) + (2×12) + (6×16) = 40 + 2 + 24 + 96 = 162 u

Solved Examples

Worked Sample Problems

Model answers in the CBSE style.

Example 1Hydrogen and oxygen combine 1:8 by mass to form water. Oxygen needed for 3 g of hydrogen?

Ratio H : O = 1 : 8, so 1 g H needs 8 g O. For 3 g H → 8 × 3 = 24 g of oxygen.

Example 2Find the formula of aluminium oxide

Al is 3+, O is 2−. Criss-cross the valencies: Al₂O₃. (The 3 and 2 become subscripts on the other ion.)

Example 3Ratio of atoms in CO₂ if C:O mass ratio is 3:8

C: 3 ÷ 12 = 0.25; O: 8 ÷ 16 = 0.50. Ratio 0.25 : 0.50 = 1 : 2 → formula CO₂.

Glossary

Key Terms

Quick definitions for fast revision.

Law of Conservation of Mass: In a chemical reaction, the total mass of reactants equals the total mass of products.
Law of Constant Proportions: A compound always contains the same elements in a fixed proportion by mass.
Dalton's Atomic Theory: Matter is made of indivisible atoms that combine in simple whole-number ratios.
Molecule: The smallest unit of a substance that can exist independently and keep its properties.
Chemical bond: The force of attraction that holds atoms together in a molecule or compound.
Covalent bond: A bond formed by sharing one or more pairs of valence electrons.
Ionic (electrovalent) bond: A bond formed by transfer of electrons and the attraction between the resulting ions.
Cation / Anion: A positive ion (lost electrons) / a negative ion (gained electrons).
Radical: A charged group of atoms acting as a single ion (e.g. SO₄²⁻, NH₄⁺).
Atomicity: The number of atoms in one molecule of an element.
Chemical formula: A symbolic representation of the kinds and numbers of atoms in a compound.
Molecular mass: Sum of the atomic masses of all atoms in one molecule of a covalent compound.
Formula unit mass: Sum of the atomic masses of all atoms in one formula unit of an ionic compound.

Exam Strategy

Top Exam Tips

High-yield reminders that catch most students out.

Open vessel ≠ mass lost

If a gas escapes, mass only seems to fall. Mass is always conserved — reason in a closed system.

Criss-cross the valencies

Swap the valency numbers as subscripts; bracket a radical that takes a subscript > 1, e.g. Ca(OH)₂.

Mono- dropped for the first element

CO is carbon monoxide (not monocarbon); but the second element keeps its prefix and -ide ending.

Prefixes for covalent only

Covalent compounds use di-, tri-, etc. Ionic compounds never do — the criss-cross already fixes the ratio.

Ions vs molecules conduct

Only free ions carry charge. Ionic solutions/melts conduct; covalent (sugar) solutions don't.

Atomicity notation

2O = two atoms; O₂ = one molecule (two bonded atoms); 2O₂ = two molecules.

Same numerical mass

Molecular mass and formula unit mass are calculated identically — the name differs by covalent vs ionic.

Dalton's drawbacks

Atoms ARE divisible; isotopes break "same mass", isobars break "different elements differ in mass".

Self-Evaluation

Practice Question Bank — 150 Questions

CBSE-pattern questions across all five sections with toggle answers. Section A: 50 MCQ & Assertion-Reason (1 mark) · Section B: 30 Very Short Answer (2 marks) · Section C: 30 Short Answer (3 marks) · Section D: 20 Long Answer (5 marks) · Section E: 20 Case-Based (4 marks).