Revision Notes — Chapter 3: Atoms and Molecules (NCERT)

This chapter builds the fundamental ideas about the particulate nature of matter, the historical laws that guided the development of atomic theory, how atoms combine to form molecules, and how chemists represent substances using chemical formulae and relative masses. These revision notes are strictly aligned with the NCERT syllabus for Class 9 and are written to help you revise effectively for CBSE board exams.

1. Laws of Chemical Combination — the starting point

The development of the concept of atoms was guided by several experimental laws about how substances combine.

• Law of Conservation of Mass: In a chemical reaction, the total mass of reactants equals the total mass of products. This law indicates that atoms are neither created nor destroyed in a chemical reaction.
• Law of Definite (Constant) Proportions: A given chemical compound always contains its component elements in a fixed ratio by mass, e.g., water (H₂O) always contains hydrogen and oxygen in a fixed mass ratio (about 1:8).
• Law of Multiple Proportions: When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in small whole number ratios — for example, carbon and oxygen form CO and CO₂.

Why these laws matter: They show that matter is made of indivisible building blocks (atoms) that combine in simple numerical proportions — a clue that led scientists to propose the atomic theory.

2. Dalton’s Atomic Theory — basic postulates

John Dalton (early 19th century) proposed a theory to explain these laws. Important points you must remember:

• Matter is made of extremely small particles called atoms.
• Atoms of the same element are identical in mass and properties; atoms of different elements differ.
• Compounds are formed by the combination of atoms in simple whole-number ratios.
• Atoms are indivisible in chemical processes — they are neither created nor destroyed in chemical reactions.

3. Atoms and Molecules — definitions & examples

Atom: The smallest part of an element that takes part in a chemical reaction. Atoms retain the chemical identity of elements.

Molecule: The smallest neutral particle of a substance that can exist independently and shows all the properties of that substance. A molecule may consist of two or more atoms bonded together (e.g., O₂, H₂O).

Molecules of an element: O₂, N₂, P₄, S₈. Molecules of compounds: H₂O (water), CO₂ (carbon dioxide), NaCl (lattice structure but formula unit used).

4. Chemical Formulae — representation of substances

Chemical formulae express the types and numbers of atoms in a molecule or formula unit. Examples:

• H₂O — two hydrogen atoms and one oxygen atom in a water molecule.
• CO₂ — one carbon and two oxygen atoms in a carbon dioxide molecule.
• NaCl — represents the simplest ratio (1:1) of sodium and chloride ions in sodium chloride lattice (called a formula unit).

5. Relative Atomic Mass and Relative Molecular Mass

Relative Atomic Mass (A_r): It is the weighted average mass of the atoms of an element compared with 1/12th the mass of a carbon-12 atom. Values of A_r are found in the periodic table (e.g., A_r of H ≈ 1, O ≈ 16, C ≈ 12).

Relative Molecular Mass (M_r): It is the sum of the relative atomic masses of all the atoms in the molecule. Example: M_r(H₂O) = 2×A_r(H) + A_r(O) = 2×1 + 16 = 18.

6. Mole concept — an introductory view

While the full mole concept appears later in the syllabus, at Class 9 you should understand that chemists use a relative scale (Ar and Mr) to compare masses of atoms and molecules. The mole links atomic-scale masses to real-world masses (1 mole = 6.022×10²³ particles) — the detailed mole calculations are part of higher grades, but recognising Mr and Ar is crucial for stoichiometry.

7. Writing and Interpreting Chemical Equations

Chemical equations represent reactions using formulae. Balancing equations uses the law of conservation of mass: the number of atoms of each element must be the same on both sides.

8. Counting atoms and molecules in formulae

When reading a formula, multiply subscripts by coefficients when present. Example: In 3H₂SO₄, total H atoms = 3×2 = 6; S atoms = 3×1 = 3; O atoms = 3×4 = 12.

9. Applying the concepts — common question types

• Identify laws illustrated by experiments (e.g., mass conservation in reactions).
• Distinguish between atoms, molecules and ions; name simple molecules and write formulae.
• Calculate relative molecular mass using Ar values from the periodic table.
• Balance simple chemical equations and count atoms on each side.

10. Important Tips for Exam Preparation

• Memorise the laws (conservation, constant proportions, multiple proportions) and be able to give examples.
• Practice writing formulae and counting atoms in molecules — many objective questions focus on this skill.
• Keep a small table of common Ar values memorised: H = 1, O = 16, C = 12, N = 14, Na = 23, Cl = 35.5 (approx.).
• When asked to calculate Mr, always state the Ar values used and show the addition step clearly.

11. Quick revision checklist (short form)

• Know and state the three laws of chemical combination.
• Understand and give examples for atom vs molecule.
• Be able to compute Ar and Mr for simple substances.
• Practice balancing equations and counting atoms.
• Remember common test examples: CO & CO₂, H₂O, NaCl, O₂, N₂.

12. Practice Problems (short)

1. Calculate the relative molecular mass of glucose C₆H₁₂O₆ (Ar: C=12, H=1, O=16).
2. In the reaction 2Mg + O₂ → 2MgO, verify the law of conservation of mass by counting atoms.
3. Why does Dalton’s postulate that all atoms of an element are identical not hold true today? (Hint: isotopes)

Answers: 1) Mr = 6×12 + 12×1 + 6×16 = 180. 2) Mg atoms: 2 on both sides; O atoms: 2 on both sides — law verified. 3) Isotopes are atoms of the same element with different masses, so atoms of an element are not strictly identical in mass.