Class 9 Science – Chapter 3: Atoms and Molecules (Detailed Notes)

Class 9 Science – Chapter 3: Atoms and Molecules (Detailed Notes)

1. Laws of Chemical Combination

(a) Law of Conservation of Mass

• Mass can neither be created nor destroyed in a chemical reaction.
• Total mass of reactants = Total mass of products.

(b) Law of Constant Proportions

• A pure compound always contains the same elements in a fixed proportion by mass.
• Example: Water (H₂O) always contains hydrogen and oxygen in the ratio 2:16 (or 1:8).

2. Dalton’s Atomic Theory

• All matter is made of tiny particles called atoms.
• Atoms are indivisible and indestructible.
• Atoms of a given element are identical in mass and properties.
• Atoms combine in simple whole-number ratios to form compounds.
• In a chemical reaction, atoms are rearranged but not destroyed.

3. What is an Atom?

• An atom is the smallest particle of an element which may or may not have an independent existence.
• Examples: H, O, N, etc.

Symbols of Elements (as given by Dalton and IUPAC)

• Each element is represented by a symbol (usually first one or two letters of English or Latin name).
• Examples: Hydrogen – H, Carbon – C, Sodium – Na, Potassium – K, etc.

4. Atomic Mass

• Atomic mass is the mass of an atom compared to 1/12th the mass of one atom of carbon-12 isotope.
• It is a relative quantity and has no unit.
• Example: H = 1, C = 12, O = 16, Na = 23, Cl = 35.5, etc.

5. Molecule

• A molecule is a group of two or more atoms chemically bonded together.
• Can be atoms of same element or different elements.

Types of Molecules:

• Monoatomic: He, Ne, Ar (noble gases)
• Diatomic: H₂, O₂, Cl₂, N₂
• Polyatomic: O₃, P₄, S₈

Molecules of Compounds:

• Formed by atoms of different elements.
• Examples: Water (H₂O), Ammonia (NH₃), Carbon dioxide (CO₂).

6. Ions

• Charged particles formed by loss or gain of electrons.

Types:

• Cation (+): Loss of electron(s) (e.g., Na⁺, Ca²⁺)
• Anion (–): Gain of electron(s) (e.g., Cl^–, O^2–)

Radicals:

• Groups of atoms that carry charge and behave as a single ion (e.g., OH^–, SO₄^2–)

7. Chemical Formulae

• A symbolic representation of a molecule using symbols and subscripts.
• Example: H₂O, CO₂, NaCl, Ca(OH)₂, etc.

Rules for Writing Chemical Formulae:

1. Write the symbols of elements or ions.
2. Write valency of each element.
3. Cross multiply the valencies.
4. Simplify the ratio if needed.

Examples:

• Sodium chloride: Na⁺ Cl^– → NaCl
• Magnesium chloride: Mg²⁺ Cl^– → MgCl₂
• Aluminium oxide: Al³⁺ O^2– → Al₂O₃

8. Molecular Mass

• Sum of atomic masses of all atoms in a molecule.
• Example: H₂O = 2×1 + 1×16 = 18 u

9. Formula Unit Mass

• Used for ionic compounds. It is the sum of atomic masses of all atoms in a formula unit of an ionic compound.

10. Mole Concept

• 1 mole = 6.022 × 10²³ particles (Avogadro's number)
• Molar mass: Mass of 1 mole of a substance in grams.
• Moles = Given mass / Molar mass

Examples:

• 1 mole of H₂O = 6.022 × 10²³ molecules of water = 18 g
• 1 mole of Na = 6.022 × 10²³ atoms of sodium = 23 g