Notes Class 9 Science Chapter 4 Structure of the Atom

Notes For All Chapters Science Class 9 CBSE

1. Introduction

Atoms and molecules are the fundamental building blocks of matter.

The existence of different kinds of matter is due to different types of atoms.

Scientists questioned:

(i) What makes one atom different from another?
(ii) Are atoms indivisible as Dalton proposed?

Experiments showed that atoms are made up of sub-atomic particles – electrons, protons, and neutrons.

2. Charged Particles in Matter

Rubbing two objects together can produce electrical charges.

This indicates atoms contain charged particles.

J.J. Thomson (1897) discovered the electron – a negatively charged particle.

E. Goldstein (1886) discovered canal rays, which are positively charged → led to discovery of protons.

• Proton: positive charge (+1), mass = 1 unit.
• Electron: negative charge (–1), mass = negligible.

An atom contains protons and electrons, balancing each other’s charges.

3. Structure of the Atom

(a) Thomson’s Model (Plum Pudding Model)

Atom = a sphere of positive charge with electrons embedded in it (like seeds in watermelon).

Features:

1. Atom has a positively charged sphere.
2. Electrons are embedded in it.
3. Total positive and negative charges are equal → atom is neutral.

Limitation: Could not explain later experimental results.

(b) Rutherford’s Model (Nuclear Model)

Conducted α-particle scattering experiment using thin gold foil.

Observations:

1. Most α-particles passed straight.
2. Some deflected at small angles.
3. A few rebounded at 180°.

Conclusions:

1. Most of atom is empty space.
2. Positive charge occupies very little space.
3. All positive charge and mass concentrated in nucleus.

Features of Rutherford’s model:

1. Nucleus – small, dense, positively charged center; contains most of the mass.
2. Electrons revolve around nucleus in circular paths.
3. Size of nucleus is very small compared to the atom.

Drawback:

• Electrons revolving in orbits should lose energy and fall into nucleus → atom unstable (but atoms are stable).

(c) Bohr’s Model

Neils Bohr improved Rutherford’s model (1913).

Postulates:

1. Electrons revolve only in certain discrete orbits (energy levels).
2. While revolving in these orbits, they do not radiate energy.

These orbits are called energy levels or shells: K, L, M, N (or n = 1, 2, 3, 4…).

(d) Neutrons

• J. Chadwick (1932) discovered neutron – neutral particle (no charge).
• Mass of neutron ≈ mass of proton.
• Neutrons are present in the nucleus (except in hydrogen).
• Mass of atom = protons + neutrons (called nucleons).

4. Distribution of Electrons in Orbits (Bohr–Bury Scheme)

Rules:

• Maximum electrons in a shell = 2n² (n = shell number).
  • K = 2, L = 8, M = 18, N = 32.
• Outermost shell can have maximum 8 electrons.
• Inner shells must be filled first.

Example:

• Carbon (Z = 6): 2,4
• Sodium (Z = 11): 2,8,1

5. Valency

Valence electrons: electrons in the outermost shell.

Valency: combining capacity of an atom.

Atoms with full outer shells (like noble gases) are inert (valency = 0).

Atoms react to achieve a full outer shell (octet).

• By losing, gaining, or sharing electrons.

Examples:

• Hydrogen, lithium, sodium → lose 1 e⁻ → valency = 1.
• Magnesium → loses 2 e⁻ → valency = 2.
• Aluminium → loses 3 e⁻ → valency = 3.
• Fluorine → gains 1 e⁻ → valency = 1.
• Oxygen → gains 2 e⁻ → valency = 2.

6. Atomic Number and Mass Number

Atomic Number (Z) = number of protons in nucleus.

• Defines the element.
• Example: H → Z = 1, C → Z = 6.

Mass Number (A) = number of protons + neutrons.

• Example: Carbon → 6 + 6 = 12.

Representation:

7. Isotopes

Atoms of the same element having same atomic number but different mass numbers.

Examples:

• Hydrogen → ¹H (Protium), ²H (Deuterium), ³H (Tritium).
• Carbon → ¹²C, ¹⁴C.
• Chlorine → ³⁵Cl, ³⁷Cl.

Properties:

• Same chemical properties.
• Different physical properties.

Average atomic mass = weighted mean of isotopic masses.

• Chlorine example: 35.5 u.

Uses of Isotopes:

1. Uranium isotope → nuclear fuel.
2. Cobalt isotope → cancer treatment.
3. Iodine isotope → treatment of goitre.

8. Isobars

Atoms of different elements having same mass number but different atomic numbers.

Example:

Calcium (⁴⁰₂₀Ca) and Argon (⁴⁰₁₈Ar).

9. Key Points 

Electron: discovered by J.J. Thomson (– charge).

Proton: discovered by E. Goldstein (+ charge).

Neutron: discovered by J. Chadwick (no charge).

Thomson’s Model – electrons embedded in positive sphere.

Rutherford’s Model – nucleus discovered; atom mostly empty.

Bohr’s Model – electrons revolve in fixed energy orbits.

Atomic Number (Z) = no. of protons.

Mass Number (A) = protons + neutrons.

Valency = combining capacity.

Isotopes = same Z, different A.

Isobars = same A, different Z.