Class 9 Science – Chapter 4: Structure of the Atom (Detailed Notes)

Class 9 Science – Chapter 4: Structure of the Atom (Detailed Notes)

1. Charged Particles in Matter

• Atom consists of subatomic particles – electrons, protons, and neutrons.
• Electrons are negatively charged and discovered by J.J. Thomson.
• Protons are positively charged and discovered by Goldstein.
• Neutrons are neutral particles discovered by James Chadwick.

2. Discovery of Electron

• J.J. Thomson discovered the electron using cathode ray experiment.
• Cathode rays are negatively charged particles that travel from cathode to anode.

3. Discovery of Proton

• E. Goldstein discovered protons using canal ray experiment.
• Canal rays are positively charged particles present in the atom.

4. Discovery of Neutron

• James Chadwick discovered neutrons in 1932.
• Neutrons are neutral particles found in the nucleus of an atom.

5. Thomson’s Model of Atom

• Atom is a positively charged sphere with negatively charged electrons embedded in it (like seeds in a watermelon).
• Fails to explain the stability of atom and results of scattering experiment.

6. Rutherford’s Model of Atom

• Performed alpha-particle scattering experiment on gold foil.
• Observations:
  • Most alpha particles passed through without deflection.
  • Some were deflected at small angles.
  • Very few bounced back.
• Conclusions:
  • Most of the atom is empty space.
  • There is a small, dense, positively charged center called the nucleus.

Drawbacks of Rutherford’s Model:

• Could not explain the stability of atom as accelerating electrons should lose energy and fall into the nucleus.

7. Bohr’s Model of Atom

• Electrons revolve around the nucleus in fixed circular paths called orbits or shells.
• Each orbit has a fixed energy; hence, also called energy levels.
• Electrons do not radiate energy while revolving in these energy levels.

Bohr's Postulates:

• Only certain special orbits known as discrete orbits are allowed.
• Energy of an electron remains constant in these orbits.

8. Atomic Structure

• Atoms are made up of electrons, protons, and neutrons.
• Protons and neutrons reside in the nucleus.
• Electrons revolve around the nucleus in specific energy levels.

9. Distribution of Electrons (Bohr-Bury Rules)

• Maximum electrons in a shell = 2n² (n = shell number)
• Maximum electrons in:
  • K shell = 2
  • L shell = 8
  • M shell = 18
• Outer shell cannot have more than 8 electrons.
• Electrons are filled in a step-wise manner, starting from the inner shell.

10. Valency

• The combining capacity of an atom is called valency.
• Valency = Number of electrons lost, gained, or shared to complete an octet.
• Example: Valency of oxygen = 2, sodium = 1.

11. Atomic Number (Z)

• Number of protons in the nucleus = Atomic number (Z).
• Atomic number also equals the number of electrons in a neutral atom.
• Example: Hydrogen (Z = 1), Carbon (Z = 6).

12. Mass Number (A)

• Mass Number = Number of protons + Number of neutrons.
• Example: Carbon-12 → A = 6 + 6 = 12.

13. Isotopes

• Atoms of the same element with the same atomic number but different mass numbers.
• Examples:
  • Hydrogen: Protium (¹H), Deuterium (²H), Tritium (³H)
  • Carbon: ¹²C and ¹⁴C
• Uses:
  • Isotope of Iodine used in treatment of goitre.
  • Isotope of Cobalt used in cancer treatment.

14. Isobars

• Atoms with same mass number but different atomic numbers.
• Example: Calcium (²⁰Ca⁴⁰) and Argon (¹⁸Ar⁴⁰).