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

Alright class, let's delve into Chapter 4, 'Structure of the Atom'. Understanding the atom is fundamental not just for chemistry, but for much of science. This chapter lays the groundwork, and these concepts often appear in various government exams. Pay close attention to the definitions, discoveries, and models.

Chapter 4: Structure of the Atom - Detailed Notes for Exam Preparation

1. Introduction: Charged Particles in Matter

• Matter is composed of tiny particles called atoms.
• Atoms are not indivisible as proposed by Dalton. They contain even smaller particles.
• Experiments revealed the presence of charged particles within the atom. Opposite charges attract, and like charges repel.

2. Discovery of Sub-atomic Particles

• (a) Electron (e⁻):
  • Discovery: J.J. Thomson (1897) through Cathode Ray experiments.
  • Cathode Rays: Streams of negatively charged particles moving from cathode (-) to anode (+) in a discharge tube at low pressure and high voltage.
  • Properties:
    • Charge: Negative (-1.602 × 10⁻¹⁹ Coulomb). Relative charge: -1.
    • Mass: Very small (9.1 × 10⁻³¹ kg), considered negligible for calculating atomic mass. Relative mass: ~1/1840 u.
    • Location: Orbit the nucleus in specific energy levels.
• (b) Proton (p⁺):
  • Discovery: E. Goldstein (1886) observed Canal Rays (Anode Rays) – positively charged radiations in a modified discharge tube. These led to the discovery of the proton.
  • Properties:
    • Charge: Positive (+1.602 × 10⁻¹⁹ Coulomb). Relative charge: +1.
    • Mass: Approximately 1.672 × 10⁻²⁷ kg. Relative mass: ~1 atomic mass unit (u).
    • Location: Reside inside the nucleus.
• (c) Neutron (n⁰):
  • Discovery: J. Chadwick (1932) by bombarding Beryllium with alpha particles.
  • Properties:
    • Charge: No charge (Neutral). Relative charge: 0.
    • Mass: Slightly greater than a proton (1.674 × 10⁻²⁷ kg). Relative mass: ~1 u.
    • Location: Reside inside the nucleus (except in Hydrogen-1).

3. Atomic Models: Understanding the Arrangement

• (a) Thomson's Model (1904) - "Plum Pudding" or "Watermelon" Model:

  • Postulates:
    • An atom consists of a positively charged sphere.
    • Electrons (negatively charged) are embedded within this sphere, like plums in a pudding or seeds in a watermelon.
    • The total positive charge equals the total negative charge, making the atom electrically neutral.
  • Limitation: Failed to explain the results of Rutherford's alpha-scattering experiment.

• (b) Rutherford's Model (1911) - "Nuclear Model":

  • Alpha (α)-Particle Scattering Experiment:
    • Setup: Fast-moving alpha particles (Helium nuclei, He²⁺) were directed at a very thin gold foil. A fluorescent screen detected the scattered particles.
    • Observations:
      1. Most α-particles passed straight through the foil undeflected.
      2. A small fraction of α-particles were deflected by small angles.
      3. Very few (about 1 in 12,000) bounced back (deflected by 180°).
    • Conclusions:
      1. Most of the space inside an atom is empty.
      2. There is a small, dense, positively charged centre called the nucleus.
      3. The nucleus contains almost all the mass of the atom.
      4. Electrons revolve around the nucleus in circular paths.
      5. The size of the nucleus is very small compared to the size of the atom.
  • Features of the Model: Positively charged nucleus at the centre; electrons orbiting the nucleus; atom is mostly empty space.
  • Drawback (Major Limitation): According to classical electromagnetic theory, any charged particle undergoing acceleration (like an electron orbiting a nucleus) should radiate energy continuously. This would cause the electron to lose energy, spiral inwards, and eventually fall into the nucleus, making the atom unstable. However, atoms are known to be stable.

• (c) Bohr's Model (1913) - Addressing Rutherford's Drawback:

  • Postulates:
    1. Electrons revolve around the nucleus only in certain specific, discrete orbits called energy levels or shells.
    2. While revolving in these discrete orbits, electrons do not radiate energy.
    3. Energy is emitted or absorbed only when an electron jumps from one energy level to another.
  • Energy Levels (Shells): Designated by letters K, L, M, N... or numbers n = 1, 2, 3, 4... starting from the nucleus outwards.
    • K-shell (n=1): Lowest energy level.
    • L-shell (n=2)
    • M-shell (n=3)
    • N-shell (n=4), and so on. Energy increases as we move away from the nucleus.

4. Atomic Number (Z) and Mass Number (A)

• Atomic Number (Z):
  • Definition: The number of protons present in the nucleus of an atom.
  • Significance: It determines the identity of an element. All atoms of a particular element have the same atomic number.
  • In a neutral atom: Number of protons (Z) = Number of electrons.
  • Represented as a subscript before the element symbol (e.g., ₆C).
• Mass Number (A):
  • Definition: The total number of protons and neutrons present in the nucleus of an atom. (Nucleons = Protons + Neutrons).
  • Calculation: Mass Number (A) = Number of protons (Z) + Number of neutrons (n).
  • Therefore, Number of neutrons (n) = A - Z.
  • Represented as a superscript before the element symbol (e.g., ¹²C).
• Notation: An atom is represented as ^A_Z X, where X is the element symbol. Example: ¹²₆C represents Carbon with Z=6 and A=12.

5. Distribution of Electrons in Shells (Bohr-Bury Scheme)

• The arrangement of electrons in different energy levels is called the electronic configuration.

• Rules:

  1. The maximum number of electrons that can be accommodated in a shell is given by the formula 2n², where 'n' is the shell number (n=1 for K, n=2 for L, etc.).
     • K-shell (n=1): Max 2(1)² = 2 electrons
     • L-shell (n=2): Max 2(2)² = 8 electrons
     • M-shell (n=3): Max 2(3)² = 18 electrons
     • N-shell (n=4): Max 2(4)² = 32 electrons
  2. The outermost shell (valence shell) cannot accommodate more than 8 electrons (Octet Rule), even if it has the capacity to hold more according to 2n². The first shell (K) is an exception, holding a maximum of 2 electrons (Duplet Rule).
  3. Electrons fill the shells in a stepwise manner, starting from the innermost shell (K). Inner shells are filled before outer shells are occupied.

• Examples:

  • Hydrogen (Z=1): K=1
  • Helium (Z=2): K=2 (Duplet complete)
  • Lithium (Z=3): K=2, L=1
  • Carbon (Z=6): K=2, L=4
  • Neon (Z=10): K=2, L=8 (Octet complete)
  • Sodium (Z=11): K=2, L=8, M=1
  • Argon (Z=18): K=2, L=8, M=8 (Octet complete)

6. Valency

• Definition: The combining capacity of an atom of an element. It determines how an atom will react with other atoms.
• Determination: It depends on the number of valence electrons (electrons in the outermost shell).
  • If the valence shell has 1, 2, 3, or 4 electrons: Valency = Number of valence electrons (tendency to lose electrons).
  • If the valence shell has 5, 6, 7, or 8 electrons: Valency = 8 - Number of valence electrons (tendency to gain or share electrons to achieve octet).
  • Atoms with a completely filled outermost shell (like Noble gases - He, Ne, Ar) have zero valency as they are chemically inert.
• Examples:
  • Sodium (Na: 2, 8, 1): Valence electrons = 1. Valency = 1.
  • Magnesium (Mg: 2, 8, 2): Valence electrons = 2. Valency = 2.
  • Aluminium (Al: 2, 8, 3): Valence electrons = 3. Valency = 3.
  • Carbon (C: 2, 4): Valence electrons = 4. Valency = 4.
  • Nitrogen (N: 2, 5): Valence electrons = 5. Valency = 8 - 5 = 3.
  • Oxygen (O: 2, 6): Valence electrons = 6. Valency = 8 - 6 = 2.
  • Chlorine (Cl: 2, 8, 7): Valence electrons = 7. Valency = 8 - 7 = 1.
  • Neon (Ne: 2, 8): Valence electrons = 8. Valency = 8 - 8 = 0.

7. Isotopes

• Definition: Atoms of the same element having the same atomic number (Z) but different mass numbers (A).
• Reason: They have the same number of protons but a different number of neutrons.
• Examples:
  • Hydrogen: ¹₁H (Protium, 0 neutrons), ²₁H (Deuterium, 1 neutron), ³₁H (Tritium, 2 neutrons).
  • Carbon: ¹²₆C (6 neutrons), ¹³₆C (7 neutrons), ¹⁴₆C (8 neutrons).
  • Chlorine: ³⁵₁₇Cl (18 neutrons), ³⁷₁₇Cl (20 neutrons).
• Properties:
  • Chemical properties are nearly identical (determined by electron configuration/atomic number).
  • Physical properties (like mass, density, boiling point) may differ slightly.
• Average Atomic Mass: The atomic mass of many elements is fractional because it represents the average mass of its naturally occurring isotopes, weighted by their relative abundance. Example: Chlorine exists as ³⁵Cl (75%) and ³⁷Cl (25%). Average atomic mass = (35 × 75/100) + (37 × 25/100) = 26.25 + 9.25 = 35.5 u.
• Applications:
  • Uranium isotope (U-235): Used as fuel in nuclear reactors.
  • Cobalt isotope (Co-60): Used in the treatment of cancer.
  • Iodine isotope (I-131): Used in the treatment of goitre.
  • Carbon isotope (C-14): Used in carbon dating (determining the age of fossils/archaeological samples).

8. Isobars

• Definition: Atoms of different elements having different atomic numbers (Z) but the same mass number (A).
• Reason: They have different numbers of protons (hence different elements) and different numbers of neutrons, but the sum (protons + neutrons) is the same.
• Examples:
  • ⁴⁰₁₈Ar (Argon: 18p, 22n) and ⁴⁰₂₀Ca (Calcium: 20p, 20n). Both have mass number A=40.
• Properties: Have different chemical properties (as they are different elements with different electron configurations).

Summary Table: Subatomic Particles

Multiple Choice Questions (MCQs)

1. Who is credited with the discovery of the neutron?
   a) J.J. Thomson
   b) E. Goldstein
   c) J. Chadwick
   d) Niels Bohr

2. Rutherford's alpha-particle scattering experiment concluded that:
   a) Electrons are embedded in a positive sphere.
   b) Most of the space in an atom is empty.
   c) Electrons revolve in discrete energy levels.
   d) Atoms are indivisible.

3. The maximum number of electrons that can be accommodated in the L-shell (n=2) is:
   a) 2
   b) 8
   c) 18
   d) 32

4. An atom has an atomic number (Z) of 11 and a mass number (A) of 23. How many neutrons does it have?
   a) 11
   b) 12
   c) 23
   d) 34

5. Isotopes of an element have:
   a) Same mass number, different atomic number.
   b) Same number of protons, different number of neutrons.
   c) Same number of neutrons, different number of electrons.
   d) Same mass number and same atomic number.

6. Which of the following represents a pair of isobars?
   a) ¹²₆C and ¹⁴₆C
   b) ³⁵₁₇Cl and ³⁷₁₇Cl
   c) ⁴⁰₁₈Ar and ⁴⁰₂₀Ca
   d) ¹₁H and ²₁H

7. What is the valency of Chlorine (Z=17)?
   a) 1
   b) 7
   c) 8
   d) 2

8. Which atomic model first proposed the existence of a nucleus?
   a) Thomson's Model
   b) Dalton's Theory
   c) Rutherford's Model
   d) Bohr's Model

9. The electronic configuration of Sodium (Z=11) is:
   a) 2, 9
   b) 8, 2, 1
   c) 2, 8, 1
   d) 1, 8, 2

10. An isotope of which element is used in the treatment of goitre?
    a) Uranium
    b) Cobalt
    c) Carbon
    d) Iodine

Answer Key for MCQs:

1. c) J. Chadwick
2. b) Most of the space in an atom is empty.
3. b) 8
4. b) 12 (Neutrons = A - Z = 23 - 11 = 12)
5. b) Same number of protons, different number of neutrons.
6. c) ⁴⁰₁₈Ar and ⁴⁰₂₀Ca
7. a) 1 (Electronic configuration: 2, 8, 7. Valence electrons = 7. Valency = 8 - 7 = 1)
8. c) Rutherford's Model
9. c) 2, 8, 1
10. d) Iodine

Make sure you understand the definitions, the key experiments and their conclusions, and how to calculate the number of subatomic particles, determine electronic configuration, and find valency. These form the core concepts for your exams. Good luck!