📝 Chapter Notes & Revision

Structure of the Atom

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Class 9 Science - Chapter: Structure of the Atom

MP Board Education

1. Introduction to Subatomic Particles

Atoms are divisible and made up of charged subatomic particles.

  • Electron ($e^-$):
    • Discovery: J.J. Thomson (1897) using the Cathode Ray Tube experiment.
    • Charge: Negative ($-1$ relative, $-1.6 \times 10^{-19}\text{ C}$ actual).
    • Mass: Negligible ($9.1 \times 10^{-31}\text{ kg}$ or $\frac{1}{1836}$ times the mass of a hydrogen atom).
  • Proton ($p^+$):
    • Discovery: E. Goldstein (1886) using Canal Rays / Anode Rays.
    • Charge: Positive ($+1$ relative, $+1.6 \times 10^{-19}\text{ C}$ actual).
    • Mass: $1.673 \times 10^{-27}\text{ kg}$ (taken as $1\text{ u}$).
  • Neutron ($n$):
    • Discovery: J. Chadwick (1932).
    • Charge: Neutral (no charge).
    • Mass: $1.675 \times 10^{-27}\text{ kg}$ (approximately equal to the mass of a proton).

2. Thomson's Model of the Atom (Plum Pudding Model)

  • Postulate: An atom consists of a positively charged sphere and electrons are embedded in it.
  • Analogy: Like a watermelon where the positive charge is spread all over (red edible part) and electrons are studded like seeds (black seeds).
  • Limitation: It failed to explain how protons could keep electrons from leaving and could not explain the results of later scattering experiments.

3. Rutherford's Alpha ($\alpha$) Particle Scattering Experiment

  • Setup: Fast-moving $\alpha$-particles (doubly-charged helium ions, $He^{2+}$) were made to fall on a thin gold foil.
  • Observations:
    1. Most $\alpha$-particles passed straight through the gold foil.
    2. Some particles were deflected by small angles.
    3. Very few ($\sim 1$ in $12,000$) bounced back completely.
  • Conclusions:
    1. Most of the space inside the atom is empty.
    2. Positive charge is concentrated in a very small volume inside the atom (Nucleus).
    3. The size of the nucleus is very small compared to the size of the atom.

4. Rutherford's Nuclear Model of the Atom

  • There is a positively charged center in an atom called the nucleus. Nearly all the mass of the atom resides in the nucleus.
  • Electrons revolve around the nucleus in circular paths.
  • Major Drawback: According to electromagnetic theory, revolving charged particles radiate energy. If electrons radiate energy, they would eventually fall into the nucleus, making the atom unstable. However, atoms are stable.

5. Bohr's Model of the Atom

To overcome Rutherford's objection, Neils Bohr (1913) proposed:

  1. Only certain special orbits known as discrete orbits of electrons are allowed inside the atom.
  2. While revolving in discrete orbits, electrons do not radiate energy. These orbits are called energy levels or shells.
  3. Energy shells are represented by letters ($K, L, M, N, \dots$) or numbers ($n = 1, 2, 3, 4, \dots$).

6. Distribution of Electrons in Orbits (Bohr-Bury Scheme)

  • Rule 1 (Maximum Electrons in a Shell): Given by the formula $2n^2$, where '$n$' is the shell number ($1, 2, 3,\dots$).
    • $K$ shell ($n=1$): $2(1)^2 = 2$ electrons
    • $L$ shell ($n=2$): $2(2)^2 = 8$ electrons
    • $M$ shell ($n=3$): $2(3)^2 = 18$ electrons
    • $N$ shell ($n=4$): $2(4)^2 = 32$ electrons
  • Rule 2: The maximum number of electrons that can be accommodated in the outermost orbit is 8 (Octet rule).
  • Rule 3: Electrons are not accommodated in a given shell, unless the inner shells are filled. (Shells are filled in a step-wise manner).

7. Valency (Combining Capacity)

  • Definition: The combining capacity of an atom of an element is known as its valency.
  • Calculation:
    • If the outermost shell has $1, 2, 3,$ or $4$ electrons, Valency = Number of valence electrons.
    • If the outermost shell has $5, 6, 7,$ or $8$ electrons, Valency = $8 -$ Number of valence electrons (or electrons needed to complete octet).
    • Elements with completely filled outermost shells (8 electrons, except Helium which has 2) have a valency of 0 (inert/noble gases).

8. Important Terms and Formulas

  • Atomic Number ($Z$):

    • Defined as the total number of protons present in the nucleus of an atom.
    • $Z = \text{Number of protons } (p) = \text{Number of electrons } (e)$ (in a neutral atom).

  • Mass Number ($A$):

    • Defined as the sum of the total number of protons and neutrons present in the nucleus of an atom.
    • Formula: $A = \text{Number of Protons } (p) + \text{Number of Neutrons } (n)$
    • Also called Nucleons ($p + n$).

  • Representation of an Element: $$\mathbf{_Z^A X}$$ (Where $\text{X} = \text{Symbol of element}, A = \text{Mass Number}, Z = \text{Atomic Number}$)

  • Number of Neutrons ($n$):

    • Formula: $n = A - Z$

9. Isotopes and Isobars

  • Isotopes (समस्थानिक):

    • Atoms of the same element having the same atomic number ($Z$) but different mass numbers ($A$).
    • Example: Hydrogen has three isotopes: Protium (${1}^{1}\text{H}$), Deuterium (${1}^{2}\text{H}$), and Tritium ($_{1}^{3}\text{H}$). Carbon has $^{12}\text{C}$ and $^{13}\text{C}$.
    • Application:
      • Uranium isotope is used as fuel in nuclear reactors.
      • Cobalt isotope is used in the treatment of cancer.
      • Iodine isotope is used in the treatment of goitre.

  • Isobars (समभारिक):

    • Atoms of different elements having the same mass number ($A$) but different atomic numbers ($Z$).
    • Example: Calcium (${20}^{40}\text{Ca}$) and Argon (${18}^{40}\text{Ar}$). Both have a mass number of $40$.