CBSE Class 9 Science Notes: Structure of the Atom

Subatomic Particles: Electrons, Protons, and Neutrons

Atoms, the fundamental building blocks of matter, are composed of even smaller particles called subatomic particles. These particles determine the atom’s properties and behavior.

• Electrons (e^–):
  • Negatively charged particles.
  • Orbit the nucleus in specific energy levels or shells.
  • Have a very small mass, almost negligible.
• Protons (p):
  • Positively charged particles.
  • Located within the nucleus of the atom.
  • Have a mass of approximately 1 atomic mass unit (amu).
• Neutrons (n):
  • Neutral particles (no charge).
  • Located within the nucleus of the atom, along with protons.
  • Have a mass of approximately 1 amu (similar to a proton).

Basic Atomic Models

Throughout history, scientists have proposed different models to explain the structure of the atom. Here are some key models:

• Dalton’s Atomic Theory (Early 1800s):
  • Proposed that all matter is made of indivisible atoms.
  • Atoms of a given element are identical in mass and properties.
  • Atoms combine in simple whole-number ratios to form compounds.
• Thomson’s “Plum Pudding” Model (Late 1800s):
  • Proposed a sphere of positive charge with electrons embedded throughout, like plums in a pudding.
  • This model accounted for the existence of electrons but didn’t explain the atom’s internal structure in detail.
• Rutherford’s Nuclear Model (Early 1900s):
  • Based on the gold foil experiment, Rutherford discovered that the atom has a small, dense, positively charged nucleus at the center.
  • Electrons orbit the nucleus.
• Bohr’s Atomic Model (Early 1900s):
  • Proposed that electrons orbit the nucleus in specific, fixed energy levels (shells).
  • Electrons can only gain or lose energy by jumping between these energy levels.

Atomic Number and Mass Number

These two numbers are crucial for identifying and understanding atoms.

• Atomic Number (Z):
  • The number of protons in the nucleus of an atom.
  • Defines the element (e.g., all atoms with 6 protons are carbon).
• Mass Number (A):
  • The total number of protons and neutrons in the nucleus of an atom.
  • Represents the approximate mass of the atom.

Formula:

Mass Number (A) = Number of Protons (Z) + Number of Neutrons (N)

Or, $A = Z + N$

Example: Carbon (C) has an atomic number of 6 (6 protons) and a mass number of 12 (6 protons + 6 neutrons). The notation is often written as ¹²C₆.

Valency and Its Simple Explanation

Valency helps us understand how atoms combine to form molecules.

Definition: Valency is the combining capacity of an atom. It represents the number of electrons an atom can gain, lose, or share to achieve a stable electronic configuration (usually resembling a noble gas).

Core Principles:

• Atoms want to achieve a stable outer electron shell (octet rule – typically 8 electrons, or a duet for Helium).
• Atoms achieve stability by:
• Losing electrons (e.g., metals).
• Gaining electrons (e.g., nonmetals).
• Sharing electrons (covalent bonds).

Examples:

• Sodium (Na) has 1 valence electron, so it has a valency of 1 (loses 1 electron).
• Chlorine (Cl) has 7 valence electrons, so it has a valency of 1 (gains 1 electron).
• Carbon (C) has 4 valence electrons, so it has a valency of 4 (shares 4 electrons).

Isotopes and Isobars

These terms describe variations in the composition of atoms of the same or different elements.

• Isotopes:
  • Atoms of the same element (same atomic number, Z) that have different mass numbers (A).
  • They have the same number of protons but different numbers of neutrons.
  • Example: Carbon-12 (¹²C) and Carbon-14 (¹⁴C). Both have 6 protons, but Carbon-12 has 6 neutrons, while Carbon-14 has 8 neutrons.
• Isobars:
  • Atoms of different elements that have the same mass number (A).
  • They have different numbers of protons and neutrons, but the sum is the same.
  • Example: Argon-40 (⁴⁰Ar) and Calcium-40 (⁴⁰Ca).