1. Introduction:

  • Atoms and molecules are the fundamental building blocks of matter.
  • Different kinds of matter are made up of different atoms.
  • Key Questions:
    • What makes atoms of different elements different from each other?
    • Are atoms indivisible, as proposed by Dalton, or do they have smaller parts?

2. Charged Particles in Matter:

  • Activity: Rubbing a glass rod with a cloth or combing hair creates an electrical charge. This happens because atoms are divisible and contain charged particles.
  • Key Particles:
    • Electrons (e–): Negative charge, very small mass.
    • Protons (p+): Positive charge, much larger mass (about 2000 times the mass of electrons).

3. Discovery of Sub-Atomic Particles:

  • Electrons: Discovered by J.J. Thomson in 1897.
  • Protons: Discovered by E. Goldstein in 1886 through canal rays (positively charged particles).
  • An atom is made of protons and electrons, which balance out the atom’s overall charge.

4. Thomson’s Model of the Atom (1898):

  • Proposed that the atom is like a Christmas pudding: a positive sphere with electrons embedded inside.
  • Key Features:
    • Positive charge is spread over the entire atom.
    • Electrons are embedded in this positive sphere.
    • The atom is electrically neutral (positive and negative charges balance).

5. Rutherford’s Model of the Atom (1911):

  • Rutherford used an experiment where alpha particles were directed at a thin gold foil.
  • Results:
    • Most alpha particles passed through without deflection.
    • Some were deflected slightly.
    • A few particles bounced back, showing that most of the atom is empty space, and its positive charge is concentrated in a small, dense nucleus.
  • Conclusion:
    • The nucleus is positively charged and holds most of the atom’s mass.
    • Electrons revolve around the nucleus in orbits.

6. Drawbacks of Rutherford’s Model:

  • Electrons moving in circular orbits would lose energy and spiral into the nucleus, making the atom unstable. But atoms are stable, so something was missing in the model.

7. Bohr’s Model of the Atom (1913):

  • To solve Rutherford’s model issue, Niels Bohr suggested:
    • Electrons revolve only in certain stable orbits (not radiating energy).
    • Electrons can jump between orbits but don’t lose energy while in a specific orbit.

8. Neutrons:

  • In 1932, J. Chadwick discovered neutrons—particles with no charge, but mass nearly equal to that of a proton.
  • Neutrons are found in the nucleus, except in hydrogen atoms.

9. Electron Distribution in Orbits:

  • Bohr’s Rule for electron distribution:
    • First orbit (K-shell): Maximum of 2 electrons.
    • Second orbit (L-shell): Maximum of 8 electrons.
    • Third orbit (M-shell): Maximum of 18 electrons.
    • Fourth orbit (N-shell): Maximum of 32 electrons.
  • Electrons fill shells from inner to outer, and the outermost shell can hold up to 8 electrons.

10. Key Concepts for Competitive Exams:

  • Sub-atomic particles: Proton (positive), electron (negative), neutron (neutral).
  • Atomic Models:
    • Thomson’s Model: Positive sphere with electrons inside (like Christmas pudding).
    • Rutherford’s Model: Atom has a tiny, dense nucleus with electrons revolving around it.
    • Bohr’s Model: Electrons in stable orbits without emitting energy.
  • Neutron Discovery: Neutrons were discovered by J. Chadwick, and they help explain atomic mass.

1. Valency

  • Valence Electrons: The electrons in the outermost shell of an atom are called valence electrons. These electrons are important because they determine how an atom reacts chemically with other atoms.
  • Bohr-Bury Scheme: According to this model, electrons are arranged in shells or orbits around the nucleus. The outermost shell can hold a maximum of 8 electrons.
  • Full Outermost Shell: When an atom’s outermost shell has 8 electrons, it is stable and not reactive. Such atoms have zero valency (e.g., noble gases like helium and neon).
  • Octet Rule: Atoms react to get 8 electrons in the outermost shell, called the octet. To achieve this, they may lose, gain, or share electrons. For example, atoms like hydrogen, lithium, and sodium each have one electron in their outer shell, so they lose that one electron to become stable. This gives them a valency of one.
  • Valency Example:
    • Magnesium (2 electrons in the outermost shell) has a valency of 2 because it loses two electrons.
    • Aluminium (3 electrons in the outermost shell) has a valency of 3 because it loses three electrons.
  • Exceptions: For elements like fluorine (7 electrons), instead of losing 7 electrons, it gains 1 electron to complete its octet, resulting in a valency of 1.
  • Calculating Valency:
    • If an element has electrons close to a full shell, subtract the number of electrons from 8 to find the valency.
    • For oxygen (6 electrons), it needs 2 electrons to complete the octet, so its valency is 2.

2. Atomic Number and Mass Number

  • Atomic Number (Z): The atomic number is the number of protons in an atom’s nucleus. It defines the element. For example, Hydrogen has an atomic number of 1, meaning it has 1 proton.
  • Mass Number (A): The mass number is the sum of protons and neutrons in the nucleus. It tells you the approximate mass of the atom. For example, Carbon (C) has 6 protons and 6 neutrons, so its mass number is 12.
  • Notation: An element’s atomic number and mass number are used to represent it. For example, Carbon is written as 12/6 C.

3. Isotopes

  • Definition: Isotopes are atoms of the same element that have the same atomic number but different mass numbers because they have different numbers of neutrons. For example, Hydrogen has three isotopes:
    • Protium (1H), with 0 neutrons
    • Deuterium (2H), with 1 neutron
    • Tritium (3H), with 2 neutrons.
  • Importance: Isotopes have similar chemical properties but different physical properties. For example, Carbon-12 and Carbon-14 are isotopes of carbon, but Carbon-14 is used in radiocarbon dating because it is radioactive.
  • Chlorine Isotopes: Chlorine exists as two isotopes with different masses (35 and 37), and they occur in a ratio of 3:1. To calculate the average atomic mass, we use the percentages and average them, giving Chlorine an average mass of 35.5 u.

4. Isobars

  • Definition: Isobars are atoms of different elements that have the same mass number but different atomic numbers. For example, Calcium (Ca) and Argon (Ar) both have mass numbers of 40, but they are different elements because they have different atomic numbers (20 for Calcium and 18 for Argon).

5. Key Points

  • Electron: Discovered by J.J. Thomson, it’s negatively charged and has a tiny mass.
  • Proton: Discovered by E. Goldstein, it’s positively charged.
  • Neutron: Discovered by J. Chadwick, it has no charge and is found in the nucleus.
  • Shells: Electrons are arranged in shells (K, L, M, N). The outer shell’s electron count determines chemical behavior.
  • Valency: The number of electrons an atom needs to lose, gain, or share to achieve a full outer shell. It explains the chemical reactivity of an element.
  • Atomic and Mass Numbers: The atomic number tells you the number of protons, while the mass number tells you the total number of protons and neutrons in an atom.
  • Isotopes and Isobars: Isotopes are atoms of the same element with different masses, while isobars have the same mass number but are different elements.

Competitive Exam Tip

  • Valency of Common Elements:
    • Sodium (Na): 1
    • Magnesium (Mg): 2
    • Oxygen (O): 2
    • Fluorine (F): 1
    • Aluminium (Al): 3

THESE ALL ARE THE NOTES OF CHAPTER 4 SCIENCE. AND AFTER SOME TIME YOU GET IMPORTANT QUESTIONS HERE. *#THANKS FOR VISITING, VISIT AGAIN#* 😊