Chapter 2 Structure of Atom Class11

Class 11 Chemistry – Chapter 2

AarishSir

Structure of Atom 

1. Introduction

• Atoms are the basic building blocks of matter.

• The idea of atoms was first given by Democritus (a Greek philosopher), but the scientific study of atoms started with Dalton’s Atomic Theory (1808).

• Over time, new experiments proved that atoms are not indivisible, but are made up of subatomic particles (electrons, protons, neutrons).

• This chapter explains how scientists discovered the structure of the atom.

2. Discovery of Subatomic Particles

(a) Electron

• Discovered by J.J. Thomson in 1897 through cathode ray tube experiments.

• Properties:

  • Negatively charged particle (charge = –1.6 × 10⁻¹⁹ C).

  • Mass = 9.1 × 10⁻³¹ kg.

• Important experiment: When high voltage was applied in a cathode ray tube, rays traveled from cathode to anode. These rays were deflected by electric and magnetic fields, proving they are made of negatively charged particles (electrons).

(b) Proton

• Discovered by E. Goldstein in 1886 using canal rays (positive rays).

• Properties:

  • Positively charged particle (charge = +1.6 × 10⁻¹⁹ C).

  • Mass ≈ 1.67 × 10⁻²⁷ kg (about 1836 times heavier than electron).

(c) Neutron

• Discovered by James Chadwick in 1932.

• Properties:

  • Neutral particle (no charge).

  • Mass ≈ equal to proton.

• Neutrons contribute to atomic mass but not to charge.

3. Atomic Models

(a) Dalton’s Atomic Theory (1808)

• Atoms are indivisible, smallest particles of matter.

• All atoms of a given element are identical in mass and properties.

• Compounds are formed when atoms of different elements combine in fixed ratios.

• Limitations: Could not explain discovery of subatomic particles, isotopes, or atomic structure.

(b) Thomson’s Model of Atom (Plum Pudding Model, 1898)

• Atom is a sphere of positive charge with electrons embedded in it like “plums in pudding.”

• Limitations: Failed to explain results of Rutherford’s experiment.

(c) Rutherford’s Nuclear Model (1911)

• Gold Foil Experiment:

  • α-particles (helium nuclei) were bombarded on thin gold foil.

  • Observations: Most passed through, some deflected, few bounced back.

  • Conclusion:

    • Atom has a small, dense, positively charged nucleus at the center.

    • Electrons revolve around the nucleus.

    • Most of the atom is empty space.

• Limitations: Could not explain stability of atom (electrons should spiral into nucleus).

(d) Bohr’s Model (1913)

• Niels Bohr proposed improvements to Rutherford’s model.

• Postulates:

  1. Electrons revolve in fixed circular orbits called energy levels/shells (K, L, M, N).

  2. Each orbit has fixed energy (quantized).

  3. Electrons neither lose nor gain energy while in these orbits.

  4. Energy is absorbed or emitted only when electron jumps between energy levels.

     • ΔE = E₂ – E₁ = hν

• Success: Explained hydrogen atom spectrum.

• Limitation: Failed for multi-electron atoms.

4. Quantum Mechanical Model of Atom

• Modern and most accepted model.

• Based on Heisenberg’s Uncertainty Principle and Schrödinger’s wave equation.

• Explains atom in terms of orbitals (regions in space where probability of finding electron is maximum), not fixed paths.

Key points:

• Electrons behave like particles + waves (dual nature).

• Orbitals are defined by four quantum numbers.

5. Quantum Numbers

Each electron in an atom is described by a set of four quantum numbers:

1. Principal Quantum Number (n):

   • Describes main shell or energy level.

   • Values: n = 1, 2, 3...

   • Higher n → larger size & energy of orbital.

2. Azimuthal/Angular Quantum Number (l):

   • Describes subshell (shape of orbital).

   • Values: l = 0 to (n – 1).

   • l = 0 (s), l = 1 (p), l = 2 (d), l = 3 (f).

3. Magnetic Quantum Number (m):

   • Describes orientation of orbital in space.

   • Values: –l to +l.

4. Spin Quantum Number (s):

   • Describes spin of electron.

   • Values: +½ or –½.

6. Shapes of Orbitals

• s-orbital: spherical.

• p-orbital: dumbbell-shaped.

• d-orbital: cloverleaf-shaped.

• f-orbital: complex shapes.

7. Electronic Configuration of Atoms

• Arrangement of electrons in orbitals of an atom.

• Rules:

  1. Aufbau Principle: Electrons fill orbitals in order of increasing energy.

     • Order: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p ...

  2. Pauli Exclusion Principle: No two electrons can have the same set of four quantum numbers (max. 2 electrons in an orbital with opposite spins).

  3. Hund’s Rule of Maximum Multiplicity: Electrons occupy orbitals singly first before pairing.

Example:

• Hydrogen (Z=1): 1s¹

• Helium (Z=2): 1s²

• Carbon (Z=6): 1s² 2s² 2p²

• Oxygen (Z=8): 1s² 2s² 2p⁴

8. Important Terms

• Isotopes: Atoms of same element with same atomic number but different mass numbers (e.g., ¹H, ²H, ³H).

• Isobars: Atoms of different elements with same mass number but different atomic numbers (e.g., ⁴⁰Ca, ⁴⁰Ar).

• Isoelectronic Species: Species with same number of electrons (e.g., Na⁺ and Ne).

9. Applications of Atomic Structure

• Explains line spectrum of hydrogen.

• Basis of periodic table arrangement.

• Explains chemical bonding.

• Helps in understanding properties of elements.

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