class 11 chemistry atomic structure notes|| atomic structure notes class 11 || class 11 chemistry unit 3

 

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🧪 Chapter: Atomic Structure

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🔹 1. Introduction to Atomic Structure

• Atom: The smallest particle of an element that retains its chemical properties.

• Atomic structure refers to the arrangement of electrons, protons, and neutrons within an atom.

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🔹 2. Subatomic Particles

Particle	Symbol	Charge	Mass (kg)
Electron	e⁻	-1.602 × 10⁻¹⁹ C	9.1 × 10⁻³¹
Proton	p⁺	+1.602 × 10⁻¹⁹ C	1.67 × 10⁻²⁷
Neutron	n⁰	0 (neutral)	1.67 × 10⁻²⁷

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🔹 3. Discovery of Subatomic Particles

• Electron: Discovered by J.J. Thomson (Cathode Ray Experiment).

• Proton: Discovered by Goldstein (Canal Ray Experiment).

• Neutron: Discovered by James Chadwick (Bombardment of beryllium with alpha particles).

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🔹 4. Atomic Models

a. Thomson’s Model (Plum Pudding Model)

• Atom is a positively charged sphere with electrons embedded in it.

• Drawback: Could not explain the results of Rutherford’s experiment.

b. Rutherford’s Nuclear Model

• Gold foil experiment (α-particle scattering).

• Major conclusions:

  • Atom is mostly empty.

  • Nucleus is dense, small, and positively charged.

  • Electrons revolve around the nucleus.

• Drawbacks:

  • Could not explain atomic stability.

  • Did not explain atomic spectra.

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🔹 5. Bohr’s Model of Atom

• Electrons revolve in discrete circular orbits (energy levels).

• Energy is quantized.

• Energy is absorbed/emitted when an electron jumps between orbits.

Postulates:

1. Electrons revolve in stable orbits without emitting energy.

2. Angular momentum is quantized:

   $$mvr = \dfrac{nh}{2\pi}$$

3. Energy is emitted/absorbed as:

   $$\Delta E = h\nu$$

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🔹 6. Atomic Number and Mass Number

• Atomic number (Z) = Number of protons = Number of electrons.

• Mass number (A) = Number of protons + Number of neutrons.

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🔹 7. Isotopes, Isobars, and Isotones

• Isotopes: Same atomic number, different mass number (e.g. ¹H, ²H, ³H).

• Isobars: Same mass number, different atomic number (e.g. ⁴⁰Ca, ⁴⁰Ar).

• Isotones: Same number of neutrons, different atomic and mass numbers.

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🔹 8. Electromagnetic Radiation

• Energy is emitted in the form of electromagnetic waves.

• Important terms:

  • Wavelength (λ): Distance between two crests/troughs.

  • Frequency (ν): Number of waves per second.

  • Speed of light (c) = 3 × 10⁸ m/s.

  • Relation:

    $$c = \lambda \nu$$

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🔹 9. Planck’s Quantum Theory

• Energy is emitted/absorbed in discrete packets (quanta or photons).

• Energy of quantum:

  $$E = h\nu$$

  where h = Planck's constant = 6.626 × 10⁻³⁴ Js.

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🔹 10. Photoelectric Effect

• Discovered by Einstein.

• Ejection of electrons from a metal surface when light of suitable frequency falls on it.

• Explained using quantum theory:

  $$h\nu = \phi + \dfrac{1}{2}mv^2$$

  (where φ = work function of the metal)

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🔹 11. Atomic Spectra

• Emission spectrum: Light emitted by atoms when electrons return to lower energy levels.

• Hydrogen spectrum: Explained by Bohr’s model.

Lyman (UV), Balmer (Visible), Paschen, Brackett, Pfund (IR)

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🔹 12. Bohr’s Energy Levels for Hydrogen Atom

• Radius of nth orbit:

  $$r_n = 0.529 \times \dfrac{n^2}{Z} \, \text{Å}$$

• Energy of nth orbit:

  $$E_n = -13.6 \times \dfrac{Z^2}{n^2} \, \text{eV}$$

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🔹 13. de Broglie’s Hypothesis

• Matter has wave-like properties.

• de Broglie wavelength:

  $$\lambda = \dfrac{h}{mv}$$

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🔹 14. Heisenberg’s Uncertainty Principle

• It is impossible to determine position and momentum of a particle simultaneously with absolute accuracy.

$$\Delta x \cdot \Delta p \geq \dfrac{h}{4\pi}$$

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🔹 15. Quantum Numbers

These describe the distribution of electrons:

Quantum No.	Symbol	Meaning
Principal	n	Shell/energy level (n = 1,2,3...)
Azimuthal	l	Subshell/shape (0 to n–1)
Magnetic	m	Orientation (–l to +l)
Spin	s	Spin direction (+½ or –½)

Subshells and Orbitals:

• s → l=0 → 1 orbital

• p → l=1 → 3 orbitals

• d → l=2 → 5 orbitals

• f → l=3 → 7 orbitals

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🔹 16. Aufbau Principle

• Electrons are filled in increasing order of energy.

• Order:
  1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s ...

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🔹 17. Pauli’s Exclusion Principle

• No two electrons in an atom can have the same set of four quantum numbers.

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🔹 18. Hund’s Rule of Maximum Multiplicity

• Electrons occupy all orbitals singly before pairing begins, to maximize total spin.

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🔹 19. Electronic Configuration

• Shows the arrangement of electrons.

• Example:

  • Hydrogen: 1s¹

  • Oxygen (Z=8): 1s² 2s² 2p⁴

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🔹 20. Limitations of Bohr’s Model

• Could not explain:

  • Zeeman effect (effect of magnetic field),

  • Stark effect (effect of electric field),

  • Spectrum of multi-electron atoms,

  • Dual nature of electron.

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