Lecture Notes on Atomic Theory and Subatomic Particles

Introduction to Matter

• Definition: Any substance that has mass and occupies space is called matter.
• Atoms: Small particles called atoms that cannot be further divided.
  • Initially proposed by Democritus (460 BC).
  • Ignored for ~2000 years due to lack of scientific evidence.

Dalton's Atomic Theory (1808)

• John Dalton proposed the atomic theory of matter.
• Postulates:
  • All matter is composed of atoms that cannot be created or destroyed.
  • All atoms of a given element are identical.
  • Different elements have different atoms.
  • Chemical reactions involve rearrangement of atoms.
  • Compounds are formed from atoms of different elements.
• Limitations: Couldn’t explain the electric nature of matter and results of later experiments.

Discovery of Electric Nature of Matter

• Michael Faraday (1830): Proved electricity consists of charged particles through electrolyte experiments.

Cathode Ray Experiments

• William Crookes (1879): Studied electrical discharge through gases using cathode ray tubes.
  • Observed that cathode rays move from the cathode to the anode.
• J.J. Thomson (1897): Studied properties of cathode rays and concluded they are negatively charged particles (electrons).
  • Calculated charge/mass ratio of electron: 1.758820 × 10¹¹ C/kg.

Discovery of Electron Charge and Mass

• Robert Millikan (1909): Oil-drop experiment to calculate the charge of an electron (1.6 × 10⁻¹⁹ C) and mass (9.1094 × 10⁻³¹ kg).

Discovery of Protons

• Eugen Goldstein (1886): Discovered canal rays (protons) through perforated cathode experiments. Proton charge: +1.6 × 10⁻¹⁹ C, Mass: ~1.67 × 10⁻²⁴ g.

Discovery of Neutrons

• James Chadwick (1932): Proved presence of neutrons (neutral particles, mass similar to protons).

Atomic Models

Thomson’s Model (1898)

• Plum pudding model: Atom is a sphere of positive charge with embedded electrons.
• Limitations: Couldn’t explain the scattering experiment results and atomic stability.

Rutherford's Nuclear Model (1911)

• Gold Foil Experiment:
  • Most alpha particles pass through foil undeflected.
  • Some deflected by small/large angles, a few bounced back.
  • Concluded that atoms have a dense, positively charged nucleus.
• Rutherford's Model: Electrons orbit nucleus like planets around the sun.
  • Limitations: Couldn’t explain atomic stability or atomic spectra.

Isotopes and Isobars

• Isotopes: Atoms of the same element with different numbers of neutrons (e.g., Hydrogen isotopes: Protium, Deuterium, Tritium).
• Isobars: Atoms of different elements with the same mass number but different atomic numbers.

Bohr’s Atomic Model

• Postulates by Niels Bohr solved atoms' stability issue using quantized orbits.

Nature of Electromagnetic Radiations

• Electromagnetic waves: Oscillating electric and magnetic fields.
• Properties:
  • Wavelength (λ)
  • Frequency (ν)
  • Velocity (c)
  • Relationship: c = λν
  • Wave number (ν̄): 1/λ

Electromagnetic Spectrum

• Arrangement of EM radiations from radio waves to gamma rays.
• Characteristics of visible spectrum, UV, IR, etc.

Blackbody Radiation (Max Planck's Theory)

• Radiation emitted/absorbed in discrete packets (quanta).
• Energy of quantum: E = hν

Photoelectric Effect (Albert Einstein)

• Observations:
  • Emission of electrons when metals are exposed to light of sufficient frequency.
  • Kinetic energy of ejected electrons depends on light frequency, not intensity.
  • Equation: hν = hν₀ + 1/2 mv²
  • Work function (W₀): Minimum energy to eject an electron.

Atomic Spectra

• Emission Spectra: Produced when atoms emit energy as they return to ground states.
• Absorption Spectra: Produced when atoms absorb specific wavelengths.
• Hydrogen Spectrum: Multiple series (Lyman, Balmer, Paschen, etc.).
• Rydberg Formula: Used to predict wavelengths of hydrogen spectral lines.

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These notes provide a comprehensive summary of atomic theory and the discovery and study of subatomic particles, which are integral to understanding the structure of matter.