Overview
This lecture covers the structure of atoms, discovery and characteristics of subatomic particles, historical atomic models, quantum mechanical concepts, rules for electron arrangement, and electronic configuration of atoms.
Discovery of Subatomic Particles
- Atoms are made of smaller particles: electrons (negative), protons (positive), and neutrons (neutral).
- Cathode ray experiments showed electrons are present in all atoms.
- J.J. Thomson measured electron's charge-to-mass ratio (e/me = 1.76 × 10¹¹ C/kg).
- Millikan oil drop experiment measured electron's charge (1.602 × 10⁻¹⁹ C).
- Protons discovered as positive particles; neutrons discovered by Chadwick (mass slightly greater than protons).
Atomic Models
- Thomson's model: Atom is a sphere of positive charge with embedded electrons (plum pudding model).
- Rutherford's model: Atom has a small, dense, positively charged nucleus with electrons orbiting around it.
- Rutherford couldn't explain atom's stability and electron arrangement.
Atomic Number, Mass Number, Isotopes, and Isobars
- Atomic number (Z): Number of protons, equals electrons in a neutral atom.
- Mass number (A): Total protons + neutrons.
- Isotopes: Same atomic number, different mass numbers (e.g., ¹H, ²H, ³H).
- Isobars: Same mass number, different atomic numbers.
Electromagnetic Radiation & Quantum Theory
- Electromagnetic waves consist of oscillating electric and magnetic fields, travel at speed of light (c = 3 × 10⁸ m/s).
- Frequency (ν), wavelength (λ), and energy related by c = λν.
- Planck: Energy is quantized (E = hν), where h = 6.626 × 10⁻³⁴ Js.
- Photoelectric effect: Electrons emitted from metals if light exceeds threshold frequency; explained by Einstein using photons.
Atomic Spectra and Bohr's Model
- Atoms emit absorption and emission spectra; hydrogen shows distinct line spectra (Balmer, Lyman series, etc.).
- Bohr's postulates: Electrons orbit nucleus in quantized energy levels; energy emitted/absorbed when electrons jump between orbits.
- Energy of nth orbit: Eₙ = -2.18 × 10⁻¹⁸ J/n².
- Bohr's model explains hydrogen's spectrum but fails for multi-electron atoms.
Quantum Mechanical Model
- Electrons have dual wave-particle nature (de Broglie relation: λ = h/mv).
- Heisenberg Uncertainty Principle: Cannot know position and momentum of electron simultaneously.
- Quantum mechanics uses wave functions (ψ), probability densities (|ψ|²), and quantum numbers (n, l, mₗ, mₛ) to describe electrons.
- Orbitals defined by quantum numbers; s, p, d, f shapes/energies.
Arrangement and Configuration of Electrons
- Aufbau principle: Electrons fill orbitals from lowest to highest energy.
- Pauli Exclusion Principle: No two electrons in an atom have same set of four quantum numbers; max two electrons per orbital with opposite spins.
- Hund's Rule: Electrons singly occupy all degenerate orbitals before pairing.
- Stability observed in completely or half-filled subshells (extra symmetry/exchange energy).
Key Terms & Definitions
- Electron — Negatively charged subatomic particle found in all atoms.
- Proton — Positively charged particle in atom's nucleus.
- Neutron — Neutral particle in nucleus, mass slightly more than proton.
- Atomic Number (Z) — Number of protons in the nucleus.
- Mass Number (A) — Sum of protons and neutrons.
- Isotope — Same atomic number, different mass numbers.
- Isobar — Same mass number, different atomic numbers.
- Quantum Numbers — Set of numbers (n, l, mₗ, mₛ) describing unique electron states.
- Aufbau Principle — Order of electron filling in orbitals, lowest first.
- Pauli Exclusion Principle — No two electrons can share all four quantum numbers.
- Hund's Rule — Maximum multiplicity: electrons occupy orbitals singly before pairing.
Action Items / Next Steps
- Solve provided exercises to practice calculations on atomic structure, spectra, and quantum numbers.
- Review the tables of electron configurations for elements.
- Read next unit for further details on chemical bonding and periodic trends.