Overview
This lecture covers the structure of the atom, focusing on the discovery and properties of subatomic particles, atomic models, quantum numbers, electron configurations, and related principles important for understanding atomic structure.
Discovery of Subatomic Particles
- Dalton's atomic theory described atoms as indivisible, but later discoveries revealed subatomic particles.
- Electron discovered by J.J. Thomson using the cathode ray tube experiment; electron is negatively charged.
- Protons are positively charged particles found using canal rays (anode rays); discovered by Goldstein.
- Neutron discovered by James Chadwick by bombarding beryllium with alpha particles; neutrons are neutral.
Atomic Models
- Thomson's "Plum Pudding" model: atom is a sphere of positive charge with embedded electrons.
- Rutherford's gold foil experiment revealed atoms have a small, dense, positively charged nucleus; most of the atom is empty space.
- Rutherford's model couldn't explain atomic stability or electron arrangement.
- Bohr's model: electrons revolve in fixed orbits with quantized energy levels; explained hydrogen’s spectrum.
Quantum Numbers & Atomic Orbitals
- Four quantum numbers (n, l, m, s) uniquely describe each electron in an atom.
- Principal quantum number (n): main energy level (shell).
- Azimuthal (l): subshell shape (s, p, d, f).
- Magnetic (m): orbital orientation.
- Spin (s): direction of electron spin.
- Orbitals are regions with high probability for finding an electron; S—spherical, P—dumbbell, D—double-dumbbell shapes.
Electron Configuration & Rules
- Electrons fill orbitals in increasing energy (Aufbau principle).
- No two electrons can have same set of quantum numbers (Pauli exclusion).
- Electrons enter degenerate orbitals singly before pairing (Hund’s rule).
- Half-filled and fully-filled subshells (d5, d10) are especially stable due to exchange/symmetry energy.
Electromagnetic Radiation & Dual Nature
- Electromagnetic radiation has dual nature: wave and particle (Planck’s quantum theory).
- Energy of photon: E = hν (h = Planck’s constant, ν = frequency).
- Photoelectric effect: electrons ejected from metal surface when light of threshold frequency shines.
Atomic Spectra & Series
- Atoms emit/absorb light at specific wavelengths—lines in spectra (Lyman—UV, Balmer—visible, Paschen/Brackett/Pfund—infrar).
- Spectral lines explained by electron transitions between quantized orbits.
De Broglie Hypothesis & Heisenberg Uncertainty
- Matter has wave-like properties (λ = h/mv).
- Heisenberg uncertainty principle: it's impossible to know both position and momentum of electron exactly.
Key Terms & Definitions
- Electron — Negatively charged subatomic particle.
- Proton — Positively charged subatomic particle in the nucleus.
- Neutron — Electrically neutral subatomic particle in the nucleus.
- Quantum Numbers — Set of four numbers describing electron position and energy.
- Photoelectric Effect — Ejection of electrons from a material by light.
- Aufbau Principle — Electrons fill lower-energy orbitals first.
- Pauli Exclusion Principle — No two electrons can have same quantum numbers.
- Hund’s Rule — Electrons fill degenerate orbitals singly before pairing.
- Isotopes — Atoms of same element with different numbers of neutrons.
- Orbital — Region of maximum probability for finding an electron.
- Ionization Energy — Energy required to remove an electron from an atom.
Action Items / Next Steps
- Review all quantum numbers and their significance.
- Practice electron configuration for various elements.
- Complete assigned homework problems, especially on de Broglie wavelength, spectra, and quantum numbers.
- Prepare summary notes for exceptions in electronic configuration (Cr, Cu).
- Attempt all practice and past exam questions provided in the lecture.