Overview
This lecture covers atomic electron configuration, sublevels, quantum numbers, orbital filling rules, magnetic properties of elements, and how to interpret and calculate related questions.
Electron Configuration & Sublevels
- Electron configuration represents how electrons are arranged in an atom's orbitals.
- Each energy level (n) contains sublevels (s, p, d, f) whose number increases with n.
- s holds 2 electrons, p holds 6, d holds 10, and f holds 14 electrons.
- Use the periodic table to determine electron configuration (e.g., fluorine: 1s² 2s² 2p⁵).
Orbital Diagrams & Filling Rules
- Orbitals are represented by boxes; each can hold 2 electrons with opposite spins.
- Hund’s Rule: Add electrons to degenerate (same energy) orbitals singly with parallel spins before pairing.
- Aufbau Principle: Fill the lowest available energy levels first.
- Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.
Paramagnetism & Diamagnetism
- Atoms with unpaired electrons are paramagnetic (attracted to magnetic fields).
- Atoms with all electrons paired are diamagnetic (weakly repelled by magnetic fields).
- Example: Fluorine is paramagnetic (unpaired electron), phosphide ion is diamagnetic (all paired).
Quantum Numbers
- Principal quantum number (n): main electron energy level.
- Angular momentum quantum number (l): sublevel (0=s, 1=p, 2=d, 3=f; l ranges from 0 to n-1).
- Magnetic quantum number (ml): specifies orbital within sublevel (-l to +l).
- Spin quantum number (ms): electron spin (+½ or -½).
Identifying Elements & Electron Counts
- You can identify elements by summing exponents in electron configurations to get atomic numbers.
- To find paired/unpaired electrons, examine orbital diagrams or subtract unpaired from total electrons.
Valence & Core Electrons
- Valence electrons are in the highest energy level; core electrons are all others.
- Groups in the periodic table can be found by adding the exponents in the outer configuration (e.g., ns²np⁵ = Group 7A Halogens).
Maximum Electrons & Orbitals in Energy Levels
- Maximum electrons per level: 2n²; maximum orbitals: n².
- Sublevels and their orbitals per level: s(1), p(3), d(5), f(7), etc.
Key Terms & Definitions
- Aufbau Principle — electrons fill lowest energy orbitals first.
- Hund’s Rule — electrons occupy equal-energy orbitals singly before pairing.
- Pauli Exclusion Principle — no two electrons have identical quantum numbers in an atom.
- Quantum Numbers — set of four numbers (n, l, ml, ms) specifying an electron’s state.
- Valence Electrons — electrons in the outermost energy level.
- Paramagnetism — property of materials with unpaired electrons, attracted to magnets.
- Diamagnetism — property of all electrons paired, weakly repelled by magnets.
Action Items / Next Steps
- Practice writing electron configurations and orbital diagrams for various elements.
- Memorize Quantum number constraints and relationships.
- Solve practice problems on paramagnetism/diamagnetism and quantum numbers.
- Review periodic table groups and electron configuration patterns.