Electron Configuration Rules
Principles for filling atomic orbitals
Understanding Electron Configuration
Electron configuration describes the distribution of electrons among atomic orbitals in an atom. It's fundamental to understanding chemical bonding, periodic properties, magnetic behavior, and spectroscopic characteristics. The arrangement of electrons determines how atoms interact with light and form bonds with other atoms, making it central to all of chemistry.
Developed from quantum mechanics in the early 20th century, electron configuration follows specific rules derived from the physics of atomic orbitals. Niels Bohr's model introduced energy levels, while later quantum mechanical treatments by Schrödinger, Heis enberg, and Pauli provided the mathematical foundation for modern orbital theory and the principles governing electron arrangement.
The electron configuration notation uses numbers (principal quantum number n) and letters (orbital type: s, p, d, f) with superscripts showing electron count. For example, 1s² 2s² 2p⁶ represents neon's configuration with completely filled first and second shells, explaining its chemical inertness as a noble gas.
Three Fundamental Principles
1. Aufbau Principle
Fill orbitals from lowest to highest energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p...
The n+l rule predicts ordering: lower (n+l) fills first; if tied, lower n fills first. Example: 4s (n+l=4) before 3d (n+l=5).
2. Pauli Exclusion Principle
No two electrons can have identical quantum numbers. Maximum 2 electrons per orbital with opposite spins (↑↓).
Wolfgang Pauli's 1925 principle explains orbital capacity limits and why matter has volume.
3. Hund's Rule
Fill degenerate orbitals singly with parallel spins before pairing. Minimizes electron-electron repulsion.
Example: p³ is ↑ ↑ ↑ (all singly occupied) not ↑↓ ↑ (one paired).
Detailed Example: Oxygen (Z = 8)
Step 1: Total electrons = 8
Step 2: Fill by Aufbau: 1s(2), 2s(2), 2p(4)
Step 3: Apply Hund's to 2p⁴: ↑↓ ↑ ↑
Configuration: 1s² 2s² 2p⁴
Answer: 1s² 2s² 2p⁴
Oxygen has 2 unpaired electrons (paramagnetic)
Orbital Capacity
| Subshell | # Orbitals | Max Electrons | Shape |
|---|---|---|---|
| s | 1 | 2 | Spherical |
| p | 3 | 6 | Dumbbell |
| d | 5 | 10 | Complex |
| f | 7 | 14 | Very complex |
Notable Exceptions
Chromium (Cr, Z=24)
Expected: [Ar] 4s² 3d⁴
Actual: [Ar] 4s¹ 3d⁵
Half-filled d⁵ provides extra stability
Copper (Cu, Z=29)
Expected: [Ar] 4s² 3d⁹
Actual: [Ar] 4s¹ 3d¹⁰
Fully filled d¹⁰ more stable