S1.3 Electron configurations

Energy level (shell)

A region around the nucleus where electrons of similar energy are found, given a principal quantum number n

Sub-level

A division within an energy level containing orbitals of the same energy and shape (s, p, d, f)

Orbital

A region of space where there is a high probability of finding an electron; each orbital holds a maximum of 2 electrons

Aufbau principle

Electrons fill orbitals starting from the lowest energy level available, before filling higher energy levels

Pauli exclusion principle

No two electrons in the same atom can have the same set of four quantum numbers; two electrons in the same orbital must have opposite spins

Hund's rule

Electrons fill degenerate (same-energy) orbitals singly, with parallel spins, before any orbital is doubly occupied

Number of orbitals per sub-level

s = 1, p = 3, d = 5, f = 7

Maximum electrons per sub-level

s = 2, p = 6, d = 10, f = 14

Full electron configuration

Written using the format showing energy level, sub-level, and number of electrons in each, e.g. 1s² 2s² 2p⁶ (filled in order of increasing energy)

Condensed (noble gas) electron configuration

Uses the symbol of the preceding noble gas in square brackets to represent its full configuration, followed by the remaining electrons, e.g. Na = [Ne] 3s¹

Electron configuration of ions

For cations, remove electrons from the outermost (highest n) sub-level first; for anions, add electrons following the normal filling order

Diamagnetic vs paramagnetic (HL)

Diamagnetic species have all electrons paired (weakly repelled by a magnetic field); paramagnetic species have one or more unpaired electrons (attracted into a magnetic field)

Exceptions to the Aufbau principle (HL)

Chromium and copper (and their ions in the same groups) have configurations ending in …3d⁵4s¹ and …3d¹⁰4s¹ respectively, because a half-filled or fully-filled d sub-level is more stable