S1.3 Electron configurationsIB Chemistry HL: Revision notes

Shorter wavelength means higher frequency and higher photon energy (c = λf, E = hf). A continuous spectrum contains all wavelengths; a line spectrum contains only specific wavelengths. Emission lines are produced when excited electrons fall to lower discrete energy levels. In hydrogen, transitions to n = 1 give the ultraviolet series, to n = 2 the visible series and to n = 3 the infrared series. The levels converge at higher energy, so lines in a series converge at high frequency.

Level n holds 2n² electrons. Sublevels s, p, d, f have 1, 3, 5, 7 orbitals, each holding two electrons of opposite spin (Pauli). s orbitals are spherical; the three p orbitals are dumbbell-shaped along x, y and z. Fill lowest energy first (Aufbau: 1s 2s 2p 3s 3p 4s 3d 4p) and occupy equal-energy orbitals singly with parallel spins before pairing (Hund). Orbital (arrow-in-box) diagrams show each orbital as a box holding up to two opposite half-arrows.

Write full (1s²2s²2p⁶…) or condensed ([Ar]4s²3d⁶) configurations up to Z = 36. Exceptions: Cr = [Ar]4s¹3d⁵ and Cu = [Ar]4s¹3d¹⁰. When transition metals form ions, 4s electrons are lost first: Fe³⁺ = [Ar]3d⁵, Cu⁺ = [Ar]3d¹⁰, Mn²⁺ = [Ar]3d⁵.

The first ionisation energy is the energy needed to remove one mole of electrons from one mole of gaseous atoms: X(g) → X⁺(g) + e⁻. In hydrogen's spectrum the convergence limit of the series ending at n = 1 corresponds to an electron moving between n = 1 and n = ∞, i.e. ionisation.

Calculation: from the limit wavelength, f = c ÷ λ, then E = hf per atom, then multiply by Avogadro's constant and divide by 1000 for kJ mol⁻¹. For λ = 91.2 nm: f = 3.29 × 10¹⁵ Hz, E = 2.18 × 10⁻¹⁸ J, IE = 1.31 × 10³ kJ mol⁻¹.

Across a period IE generally increases: nuclear charge rises while electrons enter the same level with similar shielding, so the radius decreases and attraction increases.

Down a group IE decreases: the outer electron is in a higher level, further from the nucleus and more shielded, which outweighs the greater nuclear charge.

Discontinuities in period 3: Mg > Al because Al's electron is removed from 3p (higher energy, shielded by 3s); P > S because S's electron is removed from a doubly occupied 3p orbital and is repelled by its partner. The same pattern appears at Be/B and N/O in period 2.

Successive IEs always increase, because each electron is removed from an increasingly positive ion with less repulsion. A large jump shows an electron being removed from a new main energy level closer to the nucleus. The number of electrons removed before the first large jump equals the number of outer electrons, so it gives the group: a jump after IE2 means group 2; after IE3, group 13. Smaller jumps within a level can reveal sublevels (e.g. aluminium's IE1 from 3p, IE2 and IE3 from 3s).