Shells (2, 8, 8) become subshells and orbitals

Electrons occupy shells that hold 2, 8, 8, (18). You draw concentric circles with crosses. Electron configuration for calcium is 2,8,8,2.

Main energy level n holds a maximum of 2n² electrons (Structure 1.3.3), but is subdivided into s, p, d, f sublevels of successively higher energy (1.3.4), and each sublevel consists of orbitals — regions of probability, each holding two electrons of opposite spin (1.3.5). Configurations follow the aufbau principle, Hund's rule and the Pauli exclusion principle. Calcium is 1s²2s²2p⁶3s²3p⁶4s². 4s fills before 3d — and then, crucially, in transition-metal ions, the 4s electrons are removed first (Fe²⁺ is [Ar]3d⁶, not [Ar]3d⁴4s²). Cr and Cu are anomalous ([Ar]3d⁵4s¹, [Ar]3d¹⁰4s¹). Evidence for this structure comes from emission spectra (1.3.1–1.3.2, converging lines) and successive ionisation energy data (1.3.7).

Three things at once. (a) Structure: shells split into sublevels, sublevels into orbitals. (b) Ordering: the 4s/3d crossover breaks the naive "fill outwards" rule. (c) Epistemology: the model is now justified by experimental evidence (spectra, IEs) rather than asserted — and IB Paper 1 will ask you to read ionisation-energy data and deduce a configuration or a group number.

2,8,8 is enough to predict ionic charges for groups 1, 2, 6, 7 and to place elements in the periodic table, which is all GCSE needs. Orbitals require quantum ideas GCSE cannot support.

Cannot write configurations for transition metals or their ions (which kills Structure 3.1.8–3.1.10 — variable oxidation states, coloured complexes, d-orbital splitting); cannot explain the small dips in first ionisation energy between groups 2→13 and 15→16 (Structure 3.1.7), which is a guaranteed exam question; cannot understand hybridisation or σ/π bonding at HL; misreads successive-IE graphs. Also a subtle one: students keep saying "shell" when they mean "sublevel", and lose marks for imprecision.