Electronic Configuration and the Periodic Table
This domain links electron configuration to the structure of the periodic table and to periodic trends. STPM tests both the rules for filling orbitals and explanations of anomalies.
Rules for filling orbitals
Use the Aufbau principle: lowest energy orbitals fill first. The n + l rule orders subshells; 4s (n + l = 4) fills before 3d (n + l = 5). Pauli exclusion limits each orbital to two electrons with opposite spins. Hund’s rule requires single occupancy of degenerate orbitals before pairing. Chromium is [Ar]3d5 4s1 and copper is [Ar]3d10 4s1 because half-filled and fully filled d subshells are more stable.
Periodic trends
Atomic radius decreases across a period because nuclear charge increases while shielding stays similar; it increases down a group because extra shells are added. First ionisation energy generally increases across a period and decreases down a group. Anomalies occur at Be > B because Be has a filled 2s subshell, and at N > O because N has a half-filled 2p subshell. Electron affinity becomes more negative across a period; electronegativity increases across a period and decreases down a group.
Blocks and transition metals
The s-block is Groups 1–2, the p-block Groups 13–18, the d-block transition metals, and the f-block lanthanides and actinides. Transition metals show variable oxidation states and form coloured complex ions.
Worked example
Iron (Z = 26) is [Ar]3d6 4s2. Fe2+ forms by removing the two 4s electrons first: [Ar]3d6. Fe3+ is [Ar]3d5. The half-filled d5 configuration helps explain the relative stability of Fe3+.
Traps that cost marks
When forming cations, remove 4s electrons before 3d electrons. Remember the Cr and Cu exceptions. The definition of ionisation energy must include gaseous atoms or ions. Down a group, increased shielding and distance outweigh increased nuclear charge, so ionisation energy falls.