Atomic Structure and the Periodic Table
What you must understand
Start from the nuclear model: protons and neutrons form the nucleus, electrons occupy quantised orbitals described by the principal quantum number n and by subshell labels s, p, d, f. Be fluent in the four quantum numbers and the rules that generate electron configurations: the Aufbau order, the Pauli exclusion principle, and Hund's rule. Know how to write configurations in full and in noble-gas shorthand, and how to derive the configuration of ions by removing electrons from the highest-n subshell first (4s before 3d for transition metals).
Periodic trends follow from effective nuclear charge and shielding: atomic radius decreases across a period and increases down a group; ionisation energy generally rises across a period with characteristic dips at group 13 and group 16; electron affinity and electronegativity follow similar patterns. Metallic character decreases across a period.
Traps to watch for
- Claiming 4s electrons are removed after 3d when forming transition-metal cations; the opposite is true.
- Confusing the number of orbitals (one s, three p, five d) with the number of electrons they can hold (double those values).
- Forgetting that chromium and copper adopt anomalous configurations, and that half-filled and fully filled subshells gain extra stability.
- Treating ionisation energy as strictly monotonic across a period; the group 13 and 16 dips are real.
Worked example
Write the configuration of Fe and Fe2+. Fe is [Ar]3d6 4s2, since 26 electrons fill 4s before 3d. For Fe2+, remove the two electrons of highest n first: the 4s pair goes, leaving [Ar]3d6. A common error is writing [Ar]3d4 4s2, which is incorrect.
Rule of thumb
For cations of d-block metals, empty the s subshell before touching the d subshell; then check the total electron count adds up.