Atomic Structure and the Periodic Law
Core concepts
Every candidate must move fluently between the nuclear model (protons, neutrons, isotopes), the electron-configuration model (shells, subshells, orbitals, the aufbau order, Hund's rule, Pauli exclusion), and the periodic table as a map of those configurations. Know that group number follows valence-shell count for main-group elements, that period number equals the highest occupied shell, and that periodic trends (atomic radius, ionisation energy, electron affinity, electronegativity, metallic character) all follow from effective nuclear charge and shielding.
Traps that appear in questions
- Isotopes vs ions: isotopes differ in neutrons only; an ion differs in electrons. Questions often ask for the particle count of a specific nuclide, such as protons, neutrons and electrons in a sulfide ion.
- Exceptions to the aufbau order: chromium and copper-type anomalies are asked; justify them by the extra stability of half-filled or filled subshells.
- Ionisation energy discontinuities: a large jump marks the shell from which the electron is removed, a classic way to identify an element's group.
- Radius comparisons across a period vs down a group: across a period the radius contracts because nuclear charge grows while shielding stays roughly constant; down a group it expands because new shells are added.
- Quantum numbers: know which combinations are allowed (n, l, m, spin) and that l ranges from 0 to n−1.
Worked example
Identify the group of an element whose successive ionisation energies (kJ/mol) are 590, 1150, 4940, 6480. The first two are moderate, the third is roughly four times the second, so the third electron comes from a new, inner shell. The element has two valence electrons, so it is in group 2.
Rule of thumb
Electron configuration first, then everything else: trends, ion charges and bonding all follow from the valence configuration.