Atomic Structure and the Periodic Table
The Atomic Structure and the Periodic Table domain tests the nuclear model, isotopes, electron configuration, quantum numbers, orbital shapes, and periodic trends. Expect items that connect subatomic structure to position on the periodic table and to observable properties such as radius, ionization energy, and electronegativity.
Core rules and concepts
- Atomic number Z equals protons and identifies the element; mass number A equals protons plus neutrons; isotopes differ in neutrons.
- Average atomic mass is a weighted average of isotope masses by natural abundance.
- Electrons fill orbitals by Aufbau order, Pauli exclusion, and Hund’s rule; Cr and Cu are common exceptions.
- Quantum numbers: n (shell), l (subshell), ml (orientation), ms (spin). An orbital holds at most two electrons with opposite spins.
- Across a period, atomic radius decreases and ionization energy, electron affinity, and electronegativity generally increase. Down a group, radius increases and ionization energy generally decreases.
Worked example
Write the ground-state configuration of Fe³⁺. Neutral Fe is [Ar]4s²3d⁶. Electrons are removed from the highest-n shell first, so remove the two 4s electrons, then one 3d electron: Fe³⁺ = [Ar]3d⁵. That leaves five 3d electrons, all unpaired by Hund’s rule. The ion is therefore paramagnetic, and its 3+ charge reflects the three electrons lost, not a change in nuclear charge.
Traps that cost marks
- Filling 4s before 3d but removing 4s before 3d when forming ions.
- Confusing mass number with average atomic mass.
- Assuming periodic trends are identical within a group and across a period.
- Pairing electrons before singly occupying degenerate orbitals.
- Forgetting that noble-gas shorthand must preserve the correct valence configuration.