Atomic structure and the periodic table
Atomic structure and the periodic table underpins explanations of bonding, periodicity and reactivity. The European Baccalaureate tests your ability to move between subatomic particles, electron configurations and trends, and to explain trends rather than merely state them.
Core rules and models
- An atom is characterised by its atomic number Z (protons) and mass number A (protons + neutrons). Isotopes have the same Z but different A; they have identical chemical behaviour because chemistry depends on electrons.
- Relative atomic mass Ar is the weighted mean of isotope masses on the 12C scale: Ar = Σ(abundance × isotope mass)/100.
- Electrons occupy orbitals in order of increasing energy: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p. Apply the Pauli exclusion principle (maximum two electrons per orbital, opposite spins) and Hund's rule (parallel spins in degenerate orbitals).
- Chromium [Ar]3d5 4s1 and copper [Ar]3d10 4s1 are exceptions because half-filled and fully filled d subshells gain exchange stability.
- Across a period, nuclear charge increases while shielding changes little, so atomic radius decreases and first ionisation energy generally increases. Down a group, shells are added, shielding increases and the outer electron is farther from the nucleus, so ionisation energy decreases and atomic radius increases.
Worked example
Explain why the first ionisation energy of aluminium (577 kJ mol−1) is lower than that of magnesium (738 kJ mol−1). Mg is [Ne]3s2; Al is [Ne]3s2 3p1. The electron removed from Al is in a 3p orbital, which is higher in energy and more shielded by the 3s electrons than the 3s electron removed from Mg. Although Al has one more proton, the 3p penetration is poorer, so less energy is needed. This is a genuine periodic dip, not an anomaly.
Traps that cost marks
- Do not confuse mass number with relative atomic mass; only Ar is a weighted average and non-integer.
- State that ionisation energy refers to gaseous atoms; values differ in solution.
- When writing transition-metal configurations, remember 4s fills before 3d but is emptied first when forming ions (e.g. Fe2+ is [Ar]3d6, not [Ar]3d4 4s2).
- Explain dips at groups 13 and 16 by orbital type and electron pairing, not by nuclear charge alone.
- Successive ionisation energies increase sharply after a shell is emptied; use this to deduce group number.