Periodicity and ionisation energy · Grade 11
Elements in the periodic table are arranged by atomic number, and their properties repeat in each period: this is periodicity. Across a period atoms get smaller and hold their electrons more tightly; down a group they get bigger and lose their outer electrons more easily. Ionisation energies show this most clearly, and they are evidence for shells and sub-shells.
Key facts
- First ionisation energy: the energy needed to remove one electron from each atom in one mole of gaseous atoms, forming one mole of 1+ ions: Na(g) → Na⁺(g) + e⁻.
- Three things decide it: the nuclear charge, the distance of the outer electron from the nucleus, and shielding by the inner electrons.
- Down a group first ionisation energy decreases: the outer electron is further from the nucleus and more shielded, even though the nuclear charge is bigger.
- Across a period it generally increases: more protons, the same shell and similar shielding. There are two dips: at group 3 the electron comes from a p sub-shell, higher in energy than s; at group 6 it comes from a pair of electrons that repel each other.
- Successive ionisation energies always increase. A big jump shows the next electron comes from a shell closer to the nucleus, so the number of electrons before the jump gives the group.
- Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond. It increases across a period and up a group; fluorine is the most electronegative element (3.98 on the Pauling scale).
- Across period 3 the melting point rises from sodium to silicon (stronger metallic bonding, then a giant covalent lattice), drops for phosphorus, sulfur and chlorine (simple molecules) and is lowest for argon.