Periodicity and ionisation energy

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.

Worked examples

Example 1

The first ionisation energy of magnesium is 738 kJ/mol and that of aluminium is 578 kJ/mol. Explain why aluminium’s is lower, even though aluminium has more protons.

  1. Magnesium is [Ne] 3s². Aluminium is [Ne] 3s² 3p¹.
  2. The electron removed from aluminium is the 3p electron. A 3p sub-shell is higher in energy than 3s, and the 3p electron is slightly more shielded by the 3s electrons.
  3. So it is easier to remove, and this outweighs aluminium’s extra proton: 578 kJ/mol is less than 738 kJ/mol.
  4. The same thing happens between beryllium and boron in period 2.

Answer: Aluminium’s outer electron is in the 3p sub-shell, which is higher in energy than magnesium’s 3s, so less energy is needed to remove it.

Example 2

The first five ionisation energies of an element, in kJ/mol, are 578, 1817, 2745, 11 577 and 14 842. Which group is the element in?

  1. Each ionisation energy is bigger than the one before, because each electron is pulled from an ion that is more positive.
  2. The big jump is between the third and the fourth. 11 577 ÷ 2745 ≈ 4.2
  3. So three electrons are in the outer shell, and the fourth comes from a full inner shell, closer to the nucleus.
  4. Three outer electrons: group 3 (13). The element is aluminium.

Answer: group 3 (group 13 in the IUPAC numbering)

Example 3

Pauling electronegativities: carbon 2.55, chlorine 3.16. Which atom in a C–Cl bond is δ−, and why is chlorine more electronegative than carbon?

  1. Chlorine has the higher electronegativity, so it pulls the shared pair of electrons towards itself. 3.16 > 2.55
  2. Chlorine gets a partial negative charge, δ−, and carbon δ+: the bond is polar.
  3. Chlorine has more protons than carbon, and in group 7 it is near the top, so its bonding electrons are close to a large nuclear charge.

Answer: Chlorine is δ−. It has more protons and its bonding electrons are held more strongly.