First ionisation energy: The first ionisation energy of an element is the energy required to remove an electron from each atom in one mole of atoms in gaseous atoms

In general:

Increase across the period
- Increasing number of protons → Increasing nuclear charge → Electrons are lower in energy → More energy needed to remove them → Higher first ionisation energy
Decrease down the group
- Electrons added to new quantum shell → New shell has higher energy and electrons are more shielded by lower electrons → Higher energy electrons → Less energy needed for removal → Lower first ionisation energy

However, there are some anomalies:

Ionisation energy of Group 3 element is less than that of Group 2 element
- Example: Ionisation energy of beryllium is greater than that of boron
- Orbital energies
  - Outer electron of boron is in p orbital
  - Outer electron of beryllium is in s orbital
  - Therefore, outer electron of boron is higher in energy → Easier to remove than outer electron of beryllium → Lower ionisation energy than beryllium
- Shielding
  - Outer electron of boron experiences shielding from two inner electrons subshells
  - Outer electron of beryllium experiences shielding from only one inner electron subshell
  - Therefore, outer electron of beryllium is lower in energy→ Electron on beryllium feels more attractive force of nucleus → Beryllium has higher ionisation energy
Ionisation energy of Group 6 element is less than that of Group 5 element
- Example: Ionisation energy of nitrogen is greater than that of oxygen
- Paired electrons
  - Presence of two electrons in one orbital increases electron-electron repulsion → Electrons will be higher in energy
  - First electron removed from oxygen comes from a pair → Higher in energy → Requires less energy to remove than unpaired outer electron of nitrogen

Periodicity