Ionisation energies

• Ionisation energy measures how easily an atom loses electrons to form positive ions.
• The first ionisation energy is the energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.
• Three factors affect the attraction between the nucleus and the outer electrons of an atom and therefore ionisation energy:

  • atomic radius
  • nuclear charge
  • electron shielding
• Effect of atomic radius

  The greater the distance between the nucleus and the outer electrons the less the nuclear attraction.
• Effect of nuclear charge

  The more protons that are in the nucleus of an atom, the greater the attraction between the nucleus and the outer electrons.
• Effect of electron shielding

  Electrons are negatively charged and so inner-shell electrons repel outer-shell electrons. This repulsion, called the shielding effect, reduces the attraction between the nucleus and the outer electrons.
• Trend in first ionisation energy down a group

  First ionisation energy decreases down every group in the periodic table as:

  • atomic radius increases
  • more inner shells so shielding increases
  • nuclear attraction on outer electrons decreases
• Trend in first ionisation energy across a period

  The first ionisation energy across periods increases as:

  • nuclear charge increases
  • same shell so similar shielding
  • nuclear attraction increases
  • atomic radius decreases
• Comparing beryllium to boron

  The fall in the first ionisation energy from beryllium to boron marks the start of filling the 2p sub-shell. The 2p subshell in boron has a higher energy than the 2s subshell in beryllium. Therefore, in boron the 2p electron is easier to remove than one of the 2s electrons in beryllium. The first ionisation energy of boron is less than the first ionization energy of beryllium.
• Comparing nitrogen and oxygen

  The fall in first ionisation energy from nitrogen to oxygen marks the start of electron pairing in the p-orbitals of the 2p subshell. They both have the highest energy electrons in the 2p subshell. In oxygen the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from an oxygen atom than a nitrogen atom. Therefore the first ionisation energy of oxygen is less than the first ionisation energy of nitrogen.