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.

  For example:

  $Na(g) -> Na^+$$(g) +$$e^-$ First ionisation energy = +496 KJ/mol
• The three factors that affect ionisation energy are: atomic radius, nuclear charge, electron shielding. These affect the attraction between the nucleus and the outer electrons which is why they affects ionisation energy.
• The larger the atomic radius, the less the nuclear attraction because the nucleus is further away from the outer electrons so the ionisation energy decreases.

  The force of attraction falls off sharply with increasing distance, so atomic radius has large affect on ionisation energy
• The more protons in a nucleus of an atom (greater nuclear charge), the greater the attraction between the nucleus and outer electrons so the ionisation energy is greater.
• Electrons are negatively charged so inner shell electrons repel outer shell electrons. This repulsion, called the shielding effect, reduces the attraction between the outer shell electrons and the nucleus.
• The second ionisation energy is the energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions
• Successive ionisation energies provide important evidence for the different electron energy levels in an atom.
• A large increase in ionisation energies indicates a change in the shell that the electron is being removed from. The electron is now being removed from a shell that is closer to the nucleus and has less shielding.
• Successive ionisation energies allow predictions to be mad about:

  • the numbers of electrons in the outer shell
  • the group of the element in the periodic table
  • the identity of an element
• Across each period there is a general increase in first ionisation energy because shielding stays the same, as the outer electrons are in the same shell, but nuclear charge increases so nuclear attraction increases. Also, the atomic radius decreases because of the increasing nuclear charge pulling outer shells in.

  However there is a sharp decrease in first ionisation energy between the end of one period to the start of the next period.
• First ionisation energies decrease going down a group. This is because atomic radius increases, more inner shells so shielding increases, so nuclear attraction on outer electrons decreases.
• Although the first ionisation energy shows a general increase across both period 2 and period 3, it does fall in two places in each period. The drops occur in the same positions in each period, suggesting a periodic cause. The reason is linked to the existence of sub-shells, their energies, and how orbitals fill with electrons.
• Across period 2, the first ionisation energy graph shows three and two falls:

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

  So, the first ionisation energy of boron is less than of beryllium.
• The fall in the first ionisation energy from nitrogen to oxygen marks the start of electron pairing in the p-orbitals of the 2p sub-shell.
• In nitrogen and oxygen, the highest energy electrons are in the 2p sub-shell.

  In oxygen, the paired electrons in one of the p-orbitals repel one another, making it easier to remove an electron from an oxygen atom than from a nitrogen atom.

  Therefore the first ionisation energy of oxygen is less than the first ionisation energy of nitrogen.