• Ionisation energies show periodicity
• As could be expected from their electronic configuration, the group I metals show low IE whereas the noble gases have very high IE’s
• The first ionisation energy increases across a period and decreases down a group caused by four factors that influence the ionisation energy
  • Size of the nuclear charge: the nuclear charge increases with increasing atomic number, which means that there are greater attractive forces between the nucleus and electrons, so more energy is required to overcome these attractive forces when removing an electron
  • Distance of outer electrons from the nucleus: electrons in shells that are further away from the nucleus are less attracted to the nucleus so the further the outer electron shell is from the nucleus, the lower the ionisation energy
  • Shielding effect of inner electrons: the shielding effect is when the electrons in full inner shells repel electrons in outer shells preventing them to feel the full nuclear charge so the greater the shielding of outer electrons by inner electron shells, the lower the ionisation energy
  • Spin-pair repulsion: electrons in the same atomic orbital in a subshell repel each other more than electrons in different atomic orbitals which makes it easier to remove an electron (which is why the first ionization energy is always the lowest)

A graph showing the ionisation energies of the elements hydrogen to sodium

Ionisation energy across a period

• The ionisation energy over a period increases due to the following factors:
  • Across a period the nuclear charge increases
  • The distance between the nucleus and outer electron remains reasonably constant
  • The shielding by inner shell electrons remain reasonably constant
• There is a rapid decrease in ionisation energy between the last element in one period and the first element in the next period caused by:
  • The increased distance between the nucleus and the outer electrons
  • The increased shielding by inner electrons
  • These two factors outweigh the increased nuclear charge
• There is a slight decrease in IE₁ between beryllium and boron as the fifth electron in boron is in the 2p subshell which is further away from the nucleus than the 2s subshell of beryllium
  • Beryllium has a first ionisation energy of 900 kJ mol^-1 as its electron configuration is 1s² 2s²
  • Boron has a first ionisation energy of 800 kJ mol^-1 as its electron configuration is 1s² 2s² 2p_x¹
• There is a slight decrease in IE₁ between nitrogen and oxygen and phosphorus due to spin-pair repulsion in the 2p_x orbital of oxygen
  • Nitrogen has a first ionisation energy of 1400 kJ mol^-1 as its electron configuration is 1s² 2s² 2p_x¹ 2p_y¹ 2p_z¹
  • Oxygen has a first ionisation energy of 1310 kJ mol^-1 as its electron configuration is 1s² 2s² 2p_x² 2p_y¹ 2p_z¹

Ionisation energy down a group

• The ionisation energy down a group decreases due to the following factors:
  • Across a period the nuclear charge increases
  • The distance between the nucleus and outer electron increases
  • The shielding by inner shell electrons increases

Ionisation energy trends across a period & going down a group table

Successive ionisation energies of an element

• The successive ionisation energies of an element increase as removing an electron from a positive ion is more difficult than from a neutral atom
• As more electrons are removed the attractive forces increase due to decreasing shielding and an increase in the proton to electron ratio
• The increase in ionisation energy, however, is not constant and is dependent on the atom’s electronic configuration
• Taking calcium as an example:

Ionisation energies of calcium table

• The first electron removed has a low IE₁ as it is easily removed from the atom due to the spin-pair repulsion of the electrons in the 4s orbital
• The second electron is more difficult to remove than the first electron as there is no spin-pair repulsion
• The third electron is much more difficult to remove than the second one corresponding to the fact that the third electron is in a principal quantum shell which is closer to the nucleus (3p)
• Removal of the fourth electron is less difficult as the orbital is no longer full and there is less spin-pair repulsion