• The?ionisation?energy?(IE) of an element is the amount of energy required to remove?one mole?of electrons from?one mole?of atoms of an element in the gaseous state to form?one mole?of gaseous ions
• Ionisation energies are measured under?standard conditions?which are 298 K and 100 kPa
• The units of IE are?kilojoules?per?mole?(kJ mol^-1)
• The?first ionisation energy?is the energy required to remove the one mole of electrons from one mole the gaseous atoms
  • E.g. the first ionisation energy of calcium:

Ca(g) → Ca⁺?(g) + e^{-? ? ? ? ? ??}1st? ??H I.E. = +590 kJ mol^-1

• 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 IEs
• The?first?ionisation?energy?increases?across a period and?decreases?down a group and is 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 outer 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: paired electrons in the same atomic orbital in a subshell repel each other more than electrons in different atomic orbitals; this 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 across 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 remains the same
• 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 1st I.E. 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?801 kJ mol^-1?as its electron configuration is?1s²?2s²?2p¹
• There is a slight?decrease?in 1st I.E. between?nitrogen?and?oxygen?due to?spin-pair repulsion?in the 2p subshell of oxygen
  • Nitrogen?has a first ionisation energy of?1402 kJ mol^-1?as its electron configuration is 1s²?2s²?2p³
  • Oxygen?has a first ionisation energy of?1314 kJ mol^-1?as its electron configuration is 1s²?2s²?2p⁴

Ionisation energy down a group

• Although going down a group the?nuclear charge?increases, the ionisation energy down a group decreases and it is due to the following factors:
  • The?distance?between the nucleus and outer electron?increases
  • The?shielding?by inner shell electrons?increases
  • The?effective nuclear charge?is decreasing as?shielding?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 values become very large and difficult to represent meaningfully, so it is more convenient to show the logarithm of the ionisation energies
• This helps us to see significant jumps in I.E.

Successive ionisation energies for the element calcium

• The?first?electron removed has a low ionisation energy 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 a little more difficult to remove than the first electron as you are removing an electron from a positively charged ion
• 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)
• The graph shows there is a large increase in successive ionisation energy as the electrons are being removed from an increasingly positive ion
• The big jumps on the graph show the change of?shell?and the small jumps are the change of?subshell