Ionisation Energy (OCR A Level Chemistry A): Revision Note

Periodic Trends In Ionisation Energy

• The first ionisation energy (IE₁) is the 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

  • E.g. the first ionisation energy of Na is:

Na (g) → Na⁺ (g) + e^-          First ionisation energy = +496 kJ mol^-1

Factors affecting ionisation energy

• Ionisation energies show periodicity

• The first ionisation energy increases across a period and decreases down a group and is caused by three factors that influence the ionisation energy:

  1. Atomic radius: 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

  2. 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

  3. Electron shielding: 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

• These three factors affect the attraction between the nucleus and the outer electrons of an atom, and therefore, the ionisation energy.

Trend in first ionisation energy down a group

• As you move down a group, the nuclear charge increases but the ionisation energy decreases

• This is due to the following factors:

  • The atomic radius increases

  • The shielding (by inner shell electrons) increases

  • Therefore, the attraction between the nucleus and the outer electrons decreases 

Trend in first 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 (no significant change in atomic radius)

  • 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 first ionisation energy 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 first ionisation energy between nitrogen and oxygen as the paired electrons in the 2p subshell of oxygen repel each other, making it easier to remove an electron in oxygen than nitrogen.

  • 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

Table explaining the Ionisation Energy Trends Across a Period & Down a Group

Predicting Ionisation Energy

• 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

• Take 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 ionisation energies

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  repulsion of the paired 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 shell that 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 sub-shell