Ionisation Energy Trends (Cambridge (CIE) AS Chemistry)

Trends in Ionisation Energy

• Ionisation energies show periodicity - a trend across a period of the Periodic Table

• As could be expected from their electronic configuration, the group I metals have a relatively low ionisation energy, whereas the noble gases have very high ionisation energies

• The size of the first ionisation energy is affected by four factors:

  • 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 - the nuclear attraction is weaker - 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 from feeling the full nuclear charge, so the more shells an atom has, the greater the shielding effect, and 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 ionisation energy is always the lowest)

Ionisation Energy & the Periodic Table

• So, the first ionisation energy increases across a period and decreases down a group

Graph of first ionisation energies from H to Na

There are ionisation energy trends within periods and groups

Ionisation energy across a period

• The ionisation energy over a period increases due to the following factors:

  • Across a period the nuclear charge increases

  • This causes the atomic radius of the atoms to decrease, as the outer shell is pulled closer to the nucleus, so the distance between the nucleus and the outer electrons decreases

  • The shielding by inner shell electrons remain reasonably constant as electrons are being added to the same shell

  • It becomes harder to remove an electron as you move across a period; more energy is needed

  • So, the ionisation energy increases

• There is a rapid decrease in ionisation energy between the last element in one period, and the first element in the next period because:

  • There is increased distance between the nucleus and the outer electrons as you have added a new shell

  • There is increased shielding by inner electrons because of the added shell

  • 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

• 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_z¹There is a slight decrease in IE₁ between nitrogen and oxygen due to spin-pair repulsion in the 2p_x orbital of oxygen

Ionisation energy down a group

• The ionisation energy down a group decreases due to the following factors:

  • The number of protons in the atom is increased, so the nuclear charge increases

  • But, the atomic radius of the atoms increases as you add more shells of electrons, making the atoms bigger

  • So, the distance between the nucleus and outer electron increases as you descend the group

  • The shielding by inner shell electrons increases as there are more shells of electrons

  • These factors outweigh the increased nuclear charge, meaning it becomes easier to remove the outer electron as you descend a group

  • So, the ionisation energy decreases

Table summarising ionisation energy trends across a period & down a group 

Successive Ionisation Energies of an Element

• The successive ionisation energies of an element increase

• This is because once you have removed the outer electron from an atom, you have formed a positive ion

• 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:

Table Showing the Successive Ionisation Energies of Calcium Table

Successive ionisation energies of an element

The ionisation energy increases as you remove more electrons from an element

• 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 more difficult as the orbital is no longer full, and there is less spin-pair repulsion