Successive Ionisation Energies (HL) (DP IB Chemistry): Revision Note

Successive ionisation energies

• Successive ionisation energies of an element increase

• This is because once the outer electron is removed, the atom becomes a positive ion, making further electron removal more difficult

• As more electrons are removed:

  • Shielding decreases

  • The proton-to-electron ratio increases

  • Attraction between nucleus and remaining electrons increases

• The increase is not constant, it depends on the electronic configuration

• Taking calcium as an example:

Ionisation energies of calcium table

Successive ionisation energies of calcium

• The first electron is easy to remove due to spin-pair repulsion in the 4s orbital, resulting in a low first ionisation energy (IE₁)

• The second electron is harder to remove as there’s no spin-pair repulsion

• The third electron is much harder to remove as it comes from the 3p subshell, which is closer to the nucleus

• The fourth electron is also difficult to remove as there is less spin-pair repulsion in the 3p orbital

• Successive ionisation energies increase as electrons are removed from an increasingly positive ion

• Large jumps in IE show a change in shell, while smaller jumps show changes within a subshell

• Uses of successive ionisation energy data

  • Predict or confirm the electronic configuration

  • Identify the number of outer electrons

  • Determine the group number in the Periodic Table

  • Large jumps help identify the shell where electrons are being removed

  • Used for deducing configurations of Na, Mg, Al and placing them in the Periodic Table

Successive ionisation energies table

Sodium

• A large increase between the first and second ionisation energies shows it is much easier to remove the first electron

• This means the first electron removed is from the valence shell, so Na is in Group 1

• The jump corresponds to removing an electron from the 3s to the full 2p subshell

  • Na: 1s² 2s² 2p⁶ 3s¹

Magnesium

• A large increase between the second and third ionisation energies suggests the first two electrons are easier to remove

• Therefore, Mg has two valence electrons, placing it in Group 2

• The jump corresponds to removing an electron from the 3s to the full 2p subshell:

  • Mg: 1s² 2s² 2p⁶ 3s²

Aluminium

• A large increase between the third and fourth ionisation energies shows the first three electrons are easier to remove

• These are the 3p and 3s electrons, which are farther from the nucleus and experience less nuclear charge than 2p electrons

• This suggests Al has three valence electrons, so it belongs to Group 13 (Group III)

• The jump corresponds to removing an electron from the third shell to the second shell:

  • Al: 1s² 2s² 2p⁶ 3s² 3p¹

• Ionisation energy shows periodicity; a repeating trend across a period in the Periodic Table

• Group 1 metals have low first ionisation energies; noble gases have very high values due to stable configurations

• Four key factors affect first ionisation energy:

  • Nuclear charge (more protons = stronger attraction)

  • Distance of outer electrons from the nucleus

  • Shielding by inner electrons

  • Spin–pair repulsion between paired electrons

• Across a period:

  • first ionisation energy increases due to higher nuclear charge and same shielding

• Down a group:

  • first ionisation energy decreases due to greater atomic radius and more shielding

Ionisation energies of hydrogen to sodium 

Ionisation energy across a period

• Across a period, ionisation energy increases because:

  • Nuclear charge increases, pulling electrons closer

  • Atomic radius decreases, reducing the distance to outer electrons

  • Shielding stays roughly constant, since electrons are added to the same shell

  • Outer electrons are held more tightly, so more energy is needed to remove them

Dips in the trend

• 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 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

  • In oxygen, there are 2 electrons in the 2p_x orbital, so the repulsion between those electrons makes it slightly easier for one of those electrons to be removed

From one period to the next

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

• This is 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