Ionisation Energy Trends (Cambridge (CIE) AS Chemistry): Revision Note

Trends in Ionisation Energy

• Ionisation energies show periodicity, a recurring trend across each period of the Periodic Table

  • As you move left to right across a period, first ionisation energy generally increases

  • This trend is linked to the electronic configuration of elements:

    • Group 1 metals (e.g. Na) have low first ionisation energies because they have a single outer electron that is easy to remove

    • Noble gases (e.g. Ne, Ar) have very high ionisation energies due to their full outer shells and strong nuclear attraction

• There are four factors that affect ionisation energy

  • Nuclear charge

  • Distance from the nucleus

  • Electron shielding

  • Spin-pair repulsion

Nuclear charge (number of protons)

• As atomic number increases, so does the positive charge of the nucleus

• A higher nuclear charge creates a stronger electrostatic attraction between the nucleus and outer electrons

• Therefore more energy is required to remove the electron means a higher ionisation energy

Distance from the nucleus (atomic radius)

• Outer electrons that are further from the nucleus experience weaker attraction

• The greater the distance, the easier it is to remove the electron

• Ionisation energy decreases with increased distance from the nucleus

Electron shielding (inner shell repulsion)

• Electrons in inner shells repel electrons in outer shells and partially block the attraction of the nucleus

• This is known as the shielding effect

• More inner electron shells means greater shielding and therefore a lower ionisation energy

Spin-pair repulsion (electron–electron repulsion)

• In orbitals that contain two electrons, the electrons experience repulsion due to their like charges

• This repulsion slightly reduces the energy needed to remove one of the electrons

• This explains small drops in ionisation energy across a period, such as:

  • Oxygen vs nitrogen (O’s paired 2p electrons are easier to remove)

  • Aluminium vs magnesium (Al’s 3p¹ electron is higher in energy than Mg’s 3s²)

Ionisation Energy & the Periodic Table

• 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

Graph of first ionisation energies from H to Na

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

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

• 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:

Table showing the successive ionisation energies of calcium table

Successive ionisation energies of calcium

• The first ionisation energy is relatively low due to spin-pair repulsion in the 4s orbital

• The second electron is harder to remove because this repulsion is no longer present

• The third ionisation energy increases sharply as the electron is removed from the 3p subshell, which is closer to the nucleus

• The fourth electron is also more difficult to remove due to reduced spin-pair repulsion within the 3p orbital

• Successive ionisation energies always increase because electrons are being removed from an increasingly positive ion

• A large jump in ionisation energy indicates a change in shell, while smaller jumps reflect changes within the same shell or subshell

Using successive ionisation data

• Helps predict or confirm electronic configuration.

• Identifies the number of outer-shell electrons.

• Indicates the group number by locating the position of a significant jump in ionisation energy.

• Commonly applied to elements like sodium (Na), magnesium (Mg), and aluminium (Al) to deduce their place 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