Periodicity - Ionisation Energy (Edexcel International A Level Chemistry)

• Elements in the periodic table are arranged in order of increasing atomic number and placed in vertical columns (groups) and horizontal rows (periods)
• The elements across the periods show repeating patterns in chemical and physical properties
• This is called periodicity

All elements are arranged in the order of increasing atomic number from left to right

• Ionisation energies show a trend across a period of the Periodic Table
• As could be expected from their electron configuration, the group 1 metals have a relatively low ionisation energy, whereas the noble gases have very high ionisation energies

Graph showing first ionisation energies of elements 1 to 11 

Ionisation energy across period 2 and 3

• The ionisation energy across a period generally 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

Dips in the trend for period 2

• 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 3p_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

Dips in the trend for period 3 

• There is again a slight decrease between magnesium and aluminium as the thirteenth electron in aluminium is in the 3p subshell, which is further away from the nucleus than the 3s subshell of magnesium
  • Magnesium has a first ionisation energy of 738 kJ mol^-1 as its electron configuration is 1s² 2s² 2p⁶ 3s² 
  • Aluminium has a first ionisation energy of 578 kJ mol^-1 as its electron configuration is 1s² 2s² 2p⁶ 3s² 3p_x¹
• There is a slight decrease in IE₁ between phosphorus and sulfur due to spin-pair repulsionin the 3p_x orbital of oxygen
  • Phosphorus has a first ionisation energy of 1012 kJ mol^-1 as its electron configuration is 1s² 2s² 2p⁶ 3s² 3p_x¹ 3p_y¹ 3p_z¹
  • Sulfur has a first ionisation energy of 1000 kJ mol^-1 as its electron configuration is 1s²2s² 2p⁶ 3s² 3p_x² 3p_y¹ 3p_z¹
• In sulfur, there are 2 electrons in the 3p_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 are adding 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

Summary Explanation for Trend in Ionisation Energy Across a Period and Down a Group