Ionisation Energy Trends Across a Period (HL) (DP IB Chemistry): Revision Note

Ionisation energy trends across a period

• The trends in ionisation energy across a period and down a group have been discussed in our revision note on Periodicity 

  • Trends in ionisation energy across a period provide evidence for the existence of energy sublevels

Graph showing ionisation energies From H to Ne

• 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

  • 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

Exceptions to the general trend in ionisation energy

• There are two key exceptions to the general trend across Period 2:

  • Beryllium and boron

  • Nitrogen and oxygen

• Both exceptions provide evidence for the existence of energy sublevels

Beryllium and boron

• Beryllium has:

  • A first ionisation energy value of 900 kJ mol^-1

  • An electron configuration of 1s² 2s²

• Boron has:

  • A first ionisation energy value of 801 kJ mol^-1

  • An electron configuration of 1s² 2s² 2p¹

• There is a slight decrease in 1st I.E. between beryllium and boron, because:

  • The electron removed from boron is in a 2p orbital

    • 2p orbitals are higher in energy than 2s orbitals

  • The 2p electron experiences slightly more shielding than the 2s electron in beryllium

  • This means that the outer electron of boron:

    • Is further from the nucleus

    • Experiences a weaker attraction to the nucleus

    • Requires less energy to remove

  • The increased nuclear charge is not enough to outweigh these factors

Nitrogen and oxygen

• Nitrogen has:

  • A first ionisation energy of 1402 kJ mol^-1

  • An electron configuration of 1s² 2s² 2p³

• Oxygen has:

  • A first ionisation energy of 1314 kJ mol^-1

  • An electron configuration of 1s² 2s² 2p⁴

• There is a slight decrease in 1st I.E. between nitrogen and oxygen, because:

  • In nitrogen, all three 2p electrons occupy separate orbitals

    • There is no electron pairing

  • In oxygen, the fourth electron is in a 2p orbital that contains a pair of electrons

  • The electron pairing in oxygen increases electron–electron repulsion

  • This means that the outer 2p electron of oxygen:

    • Is slightly higher in energy

    • Requires less energy to remove

  • The increased nuclear charge is not enough to outweigh the electron–electron repulsion

Summary of ionisation energy trends

Across a period

• Nuclear charge increases

  • So, there is a stronger attraction between the nucleus and the outer electrons

• Number of electron shells is constant

  • So, the outer electrons are approximately the same distance from the nucleus

• Shielding remains reasonably constant

• Atomic radius decreases

  • This is due to the increased nuclear charge pulling electrons closer

• The outer electron is held more tightly and requires more energy to remove

Down a group

• Nuclear charge increases

  • But, this effect is outweighed by other factors

• Number of electron shells increases

  • So, the outer electrons are further from the nucleus

• Shielding increases

  • There are more inner shells repelling the outermost electron

• Atomic radius increases

• The outer electron is held more loosely and requires less energy to remove