Ionisation energies

• Define first ionisation energy


  The energy required to remove 1 electron from each atom in 1 mole of gaseous atoms of an element to form 1 mole of gaseous 1+ ions
• What are the 3 factors affecting ionisation energy?

  • Atomic radius
  • Nuclear charge
  • Electron shielding
• How does atomic radius affect ionisation energy?

  • The greater the distance between the nucleus and outer electrons, the less nuclear attraction
  • So is easier to lose an electron
• How does nuclear charge affect ionisation energy?

  • The more protons there are in the nucleus, the higher the nuclear charge
  • Therefore the greater the attraction between the nucleus and outer electrons
  • So is harder to lose an electron
• How does electron shielding affect ionisation energy?

  • Inner shell electrons repel outer shell electrons -> Shielding effect
  • The more shells, the more shielding between the nucleus and outer electrons, so nuclear attraction weakens
  • So is easier to remove an electron
• What is the equation for the first ionisation energy?

  Mg (g) -> Mg+ (g) + e-
• What is the equation for the second ionisation energy?

  Mg+ (g) -> Mg 2+ (g) + e-
• What is the trend in the first ionisation energy down a group?

  • Atomic radius increases, as more shells are added
  • Nuclear charge increases, as there are more protons in the nucleus
  • Electron shielding increase, as more inner shells are added
  • Nuclear attraction between nucleus and outer electrons decrease
  • Therefore, first ionisation energy decreases down a group
• What is the trend in first ionisation energy across a period?

  • Atomic radius decreases, as the nuclear attraction increases
  • Nuclear charge increases, as there are more protons in the nucleus
  • Electron shielding is similar, as they all have the same number of shells
  • Therefore, first ionisation energy increases across a period
• What are some exceptions to first ionisation energy trend in period 2?

  • A fall between beryllium to boron
  • A fall between nitrogen to oxygen
• Explain why there is a slight fall in first ionisation energy from Be to B

  • Because in boron we see the first 2p orbital being filled
  • The 2p sub-shell has higher energy than the 2s sub-shell so the 2p electron is easier to remove as nuclear attraction is lower
  • This marks the start of the 2p sub-shell filling
• Explain why there is a slight fall in first ionisation energy from N to O


  • In nitrogen and oxygen the highest energy electrons are in a 2p sub-shell
  • In oxygen, the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from an oxygen atom than a nitrogen atom
  • This marks the start of electron pairing in the p-orbitals of the 2p sub-shell
  • Therefore the first ionisation energy of oxygen is less than the first ionisation energy of nitrogen
• Explain why there's a drop in first ionisation energy between groups 2 and 3

  • In group 3, the electron is removed from a p-orbital rather than an s-orbital in group 2
  • P-orbitals have slightly higher energy than s-orbitals, so the outer electron is further away from nucleus
  • The p-orbital also experiences shielding from the nucleus rpovided by s electrons
  • Less energy is required to remove the outermost p electron from group 3 element compared to s electron
• Why are there large jumps in success ionisation energy between shells


  • When reaching a new inner electron shell, there is a big increase in the ionisation energy needed to remove the first electron in the new shell
  • This happens because the attraction to the nucleus is much greater for inner shell electrons closer to nucleus