Ionisation Energies

• Ionisation energy measures how easily an atom loses electrons to form positive ions
• First ionisation energy is the energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.
• Factors affecting ionisation energy

  Electrons are held in their shells by attraction from the nucleus

  First electron lost will be in the highest energy level and will experience the least attraction from the nucleus.

  Three factors affect the attraction between the nucleus and the outer electrons of an atom and therefore the ionisation energy
• Atomic radius

  The greater the distance between the nucleus and the outer shell electrons, the less the nuclear attraction. The force of attraction falls off sharply with increasing distance, so atomic radius has a large effect
• Nuclear charge

  The more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and outer electrons
• Electron shielding

  Electrons are negatively charged and so inner-shell electrons repel outer-shell electrons. This repulsion is called the shielding effect, and it reduces the attraction between the nucleus and outer electrons
• Successive Ionisation energies

  An element will have as many ionisation energies as there are electrons e.g. helium has 2 electrons and 2 ionisation energies.

  The second ionisation energy of helium is greater than the first.
• Why is the second ionisation energy of helium greater than the first?

  In a helium atom, there are two protons attracting two electrons in the 1s sub-shell - once the first electron has been lost, the single electron is pulled closer to the helium nucleus. The nuclear attraction on the remaining electron increases and so more ionisation energy will be needed to remove this second electron.
• Second ionisation energy is the energy required to remove one electron from each atom in one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions
• Successive ionisation energies can provide important evidence for different electron energy levels in an atom.

  For example, on a graph, a large increase between the 7th and 8th ionisation energies suggests that the 8th electron must be removed from a different shell, closer to the nucleus and with less shielding.

  n=1 contains 2 electrons

  n=2 contains 7 electrons (outer shell)
• Making predictions from successive ionisation energies

  Number of electrons in the outer shell

  The group of the element in the periodic table

  The identity of an element

  e.g. If ionisation energies steadily increase but then there is a large increase between the 3rd and 4th ionisation energies - the 4th electron must be being removed from an inner shell - therefore there are three electrons in the outer shell and the elemt must be in group 3
• Trends in ionisaton energies

  Periodic trends in ionisation energies provide important evidence for the existence of shells and sub-shells.

  e.g. general increase in 1st ionisation energy across each period and a sharp decrease in 1st ionisation energy between the end of one period and the start of the next period.
• Trend in first ionisation energy down a group

  Atomic radius increases

  More inner shells so shielding increases

  Nuclear attraction on outer electron decereases

  First ionisation energy decreases

  (Although nuclear charge increases, its effect is outweighed by the increased radius, and to a lesser extent, the increased shielding)
• Trend in first ionisation energy across a period

  Nuclear charge increases

  Same shell = similar shielding

  Nuclear attraction increases

  Atomic radius decreases

  First ionisation energy increases
• Sub-shell trends in first ionisation energy

  Although 1st ionisation shows a general increase across periods 2 and 3, it falls in two places in each period. These drops occur in the same positions in each period suggesting there may be a periodic cause - the reason is linked to the existence of sub-shells, their energies and how orbitals fill with electrons
• Across period 2, 1st ionisation energy shows 3 rises and 2 falls

  A rise from lithium to beryllium

  A fall to boron followed by a rise to carbon and nitrogen

  A fall to oxygen followed by a rise to fluorine and neon
• Comparing beryllium and boron

  The first fall in 1st ionisation energy from beryllium to boron marks the filling of the 2p sub-shell

  The 2p sub-shell in boron has a higher energy than the 2s sub-shell in beryllium.

  Therefore, in boron the 2p electron is easier to remove than one of the 2s electrons in beryllium.

  So the 1st ionisation energy of boron is less than 1st ionisation energy of beryllium
• Comparing nitrogen and oxygen

  The fall in 1st ionisation energy from nitrogen to oxygen marks the start of electron pairing in p-orbitals of the 2p sub-shell

  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

  Therefore the first ionisation energy of oxygen is less than the first ionisation energy of nitrogen
• Nitrogen?

  Three 2p electrons: 2px1, 2py1, 2pz1

  one electron in each 2p orbital

  spins are at right angles - equal repulsion as far apart as possible
• Oxygen?

  Four 2p electrons: 2px2, 2py1, 2pz1

  2 electrons in 1 2p orbital

  2p electrons start to pair

  The paired electrons repel