7.2 Ionisation Energies

7.2 Ionisation Energies

1. What is Ionisation Energy?

   • Ionisation energy measures how easily an atom loses electrons to form positive ions
   • The 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.

2. Factors Effecting Ionisation Energy

   • Atomic Radius
     • The greater the distance between the nucleus and the outer electrons, the less the nuclear attraction. The force of attraction falls off sharply with increasing distance, so atomic radius has a large effect
   • Electron Shielding
     • Electrons are negatively charged and so inner-shell electrons repel outer-shell electrons. This repulson, called the shielding effect, reduced the attraction between the nucleus and outer electrons
   • Nuclear Charge
     • The more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and the outer shell electrons

3. Successive Ionisation Energies

   • Successive Ionisation Energies and Shells
     • In fluorine for example, there is a large increase between the 7th and 8th ionisation energies. This suggests that the 8th electron must be removed from a different shell, closer to the nucleus with less shielding
       • The first shell (n=1, closer to the nucleus) contains 2 electrons
         • The second shell (n=2, the outer shell) Contains 7 electrons
   • Making predictions from successive ionisation energies
     • successive ionisation energies allow predictions to be made about the number of electrons in the outer shell, the groups of the element in the periodic table and the identity of the element
       • An element in period 3 has successive ionisation energies that steadily increase, but then there is a large increase between the third and fourth ionisation energies. This shows that the fourth electron is being removes from an inner shell. Therefore there must be three electrons in its outer shell and the element must be in group 13 (3). Since it is period 3, the element must be Aluminium.

4. Trends in first ionisation energies

   • For the first 20 elements in the periodic table, there are two key patterns
     • A general INCREASE in first ionisation energy across each period (H --> He, Li --> Ne, Na --> Ar)
     • A sharp DECREASE in first ionisation energy between the end of one period and the start of the next period (He --> Li, Ne --> Na, Ar --> K)
   • Trend in first ionisation energy down a group
     • First ionisation energies DECREASE down a group.
       • This is because although the nuclear charge increases, its effect is outweighed by the increased atomic radius and, to a lesser extent, increased shielding. Therefore first ionisation energies decrease.
   • Trend in first ionisation energies across a period
     • Although first ionisation energy shows a general increase across both periods, it does fall in two places in each period. The drops also happen at the same positions in each period. The reason is linked to the existance of sub-shells, their energies and how orbitals fill with electrons
       • Across Period 2 there are three rises and two falls.
         • A rise from Li to Be
           • A fall from Be to B
             • Rise from B to C to N
               • Fall from N to O
                 • Rise from O to F to Ne
                 • The fall marks the start of electron pairing in the p-orbitals of the 2p sub-shell
                   • In N and O the highest energy electrons are in a 2p sub-shell
                     • In O, the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electrons from an oxygen atom than a nitrogen atom.
                       • Therefore the first ionisation energy of O is less than the first ionisation energy of nitrogen
             • This fall marks the start of filling the 2p sub-shell. The 2p sub-shell 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. The first ionisation energy of boron is less than the first ionisation energy of beryllium