What is Ionisation?
When electrons have been removed from an atom or molecule, it's been ionised. The energy you need to remove the first electron is called the first ionisation energy (or often just ionisation energy).
The first ionisation energy is the energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions.
You can write equations for this process - here's the equation for the first ionisation of oxygen:
O (g) --> O+ (g) + e-
Here are a few rather important points about ionisation energies:
1) You must use the gas state symbol (g), because the ionisation energies are measured for gaseous atoms.
2) Always refer to 1 mole of atoms, as stated in the definition, rather than to a single atom.
3) The lower the ionisation energy, the easier it is to form an ion.
Factors affecting ionisation energy
Nuclear Charge: The more protons there are in the nucleus, the more positively charged the nucleus is and the stronger the attraction for the electrons.
Distance from Nucleus: Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away.
Shielding: As the number of electrons between the outer electrons and the nucleus increases, the outer electrons feel less attraction towards the nuclear charge.This lessening of the pull of the nucleus by inner shells of electrons is called shielding (or screening).
A high ionisation energy means there's a high attraction between the electron and the nucleus.
Ionisation Energy down a group
1. This provides evidence that electron shells really do exist.
2. If each element down a group has an extra electron shell compared to the one above, the extra inner shells witl shield the outer electrons from the attraction of the nucleus.
3. Also, the extra shell means that the outer electrons further away from the nucleus, so the nucelus' attraction will be greatly reduced.
It makes sense that both of these factor will make it easier to remove outer electrons, resulting in a lower ionsation energy.
Ionisation energy across a period
1. As you move across a period, the general trend is for the ionisation energies to increase- i.e. it gets harder to remove the outer electron.
2. This can be explained because the number of protons is increasing, which means a stonger nuclear attraction.
3. All the extra electrons are at roughly the same energy level, even if the outer electrons are in different orbital types.
4. This means there's generally little extra shielding effect or extra distance to lessen the attraction from the nucleus.
5. But, there are small drops between Groups 2 and 3, and 5 and 6. Tell me more, I hear you cry. Well, alright then...
Groups 2 and 3
1. Aluminium's outer electron is in a 3p sub-shell rather than a 3s. The 3p orbital has a slightly higher energy than the 3s orbital, so the electron is, on average, to be found further from the nucleus.
2. The 3p orbital has additional shielding provided by the 3s2 electrons.
3. Both these factors together are strong enough to override the effect of the increased nuclear charge, resulting in the ionisation energy dropping slightly.
4. This pattern in ionisation energies provides evidence for the theory of electron sub-shells.
Groups 5 and 6
1. The shielding is identical in the phosphorus and sulphur atoms, ant the electrons is beings removed from an identical orbital.
2. In phosphorus' case, the electron is being removed from a singly-occupied orbital. But in sulphur, the electron is being removed from an orbital containing two electrons. The repulsion between the two electrons means that electrons are easier to remove from shared electrons.
3. Yup, yet more evidence for the electronic structure model.