Ionisation Energy |
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At the end of this topic you should be able to:
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Recall the definitions of first and successive ionisation energies.
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Understand that successive ionisation energies provide evidence of quantum shells.
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Understand that electronic structure determines chemical properties.
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Define first and second electron affinities, and explain why the second electron affinity is endothermic.
Definitions
Ionisation Energy
This is the energy required to remove one mole of electrons from the outermost shell of an atom to form a positively charged ion.
M(g) 
M+(g)
+ e-
This process can be repeated again to give the second ionisation energy. This is the removal of one mole of electrons from gaseous monopositive ions. This is more difficult than the first ionisation energy because we are removing a negative electron from a positive ion.
M+(g) M2+(g)
+ e-
And again......
M2+(g) M3+(g)
+ e-
It is possible to continue in this way until al of the electrons on an atom have been removed.
Electron Affinity
This is the enthalpy change when 1 mole gaseous atoms gains 1 mole of electrons under standard conditions.
Cl (g) + e- 
Cl-(g)
The elements in group 7 have the highest electron affinities, they form negative ions easily, as go down the group the electron affinity decreases so reactivity decreases.
The second electron affinity is the energy needed to to add an electron to 1 mole of gaseous 1- ions to form 1 mole of gaseous 2- ions under standard conditions (where standard conditions are 100kpa and 298K).
Cl- (g) + e-
Cl2-(g)
This process involves adding a negatively charged electron to a negative ion - naturally this process is endothermic since energy needs to be supplied to overcome the repulsive forces between the negative ion and the negative incoming electron.
Successive Ionisation Energies of an Atom
The successive ionisation energies of an element can tell us what group the element is in and also how many electron shells the atom has. A graph of the successive ionisation energies of sodium are shown below:
The electronic configuration of sodium show that it has electrons in three different energy levels. The electrons closest to the nucleus (first shell) are the 1s2 electrons. The next eight electrons are the 2s2 and 2p6 electrons and the last electron in the outer shell is the 3s2 electron.
There is a big jump between the first and second ionisation energies because the second electron is removes from a shell closer to the nucleus. There is a big jump between the 9th and 10th ionisation energies for the same reason. Successive ionisation energies generally tend to increase because the nuclear attraction of the protons is increasing as more and more electrons are removed.
Trends in Ionisation Energy
How Does the First Ionisation Energy Change Going Across a Period?
Before we explain this we first need to know the electronic configuration of the elements in the period. In this example we will use period three.
| Element | Electronic Configuration |
| Lithium | 1s2, 2s1 |
| Beryllium | 1s2, 2s2 |
| Boron | 1s2, 2s2, 2p1 |
| Carbon | 1s2, 2s2, 2p2 |
| Nitrogen | 1s2, 2s2, 2p3 |
| Oxygen | 1s2, 2s2, 2p4 |
| Fluorine | 1s2, 2s2, 2p5 |
| Neon | 1s2, 2s2, 2p6 |
There are two basic rules which need to be applied to explain the trend
1. It takes more energy to remove an electron from a full or half full sub-shell because it is more stable.
2. As the number of protons in the nucleus increase, the nuclear attraction increases, therefore making it more difficult to remove an electron.
The following graph shows how ionisation energy varies across period two.
Main Points
1. The lowest ionisation energy is for lithium because
a) it has only 3 protons in the nucleus
b) it can lose its outer 2s1 electron easily because it will then have a full 1s2 sub-shell which is very stable.
2. The highest ionisation energy is for Argon because
a) it has a full 2p6 sub-shell
b) it has 10 protons in its nucleus
3. Boron has a lower ionisation energy than Beryllium even though there is a greater nuclear attraction in Boron because
a) Boron loses it's 2p1 electron from the sub-shell easily to gain a full 2s2 sub-shell
b) Beryllium has a full 2s2 sub-shell giving it extra stability and therefore making it more difficult to remove an electron
4. Oxygen has a lower ionisation energy than Nitrogen because
a) Oxygen loses it's 2p4electron in its sub-shell easily to gain a half full 2p3 sub-shell which is more stable.
b) Nitrogen has it's outer electron in a half full 2p3 sub-shell giving it extra stability and therefore making it more difficult to remove an electron.
How does the first ionisation energy change going down a group?
The outer electrons are held in their shells by the attractive force of the positive protons in the nucleus, the nuclear attraction. As more and more electron shells are added this force gets weaker because
- the distance between the outer electrons and the nucleus is increasing
- The inner electrons shield the nuclear electrons from the outer electrons , electronic shielding.
The lower the ionisation energy the easier it is to remove electrons from the outermost shell of the atom. As you go down a group the ionisation energy decreases. This also explains why metals get more reactive as you go down a group. It gets easier for them to give up electrons to form bonds.