Trends in Ionisation Energy

Electrons are logical particles. They’re governed by simple laws of attraction and energy states, and ionisation is no different - there are clear trends in ionisation energies. Let's study these trends step by step.

Ions are everywhere in nature. Many elements form ions instead of remaining as atoms. For example, copper wire consists of positive copper ions suspended in a sea of delocalised electrons. Calcium tends to form 2+ ions, whereas potassium generally forms 1+ ions - you’ll rarely find potassium with any other charge. But why do some atoms form ions more readily than others, and why do they only form ions with certain charges? To understand this, let’s first recap ionisation energy.

In general, ionisation energy shows periodicity. This means that it has a trend that repeats across each period in the periodic table. The general pattern shown in period 2, for example, repeats in period 3. There are two other main trends you should know:

• Ionisation energy increases across a period.
• Ionisation energy decreases down a group.

Ionisation energy increases across a period

Ionisation energy increases across a period due to increasing nuclear charge. This increases the attraction between the nucleus and the outermost electron.

Take carbon and boron, for example. Carbon has six protons whilst boron has five, and so carbon’s nuclear charge is larger. Therefore, carbon’s first ionisation energy is higher than boron’s.

Fig. 2 - Boron, left, and carbon, right. Notice that carbon has one more proton in its nucleus than boron

Ionisation energy decreases down a group

Ionisation energy decreases down a group because the distance between the nucleus and the outermost electron increases. In addition, the outer shell electron is shielded by more shells of inner electrons. This decreases the attraction between the electron and the nucleus. These factors negate the impact of increased nuclear charge.

Exceptions to the trends in ionisation energy

Look at the following graph. It shows the first ionisation energy of each of the elements in period 3.

Fig. 3 - A graph showing the first ionisation energies of the elements in period 3. Note the upward trend

Overall it follows the trends described above - an increase in ionisation energy across the period as nuclear charge increases - but there are some exceptions. To understand why, we need to look more closely at electron configuration.

Groups 5 and 6

Elements in Group 6 in the periodic table have lower first ionisation energies than those in group 5, despite their increased nuclear charge. This can be explained by looking at their electron configuration.

For example, nitrogen (with seven electrons) has the structure while oxygen (with eight electrons) has the structure . Their electron configurations are shown below.

Fig. 4 - The electron configurations of nitrogen and oxygen

You’ll notice that nitrogen only has three electrons in the 2p subshell. According to Hund’s rule, electrons within a subshell will prefer to fill empty orbitals, and so there is a single electron located in each of 2p’s three orbitals. However, oxygen has four electrons in the 2p subshell. This means that one orbital must contain two electrons, which repel each other quite strongly. The electron-electron repulsion means that the outermost electron is easier to remove - it is already partially repelled by the other electron in its orbital.

For a reminder on Hund’s rule, see Electron Configuration.

Groups 2 and 3

Group 3 elements also have lower first ionisation energies than group 2, despite their increased nuclear charges. Again, it is to do with their electron configuration.

Fig. 5 - The electron configurations of beryllium and boron

For example, beryllium has the structure and boron has the structure . Boron’s outermost electron is located in the 2p subshell, whereas beryllium’s is in the 2s subshell. 2p is of a slightly higher energy level than 2s and is also slightly further away from the nucleus. This reduces the attraction between the nucleus and the outermost electron and makes it easier to lose, reducing the ionisation energy.

Trends in successive ionisation energy

Take a look at the successive ionisation energies of sodium.

Fig. 6 - A table showing the successive ionization energies of sodium

The first ionisation energy is relatively small. The second, however, is much larger. We’d expect successive ionisation energies to increase slightly, as it requires more energy to remove a negative electron from an increasingly positive ion, but not to this magnitude. This means that sodium will tend to form ions with a charge of +1, unless supplied with a significant amount of energy. Why is this?

Sodium has the electron configuration . Its outermost electron comes from the 3s subshell and this is the electron it will lose first when ionised. Sodium’s outermost electron is now in the 2p subshell, part of the second electron shell. The outermost electron is much nearer the nucleus and so experiences a stronger attraction to the positive protons. It is also shielded by fewer electron shells - just the first shell - and so the relative charge felt by the electron is stronger. The second shell is now complete, and for all the above reasons, we say it is energetically more stable. Therefore, it will take a lot of energy to remove this second electron from a sodium ion, which is why the second ionisation energy of sodium is so much larger than the first.

Using our knowledge of electron configuration and shells, we can determine what group an element is in by examining its successive ionisation energies. Let’s look at aluminium.

There is a large jump between the third ionisation energy and the fourth. This is because once you have removed the third electron from aluminium, it has a full outer shell. From this we can infer that aluminium is in group 3.