Ionisation energy trends 

In a nutshell

The first ionisation energies across the periodic table generally increase across a period and decrease down a group. There are some exceptions to this rule, which occur due to the different energy levels occupied by the electron orbitals.

First ionisation energy

The first ionisation energy is the energy required to convert one mole of gaseous atoms into one mole of singly positively charged gaseous atoms (the energy required to remove one electron from each atom). The general formula for first ionisation energy, where X represents an element is:

​$$X (g) \longrightarrow X^{+} (g) + e^{-} (g)$$​​

Trends in first ionisation energy

There are general trends seen across the periodic table. Below is the periodic table, with arrows showing the general trends in first ionisation energy.

Trends across periods and groups

A horizontal row in the periodic table is called a period. Across a period, the ionisation energy generally increases. This is because, as you go across the periodic table (left to right) the increased charge on the nucleus makes it more difficult to remove an electron. The difficulty results in higher first ionisation energy. 

A vertical column in the periodic table is called a group. The first ionisation energy decreases as you go down a group in the periodic table. This is for two main reasons.

The first reason is that the outer electron experiences more shielding as you go down the group. The second reason is that as you go down the group, the outer electron is further away from the nucleus and therefore held less strongly. These two factors result in less energy being needed to remove the outer electron and hence lower first ionisation energy. 

Across the periodic table, there is a general drop in ionisation energy from one period to the next, for example, Period $$0$$​ has higher ionisation energy than Period $$1$$​. This is because there is an increase in atomic radius as you go from one period to the next. This, as previously mentioned, means that the outer electron is further away from the nucleus, held less strongly and hence requires lower ionisation energy to remove that electron.

Exceptions to the rule

One exception to the trend is the drop in first ionisation energy seen from Group $$2$$ to Group $$3$$, and then another drop seen between Group $$5$$ and Group $$6$$. These exceptions are due to electron orbital arrangement. Two rules are key here:

1. Aufbau principle: Electrons will fill up lower energy orbitals before higher energy one
2. Hund's rule: When considering a particular energy level, electrons will occupy orbitals singly before doubling up.
   ​ ​

The drop in ionisation energy from Group $$2$$ to Group $$3$$​:

As shown by the graph, the first ionisation energy drops between Group $$2$$ and Group $$3$$​ and again between Group $$5$$​ and Group $$6$$.  This is due to the energy levels occupied by electron orbitals. The higher the energy level of the orbital, the less energy is required to remove the electron within it - hence lower ionisation energy.  

As shown by the diagram, $$2p$$​-orbitals have a higher energy level than electrons in the $$2s$$​-orbital. This means that atoms losing their outer electron from the $$2p$$​-orbital will have lower ionisation energy. 

Example

Compare Boron ($$B$$) and Beryllium ($$Be$$) in terms of their ionisation energy. ​

You might expect Boron to have the higher ionisation energy as it has an extra electron. However, that extra electron occupies a $$2p$$​-orbital, whereas Beryllium's outer electron occupies a $$2s$$​-orbital. The higher energy of the $$2p$$-orbital means that Boron will have lower first ionisation energy. 

The drop in first ionisation energy from Group $$5$$​ to Group $$6$$​:

When electrons occupy the same orbital they repel each other, so it is easier to remove one electron from that orbital. Elements in Group $$5$$ such as Silicone ($$Ne$$​) have their outer electrons in different orbitals, without repulsion ($$1s^{2}2s^{2}2p^{6}3s^{2}3p^{2}$$). 

In Group $$6$$​, elements such as Sulfur have their outer electrons paired in an orbital ($$1s^{2}2s^{2}2p^{6}3s^{2}3p^{4}$$​).

This repulsion between the electrons means that less energy is required to remove the outer electron and therefore, lower the first ionisation energy.​