Ionization potential, factors affecting ionization potential and its periodic variation

When heat is applied to an atom, the electrons absorb energy, and jump to the higher energy levels.
If this process is continued, a stage will come when the electron goes completely out of the influence of the nucleus resulting in the formation of the positive ion.
The amount of energy required to remove the electron from an isolated gaseous neutral atom to give a positive ion is called ionization energy or ionization potential.
Atoms having less ionization energy ionize very easily.
This is an endothermic process and hence ionization energy is represented by a positive sign.
Ionization potential is measured in electron volts (ev) or kilojoules.

1 ev= 1.6022 x 10^-19 J

The electrons are removed in stages one by one. The first ionization potential or first ionization energy (I₁), removes only one electron from its valence shell. The removal of the second electron from the +1 ion is called second ionization potential (I₂) and so on. Similarly, the removal of the third electron from the +2 ion is called third ionization energy (I₃).
The more the removal of electrons from an atom, the greater is the nuclear charge. As the nuclear charge increases, the attractive force of the nucleus for the valence electrons also increases. Hence, the ionization energy increases from first (I₁) to third (I₃).
I₁<I₂<I₃

The ionization energy above the third or ions with charges higher than +3 is rarely known under ordinary condition. The ionization potential of metals is lower than those of non-metals.
If a metal has higher ionization potential, it is difficult for it to form an ionic compound with a non-metal.

Factors on which ionization potential depends and its periodic variation:

The magnitude of ionization potential depends upon the following factors:

completely filled and half-filled orbitals

1. Charge on the nucleus

Let’s take the example of Li to Ne in the second period of the periodic table.
Li has one electron in its outermost shell and the number of valence electrons goes on increasing in the same shell until it possesses eight electrons in Ne.
As the number of protons (charge on the nucleus) increases, the attraction between the nucleus and the valence electrons increases thereby drawing the valence shell inward towards the nucleus and consequently the atomic size decreases.
Hence, the removal of valence electrons becomes more and more difficult as the atomic size decreases.
In other words, when we move from left to right in the same period across the periodic table, then number of protons or nuclear charge increases whereby the electrostatic attraction between the nucleus and the valence electrons increases.
This leads to the decrease in atomic size and hence ionization potential increases.

2. Shielding effects

It is known that the electrons in the valence shell are attracted by the nucleus and repelled by the electrons of the inner shells.
The combined effect of these two effects in the valence electrons experience more repelling effect which is known as shielding effect or screening effect.
The larger the number of electrons in the inner shells, the more is the shielding effect.
When we move from top to bottom in any group in the periodic table, we see that the size of atom increases.
For example, Elements Li, Na, K, Rb and Cs fall in group IA of the periodic table. Here the number of shielding electrons increases from top to bottom (Li to Cs). As a result, the atomic size increases from Li to Cs and hence the attractive force between the nucleus and the valence electrons decreases.
That is to say, the ionization potential decreases from Li to Cs.

3. Atomic radius

As the atomic size increases, the atomic radius also increases and the ionization potential decreases.
That is why, the ionization potential decreases gradually from Lithium to Caesium.

4. Principal quantum number

Principal quantum number (n) is any positive whole number (except zero) that indicates energy levels as well as relative distance from the nucleus.
It indirectly describes the size of the electron orbital.
The greater the value of n, the larger is the size of an atom.
The valence electrons in the bigger atoms will be far from the nucleus and experience less attractive force.
Hence, a little energy will be sufficient to remove the valence electrons from these atoms.
Therefore, the ionization potential decreases with the increase in the value of principal quantum number (n).
That is why, when we go down from top to bottom in any group in the periodic table, the ionization potential decreases.

5. s, p, d and f electrons

The ‘d’ electrons approach more closely to the nucleus than the ‘f’ electrons, the ‘p’ electrons closer than the ‘d’ electrons, and the ‘s’ electrons still closer than the ‘p’ electrons.
Therefore, the ‘s’ electrons are more difficult to remove from an atom than the ‘p’ electrons, the ‘p’ electrons than the ‘d’ electrons and so on.
Thus, the ionization energy of ‘Al’ is lesser than that of ‘Mg’.

Mg: 1s² 2s²2p⁶ 3s²

Al: 1s² 2s²2p⁶ 3s²3p¹

6. Completely filled and half-filled orbitals

The completely filled and half-filled orbitals are comparatively more stable, according to Hund’s rule.
Therefore, more energy is required to remove the electrons from such atoms. This shows some irregularities in the expected values of the ionization potential.
For example, according to the nuclear charge concept, the ionization potential value of Aluminium should higher than that of Magnesium and of Phosphorus than that of Sulfur when we move from left to right in the third period across the periodic table.
However, Magnesium has a higher value of ionization potential than Aluminium due to the completely filled ‘s’ orbital 1s² 2s²2p⁶ 3s² for Magnesium and Phosphorus has a higher value of ionization potential than Sulfur owing to the half-filled ‘p’ orbital 1s² 2s²2p⁶ 3s²3p³ for Phosphorus.

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