Periodicity of ionization energy

The amount of energy required to remove the most loosely bound electron from an isolated gaseous atom is called ionization energy (IE).

ionization potential

Ionization energy is also called as ionization potential because it is measured as the minimum potential required to remove the most loosely held electron from the rest of the atom. It is measured in the units of electron volts (eV) per atom or kilo joules per mole of atoms (kJ mol-1).

1 eV per atom = 96.64 kJ mol-1 = 23.05 k cal mol-1

Thus, the ionization energy gives the ease with which the electron can be removed from an atom. The smaller the value of the ionization energy, the easier it is to remove the electron from the atom.

Factors Governing Ionization Energy

Size of the atom

As the size of the atom increases the outermost electrons are held less tightly by the nucleus (attractive force between the electron and the nucleus is inversely proportional to the distance). As a result it becomes easier to remove the electron and therefore the ionization energy decreases with the increase in atomic size.

Charge on the nucleus

The attractive force between the nucleus and the electron increases with the increase in nuclear charge making it more difficult to remove an electron. The ionization energy thus increases with the increase in the nuclear charge.

Screening effect

In multielectron atoms, the outermost electrons are shielded or screened from the nucleus by the inner electrons. This is known as shielding or screening effect. The outer most electrons do not feel the complete charge of the nucleus and the actual charge felt is called the effective nuclear charge. When the inner electrons are more, the screening effect will be large, the nuclear attraction will be less. Thus when the inner electrons increase the ionization energy will decrease.

Penetration effect

The 's' electrons are more penetrating (maximum probability of finding near the nucleus) towards the nucleus than the p electrons. The order of penetration power in a given shell is s > p > d > f

If the penetration power of the electron is more, it will be closer to the nucleus and will be held more firmly. Thus ionization energy will increase with the increase in the penetration power of the electrons. For the same sub shell the ionization energy would be more to remove an 's' electron than to remove a 'p' electron which in turn will be more than that for removing a 'd' electron.

Electronic arrangement

Certain electronic configuration like half-filled and completely-filled shells have extra stability. It is more difficult to remove electron from these stable configuration and the ionization energy is very high. For example, the noble gases have the most stable configuration and so have high ionization energy; elements like Be and Mg have completely filled orbitals while N and P have exactly half-filled sub shells. Thus, their ionization energies are high. The more stable the electronic configuration, the higher is the ionization energy.

Variation along a period

The ionization energy increases with increasing atomic number in a period. This is because

  • The nuclear charge increases on moving across a period from left to right.
  • The atomic size decreases along a period though the main energy level remains the same.
Due to the increased nuclear charge and simultaneous decrease in atomic size, the valence electrons are more tightly held by the nucleus. Therefore more energy is needed to remove the electron and hence ionization energy keeps increasing. However some irregularities have been noticed due to the extra stability of the half filled and completely filled configurations.

For example, the nuclear charge on Boron is more than Beryllium, yet there is slight decrease in ionization energy from Be to B. This is because, in boron the last electron goes to '2p' orbital which is at a slightly higher energy than '2s' orbital. Also, the electronic configuration of B is less stable than that of Be (has completely filled orbitals). Hence the ionization energy is less than that of Be. Similarly, nitrogen, which has half filled '2p' orbitals, is more stable than oxygen. Therefore the ionization energy of nitrogen is more than that of oxygen.

variation of ionization energy with atomic number along a period

Fig: 4.7 - Variation of ionization energy with atomic number

Variation down a group

The ionization energy gradually decreases in moving from top to bottom in a group. This is due to the fact that:

  • The nuclear charge increases in going from top to bottom in a group.
  • An increase in the atomic size due to an additional energy shell (level) 'n'.
  • Due to the increase in the number of inner electrons there is an increase in the shielding effect on the outer most electron. The effect of increase in atomic size and the shielding effect is much more than the effect of increased nuclear charge.

As a result , the electron becomes less firmly held to the nucleus and so the ionization energy decreases as we move down the group.

variation of ionization energy with atomic number along a group

Fig: 4.8 - The variation of ionization energy with atomic number

Successive ionization energies

The energies required to remove subsequent electrons from a gaseous atom is called as successive ionization energies. They are termed as first ,second, third …… ionization energy depending on the removal of the first, second, third electron respectively.

Successive ionization energies

Successive ionization energies

Successive ionization energies

The second ionization energies are higher than the first due to the fact that after the removal of the first electron the atom changes into a monovalent positive ion. In this ion, the number of electrons decreases but the nuclear charge remains same and so the remaining electrons are held more tightly by the nucleus and it becomes difficult to remove the second electron. Hence the value of the second ionization energy (IE2) is higher than the first (IE1). In the same way the removal of the second electron will result in the formation of di-positive ion making the attraction between the nucleus and the remaining electrons stronger. This results in higher value of third ionization energy (IE3).


6. Calcium (Z = 20) loses electrons successively to form Ca+, Ca2+ and Ca3+ ions. Which step will have highest ionization energy?


The step, which involves the formation of Ca3+ from Ca2+� would have highest ionization energy.

formation of ca ion

7. Consider the ground state electronic configurations given below:

(A) 1s2 2s2 2p6 (B) 1s2 2s2 2p4 (C) 1s2 2s2 2p6 3s2 (D) 1s2 2s2 2p6 3s1 (E) 1s2 2s2 2p5.

Which of the above configuration is associated with the lowest and which is associated with highest ionization energy?


Lowest ionization energy = D

Highest ionization energy = A.

8. Why is the ionization energy of nitrogen unexpectedly high?


The electronic configuration of nitrogen (1s22s22p3) shows that it possesses exactly half-filled 'p' orbitals as its outershell configurations. Since, half-filled orbitals are extraordinarily stable, it is more difficult to remove an electron from this atom. Thus nitrogen has unexpectedly high ionization energy.


Deirdre Hanvey said...

Thank you for this clearly written explanation.!
New chem student

Deirdre Hanvey said...

Could you explain the descrepancy in ionization trend from Aluminum to Gallium?