Ionization energy:

The energy required to remove the outermost shell electron in isolated gaseous atom when present in ground state is called Ionization energy (I.E). It is an elemental and thermodynamic property also termed as ionization potential and ionization enthalpy. It is also an endothermic process.

The ionization energy units are ev, kjmol-1, and kcalmol-1 and they are used for qualitative measurement of stability of atom.

Ionization energy relation with stability:

Ionization energy is directly proportional to stability but inversely proportional to reactivity. So as the ionization energy increases stability of element/atom also increase but it becomes less reactive.

                                                        I.E = stability = 1/reactivity

For example: noble gases have high ionization energy and stability but are low in reactivity. On the other hand, alkali metals are low in ionization energy and stability, but they are very reactive metals.

Successive ionization energy:

The energy needed to eliminate electrons one by one in stages in called successive ionization energy. The energy required to remove the first electron to form a cation is called first ionization energy, and the energy required to eliminate the second electron is called second ionization energy, and so on.

Mg –> Mg+1 + 1e  =  738kjmol-1

Mg –> Mg+2 + 2e  =  1451kjmol-1

Mg –> Mg+3 + 3e  =  7730kjmol-1

The second and third ionization energies are always higher than the first and second ones, respectively.

                                                   I.E1 < I.E2 < I.E3

The reason is that by decreasing the number of electrons, the proton number increases automatically, and the nucleus attracts the remaining electrons more powerfully. That’s why a high amount of energy is needed to remove electrons.

Helium has the highest ionization energy since it has a small size and more nuclear attraction to its electrons. Cesium has the lowest ionization energy due to its large size, and electrons are away from the nucleus, so less energy is required to remove electrons.

Factors effecting on I.E:

Effective nuclear charge

Effective nuclear charge is the nucleus attraction towards electrons, and it enhances as the size gets smaller. In easy words, it’s the nucleus grip on electrons. The ionization energy is directly proportional to the effective nuclear charge. As the charge increases, more energy is required to remove electrons.

                                                 I.E = effective nuclear charge

Shielding/Screening effect

In an atom, the inner shell electrons form a shield between the nucleus and the outer shell electrons, and as a result, the outer shell electrons face a less effective nuclear charge. This is called the shielding effect. The ionization energy is inversely proportional to the shielding effect. As the shielding effect increases, it becomes easier to remove electrons due to the less effective nuclear charge.

                                                I.E = 1/shielding effect

Atomic radius

Atomic radium is inversely proportional to I.E. By increasing the atomic radius, the shielding effect also increases, making it easier to eliminate electrons from the shell.

                                                  I.E = 1/atomic radius

Nature of the orbital

Because of the small size s-orbital requires more energy to remove electrons and due to the larger f-orbital size, requires less energy to remove electrons.

                                                       s > p > d > f

Electronic configuration:

Stable orbitals like full-filled and half-filled require the highest and highest energy to remove electrons, respectively, as compared to partially filled. Stable orbitals do not easily lose electrons as compared to unstable orbitals.

                              Fully-filled > half-filled > partially-filled


  1. To find out the index of metallic character, which is inversely proportional to I.E. Metals easily lose electrons and require less I.E. to remove them. So less I.E. value means more metallic character.
  2. To find out the strength of the reducing agent. As the reducing agents can oxidize themselves and reduce others (give electrons), it becomes easier to remove electrons. For example: alkali metals are good reducing agents, so they require less I.E. to remove electrons.
  3. To find out the reactivity of metals. Most reactive metals easily release electrons and require less I.E. So, the reactivity is inversely proportional to I.E. Note: This rule is only applicable to electro-positive elements such as alkali and alkaline earth metals but not to electro-negative elements.
  4. To predict the valency of atoms through the gap of I.E.

Periodic trends:

Top-to-bottom decreaseReasons:
Increase in shells
Increase in shielding effect
Increase in atomic size
Decrease in effective nuclear charge
Left-to-right increaseReasons:
No. of shells is the same
Shielding effects are almost same
Decrease in atomic radius
Increase in effective nuclear charge


The first three elements of groups IIIA and VIA show abnormal behavior with group IIA and VA group elements because of unstable electronic configuration. In comparison to IIA and IIIA, the IIA elements have a filled orbital, and the IIIA group elements have a partially filled orbital. Same in the case of VA and VIA, in which VA is half-filled and VIA is partially filled.

So according to stability rule: Fully-filled > half-filled > partially-filled

And I.E is directly proportional to stability/ electronic configuration

To know more about the processes and chemicals visit

Difference between Electron affinity and Ionization energy.

In simple words electron affinity is the amount of energy releases when we add an electron in gaseous atom and ionization energy is energies amount that is needed to eliminate the electron from that gaseous atom.

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