[Explained] Why do Successive Ionization Energies Increase?

You might have been wondering why the successive ionization energies of elements increase with an increase in atomic number. In this post you will learn about the reason why it is so.


The atom gets smaller as it loses electrons, the effective nuclear charge on it increases, shielding effect of the electrons is reduced, penetration of the nucleus increases and the atoms which have attained the stable electron configuration due to successive removal of electrons have higher ionization energy values.



The most accurate answer as why successive ionization energies increase would be:

The atom gets smaller as it loses electrons, the effective nuclear charge on it increases, shielding effect of the electrons is reduced, penetration of the nucleus increases and the atoms which have attained the stable electron configuration due to successive removal of electrons have higher ionization energy values.


We shall see how all this happens when an atom undergoes successive ionization in detail further in this post.

We shall first learn about what ionization energy is.


As you might know, 

The amount of energy required to remove the most loosely bound electron from an isolated gaseous atom to form a gaseous ion is called ionization enthalpy or ionization energy of that atom.

Let us also brush our knowledge about the successive ionization energy and what it means.

Successive Ionization Energies

Second Ionization Energy


Due to smaller size of the cation, the remaining electrons experience a greater pull of the nucleus. Therefore, a higher energy will be required to remove an electron from the unipositive cation. This energy is called the second ionization energy and is represented as IE2.




The removal of the most loosely bound electron from a neutral atom forms a mono-positive cation. The ionic radius of this cation is much smaller than the atomic radius of the parent atom.

Due to smaller size of the cation, the remaining electrons experience a greater pull of the nucleus. Therefore, a higher energy will be required to remove an electron from the unipositive cation. This energy is called the second ionization energy and is represented as IE2

Third Ionization Energy


The dipositive ion thus formed will be further smaller and still higher energy will be required to remove an electron from it to form a triposititve ion. This energy is called the third ionization energy and is represented as IE3.




The dipositive ion thus formed will be further smaller and still higher energy will be required to remove an electron from it to form a triposititve ion. This energy is called the third ionization energy and is represented as IE3.

General Formula for Successive Ionization Energy


                        M (g)  + IE1            → M+ (g) + e-


                        M+ (g)  + IE2          → M2+ (g) + e-


                        M2+ (g)   + IE3        → M3+ (g) + e-



The second, third and fourth etc. ionization energies are known as successive ionization energies.


                                                 IE3 > IE2 > IE1 


It is comparatively more difficult to remove electrons from completely or half filled shells. This is why helium possesses much higher value of IE1 (2372.6 kJ mol-1) as compared to that of hydrogen (1312 kJ mol-1). 

If the removal of an electron results in the removal of the valence shell, the successive ionization energy will be much higher. For example, in lithium, the removal of an electron removes its valence shell. Therefore, IE2 of Li (7297 kJ mol-1) is much higher than IE1 (520.3 kJ mol-1).

The successive increase  in ionization energies of elements can be explained using the 5 factors that affect ionization energy.

These are namely:
  • Size of the Atom
  • Nuclear Charge
  • Screening Effect
  • Penetration effect of the Nucleus
  • Electron configuration of the Atom
We shall look into each of the aforementioned factors and see how successive ionization energies are affected by the changes in them.

Size of the Atom


The smaller the size of the atom greater is the ionization energy required to remove the outermost valence electron from it.



When electrons are removed from a metal ion which generally has less than 3 valence electrons. The valence shell is removed, therefore the subsequent shell experiences greater force of attraction. Due to this increased force of attraction greater ionization energy is required to remove the electron from that shell. This is one of the reasons why alkali metals are so reactive.




The smaller the size of the atom greater is the ionization energy required to remove the outermost valence electron from it.


Nuclear Charge


Thus to conclude, if the effective nuclear charge is increased for any atom, the ionization energy, the energy required to remove the outermost electron also increases.




As the electrons in the subsequent shell are closer to the nucleus, the effective nuclear charge experienced by them is also greater. With this increase in the magnitude of nuclear charge the pull experienced by the new valence electrons is greater than the what was experienced by the previous valence electrons which have been removed.


Thus to conclude, if the effective nuclear charge is increased for any atom, the ionization energy, the energy required to remove the outermost electron also increases.

Screening effect


Therefore, with a decrease in inner electrons, the ionization energy increases for successive removal of electrons.




The more the number of shells an atom has more is the shielding effect prominent.
In terms of successive ionization energies, the shielding effect is reduced and the effective nuclear charge increases due to the removal of electrons from the atom.

Therefore, with a decrease in inner electrons, the ionization energy increases for successive removal of electrons. 

Penetration effect of the Nucleus


Due to this penetration effect, the electrons experience greater ionization energy.  Hence for the same shell, the ionization energy follows the order   s>p>d>f.




The electron clouds of various atoms do not maintain distinct boundaries in a multi electron atom. Instead the electron cloud of one electron penetrates the electron cloud of other atom.

Due to this penetration effect, the electrons experience greater ionization energy.

Hence for the same shell, the ionization energy follows the order


s>p>d>f.

Electron configuration of the Atom


Half filled and completely filled shells are found to possess extra stability. The atoms having completely filled  shells are said to possess stable electronic configuration.




When an atom of Alkali metals loses electron it achieves the electron configuration of the nearest noble gas. For example let's take Sodium. If it loses an electron it attains the electron configuration of Neon. Which has a complete octet.

It is the most stable electron configuration. Therefore to remove an electron from Neon like species it would take an enormous amount of energy.


Half filled and completely filled shells are found to possess extra stability. The atoms having completely filled  shells are said to possess stable electronic configuration.


The atoms having filled shells also show an extra stability. Such atoms possess a tendency to lose the valence electron and consequently have higher values of ionization energies.


1. Helium possesses a stable electronic configuration 1s2. In it the K shell is completely filled . This is why the ionization energy of helium is much greater than that of hydrogen. Other gases also possess completely filled shells with stable configuration ns2 np6Therefore the noble gases have very high values of ionization energy.

2. Elements like Be, Mg etc. possess electronic configuration of the type of ns2 in which orbitals are completely filled. Therefore these elements also have higher ionization energies. 


3. Elements like N, P, etc. possess configuration of the type ns2 npx1 npy1 npz1. In these electronic configurations the p orbitals belonging to the valence shell are exactly half filled. Therefore, these elements show higher stability and have relatively higher ionization energies.

Learn about: Why does Ionization Energy decrease down the group?

Also Read: Trend of Ionization Energy in the Periodic Table




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