What is Ionization Energy -[ Definition, Trend, Periodic table] & How to Calculate Ionization Energy

At the end of the article, you will able to describe- What is Ionization Energy, Definition, Trend, Periodic table and how to calculate ionization energy. Let’s start discussing one by one.

What is Ionization Energy – Definition

The amount of energy (work) required to remove one electron from the last orbit of the parent atom is term as ionization energy(I.E).

We have to utilize energy. The energy is needed to pull out electrons well. Electrons don’t simply come off. In atoms, electrons are held together by electrostatic force. The Electrostatic force comes from the positively charged nucleus and the negatively charged electrons. So, something must pull those electrons away. In other words work or energy must be inputted into our system to pull that electron off.

For Example –

  Calcium (Ca) → Ca2+ + 2e

Calcium in its neutral state. We can take away two (or more) electrons to make it calcium plus two (Ca2+).  Some atoms can pull away or can give off more than one electron.

  • Now the energy required to pull away that first electron is known as the first ionization energy.
  • While the energy required to pull away that second electron is known as the second I.E and so on.

Ionization Energy Periodic Table

We see in a periodic table that as we go across a period the ionization energy of an atom tends to increase. The effective nuclear charge (protons) increase as we go from left to right.


To explain that let’s look at Coulomb’s law. Once again, Coulomb’s law states that the force is equal to constant K times the product of two charges divided by the distance between them squared.

  • Remember the denominator (r2) decreases that means force (F) tends to increase.
  • The increase in charge causes the force to increase.
  • The decrease in the atomic radius tends to increase force.

There exist two charges one is protons and other one are electrons.  We already said that our effective nuclear charge tends to increase as we go from left to right.

Taking an example, Now what happens, as we move from lithium to fluorine. Fluorine has the highest force. In other words, the protons found in the nucleus pull those electrons on the outermost electron shell with a lot of force much more force than lithium (or beryllium or boron or carbon).

That means it’s going to require much more energy to pull those outermost electrons off and that’s exactly why as we go across a period from lithium to fluorine ionization energy tends to increase. Because as we go Left to right we have a higher effective nuclear charge which means we have a greater force.

When we go down the group, atomic radius increases. If we go back to Coulomb’s law if atomic radius increases that increases distance between two charges (protons and electrons) increases. So R also increases (denominator is increased) and that means our force is less.

So, as we go down a group our ionization energy tends to decrease.

So basically the higher ionization energy is the less likely you are to give up electrons

Ionization Energy Trend

Atomic size or radius

The larger the atomic size, smaller is the I.E.  As the size of the atom increases, the outer electrons lie farther away from the nucleus and thus exert less attraction towards the nucleus and hence can be easily removed.

What is Ionization Energy -[ Definition, Trend, Periodic table] & How to Calculate Ionization Energy

 Screening effect

The larger the number of electrons in the inner shell, the greater is the screening effect on the electrons in the outer (valence) shell which thus experiences less attraction from the nucleus and thus can be easily removed leading to the lower value of ionization potential. Now as we move down a group, the number of inner shells increases and hence the ionization potential tends to decrease.

 Nuclear charge Effect in first ionization energy

Nuclear charge is defined as the net nuclear attraction towards the valence shell electrons.

More is the effective nuclear charge, more tightly the electrons are held with the nucleus and thus more energy will be required to remove electron leading to higher I.E.

Effective nuclear charge (Zeff) = Z-S

Where Z = Nuclear charge ; S = Screening constant.

On moving along the period the charge on nucleus increases as the atomic number increases while the valence shell remains the same and thus effective nuclear charge increases which lead to higher I.E. Therefore, I.E increases along the period with some disorders, viz.  IE of elements of the 3rd group is less than that of the 2nd group elements, and similarly, IE of group 16 elements is less than that of group 15 elements.

An increase in positive charge on the ion increases the effective nuclear charge which in turn increases the I.E. On the other hand, an increase in negative charge on the ion decreases the effective nuclear charge which in turn decreases the I.E.

Half filled and fully filled electronic configuration

According to Hund’s rule, atoms having half-filled or completely filled orbitals are comparatively more stable and hence more energy is needed to remove an electron from such atoms leading to higher I.E than the usually expected value.

Arrangement of electrons(Symmetry)

Symmetry plays a vital role in first I.E. If an atom or an ion has s2p6 configuration, its I.E is extremely high due to the presence of the so-called octet arrangement (noble gas configuration). This explains why IE2 of Li(lithium) is very very high as compared to IE1

IE3 >>> IE2 > IE1

Removal of s, p, d and f electrons from the same shell

Since s-orbital is more closet the nucleus than the p-orbital of the same orbit, its electrons experience more attraction than that of p and hence their removal is difficult leading to higher IE. In general, the I.E follows the following order s> p > d > f orbital of the same orbit.

s> p > d > f orbital of the same orbit.

Remember: I.E. measured in a unit of an electron volt (eV) per atom or kilojoule per mole or kilocalorie per mole.
1 eV/atom kcal.

How to Calculate Ionization Energy

Actually, the electrons are held by the nucleus of the metal atoms by a certain force called binding force. In order to escape electrons we to supply the energy to overcome the binding force. This job is performed by the photon which contains minimum energy be called as threshold energy to break down the binding energy. The threshold energy is also known as work function.

Ionization Energy=Work function + kinetic energy


  • Energy E = hv.
  • Work function E0 =hv0
  • kinetic energy K.E = 1/2mv2.

  1. If the energy of incident Photon is <(less than) threshold energy no electron will be emitted.
  2. If the incident Photon has the energy =(equal) to the threshold energy electron will just release from the metal surface.
  3. In case, an incident photon has the energy >(greater than) than threshold energy emit electron will acquire some kinetic energy.
  4. The kinetic energy of electrons is directly proportional to the frequency of striking Photon and it is quite independent of intensity.
  5. The number of electrons is ejected per second depends upon the intensity of striking Photon, not upon their frequency.

This is all about the basics – What is Ionization Energy, Definition, Trend, Periodic table and how to calculate ionization energy.

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