# What are First and second Ionization potentials

First and second Ionization potentials :

Ionization potential is the amount of energy required to remove the most loosely held electron from isolated a neutral gaseous atom to convert it into gaseous ion. It is also known as first ionization potential because it is the energy required to remove the first electron from atom.

It is denoted as $I_{1}$ and is expressed in electron volts per atom, kilo calories (or) kilo joules per mole.

$M_{(g)}&space;+&space;I_{1}&space;\xrightarrow{\hspace*{1cm}}$$Mg_{(g)}^{+}&space;+&space;e^{-}$

$I_{1}$ is first ionization potential.

The energy required to remove another electron from the unipositive ion is called the Second ionization potential. It is denoted by $I_{2}$.

$Mg_{(g)}^{+}&space;+&space;I_{2}$ $\xrightarrow{\hspace*{1cm}}$ $Mg_{(g)}^{2+}&space;+&space;e^{-}$

The second ionization potential is greater than the first ionization potential. On removing an electron from an atom, the unipositive ion formed will have more effective nuclear charge than the number of electrons. As a result, the effective nuclear charge increases over the outermost electrons. Hence more energy is required to remove the second electron. This shows that the second ionization potential is greater than the first ionization potential.

For sodium, $I_{1}$ is 5.1 eV and $I_{2}$ is 47.3 eV.

$I_{1}&space;<&space;I_{2}&space;<&space;I_{3}&space;\;......\;&space;I_{n}$

Factors affecting ionization potentials :

As the size of the atom increases the distance between the nucleus and the outermost electrons increases. So the effective nuclear charge on the outermost electrons decreases. In such a case the energy required to remove the electrons also decreases. This shows that with an increase in atomic radius the ionization potential decreases.

Nuclear charge :

As the positive charge of the nucleus increases, its attraction increases over the electrons. So it becomes more difficult to remove the electrons. This shows that the ionization potential increases as the nuclear charge increases.

Screening effect or shielding effect :

In multielectron atoms, valence electrons are attracted by the nucleus as well as repelled by electrons of inner shells. The electrons present in the inner shells screen the electrons present in the outermost orbit from the nucleus. As the number of electrons in the inner orbits increases, the screening effect increases. This reduces the effective nuclear charge over the outermost electrons. It is called screening effect or shielding effect. With the increase of screening effect the ionization potential decreases. Screening efficiency of the orbitals falls off in the order $s&space;>&space;p>&space;d>&space;f.$

(Magnitude of screening effect) $\propto$ $\frac{1}{(Ionzation&space;\;enthalapy)}$

Trend in a group:

The ionization potential decreases in a group, gradually from top to bottom as the size of the elements increases down a group.

Trend in a period:

In a period from left to right I.P. value increases as the size of the elements decreases along the period.