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Ionization Energy

In long form of the periodic table, elements are arranged in 9 periods and 18 groups. In a group, elements have the same number of valence electrons in their outermost shell. Valence electrons refer to those electrons of an atom that can form chemical bonds with other atoms. The arrangement of electrons across the shells is known as the electronic structure of the element.

In the electronic structure, the number of shells increases steadily as we compare elements of one group. Or we can say that the number of shells increases down the group. Now observe the electronic structure of the elements in the second period. What do you observe? The elements have different valence electrons but same number of shells, that is, 2. Notice that the atomic number of these elements increase by 1 unit, from left to right, the number of valence shell electrons also increases by 1 unit.

The number of elements in each period is based on how electrons are filled in various shells. For example,
K Shell – 2 × (1)2 = 2, hence the first period has 2 elements.
L Shell – 2 × (2)2 = 8, hence the second period has 8 elements and so on.

From all these observations, we can conclude trend of various properties of elements such as ionisation potential, electron affinity, electron negativity etc. Let’s discuss the ionisation potential of elements in the periodic table.

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Ionization Energy Definition

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The electrons are raised to higher energy levels by absorption of energy from an external source. If this process is continued, a stage comes when the electron goes completely out of the influence of the nucleus and a positive ion is produced. The electron in an atom is attracted by the positively charged nucleus.
In order to remove an electron from an atom, energy has to be supplied to it to overcome the attractive force. This energy is referred to as ionization energy or ionization potential.
"The amount of energy required to remove the most loosely bound electron or the outermost electron from an isolated gaseous atom of an element in its lowest energy state or the ground state to produce a cation is known as ionization potential or ionization energy of that element."

A(g) + Energy $\to$ A+ (g) + electron
It is generally represented as I or IP and is measured in electron volts (eV) or kilo calories (K calories) per gram of atom.

1eV atom = 23.06 kcal / mole = 96.3 kj/mole

Ionization potential is a very important property which gives an idea about the tendency of an atom to form a gaseous positive ion. The smaller the value of ionization energy, the easier it is to remove the electron from the atom.

The process by which the element loses an electron, to convert itself into a cation is called its ionization. This process is an endothermic process, since the energy is supplied to effect it.

Ionization Energy Trend

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In a group, as we move from top to bottom, ionization potential decreases. As we move down the group, with the increase in atomic number, new shells ass on and the size of the atom gradually increases. The distance between the nucleus and the outermost shell increases and the screening effect also increases.

As a result, the attractive force of the nucleus on the outermost electron decreases. Thereby, the amount of energy required to remove an electron from the atom decreases. Ionization potential, thus, decreases down a group.

Element kJ/mol eV/atom
H 1312.0 13.595
Li 520.1 5.390
Na 459.2 5.138
K 418.7 4.339
Rb 403.0 4.176
Cs 375.7 3.893

The element with highest ionization potential in the periodic table is Helium (2372.1 kj/mol), while the element with the lowest ionization potential is Caesium (375.7 kj/mol).

First Ionization Energy

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First ionization energy of an atom is the energy required to remove the first electron from the outermost shell of an atom.

Once the first electron has been removed from the gaseous atom, it is possible to remove second and successive electrons from positive ions one after the other.

Ionization Energy Periodic Table

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In a period, as we move from left to right, the ionization energy increases.
As we move along a period, the nuclear charge increases and the electrons are added into the same shell. Thereby the effective nuclear charge increases and size decreases. Therefore, the energy required to remove an electron increases.

Elements
Li Be B C N O F Ne
Ionization Potential 520 900 802 1085
1398
1315
1680
2080

However, some irregularities in the general trend have been noticed. These are due to half - filled and completely filled configurations which have extra stability.

To illustrate this, let us consider the ionization energies beryllium and boron. The ionization energy of boron is slightly less than that of beryllium even though boron has a greater nuclear charge. This can be understood by comparing the electronic configurations of Be (1s2, 2s2 ) and B (1s2, 2s2, sp1).

In the case of Beryllium, the electron removed during ionization is an s- electron whereas the electron removed in the case of boron is p-electron. The penetration of a 2s electron to the nucleus is more than that of a 2p electron, hence the 2p - electron of boron is more shielded than the 2s electron of Beryllium. Therefore, ionization energy of Beryllium is more than that of Boron.

Second Ionization Energy

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Once the first ionization electron is removed from the gaseous atom, it is possible to remove second and successive electrons from positive ions one after the other. Removal of a second electron from an already ionized gaseous atom is called as second ionization energy.

A(g) + IE1 → A+(g) + e- [First Ionization]
A+(g) + IE1 → A2+(g) + e- [Second ionization]

The amount of energies required to remove most loosely bound electron from unipositive, dipositive, tripositive .... ions of the element in gaseous state are called second, third, fourth, etc ionization energies respectively.

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

It is also seen that IE3 > IE2 > IE1.

Ionization Energy Chart

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Ionization chart of II A group elements are as follows.

Element Number of inner shells eV/atom
Be 1 9.3
Ma 2 7.6
Ca 3 6.1
Sr 4 5.7
Ba 5 5.2

Factors Affecting Ionization Potential

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The magnitude of Ionization potential depends upon:

1. Charge in the nucleus


The greater the charge on the nucleus, the more difficult it will be to remove an outermost electron from that atom. Thus, greater the nuclear charge of an atom, greater will be its ionization potential. Halogens have highest ionization potential in their respective periods.

2. Atomic radius


With the increase in the atomic radius, the ionization potential decreases. This is because of the fact that in case of larger atoms the attraction between the nucleus and the outer most electron is less and hence it is easier to remove an electron from a larger atom than from a smaller one.

This can be seen from the fact that as we move down a group, the size increases, while ionization potential decreases.

3. Completely and partially filled in shells


The atoms having a completely filled or partially filled outer most shell are comparatively more stable than the atoms with incompletely filled outer shells. This, the ionization potential of such atoms are relatively higher than expected normally from their positions in the periodic table.

Successive Ionization Energies

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The amount of energy necessary for removing the subsequent electrons of a gaseous atom is better known as successive ionization energies. These are basically termed as 1st , 2nd, 3rd or 4th …… ionization energy and these depends completely on the removal of the 1st, 2nd, 3rd electrons respectively.

$M_{(g)} \overset{IE_1}{\to} M_{(g)}^{+} + e^-$

$M_{(g)}^{+} \overset{IE_2}{\to} M_{(g)}^{2+} + e^-$

$M_{(g)}^{2+} \overset{IE_3}{\to} M_{(g)}^{3+} + e^-$
  • 2nd ionization energy amount are higher compared to 1st because after the first electron is removed the atom changes into a positive monovalent ion.
  • The nuclear charge remains same although there is a decrease in electron number and this ultimately leads to a situation where the remaining electrons getting held more tightly by the nucleus which ultimately makes it very difficult to remove the second electron.
  • The value of the 2nd ionization energy (IE2) is more than the 1st ionization energy (IE1). Removing 2nd electron would ultimately lead to forming a di-positive ion which results in a stronger attraction between the nucleus and remaining electrons. Higher value of ionization energy results from this.
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