Electron Binding Energy – Ionisation Energy

In more technical terms, ionisation energy is defined as the minimum amount of energy that an electron in a gaseous atom or ion must absorb to escape the influence of the nucleus. It is also known as ionisation potential and is typically an endothermic process.

We can also conclude that ionisation energy provides information about the reactivity of chemical compounds. It is also helpful in determining the strength of chemical bonds. It is measured in either electron volts or kJ/mol units.

Depending on how molecules are ionised, which frequently results in changes in molecular geometry, ionisation energy can be either adiabatic ionisation energy or vertical ionisation energy.

Electron binding energy 

The ionisation energy for every ion is calculated in the periodic formula. It is first important to know the calculation behind measuring the energy required to lose electrons to calculate the ionisation energies. Ionisation energy equation can be written as- X → X+ + e-

Every time an electron is withdrawn from a molecule or an atom, a considerable amount of energy is needed to eliminate the remaining electrons from the same molecule or atom. Therefore, the change in the equation occurs. Every time there’s a change in the amount of energy needed to remove an electron, the equation changes. 

• First Ionisation energy equation

X → X+ + e- 

• Second Ionisation energy equation 

X+ → X2+ + e- 

• Third Ionisation energy equation 

X2+ →  X3+  + e- 

Units used for the measurement of Ionisation energy are not equal to each other. Due to the Ionisation energy, it is recognised how closely electrons are aligned by their atoms. The closer the electron is to its atom, the greater the potential for ionisation. 

The ionisation energy advances are directly opposite to atomic radii. As atomic radii decrease, the ionisation force mainly increases, whereas the ionisation force decreases if atomic radii increase. 

Factors Affecting Ionisation Energy

When the ionisation energy is high, removing an electron is typically more challenging. Several factors also govern the attraction forces.

• An electron is pulled towards a positively charged nucleus by force.

• When an electron is close to the nucleus, the attraction is more potent than when the electron is further away.

• The attraction forces are reduced as the number of electrons between the outer level and the nucleus increases.

• When two electrons are in the same orbital, they are repelled somehow. This now causes disturbances in the nucleus’s attraction. Because paired electrons can be easily removed, their ionisation energy will be lower.

Ionisation

Ionisation is the removal of an electron from an orbit to the outside of the atom. The difference between the energy of the electron in the initial orbit and the electron’s energy outside the atom is determined by the electrons’ orbital energy.

Ionisation energy and its Periodic trend 

In a period 

The atomic radius decreases every time you move left to the right in the periodic table. So, if the atom size decreases, the force of attraction between the outermost electrons and the nucleus increases. Because of this, after a certain point in the periodic table, the ionisation energy automatically increases. During the second period, discrepancy appears in the Ionisation energy trend starting from the boron to the beryllium. 

Ideally, it was suggested that the ionisation energy of beryllium should be less than that of boron, but sadly the opposite of it happens. This happens because of the penetration effects, and also beryllium has fully filled subshells. Another reason is that beryllium has 2s orbital only, whereas the boron deals in 2p and 2s orbitals. Therefore, it is easier to eliminate electrons from boron than beryllium with 2s orbitals. 

In groups 

As you move down in a group, the number of shells increases. As a result, the element’s ionisation energy automatically decreases. The innermost electron will be the closest to the nucleus, and vice-versa; therefore, there will be a significantly less nuclear charge. 

Conclusion :

The Ionisation energy can be defined as the amount of energy required to remove an electron from an isolated gaseous atom in its gaseous form. Every time an electron is withdrawn from an atom, a considerable amount of energy is consumed to lose it. This is known as Ionisation energy. Ionisation energy falls from top to bottom in the periodic table, whereas it rises from left to right.