# Covalent Character in Ionic Compounds for JEE

## Introduction to Covalent Character of Ionic Bonds

Ionic bonds are formed between two oppositely charged ions (cation-anion). In the process of formation of ionic bonds one element donate electrons to another element to get positive charge on itself and negative charge on another element. On the other hand covalent bonds are formed by sharing electrons between elements. However, there are some other factors such as electronegativity, charge, size, etc. that give ionic bonds some covalent character and covalent bonds some ionic character.

Generally it is a common experience that many times an ionic bond has some covalent character and a covalent bond has some ionic character. The formation of a bond intermediate between an ionic and a covalent bond occurs through a phenomenon known as polarisation of ions.

## Polarisation of Ions

When two oppositely charged ions approach each other, the attraction between the positive charge of cation and the negative charge of anion and also the simultaneous repulsion between both of their nuclei and between their electrons results in the distortion, deformation or polarisation of the electron charge cloud of the anion. The electron charge cloud of the anion that is formed, no longer remains spherical but gets distorted, i.e., polarised towards the cation, as shown in the Image below.

Polarisation of electron charge cloud of an anion by cation

The electron charge cloud of the cation also gets distorted, i.e., polarised by the anion through a similar process but the polarisation of the cation is far less  pronounced because of its small size. The high electron charge concentration between the two nuclei is a result of the polarisation of the ions. This results in the formation of a bond which is intermediate between an ionic and a covalent bond. These types of bonds are generally referred to as a polar covalent bond. A polar covalent bond is practically more stable than a pure covalent bond. The higher the degree of polarisation, the greater is the stability of the polar covalent bond. The extent of polarisation evidently depends on the polarising power of the cation and also the polarisability of the anion.

## Fajan’s Rules and Application of Fajan’s Rule

The rules regarding polarisation are referred to as Fajan's rules. The significance of Fajan’s rule of polarisation is given below.

1. The cations with smaller size have higher polarising power, i.e., they cause polarisation of electron charge clouds of an anion to a greater extent. Such cations have positive charge concentrated over a small surface area, i.e., they have a high charge density and thus distort the electron charge cloud of the anion highly efficiently. If the cation is of larger size it will  have a low polarising power.

2. The anions with large size have high polarizability, i.e., their electron charge cloud can be deformed easily. The grip of the nucleus on the orbital electrons in large anions will be weak and hence such anions will get polarised by a cation relatively easily.

3. For effective polarisation, there should be a high charge on the cation or the anion or both. The electrostatic forces which cause polarisation, would increase with increase in the charge on the ions.

4. Cations with pseudo inert gas configuration, (ns2p6d10) or with inert pair configuration, (d10(n+1)s2), have high polarising power while cations with noble gas configuration (ns2p6), have low polarising power. This is due to the greater effective nuclear charge in the former cases((ns2p6d10) and d10(n+1)s2) and the smaller effective nuclear charge in the latter cases ((ns2p6)) due to poor shielding effect of d orbital.

### Effects of Polarisation

1. The solubility of ionic compounds in polar solvents decreases with increase in the degree of polarisation, that is with an increase in the degree of covalent bonding.

2. The hardness of ionic compounds decreases with increase in the degree of polarisation i.e. with increase in the covalent character.

## Percentage Ionic Character

The percent ionic character of a polar covalent bond usually depends upon two prominent factors:

1. Electronegativity Difference of the Bonded Atoms and

2. Dipole Moment of the Compound Formed

### Electronegativity Difference and Percentage Ionic Character Formula

A pure covalent bond is formed between atoms A and B when the electronegativities of the two atoms remain the same. If A and B have different electronegativities, the resulting bond would be a polar covalent bond since the bonding electron pair is pulled more towards the more electronegative atom thereby bringing some excess negative charge on the latter.

As a result, some excess positive charge is developed on the less electronegative atom. If A is more electronegative than B, a polar covalent bond $\overset{\delta-}{A}-\overset{\delta+}{B}$ would be formed between A and B. The larger the difference in the electronegativities of A and B, the greater would be the magnitude of δ are the higher would be the polarity of the bond.

Several empirical equations have been proposed to calculate the percent ionic character of a covalent bond from the electronegativities of the bonding atoms. Two of these equations which are used are the Pauling Equation and the Hannay Smith equation.

### Pauling Equation

Pauling proposed the following empirical equation for determining the percent ionic character of a covalent bond:

Percent ionic character $=1-e^{-\dfrac{1}{4}\left(\chi_{A}-\chi_{B}\right)}$

Where $\chi_{A}$ and $\chi_{B}$ are the electronegativities of A and B, respectively.

### Hanna-Smith Equation.

Hannay and Smith proposed the below equation for the purpose :

Percent ionic character $=16\left(\chi_{A}-\chi_{B}\right)+3 \cdot 5\left(\chi_{A}-\chi_{B}\right)^{2}$

Both the equations, however, give only approximate values.

## Conclusion

We discussed what is covalent character significance in ionic bonds. Ionic compounds generally show partial covalent character. For example, the ionic compound LiCl2 shows covalent character and is soluble in organic solvents such as ethanol.

The partial covalent character of ionic bonds in ionic compounds can be explained using the concept of a phenomenon called polarisation. We know that in an ionic compound, there is an electrostatic force existing between the anion and cation. The positively charged cation attracts the valence electrons of anion while repelling the nucleus. This will be causing a distortion in the electron cloud of the anion and its electronic density drifts towards the cation, which results in some sharing of the valence electrons between these two ions. Thus, a partial covalent character is developed between them. This phenomenon is called polarisation. The extent of polarisation can be understood better by Fajan's rule.

## FAQs on Covalent Character in Ionic Compounds for JEE

1.  What is a dipole moment? Describe the dipole moment in a CO2 molecule.

The dipole moment is the term used to measure the polarity of a covalent bond, it can be defined as μ = q 2d. Where μ is the dipole moment, q is the charge and 2d is the distance between the two separate charges. The dipole moment is a vector quantity and the direction of the dipole moment vector points from the negative charge to the positive charge. In CO2, the dipole moments of two polar bonds (CO) are equal in magnitude but have opposite directions. Hence, the net dipole moment of the CO2 is $\mu=\mu_{1}+\mu_{2}=\mu_{1}+\left(-\mu_{1}\right)=0$.

2. Explain briefly about the Partial ionic character in covalent bond.

When a covalent bond is formed between two identical atoms, both of the atoms have equal tendency to attract a shared pair of electrons and hence the shared pair of electrons will be lying exactly in the middle of the nuclei of the two atoms. However in the case of covalent bond formed between the two atoms having two different electronegativities the atom with higher electronegativity will have greater tendency to attract the shared pair of electrons more towards itself than the other actors result the cloud of shared electron pair getting distorted.

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