Chemical Bonding and Molecular Structure
4.0 Covalent Bond (By Mutual Sharing of Electrons)
4.0 Covalent Bond (By Mutual Sharing of Electrons)
The covalent bond is formed when two atoms achieve stability by the sharing of an electron pair, each contributing one electron to the electron pair.
The arrangement of electrons in a covalent molecule is often shown by a Lewis structure in which only valency shells (outer shells) are depicted. For sake of clarity, the electrons on different atoms are denoted by dots and crosses.
Polarity of Bonds:
A covalent bond is set up by sharing of electrons between two atoms. It is further classified as polar or non-polar depending upon the fact whether the electron pair is shared unequally between the atoms or shared equally.
For example, the covalent bonds in $H_2$ and $Cl_2$ are called non-polar as the electron pair is equally shared between the two atoms.
In the case of hydrogen fluoride the bond is polar as the electron pair is unequally shared. Fluorine has a greater attraction for electrons or has higher electro-negativity than hydrogen and the shared pair of electrons is nearer to the fluorine atom than hydrogen atom. The hydrogen end of the molecule, therefore, appears positive with respect to fluorine.
Bond polarities affect both physical and chemical properties of compounds containing polar bond. The polarity of a bond determines the kind of reaction that can take place at that bond and even affects the reactivity at nearby bonds. The polarity of bonds can lead to polarity of molecules and affect melting point, boiling point and solubility.
Dipole Moment:
It is vector quantity and is defined as the product of the magnitude of charge on any of the atom and the distance between the atoms. It is represented by $\mu $. Magnitude of dipole moment is $$\left| \mu \right| = q \times r$$ where $q$ is charge in esu and $r$ is distance in $A$.
The unit $= 10^{–18}$ (esu) cm (D) is used in practice. In SI units charge $q$ is measured in coulombs $(C)$ and the distance, $r$ in metre, $m$.$$\begin{equation} \begin{aligned} 1C = 2.998 \times {10^9}esu\;and\;1m = {10^2}cm \\ \therefore 1Cm = 2.998 \times {10^9} \times {10^2} = 2.998 \times {10^{11}}(esu)m \\\end{aligned} \end{equation} $$Therefore in SI system, the unit of dipole moment is coulomb metre. $$\therefore 1Cm = \frac{{2.998 \times {{10}^{11}}}}{{{{10}^{ - 18}}}} = 2.998 \times {10^{29}}D$$ or $$1D = \frac{1}{{2.998 \times {{10}^{29}}}} = 3.336 \times {10^{ - 30}}$$
Dipole moment is a vector quantity and is often indicated by an arrow parallel to the line joining the point of charge and pointing towards the negative end e.g., $\overline {H - F} $.
$\%$ Ionic character of a covalent bond is calculated using $$ = \begin{equation} \begin{aligned}
\frac{Experimental Dipole moment}{Theoretical Dipole moment}\end{aligned} \end{equation} $$ assuming 100% ionic character
In general a polar bond is established between two atoms of different radii and different electronegativities while positive centres (nuclei) of different magnitudes combine to share an electron pair. Greater the values of the dipole moment, greater is the polarity of the bond.
The following points may be borne in mind regarding dipole moments:
- In case a molecule contains two or more polar bonds, its dipole moment is obtained by the vectorial addition of the dipole moments of the constituent bonds.
- A symmetrical molecule is non-polar even though it contains polar bonds.
For example, carbon dioxide, methane and carbon tetrachloride, being symmetrical molecules, have zero dipole moments.
Dipole moment of methyl chloride is a vectorial addition of dipole moments of three $C – H$ bonds and one $C – Cl$ bond.
Dipole moment gives valuable information about the structure of molecules. For example, carbon dioxide is assigned a linear structure since its dipole moment is zero.
We have seen that in a polar covalent bond between two atoms (say $A$ and $B$), there is a partial separation of charge. This bond is, therefore, said to have a partial ionic character. Greater the difference of electronegativity between $A$ and $B$, greater is the degree of ionic character (or polarity measured by dipole moment of $AB$) of the bond.
Pauling gave a fairly accurate rule by which the nature of the bond can be predicted. According to this rule,
“If the difference on the electronegativity scale between the two atoms is $1.9$, the bond is $50\%$ ionic in character. When the difference is greater than $1.9$, the bond is correspondingly more ionic”.
For example, when the electro-negativity difference is $0.8$, $1.2$, $2.2$ and $2.6$, the corresponding partial ionic character is $12\%$, $25\%$, $61\%$ and $74\%$ respectively.
