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Covalent Bond

In the formation of a covalent bond, there occurs mutual sharing of electrons between the two involved atoms. This mechanism was suggested by G. N. Lewis. Taking the example of chlorine atom, we find that it has seven electrons in the outermost shell and it can be represented as .

Lewis Structure
Considering two chlorine atoms, we find that if each atom contributes one electron to the mutual sharing of electrons, then both the atoms will acquire the stable electronic configurations of eight electrons in the outermost shells. This is represented as follows:  
The above way of representing the distribution of valence electrons in a molecule is known as Lewis formula or structure. It is important to note that the use of dots and crosses to denote electrons is for illustrative purposes only. No difference actually exists between electrons from the different atoms; they are all equivalent.

The sharing of two electrons between two atoms constitutes a single covalent bond.  A single line joining the two atoms together usually represents it. For example, the chlorine molecule is represented as Cl—Cl.
From the above example, it is clear that two electrons are required to form a single covalent bond. Sometimes atoms need sharing of four or six electrons in order to have the stable electronic configuration of eight electrons in the outermost shell. The well-known examples are the oxygen and nitrogen molecules. The Lewis structures of these molecules can be represented as follows: 


The sharing of four electrons leads to the formation of a double bond between the involved atoms. Similarly, the sharing of six electrons leads to the formation of a triple bond.

The sharing of electrons involving two different types of atoms can be represented in a similar manner. For example, the combination of hydrogen and chlorine is represented as follows:
In a molecule containing more than two atoms, the Lewis formulae are represented in a similar manner. For example, the formation of water and ammonia can be represented as follows:
Formal Charge: Formal charge of an atom in a polyatomic molecule or ion may be defined as the difference between the no. of valence electrons of that atom in an isolated or free state and the no. of electrons assigned to that atom in the Lewis structure. It is expressed as:

Formal charge = ( total number of valence electrons in the free atom) - (total no. of nonbonding electrons (lone pair) - ½ (total no. of bonding electrons (shared)


Formal charge of O marked 1
6 - 4 - ½(4) = 0

Formal charge of O marked 2
6- 6 - ½(2) = -1

Formal charge of O marked 3
6- 6 - ½(2) = -1

The formal charge is a concept based on a pure covalent view of bonding in which electrons are equally shared between the neighbouring atoms. Lowest energy structure is the one with the smallest formal charge on the atoms hence formal charge helps in locating stable structure.

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