CHEMICAL BONDING

INORGANIC CHEMISTRY

B.Sc FIRST YEAR

INTRODUCTION

The atoms are said to combine together because of the following two main
reasons:

(i) Concept of lowering of energy
It has been observed that the aggregate (or the molecules) are lower in energy than the individual atoms from which they have been formed. This means when the individual atoms combine to form molecules through a bond, the potential energy of the combining atoms decreases and the resulting molecules are more stable than the free atoms. This energy
difference between the free atoms and bonded atoms (or molecules) is generally 40kJ mol^-1 or more. It follows from this that the process of bond formation between the atoms decreases the energy of the molecule formed from these atoms and forms a system of lower energy and greater stability.
(ii) Electronic theory of valence (the octet rule)
The atoms of the noble gases-helium to radon- do not, except a few cases, react with any other atoms to form the compounds and also they do not react with themselves. Hence they stay in atomic form. These atoms are said to have low energy and cannot be further lowered by forming compounds. This low energy of noble gas atoms is associated with their outer shell electronic configuration, i.e. the stable arrangement of eight
electrons (called octet). It has also been established that the two electrons in
case of helium atom (called doublet) is as stable as an octet present in other noble gas atoms. The chemical stability of the octet of noble gases led chemists to assume that when atoms of other elements combine to form a molecule, the electrons in their outer shells are arranged between themselves in such a way that they achieve a stable octet of electrons (noble
gas configuration) and thus a chemical bond is established between the atoms. This tendency of the atoms to attain the noble gas configuration of eight electrons in their outer shell is known as octet rule or rule of eight and when the atoms attain the helium configuration, it is called doublet rule or rule of two. This octet rule was later called “Electronic Theory of Valence”.
It may be noted here that in the formation of a chemical bond, atoms interact with each other by losing, gaining or sharing of electrons so as to acquire a stable outer shell configuration of eight electrons. This means, an atom with less than eight electrons in the outer shell is chemically active and has a tendency to combine with other atoms. Accordingly, three different types of bonds may exist in the molecules/aggregates.

COVALENT BOND

A covalent bond is formed between the two combining atoms, generally of the electronegative non-metallic elements by the mutual sharing of one or more electron pairs (from their valence shell). Each of the two combining atoms attains stable noble gas electronic configuration, thereby enhancing the stability of the molecule. If one electron pair is shared between the two atoms, each atom contributes one electron towards the electron pair forming the bond. This electron pair is responsible for the stability of both the atoms. A covalent bond is denoted by the solid line (-) between the atoms. Depending on the number of shared electron pairs i.e. one, two, three etc. electron pairs between the combining atoms, the bond is known as a single, double, triple etc. covalent bond. For example,

Single Bond

H:H H-H
Cl:Cl Cl-Cl
H:Cl H-Cl

Multiple Bond
O::O O=O (double bond)
N:::N N≡N (triple bond)

In the molecules, the bond strength and bond length has been found in the following order:
Bond strength: triple bond > double bond > single bond
Bond length: triple bond < double bond < single bond
It may be noted that the covalent bond formation between multielectron atoms involves only the valence shell electrons that too, the unpaired electrons. Thus O-atom has two unpaired electrons in its valence shell and N-atom has three unpaired electrons there by forming two and three bonds with themselves or other atoms.

Polar and non-polar covalent bond
In the examples given above, most of the bonds viz. single, double and triple covalent bonds, have been shown to be formed between the like atoms such as H-H, Cl-Cl, O=O and N≡N in H2, Cl2, O2 and N2, respectively. The bonded atoms in these molecules attract the bonding or shared pair of electrons by equal forces towards themselves due to equal electronegativity of the atoms. Hence the bonding pair of electron lies at the mid point of the internuclear distance ( or bond distance). This type of bond is known as the
non-polar covalent bond.But if the covalent bond is formed between two unlike atoms of different elements, e.g.HCl, H2O, NH3 etc., the shared pair of electrons will not be equally by the bonded atoms due to electronegativity difference.It shifts towards more electronegative atom and hence moves away from less electronegative atom. This develops small negative charge on more electronegative atom and equal positive charge on less electronegative atom. Such a molecule is called a polar molecule (this is different from ionic bond) and the bond present in such molecules is known as polar covalent bond.

Valence Bond Theory (VBT)
This theory was put forward by Heitler and London in 1927 to covalent bond. They gave a theoretical treatment molecule and the energy changes taking place therein. Later Slater in 1931 to account for the directional characteristics of the covalent bond. The main points called the postulates of this theory are given below:
(i) The atoms involved in the bond formation maintain their individuality
after the bond is formed i.e. in the molecule.

(ii) The bond is formed due to the overlapping of half filled atomic orbitals (or the interaction of electron waves) belonging to the valence shell of the combining atoms as these approach each other. Thus the spins of the two electrons get mutually neutralised. The electrons in the orbital of inner shells remain undisturbed.

(iii) The filled orbitals (i.e. containing two electrons) of the valence shell do not take part in the bond formation. However, if the paired electrons can be unpaired without using much energy, they are first unpaired by promoting to the orbitals of slightly higher energy and then can take part in bonding. For example, N can form NCl₃ only retaining a lone pair while P can form both PCl₃ andPCl₅

(iv) The electrons forming the bond undergo exchange between the atoms and thus stabilize the bond.
(v) The strength of the covalent bond depends on the extent to which the two atomic orbital overlap in space.