CHEMICAL BONDING AND MOLECULAR STRUCTURE

                   1. Introduction

Matter is made up of one or different type of elements. Under normal conditions no other element exists as an  independent atom in nature, except noble gases. However, a group of atoms is found to exist together as one species having characteristic properties. Such a group of atoms is called a molecule. Obviously there must be some force which holds these constituent atoms together in the molecules. The attractive force which holds various 

constituents (atoms, ions, etc.) together in different chemical species is called a chemical bond. Since the formation of chemical compounds takes place as a result of combination of atoms of various elements in different ways, it raises many questions. Why do atoms combine? Why are only certain combinations possible? Why do some 

atoms combine while certain others do not? Why do molecules possess definite shapes? To answer such questions different theories and concepts have been put forward from time to time. These are K b ö s s e l-Lewis approach, Valence Shell Electron Pair Repulsion (V S E P R) Theory, Valence Bond (VB) Theory and Molecular Orbital (MO) Theory. The evolution of various theories of valence and the interpretation of the nature of chemical bonds have closely been related to the developments in the understanding of the structure of atom, the electronic configuration of elements and the periodic table. Every system tends to be more stable and bonding is nature’s way of lowering the energy of the system to attain stability.

 

 2. K Ö S S E L-LEWIS APPROACH TO CHEMICAL BONDING

                             

In order to explain the formation of chemical bond in terms of electrons, a number of

attempts were made, but it was only in 1916 when K o s s e l and Lewis succeeded independently in giving a satisfactory explanation. They were the first to provide

some logical explanation of valence which was based on the inertness of noble gases.

Lewis pictured the atom in terms of a positively charged ‘Kernel’ (the nucleus plus

the inner electrons) and the outer shell that could accommodate a maximum of eight

electrons. He, further assumed that these eight electrons occupy the corners of a cube

which surround the ‘Kernel’. Thus the single outer shell electron of sodium would occupy one corner of the cube, while in the case of a noble gas all the eight corners would be occupied. This octet of electrons, represents a particularly stable electronic arrangement.

Lewis postulated that atoms achieve the stable octet when they are linked by

chemical bonds. In the case of sodium and chlorine, this can happen by the transfer of an electron from sodium to chlorine thereby giving the Na+ and Cl– ions. In the case of other molecules like C l2, H2, F 2, etc., the bond is formed by the sharing of a pair of electrons between the atoms. In the process each atom

attains a stable outer octet of electrons. Lewis Symbols: In the formation of a

molecule, only the outer shell electrons take

part in chemical combination and they are

known as valence electrons. The inner shell

electrons are well protected and are generally

not involved in the combination process.

G.N. Lewis, an American chemist introduced

simple notations to represent valence

electrons in an atom. These notations are

called Lewis symbols. For example, the Lewis

symbols for the elements of second period are

as under:

Significance of Lewis Symbols : The

number of dots around the symbol represents

the number of valence electrons. This number

of valence electrons helps to calculate the

common or group valence of the element. The

group valence of the elements is generally

either equal to the number of dots in Lewis

symbols or 8 minus the number of dots or

valence electrons.

K o s s e l, in relation to chemical bonding,

drew attention to the following facts:

» In the periodic table, the highly

electronegative halogens and the highly

electropositive alkali metals are separated

by the noble gases;

» The formation of a negative ion from a

halogen atom and a positive ion from an

alkali metal atom is associated with the

gain and loss of an electron by the

respective atoms;

» The negative and positive ions thus

formed attain stable noble gas electronic

configurations. The noble gases (with the

exception of helium which has a duplet

of electrons) have a particularly stable

outer shell configuration of eight (octet)

electrons, n s 2np6.

» The negative and positive ions are

stabilized by electrostatic attraction.

The bond formed, as a result of the

electrostatic attraction between the

positive and negative ions was termed as

the electro valent bond. The electro valence

is thus equal to the number of unit

charge(s) on the ion.

Thus, calcium is assigned a positive electro  valence of two, while chlorine a negative electro valence of one. 

Kössel’s postulations provide the basis for

the modern concepts regarding ion-formation

by electron transfer and the formation of ionic

crystalline compounds. His views have proved

to be of great value in the understanding and

systematisation of the ionic compounds. At

the same time he did recognise the fact that

a large number of compounds did not fit into

these concepts.

                          3. Octet rule

K o s s e l and Lewis in 1916 developed an

important theory of chemical combination

between atoms known as electronic theory

of chemical bonding. According to this,

atoms can combine either by transfer of

valence electrons from one atom to another

(gaining or losing) or by sharing of valence

electrons in order to have an octet in their

valence shells. This is known as octet rule.

                        4. Covalent Bond

The Lewis dot structures provide a picture

of bonding in molecules and ions in terms

of the shared pairs of electrons and the

octet rule. While such a picture may not

explain the bonding and behaviour of a

molecule completely, it does help in

understanding the formation and properties

of a molecule to a large extent. Writing of

Lewis dot structures of molecules is,

therefore, very useful. The Lewis dot

structures can be written by adopting the

following steps:

» The total number of electrons required for

writing the structures are obtained by

adding the valence electrons of the

combining atoms. For example, in the C H 4

molecule there are eight valence electrons

available for bonding (4 from carbon and

4 from the four hydrogen atoms).

» For anions, each negative charge would

mean addition of one electron. For cat ions, each positive charge would result in subtraction of one electron from the total

number of valence electrons. For example,

for the CO 3 2– ion, the two negative charges

indicate that there are two additional

electrons than those provided by the

neutral atoms. For NH 4+ion, one positive

charge indicates the loss of one electron

from the group of neutral atoms.

» Knowing the chemical symbols of the

combining atoms and having knowledge

of the skeletal structure of the compound

(known or guessed intelligently), it is easy

to distribute the total number of electrons

as bonding shared pairs between the

atoms in proportion to the total bonds.

» In general the least electronegative atom

occupies the central position in the

molecule/ion. For example in the NF 3 and

CO 3 2–, nitrogen and carbon are the central

atoms whereas fluorine and oxygen

occupy the terminal positions.

» After accounting for the shared pairs of

electrons for single bonds, the remaining

electron pairs are either utilized for multiple

bonding or remain as the lone pairs. The

basic requirement being that each bonded

atom gets an octet of electrons.

        5. Molecules (the Lewis Structures)

Lewis dot structures, in general, do not

represent the actual shapes of the molecules.

In case of polyatomic ions, the net charge is

possessed by the ion as a whole and not by a

particular atom. It is, however, feasible to

assign a formal charge on each atom. The

formal charge of an atom in a polyatomic

molecule or ion may be defined as the

difference between the number of valence

electrons of that atom in an isolated or free

state and the number of electrons assigned

to that atom in the Lewis structure.

                  6.  Bond length

Bond length is defined as the equilibrium

distance between the nuclei of two bonded

atoms in a molecule. Bond lengths are

measured by spectroscopic, X-ray diffraction

and electron-diffraction techniques about

which you will learn in higher classes. Each

atom of the bonded pair contributes to the

bond length. In the case of a covalent

bond, the contribution from each atom is

called the covalent radius of that atom.

The covalent radius is measured

approximately as the radius of an atom’s

core which is in contact with the core of

an adjacent atom in a bonded situation.

The covalent radius is half of the distance

between two similar atoms joined by a

covalent bond in the same molecule. The van

der Waals radius represents the overall size

of the atom which includes its valence shell

in a non bonded situation. Further, the van

der Waals radius is half of the distance

between two similar atoms in separate

molecules in a solid. Covalent and van der

Waals radii of chlorine are depicted.