CHEM 1405 – CHAPTER 12
CHEMICAL BONDING
The chemical Bond Concept
The attractive force that holds the two atoms in a molecule together is termed as a chemical bond.
The valence electrons are responsible for holding the two atoms together.
Octet Rule
Every atom has the ability to have a stable configuration of eight electrons in the valence shell. This
tendency of atoms is the cause of chemical combination.
The Ionic Bond
An ionic bond is the electrostatic force of attraction that holds the oppositely charged ions together in
an ionic compound. Such bonds are formed between metals and non-metals.
The metal atom loses the electron from valence shell to form a cation.
The non-metal atom gains the electron to form an anion.
The number of charges on the ion is equal to the number of electrons lost or gained by the atom.
Formation of Lithium Fluoride
.
..
.. _
+
Li + .F :
Li + : F :
LiF
..
..
.
..
.. _
2+
Mg + 2 . Cl :
Mg + 2 : Cl :
MgCl2
.
..
..
The Covalent bond
A covalent bond refers to the bond formed by the mutual sharing of electrons between the atoms.
Such compounds are covalent compounds.
Covalent bonds are formed between the atoms of non- metals.
The two electrons responsible for the covalent bond are known to be the shared pair of electrons or
the bond pair of electrons. In Lewis structure of covalent compounds, a line between the bonded atoms
represents each of the shared pair of electrons. The lone pair of electrons is represented as dots
surrounding the atom. The sharing of 2 electrons (one pair) forms a single covalent bond and the
sharing of 4 electrons (two pairs) forms a double covalent bond.
Properties of ionic and covalent compounds
Ionic compounds
Covalent compounds
Good conductors of electricity in the molten
State or when dissolved in water
Do not conduct electricity as no
ions are present
Possess high melting point due to strong
electrostatic forces of attraction
Low melting points because of
weak intermolecular forces
Exist as solids at ordinary conditions
Liquids, gases or usually low melting
Solids
Soluble in polar solvents like water and
insoluble in non-polar organic solvents
Insoluble in polar solvents but soluble
in non-polar organic solvents.
Bond Length: The distance between the nuclei of two bonded atoms in a molecule is the bond length.
Bond Energy
The energy required to break a bond into its constituent atoms is the bond energy.
Electro negativity
The ability of an atom to attract the shared pair of electrons in a covalent bond is electro negativity.
In the periodic table, electro negativity increases form left to right across the period and decreases
from top to bottom in a group.
Fluorine is the most electronegative of all the elements.
Lewis Structure of Covalent Compounds ( Electron Dot Formulas of Molecules)
Ex: NH3 (Ammonia)
1. Identify the central atom and the surrounding atoms and write their symbols.
2. Find the total number of valence electrons
Element
valence e- atoms per molecule number of e- total number of valence eN
H
5
1
1
3
5
3
8
3. Represent the shared pair of electrons by drawing lines between bonded atoms.
H
H_____N_____H
Each line (covalent bond) represents two electrons
4. Position the remaining valence electrons as lone pairs (electron pairs that are not
involved in bonding)
H
H_____N_____H
..
Resonance
Resonance is the phenomenon by which some covalent molecules are represented by two or more Lewis
structures, none of which is capable of representing the actual structure of the molecule.
Ex. Carbonate ion, Nitrate ion, Benzene
Exceptions to the octet Rule
 Incomplete Octet
In some covalent compounds the number of electrons surrounding the central atom
( valence shell ) is fewer than eight.
Example. Be in BeCl2 , B in BF3 and Al in AlCl3

Expanded Octet
In some other compounds the number of electrons surrounding the central atom exceeds the
normal octet (8 electrons). This is by making use of the d sub shells.
Elements of Group 3A and beyond exhibit such behavior.
Example. S in SF6 , P in PCl5 and I in IF7
Coordinate Covalent bond
A special type of covalent bond formed by the donation of a lone pair of e- by the central atom to the
other atoms bonded to it. Example: H2O2
Polar covalent Bond
A polar covalent bond bears a partial positive charge on one end and a partial negative charge on the
other end of the bond. I.e. the bond bears a dipole. Polarity of the covalent bond is due to the
difference in the electro negativity of the bonded atoms. The more electronegative atom pulls the bond
pair of electrons towards it and bears a partial negative charge.
H  ------ F 
An ionic bond is formed when the electro negativity difference between the bonded atoms is 2
or more. Less than this value the bond formed is a polar covalent bond.
A 100 % pure covalent bond forms when the bonded atoms are of the same element, or the electro
negativity difference between the bonded atoms is zero.
The attractive forces between atoms in a polar bond, is measured by a quantity called dipole moment.
It depends on the distance between the partial charges and the magnitudes of the charges.
 =Q x r
Q = magnitude of charge
r = distance between charges
Polar molecules contain one or more polar bonds.
There can be molecules that contain two or more polar bonds and still be non-polar.
Its overall shape may be such that the dipole moments cancel each other. Thus, the geometry of the
molecule decides if the molecule polar or not.
Rules to predict a molecule polar or non polar

Symmetric molecules containing same kind of atoms around the central atom are non polar, even
if they contain polar bonds.
BeF2 and CO2 molecules are non polar though they contain polar bonds because they are linear in
geometry. Their dipole moments cancel each other.

Symmetric molecules containing different kind atoms around the central atom are polar.
Ex. CHCl3,
Asymmetric molecules are polar if they contain polar bonds.
Ex. NH3, NF3, H2O, SO2 etc
Water because of its polarity and large dipole moment is an excellent solvent for polar substances.

VSEPR Theory (Valence Shell Electron Pair Repulsion Theory)
This theory explains the 3-Diamensional molecular structure in covalent compounds.
1. The central atom in a covalent molecule is surrounded by a definite number of electron pairs.
These electron pairs include bond pair and lone pair of electrons.
2. The electron pairs arrange themselves at the maximum possible distance around the central
atom to minimize repulsion. This arrangement gives a definite geometry to the molecule.
3. The shared e- pairs are called the bond pairs and the unshared e- pairs are called lone pairs.
4. The repulsion between the e- pairs is in the order lp-lp > lp-bp > bp-bp.
Geometry of molecules
2 pairs of e3 pairs of e4 pairs of e-
1800
1200
109.50
linear
trigonal planar
tetrahedral
Structure of common covalent molecules. (Before deciding the structure, find the number of valence
electrons as described above).
Methane
H
C
H
H
H
2. BeCl2
Cl _____Be_____Cl
109.50
Tetrahedral
1800
Linear
3. BF3
F
1200
B
F
F
Trigonal planer