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.
Lewis Dot Symbols of elements. The nucleus and the inner electrons are represented by the symbol of the element and the valence electrons by equal number of dots surrounding the symbol.
.
H
. .
He
.
Li
. .
Be
.
: B. . C .
.
.
. N :
.
.
: O :
.
..
: F :
.
..
: Ne :
..
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 and 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.
.
.. .. _
Li + .F : Li
+
+ : F : LiF
.. ..
. .. .. _
Mg + 2 . Cl : Mg 2+ + 2 : Cl : MgCl
2
. .. ..
Lattice Energy of ionic compounds
This is the energy required to completely separate one mole of a solid ionic compound into its gaseous ions.
Born Haber Cycle
Born Haber Cycle helps to calculate the lattice energy of ionic compounds in the following way.
Let us take the example of LiF. The following steps are involved in the formation of LiF
(s)
Li
(s)
+ F
2(g)
LiF
(s)
H = -594.1KJ/mol
S
½ D
Li
(g)
F
(g)
IE
+
Li
(g)
+
EA
-
F
(g)
X
S - Sublimation energy, IE - Ionization energy, D - Dissociation energy, EA - Electron affinity
H = S + IE + ½ D + EA + X
Lattice energy = - X (KJ / mol)
Higher the lattice energy more stable is the ionic compound. Compounds containing ions with smaller size and high charge density possess high lattice energy.

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
Good conductors of electricity in the molten
Covalent compounds
State or when dissolved in water
Possess high melting point due to strong
Do not conduct electricity as no ions are present 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
Electro negativity
Insoluble in polar solvents but soluble in non-polar organic solvents.
The ability of an atom to attract the shared pair of electrons in a covalent bond is electro negativity.
Electro negativity increases form left to right across the period and decreases from top to bottom in a group in the periodic table.
Fluorine is the most electronegative of all the elements.
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.
Lewis Structure of Covalent Compounds
Ex: NH
3
(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 e -
N
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 electrons
Each line (covalent bond) represents two
4. Position the remaining valence electrons as lone pairs (electron pairs that are not involved in bonding)
H
H _____ N _____ H
Formal Charge
An atom’s formal charge is the electrical charge difference between the valence electrons in its isolated atom and the number of electrons assigned to that atom in a Lewis structure.
Formal Charge = # of valence electrons on free atom – ½ # of shared electrons - # of unshared electrons
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 BeCl
2
, B in BF
3
and Al in AlCl
3

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 SF
6
, P in PCl
5
and I in IF
7
Bond Enthalpy
Bond enthalpy is the enthalpy change associated with the breaking of a particular bond in one mole of the gaseous molecules.
H
2(g)
H
2
O
(g)
H
H
(g)
(g)
+ H(g)
+ OH
(g)
H
H
0
0
= 436.4 KJ/mol
= 502 KJ/mol
OH
(g)
For a chemical reaction,
H
(g)
+ O
(g)
H 0 = 427 KJ/mol
H 0 = Sum of the bond enthalpy of the reactants – Sum of the bond enthalpy of the products.