The lattice enthalpy of an ionic solid is defined as the energy required to completely separate
one mole of a solid ionic compound into gaseous constituent ions.
When anion and cation approach each other, the valence shell of anion is
pulled towards cation nucleus and thus shape of anion is deformed. This
phenomenon of deformation of anion by a cation is known as polarisation and
the ability of cation to polarize a near by anion is called as polarizing power of
The Lewis-Langmuir theory can be understood by considering the formation of the chlorine molecule, Cl2. The
Cl atom with electronic configuration, [Ne]10 3s2 3p5, is one electron short of the argon configuration. The formation
of the Cl2 molecule can be understood in terms of the sharing of a pair of electrons between the two chlorine
atoms, each chlorine atom contributing one electron to the shared pair. In the process both
The bond formed between two atom in which contribution of an electron pair is made by one of them while the
sharing is done by both.
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. It is expressed as
As we know that Lewis approach helps in writing the structure of molecules but it fails to explain the formation
of chemical bond. It also does not give any reason for the difference in bond dissociation enthalpies and bond
lengths in molecules like H2 (435.8 kJ mol –, 74 pm) and F2 (150.6 kJ mol– , 42 pm); although in both the cases
a single covalent bond is formed by the sharing of an electron pair between the respective atoms. It also gives no
idea about the shapes of polyatomic molecules. Similarly the VSEPR theory gives the geometry of simple
molecules but theoretically, it does not explain them and also it has limited applications.
When two atoms come close to each other there is overlapping of atomic orbitals. This overlap may be positive,
negative or zero depending upon the properties of overlapping of atomic orbitals. The various
arrangements of s and p orbitals resulting in positive, negative and zero overlap are depicted in the following
This type of covalent bond is formed by the end to end (hand-on) overlap of bonding orbitals along the internuclear
axis. This is called as head on overlap or axial overlap. This can be formed by any one of the following types of
combinations of atomic orbitals.
In the formation of bond the atomic orbitals overlap in such a way that their axes remain parallel to each other
and perpendicular to the internuclear axis. The orbitals formed due to sidewise overlapping consists of two
saucer type charged clouds above and below the plane of the participating atoms
Lewis concept is unable to explain the shapes of molecules. This theory provides a simple procedure to predict
the shapes of covalent molecules. Sidgwick and Powell in 1940, proposed a simple theory based on the repulsive
interactions of the electron pairs in the valence shell of the atoms. It was further developed and redefined by
Nyholm and Gillespie (1957).
Shape (molecular geometry) of Some Simple Molecules / ions with central atom / ion having
no Lone Pairs of Electrons (E).
Shape molecule with Lone Pairs
Shapes of Molecules containing Bond Pair and Lone Pair
1. The number of hybrid orbitals is equal to the number of the atomic orbitals that get hybridised.
2. The hybridised orbitals are always equivalent in energy and shape.
Steric No. of an atom = number of atom bonded with that atom + number of lone pair(s) left on that atom.
This type of hybridisation can be explained by taking the example of CH4 molecule in which there is mixing of
one s-orbital and three p-orbitals of the valence shell to form four sp3 hybrid orbital of equivalent energies and
shape. There is 25% s-character and 75% p-character in each sp3 hybrid orbital. The four sp3 hybrid orbitals so
formed are directed towards the four corners of the tetrahedron. The angle between sp3 hybrid orbital is 109.5° as
shown in figure.
The molecular orbital theory was developed by F. Hund and R.S. Mulliken in 1932. The salient features are:
The energy levels of molecular orbitals have been determined experimentally from spectroscopic data for
homonuclear diatomic molecules of second row elements of the periodic table. The increasing order of energies
of various molecular orbitals for O2 and F2 is given below
In reality no bond or a compound is either completely covalent or ionic. Even in case of covalent bond between
two hydrogen atoms, there is some ionic character
It is often observed that a single Lewis structure is inadequate for the representation of a molecule in conformity
with its experimentally determined parameters. For example, the ozone, O3 molecule can be equally represented
by the structures I and II shown below
Nitrogen, oxygen and fluorine are the highly electronegative elements. When they are tied to a hydrogen atom
to form covalent bond , the electrons of the covalent bond are shifted towards the more electronegative atom.
This partially positively charged hydrogen atom forms a bond with the other more electronegative atom. This
bond is called as hydrogen bond and is weaker than covalent bond. For example, in HF molecule, the hydrogen
bond exists between hydrogen atom of one molecule and fluorine atom of another molecule as given below
This type of H-bonding occurs when polar H and electronegative atom are present in the same molecule i.e., it
is formed when hydrogen atom is present in between the two highly electronegative (F, O, N) atoms within the
Exists between the negative and positive ends of different molecules of the same or different substances i.e., it
is formed between two different molecules of the same or different compounds
Van der Waal’s Forces
Most metals crystallise in close-packed structures. The ability of metals to conduct electricity and heat must
result from strong electrons interactions among 8 to 12 nearest neighbours (which is also called coordination
number). Bonding in metals is called metallic bonding. It results from the electrical attractions among positively
charged metal ions and mobile, delocalised electrons belonging to the crystal as a whole.
Two models are considered to explain metallic bonding
Back bonding generally takes place when out of two bonded atoms one of the atom has vacant orbitals
(generally this atom is from second or third period) and the other bonded atom is having some non-bonded
electron pair(generally this atom is from the second period). Back bonding increases the bond strength and
decreases the bond length. For example, in BF3 the boron atom completes its octet by accepting two 2pelectrons
of fluorine into 2p empty orbital.