Chemical Bonding

Ionic Bond
Covalent Bond
Co-ordinate Bond Or Dative Bond
VSEPR Theory
Molecular Orbital Theory
Hydrogen Bond

Electronic Theory of Valency

The chemical behaviour of an atom is determined to a large extent by the number and the arrangement of electrons in the outer orbitals of the atom. Only these electrons are involved in chemical combination and so these are called the valence electron .

Completed Electron Octet Or Duplet

Group zero of the periodic table contains the nobal gases. With the exception of helium which has a 1s2 electron arrangement others have ns2 np6 configuration in the outer orbitals.

He 1s2
Ne 1s2 2s2 2p6
Ar 1s2 2s2 2p6 3s2 3p6
Kr 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6
Xe 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p6

Since the atoms of the noble gases were not known to form chemical bond, it was argued that the presence of 8 electrons (an electron octet) in the valence shell makes the atom stable. Therefore all other atoms must undergo bonding by gaining or losing or sharing electrons so as to aquire the electronic configuration of the nearest inert gas .

Ionic Bond (Or Electrovalent Bond)

Ionic bond or Electrovalent bond is a bond formed by the complete transfer of electrons from one atom to another, so as to complete their valence shell with eight electrons (i.e. octet) or two electrons (i.e. duplet). Atoms of metals generally lose electrons and those of non-metals gain electrons.

Covalent Bond

A Covalent bond is formed by the mutual sharing of electrons between two atoms, each atom contributing one electron to the shared pair . A covalent bond is usually represented by a short line
(i.e, a dash) between the atoms.

Lewis Structure

A Lewis symbol is a symbol in which the electrons in the valence shell of an atom or simple ion are represented by dots placed around the letter symbol of the element. Each dot represents one electron .


  1s2 2s2 2p4

Co-ordinate Covalent Bond Or Dative Bond

We have seen that the formation of a covalent bond between two atoms, each atom contributes one electron to the shared pair. Sometimes both the electrons of the shared pair may come from one of the atoms. The covalent bond thus formed is called a co-ordinate bond or dative bond .

Polar Covalent Bonds

Covalent bonds in which the sharing of the electron pair is unequal , with the electrons spending more time around the more nonmetallic atom, are called polar covalent bonds . In such a bond there is a charge separation with one atom being slightly more positive and the other more negative, i.e., the bond will produce a dipole moment. The ability of an atom to attract electrons in the presense of another atom is a measurable property called electronegativity .


The phenomenon where a molecule can be represented by more than one structural formula is known as resonance .

Valence Shell Electron Pair Repulsion (VSEPR) Theory

VSEPR theory is used to predict the shapes of molecules. This theory predicts that bonding (sigma bonds only) and non-bonding electron pairs in a molecule will adopt a geometry in which the distance between the electron pairs is maximized from one another in order to minimize the repulsions. This will result in a molecular geometry with the lowest possible energy. The theory also allows us to predict which hybridization the central atom takes in bonding to other atoms.

For this, we count how many pairs of electrons (triple bonds are counted as one pair) are around the central atom. If there are two pairs of electrons, they must be positioned 180° apart from each other and the shape is therefore linear. Three pairs are best positioned 120° apart and the shape is thus trigonal planar. Four pairs of electrons are best positioned as tetrahedral shape. For five pairs of electrons, the shape is predicted to be trigonal bipyramidal. Last, the octahedral is the shape predicted for six pairs of electrons.


Electron pairs

Shapes with, and without non-bonding e- pair

Hybridization of central atom



linear, linear




trigonal planar, trigonal planar




tetrahedral, tetrahedral




tetrahedral, trigonal pyramidal




tetrahedral, bent




trigonal bipyramidal, trigonal bipyramidal




trigonal bipyramidal, T-shaped




trigonal bipyramidal, Seesaw




octahedral, octehedral




octahedral, square planar




trigonal bipyramidal, linear



Molecular Orbital Theory

In the molecular orbital approach, all of the electrons are present in the molecular orbitals. These molecular orbitals are formed by linear combination of atomic orbitals (LCAO) . Thus it is also known as LCAO -MO method.

The following are the essential features of the M.O Theory -

(i) In the M.O. model, all the electron are taken together and considered as moving in the field of all the nuclei.

(ii) The atomic orbitals are combined to form, which are called molecular orbitals and electrons are fed into these orbitals.

(iii) The number of combining atomic orbitals is equal to the number of molecular orbitals formed.

(iv) When two atomic orbitals combine, two M.O's are formed, of which one has a lower energy, while the other has a higher energy. The former one is known as the bonding orbital and the latter one is known as the antibonding orbital .


Hybridisation is the combination of certain number of atomic orbitals of slightly different energies to form the same number of new (hybrid) orbitals of equal energy.

Type of hydridization Shape of molecule Examples
sp Linear BeCl2, BeH2, C2H2
sp2 Triangular planar BF3, BCl3, C2H4
sp3 Tetrahedral CH4, NH4+, CCl4
dsp2 Square planar Ni(CN)4]2–, [PtCl4]2–
sp3d Trigonal bipyramidal PF5, PCl5
sp3d2 Octahedral SF6, [Co(NH3)6]2+

Hydrogen Bond

A hydrogen atom normally forms a single bond . In some compounds, however, the hydrogen atom may be located between two atoms acting as a bridge between them. Hydrogen atom is now involved in two bonds, one a normal covalent bond , the other a hydrogen bond. A hydrogen bond is always formed between two small, strongly electronegative atoms such as fluorine, oxygen and nitrogen.

There are two types of hydrogen bond :-

(i) Intermolecular Hydrogen Bond .

(ii) Intramolecular Hydrogen Bond .
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