Monday, February 4, 2008


JEE Syllabus

Orbital overlap and covalent bond;
Hybridisation involving s, p and d orbitals only;
Orbital energy diagrams for homonuclear diatomic species;
Hydrogen bond;
Polarity in molecules, dipole moment (qualitative aspects only);
VSEPR model and shapes of molecules (linear, angular, triangular, square planar, pyramidal, square pyramidal, trigonal bipyramidal, tetrahedral and octahedral).


Covalent bonds are formed as a result of the sharing of one or more pairs of bonding electrons. Each atom donates half of the electrons to be shared.

This sharing of electrons is as a result of the electronegativity(electron attracting ability) of the two bonded atoms are either equal or the difference is no greater than 1.7.

Atomic Orbital Approach to Bonding: The basic premise of this theory is that bonds are formed when atoms get close enough so that atomic orbitals on the individual atoms will be able to overlap so that the three dimensional probability regions share a common volume.


In terms of chemistry we refer to the hybridisation of atomic orbitals to explain the change that seems to happen between the atomic orbitals in an uncombined atom and the orbitals used by the same atom when bonding.

The classic molecule to consider is methane CH4. In this molecule the bond angles indicate that the shape of the molecule is a perfect tetrahedron with bond angles of 109º 28' (approximately 109.5º)

It seems that the orbitals used for bonding in CH4 are arranged as far apart as possible suggesting that they have the same energy (degenerate). We know that the orbitals on the carbon atom do not have the same energy - the 2s orbital is of lower energy that the three 2p orbitals. Hybridisation is a model that allows us to combine the atomic orbitals and then produce four degenerate orbitals to be used for bonding.

Hydrogen Bonding

The attractive force which binds hydrogen atom of one molecule with electronegative atom (F,O or N) of another molecule is known as hydrogen bond or hydrogen bonding.


Bonds in which the electron density is symmertically distributed between the nuclei are called non-polar bonds, while those in which the electron density is unsymmetrically distributed are called polar bonds.


The bonding pairs and lone pairs around any particular atom in a molecule adopt positions in which their mutual interactions are minimized. The logic here is simple. Electron pairs are negatively charged and will get as far apart from each other as possible.

Shapes of Molecules

linear - 2 electron pairs

triangular,- 3 electron pairs

tetrahedral - 4 electron pairs

trigonal bipyramidal 5 bond pairs,

octahedral – 6 bond pairs.

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