Electronegativity is defined as a measure of the ability of an aotme in a molecule aof attracts the shared electron pair towards itself. For example, in a homonuclear diatomic molecule such as H₂ and CI₂. the shared electron pair is attracted equally by both the bonding atom and so there is no net polarity I the molecule. This is because the electronegativity of both the atoms is the same. However, in a heteronuclear molecule such as HCI, the shred electron pair is not equally attracted by the bonding atoms, i.e., hydrogen and chlorine. The electron pair is more towards chlorine than hydrogen and the molecule acquires polarity ( Hδ ⁺– CIδ ^- ). These changes are of equal magnitude but opposite sign, and the overall charge is zero. Such small charges are denoted by the symbol δ (delta). This implies that chlorine is more electronegative than hydrogen.

Electronegativity is a qualitative concept. There is no direct way of measuring this ability experimentally though a number of indirect methods have been suggested. Pauling was the first to given a general method of computing electronegativities of elements. He proposed that if two atoms A and B had the same electronegativity value, then the bond energy of the A-B bond would simply be the geometric mean of the A-A, and B-B bond energies, because all the three molecules (A-B, A-A and B-B) are bonded by pure covalent bonds. However, if the two atoms A and B have different electronegativities, then the energy of the A-B bon would be greater than the geometric means of the energies of A-A and B-B bonds. This excess energy is due to the contribution of ionic structure in addition to the covalent structure.

A-B, A⁺B^- (Say B is more electronegative than A )

Pauling, therefore, suggested that this excess bond energy in molecules (containing atoms of different electronegativities) can be used as a measure of the electronegativity differences.

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