## Dipole Moments

When a covalent bond is formed between two identical atoms, e.g. H-H, Cl-Cl, etc., the two electrons forming the covalent bond may be regarded as being symmetrically disposed between the two atoms. The centres of gravity of the electrons and nuclei therefore coincide. With two dissimilar atoms the two electrons are no longer symmetrically disposed, because each atom has a different electronegativity, i.e., attraction for electrons. Chlorine has a much greater electronegativity than hydrogen that when chlorine and hydrogen combine to form covalent hydrogen chloride, the electrons forming the covalent bond are displaced towards the chlorine atom without any separation of the nuclei :
The hydrogen atoms will, therefore, be slightly positively charged, and the chlorine atom slightly negatively charged. Thus, owing to the greater electronegativity of the chlorine atom, the covalent bond in hydrogen chloride is characterized by the separation of small charges in the bond. A covalent bond such as this, in which one atom has a larger share of the electron-pair, is said to possess partial ionic character.

In analogy with a magnet, such a molecule is called a dipole, and the product of the electronic charge, e, and the distance d, between the charges (positive and negative centres) is called the dipole moment, μ ; i.e.,  μ= e✕d. e is the order of 10-10  e.s.u. ; d, 10-8 cm. Therefore μ is of the order of  10-18 e.s.u., and this unit is known as the Debye (D), in honour of Debye, who did a large amount of work on dipole moments.

The dipole moment is a vector quantity, and its direction is often indicated by an arrow parallel to the line joining the points of charge, and pointing towards the negative end.

e.g., H-Cl. The greater the value of the dipole moment, the greater is the polarity of the bond. The terms polar and non-polar are used to describe bonds, molecules and groups, and the reader is advised to make sure he appreciates how the terms are applied in each case under consideration.

Not only do polar bonds contribute to the dipole moment of a molecule, but so do long pairs, e.g., the lone pair in ammonia makes a large contribution to the dipole moment.

In the table given below gives some electronegativity values, and it will be seen that electronegativity increases from left to right and decreases from top to bottom of the periodic table. The values are not absolute values; they are relative values and this is a satisfactory scheme since their use involves differences in electronegativities. Furthermore, it should be noted that electronegativity deals with electron attraction within a molecule, whereas electron affinity deals with the attraction of an electron from outside the atom and has an absolute value.