Bond Polarity and the Dipole Moment

If the molecule has several polar bonds, the average polarity corresponds to the vector sum of bond dipoles called the dipole moment and labeled by the symbol p. 

The dipole moment is defined as the product of charges and distance between centers of the same charges. The unit for measuring the dipole moment is debye (D), in the honor of Peter Debye, the scientist who made the greatest contribution to research on the polarity of molecules. The molecules with nonzero dipole moments are called polar molecules. 

For the example, if the proton and electron are separated by 100 pm the dipole moment is 4.80 D. This is the limiting value which can be used as a standard for the chemical bond with 100 % ionic character [(1.60 x 10 C m) (lD/3.336 X 10-30 C m) = 4.80 D]. The C-Cl bond-length is 178 pm and the dipole moment is 1.87 D. If we assume that the C-Cl bond is 100 % ionic, the expected dipole moment would be / = (178/100) (4.80 D) = 8.54 D. Since the experimental dipole moment is much smaller, 1.87 D, the character of the C-Cl bond is only 22 % ionic (% of ionic character = (1.87/8.54) x 100 = 22 %). In this way the polarity can be related to the ionic character of the chemical bond. 

Since the dipole moment is the vector sum of the polarities of the chemical bonds, it depends on the molecular geometry. Consequently, molecules such as CO2 or CCI4, with symmetrically distributed polar bonds, are nonpolar because their total dipole moment equals zero. 

As already mentioned, chlorides of methane are very useful solvents. From the knowledge about the polarity of molecules, i.e. their dipole moments, we can distinguish between polar and nonpolar solvents. The polar solvents are for example CH2CI2 and CHCI3, while tetrachloromethane, CCI4, is a nonpolar solvent. Polarity of solvents is one of the most important properties, not only for practical laboratory applications, but also for the theory of reaction mechanisms, reactivity and selectivity of organic compounds. 

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