Bond dipole moment
The bond dipole moment uses the idea of electric dipole moment to measure the polarity of a chemical bond within a molecule. It occurs whenever there is a separation of positive and negative charges. The bond dipole μ is given by:
The bond dipole is modeled as +δ — δ- with a distance d between the partial charges +δ and δ-. It is a vector, parallel to the bond axis, pointing from minus to plus, as is conventional for electric dipole moment vectors. (Some chemists draw the vector pointing from plus to minus, but only in situations where the direction is not important.) This vector can be physically interpreted as the movement undergone by electrons when the two atoms are placed a distance d apart and allowed to interact, the electrons will move from their free state positions to be localised more around the more electronegative atom.
The SI unit for electric dipole moment is the coulomb-meter. This is too large to be practical on the molecular scale. Bond dipole moments are commonly measured in debyes, represented by the symbol D, which is obtained by measuring the charge in units of 10−10statcoulomb and the distance d in Angstroms. Note that 10−10 statcoulomb is 0.208 units of elementary charge, so 1.0 debye results from an electron and a proton separated by 0.208 Angstrom. Another useful conversion factor is 1 C m = 2.9979×1029 D.
For diatomic molecules there is only one (single or multiple) bond so the bond dipole moment is the molecular dipole moment, with typical values in the range of 0 to 11 D. At one extreme, a symmetrical molecule such as chlorine, Cl
2, has zero dipole moment, while near the other extreme, gas phase potassium bromide, KBr, which is highly ionic, has a dipole moment of 10.5 D.
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