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 the other way around, pointing from plus to minus, but only in situations where the direction doesn't really matter.)
The SI unit for electric dipole moment is the coulomb-meter, but that is much too large to be practical on the molecular scale. Bond dipole moments are commonly measured in debyes, represented by the symbol D, which is what you get if you measure the charge in units of 10-10 statcoulomb and measure the distance d in Angstroms. Note that 10-10 statcoulomb is 0.48 units of elementary charge. Another useful conversion factor is 1 C m = 2.9979 D.
Typical dipole moments for simple diatomic molecules are in the range of 0 to 11D. At one extreme, a symmetrical molecule such as chlorine, Cl2, has zero dipole moment, while near the other extreme, gas phase potassium bromide, KBr, which is highly ionic, has a dipole moment of 10.5D.
For a complete molecule the total molecular dipole moment may be approximated as the vector sum of individual bond dipole moments. Often bond dipoles are obtained by the reverse process: a known total dipole of a molecule can be decomposed into bond dipoles. The reason for doing this is the transfer of bond dipole moments to molecules that have the same bonds, but for which the total dipole moment is not yet known. The vector sum of the transferred bond dipoles gives an estimate for the total (unknown) dipole of the molecule.
The Bond Dipole is two atoms in a bond, such that the electronegativity of one atom changes and draws the electrons towards the other, causing a partial negative charge. There is an increase difference in polarity, and an increase in dipole moment.