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How does a dipole moment measure bond polarity?

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Diagram showing the formation of a polar covalent bond in a water molecule, with two hydrogen atoms bonding to one oxygen atom to form H2O, illustrating shared electrons and chemical reaction. — covalent bond In polar covalent bonds, such as that between hydrogen and oxygen atoms, the electrons are not transferred from one atom to the other as they are in an ionic bond. Instead, some outer electrons merely spend more time in the vicinity of the other atom. The effect of this orbital distortion is to induce regional net charges that hold the atoms together, such as in water molecules.

The dipole moment, denoted by $μ$ , is a quantitative measure of bond polarity that reflects the distribution of electrical charge between two atoms.12

When atoms of different elements form a covalent bond, their differing electronegativities cause the more electronegative atom to attract the shared electron pair more strongly. This creates partial positive and negative charges on the atoms, resulting in an electric dipole.12

The magnitude of this polarity is defined as the dipole moment, calculated as the product of the magnitude of the partial charges ( $q$ ) and the distance of their separation ( $d$ ):
$μ = q \times d$

A larger dipole moment indicates a greater degree of charge separation, signifying a more polar bond. If the bonded atoms are identical, the charges are uniform, the dipole moment is zero, and the bond is nonpolar. As the electronegativity difference between atoms increases, the dipolar character grows until the bond may be characterized as ionic.12