Molecular Dipole Moment (College Board AP® Chemistry): Study Guide

Molecular dipole moment

• Molecular dipole moments give information about the separation of charges in a molecule

• The dipole moment (μ) is a measure of the net polarity of a molecule and is measured in debyes (D)

  • It is the product of the charge, Q, at each end of the dipole and the distance, r, between the charges

• The greater the difference in electronegativity between two bonded atoms, the more polar the bond

• A molecule’s dipole moment also depends on its shape: in CO₂ and CCl₄ the bond dipoles cancel, so these molecules have no overall dipole moment

• For example, in chloromethane, the C—Cl bond has a dipole moment because chlorine is more electronegative than carbon

  • The chlorine atom develops a partial negative (δ^-) charge

  • The carbon atom develops a partial positive charge (δ⁺)

  • Therefore, the C—Cl bond is described as a polar bond

• The dipole moment in the molecule is represented using an arrow with a plus at the end

  • The plus end of the arrow represents the positive atom or end of the polar bond / dipole moment

  • The arrowhead represents the negative atom or end of the polar bond / dipole moment

• The signs of the partial charges explain how polar molecules interact: the δ^- end of one molecule attracts the δ⁺ end of a neighbor (dipole-dipole), a cation attracts the δ^- end and an anion attracts the δ⁺ end (ion-dipole)

• These attractions are strongest when opposite partial charges line up close together, so dipole-dipole and ion-dipole forces depend on how the molecules are oriented

Bond dipole in chloromethane

The C—Cl bond has a dipole moment because chlorine is more electronegative than carbon.