The difference between the electronegativities of chlorine (EN = 3.16) and hydrogen (EN = 2.20) is large enough to assume that the bond in HCl is polar.

+		-
H		Cl

Because it contains only this one bond, the HCl molecule can also be described as polar.

The polarity of a molecule can be determined by measuring a quantity known as the dipole moment, which depends on two factors: (1) the magnitude of the separation of charge and (2) the distance between the negative and positive poles of the molecule. Dipole moments are reported is units of debye (d). The dipole moment for HCl is small: µ = 1.08 d. This can be understood by noting that the separation of charge in the HCl bond is relatively small (EN = 0.96) and that the H-Cl bond is relatively short.

C-Cl bonds (EN = 0.61) are not as polar as H-Cl bonds (EN = 0.96), but they are significantly longer. As a result, the dipole moment for CH₃Cl is about the same as HCl: µ = 1.01 d. At first glance, we might expect a similar dipole moment for carbon tetrachloride (CCl₄), which contains four polar C-Cl bonds. The dipole moment of CCl₄, however, is 0. This can be understood by considering the structure of CCl₄ shown in the figure below. The individual C-Cl bonds in this molecule are polar, but the four C-Cl dipoles cancel each other. Carbon tetrachloride therefore illustrates an important point: Not all molecules that contain polar bonds have a dipole moment.

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