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The dipole moment of ethyl chloride is 2.10 D (J. Mol. Struct. 1990, 216, 9–26), and that of vinyl chloride is 1.42 D (J. Chem. Phys. 1960, 32 (1), 205–209).
Why is this the case? I would have thought that hyperconjugation would lead to vinyl chloride possessing a larger dipole moment instead.
I see two explanations for why the dipole in vinyl chloride is smaller than ethyl chloride.
Consider a resonance structure where the chlorine donates electrons to the alkene. This puts a formal positive charge on chlorine:
This would be a very minor contributor given the electronegativities and the mismatch in orbital size. However, this is not possible in ethyl chloride at all.
The more likely explanation is that an sp2-hybridized carbon is more electronegative than an sp3-hybridized carbon. Since the electronegativity of chlorine will be closer to the electronegativity of the carbons in vinyl chloride than the carbons of ethyl chloride, the C–Cl bond will be less polarized in vinyl chloride, giving a smaller dipole.
There are two factors:
The bonding orbital of carbon in $\ce{C2H3Cl}$ is an $\mathrm{sp^2}$ hybrid orbital while $\ce{C2H5Cl}$’s is an $\mathrm{sp^3}$ one. The former has a higher s-content and a lower p-content resulting in it being closer to the atom resulting in a shorter bond. A shorter bond means less charge separation.
$\mathrm{sp^2}$ carbons are also more electronegative than $\mathrm{sp^3}$ carbons; see the discussion on how to interpret the high acidity of phenol.
Don't forget that dipole moment depends not only on electronegativity differences but also bond lengths. All other factors held equal, the longer the bond, the greater the dipole moment.
Why would dipole moment depend on bond length? Consider an extreme example in which, say, sodium ion were in exactly the same place as fluoride ion. Obviously the two differ tremendously in electronegativities. However, there cannot be a dipole moment; there are no poles of opposite charge since the two are directly on top of each other.
