Question

Which of the following has diopole moment

Answer 1

Using the equation above, the dipole moment is calculated to be 1.85 D by multiplying the distance between the oxygen and hydrogen atoms by the charge difference between them and then finding the components of each that point in the direction of the net dipole moment (the angle of the molecule is 104.5˚).

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Answer 2

Answer

In the case of #Cl_2#, the 2 atoms are identical, so no polarization of the bond is possible, and the dipole moment is zero.

In every other case except #H_2S#, the polarization of charge associated with each bond is exactly cancelled by the other bonds, resulting in no net dipole moment.

For #CO_2#, each C-O bond is polarized (with oxygen taking on a partial negative charge, and carbon a positive charge). However, #CO_2# is a linear molecule, so the two C-O bonds are polarized in equal and opposite directions, and exactly cancel each other out. Therefore, #CO_2# has no dipole moment.

For #H_2S# the bonds are both polarized, but #H_2S# is a bent molecule, not linear, so the polarizations do not cancel, and #H_2S# has a net dipole moment.

For #BCl_3#, the geometry is an equilateral triangle of Cl atoms, with the boron atom in the center of the triangle. The polarization of the 3 B-Cl bonds exactly cancels out, so #BCl_3# has no dipole moment.

Similarly, the 4 C-Cl bonds in CCl4 are oriented to point at the vertices of a regular tetrahedron, and they cancel each other out exactly, so CCl4 has no dipole moment.