Based on symmetry alone, we know that #H_2S# is the only one of these molecules that has a dipole moment.
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.
