To predict whether or not a polyatomic molecule will have a dipole moment (i.e, be a polar molecule), we note that molecules in which the central atom is symmetrically surrounded by identical atoms are nonpolar. For molecules of the general formula ABn, certain geometries - linear (AB2), trigonal planar (AB3), tetrahedral and square planar AB4, trigonal bipyramidal AB5, and octahedral (AB6) - will lead to nonpolar molecules no matter how polar the individual bonds are. In these cases, the individual bond dipoles will cancel each other out.
This species is polar - there are only two atoms in this molecule, and they differ in electronegativity, so by definition the molecule is polar.
Draw out the Lewis structure; there are two double bonds on the central C atom, and we count these as two electron pairs. The molecular geometry is linear (two pairs, both bonding) - this is the case of linear AB2, and the molecule is nonpolar.
Draw the Lewis structure - the electron pair geometry is trigonal planar (the S-O double bond counts as one pair.) The molecular geometry is trigonal planar - we therefore have an AB3molecule with a trigonal planar geometry. This species is nonpolar.
This molecule is of the general formula AB3; its molecular geometry is trigonal pyramidal (the only geometry for AB3 which leads to a nonpolar molecule is trigonal planar.) This species is polar.
Here we have an AB6-type molecule with an octahedral molecular geometry. It is nonpolar.
Draw the Lewis structure of this species; its electron pair geometry is octahedral, but its molecular geometry is square pyramidal (5 bonding/1 nonbonding.) The only geometry for formula AB5 which leads to a nonpolar molecule is trigonal bipyramidal; this species is therefore polar.