There are a few ways of determining the degree of polarity from an external view; for example, you could see if the molecule's orientation changes in the presence of polar molecules, or in an induced magnetic field, or you could take a diagrammatical approach and draw Lewis structures. However, speaking strictly in terms of the molecule's outward appearance, with no other information, the best indication is the size of its component atoms, and to a lesser extent, its bond angles.
We can be sure of a few general rules; primarily, that when two identical atoms are bonded to each other, there will be no polarity in that bond. This changes the moment we add different substituents to each molecule, or if we get fancy and start considering things like London dispersion forces, but in general this rule should hold. Second, we know that polarity will increase proportionally to the electronegativity difference between two atoms in a bond. We also know that atoms are more electronegative the further to the right and the higher up the periodic table they are, and that this same trend also corresponds to the atom's radius. We also know that almost all atoms have different sizes from any other atom.
Therefore, we can conclude that, as viewed from the outside, any constituent atom that is different in diameter from the others will probably be the site of a polar bond.
We can also assume that any molecule with a bent linear configuration stands a good chance of being polar, because this is one of the few configurations in which it is impossible for the electrical charges to be evenly distributed over the surface of the molecule.
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