AP Chemistry Molecular Polarity

Telling polar from nonpolar molecules by shape and bond dipoles, with four worked examples and a diagram.

A bond can be polar while the molecule it sits in is not, and that gap is exactly where exam questions on polarity hide. Whether a molecule is polar decides which intermolecular forces it has, what it dissolves in, and how it boils, so it links the bonding in Unit 2 to the forces in Unit 3. The method below uses only two things: the shape and the direction of each bond's pull on electrons.

How to Decide Whether a Molecule Is Polar in Four Steps

Draw the Lewis structure and name the shape. Count the electron domains on the central atom and use the VSEPR chart; the Lewis structures guide covers the first half.
Decide whether each bond is polar. If the two atoms differ in electronegativity, the bond is polar, with its dipole pointing toward the more electronegative atom. Use periodic trends: electronegativity rises toward the upper right, excluding noble gases.
Add the bond dipoles as vectors. Equal dipoles in a symmetric arrangement cancel. Different atoms on the central atom, or a lone pair on it, break the symmetry.
Classify. No net dipole means nonpolar; a net dipole means polar.

A quick check for the common shapes: a central atom with no lone pairs and identical outer atoms gives a nonpolar molecule. Anything that breaks one of those conditions is usually polar, as the examples show.

Worked Example 1: Comparing CO₂ and H₂O

Both molecules contain polar bonds. Which are polar molecules?

CO₂: carbon has two electron domains (two double bonds) and no lone pairs, so the shape is linear. Oxygen is more electronegative than carbon, so each C=O dipole points toward oxygen. The two dipoles are equal and point in opposite directions, so they cancel: CO₂ is nonpolar.
H₂O: oxygen has four electron domains (two bonds, two lone pairs), so the shape is bent. Each O–H dipole points toward oxygen, and because the bonds are not in a straight line the dipoles add: H₂O is polar, with the negative end at the oxygen.

Worked Example 2: Comparing BF₃ and NH₃

Explain why BF₃ is nonpolar while NH₃ is polar.

BF₃: boron has three electron domains and no lone pairs, giving a trigonal planar shape. The three B–F dipoles are identical and spaced 120° apart, so they cancel completely: nonpolar.
NH₃: nitrogen has four electron domains (three bonds, one lone pair), giving a trigonal pyramidal shape. The N–H dipoles point toward nitrogen, and the lone pair makes the shape asymmetric, so they add up rather than cancel: polar.

Both examples have the same number of identical bonds on the central atom. What separates them is the lone pair, which is why you must look at the shape and not just the formula.

Worked Example 3: Comparing CCl₄ and CH₂Cl₂ With the Same Shape

Both molecules are tetrahedral around carbon. Classify each.

CCl₄: four identical C–Cl bonds, each polar, arranged symmetrically in a tetrahedron. The dipoles cancel: nonpolar.
CH₂Cl₂: the same tetrahedral shape, but two of the outer atoms are hydrogen and two are chlorine. The C–Cl bonds are much more polar than the C–H bonds, so the pulls are no longer equal in all four directions and a net dipole remains: polar.

The lesson is that identical geometry does not guarantee identical polarity. Replacing even one outer atom with a different one breaks the cancellation.

Worked Example 4: Using Polarity to Name the Strongest Intermolecular Force

Name the strongest intermolecular force in a pure sample of CO₂, H₂O, CH₂Cl₂, and CCl₄.

CO₂: nonpolar, so only London dispersion forces.
H₂O: polar, with H bonded to O, so hydrogen bonding.
CH₂Cl₂: polar with no H–F, H–O, or H–N bond, so dipole-dipole forces (plus dispersion).
CCl₄: nonpolar, so only London dispersion forces.

Having no dipole-dipole forces does not make CCl₄ a weakly bound liquid: its large, polarizable electron cloud gives strong dispersion forces, and it actually boils higher than the polar CH₂Cl₂ (about 77°C against 40°C). The reasoning behind that trade-off is worked through in London dispersion forces.

Common Molecular Polarity Mistakes

Assuming polar bonds mean a polar molecule. CO₂, BF₃, and CCl₄ all have polar bonds and are nonpolar.
Judging polarity from the formula alone. BF₃ and NH₃ look similar but differ in shape, and so in polarity.
Forgetting to draw the lone pairs. They set the shape, and the shape decides whether the dipoles cancel.
Treating C–H as polar. The C–H bond is effectively nonpolar, so hydrocarbons are nonpolar.
Pointing the dipole the wrong way. The arrow points toward the more electronegative atom, the partial negative end.

Once you can classify a molecule, a lot of other predictions follow, including what it dissolves in (like dissolves like) and whether it forms the bond type you expect. For help turning practice results into a score estimate, the AP Chem Score Calculator is a good next stop.

Frequently Asked Questions

What makes a molecule polar?

A molecule is polar when its bond dipoles do not cancel, leaving a net dipole: one side of the molecule is partially negative and the other partially positive. This needs both polar bonds and an asymmetric shape.

Can a molecule have polar bonds and still be nonpolar?

Yes. In CO2, BF3, and CCl4 each bond is polar, but the molecule's symmetric shape makes the bond dipoles cancel, so the molecule as a whole is nonpolar.

Is CO2 polar or nonpolar?

Nonpolar. The two C=O bonds are polar, but CO2 is linear, so the two bond dipoles point in opposite directions with equal size and cancel.

Why is water polar?

Water has two polar O-H bonds in a bent shape (two bonding domains and two lone pairs on oxygen), so the bond dipoles add up to a net dipole pointing toward the oxygen end.

Does a lone pair on the central atom make a molecule polar?

Usually, because it makes the shape asymmetric, as in NH3 and H2O. But a few molecules whose lone pairs are arranged symmetrically are nonpolar, so check the symmetry of the whole shape rather than counting lone pairs alone.

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