How to Tell If a Molecule Is Polar

“Is this molecule polar or non-polar?” sounds like a simple question, but it needs two separate pieces of information: whether the bonds are polar, and whether the molecule’s shape lets those bond polarities cancel. Getting one without the other is the most common source of wrong answers. This guide gives you a reliable four-step method, a few shortcuts, and twelve worked examples.

Why it matters

A polar molecule has an overall separation of charge — one end slightly positive, the other slightly negative — and so a dipole moment (see dipole moments). Polarity decides:

• Solubility: polar substances dissolve in polar solvents such as water; non-polar substances dissolve in non-polar solvents such as hexane (“like dissolves like”).
• Boiling points: polar molecules attract each other through dipole–dipole forces, and often hydrogen bonds (see intermolecular forces).
• Behaviour in electric fields and in microwave ovens.

The four-step method

Step 1: Are any bonds polar?

Look up the electronegativities (χ) of the bonded atoms (see bond polarity):

• Δχ below about 0.4 → essentially non-polar bond (e.g. C–H, 0.35).
• Δχ above about 0.4 → polar bond (e.g. O–H, C–Cl, C=O).

If no bond is polar, the molecule is non-polar. Stop here. (Examples: H₂, Cl₂, and — to a good approximation — hydrocarbons.)

Step 2: What is the shape?

Draw the Lewis structure and use VSEPR to find the molecular shape (see VSEPR and molecular geometry). Include lone pairs on the central atom — they’re essential.

Step 3: Do the bond dipoles cancel?

Treat each polar bond as an arrow pointing towards the more electronegative atom. Then ask: do the arrows cancel out by symmetry?

• They cancel if the molecule has a symmetrical shape and identical atoms in all outer positions. Then the molecule is non-polar.
• They don’t cancel if the shape is unsymmetrical or the outer atoms differ. Then the molecule is polar.

Step 4: Check for lone pairs on the central atom

Lone pairs on the central atom usually make a molecule polar, because they create an unsymmetrical distribution of electrons. The exceptions are shapes where the lone pairs themselves sit symmetrically — linear XeF₂ and square planar XeF₄, for example — which are non-polar.

Shortcut table

Shape (with identical outer atoms)	Lone pairs on centre	Polar?
Linear (e.g. CO₂)	0	No
Trigonal planar (e.g. BF₃)	0	No
Tetrahedral (e.g. CCl₄)	0	No
Trigonal bipyramidal (e.g. PCl₅)	0	No
Octahedral (e.g. SF₆)	0	No
Linear from 3 lone pairs (XeF₂)	3	No
Square planar (XeF₄)	2	No
Bent (H₂O, SO₂)	1 or 2	Yes
Trigonal pyramidal (NH₃, PCl₃)	1	Yes
Seesaw (SF₄)	1	Yes
T-shaped (ClF₃)	2	Yes
Square pyramidal (BrF₅)	1	Yes

And the golden rule: if the outer atoms are not all the same, a symmetrical shape usually becomes polar.

Worked examples

1. CO₂

• Bonds: C=O is polar (Δχ = 0.89).
• Shape: linear.
• Two identical dipoles point in opposite directions and cancel.
• Non-polar.

2. H₂O

• Bonds: O–H polar (Δχ = 1.24).
• Shape: bent (two lone pairs on O).
• Dipoles don’t cancel.
• Polar (see the shape of a water molecule).

3. BF₃

• Bonds: B–F very polar (Δχ = 1.94).
• Shape: trigonal planar, three identical F atoms at 120°.
• Dipoles cancel.
• Non-polar.

4. NH₃

• Bonds: N–H polar.
• Shape: trigonal pyramidal (lone pair on N).
• Polar.

5. CH₄

• Bonds: C–H barely polar; shape tetrahedral; identical atoms.
• Non-polar.

6. CH₃Cl (chloromethane)

• Bonds: C–Cl polar; C–H nearly non-polar.
• Shape: tetrahedral, but the outer atoms are not identical.
• The C–Cl dipole isn’t balanced.
• Polar (1.87 D).

7. CCl₄

• Bonds: C–Cl polar.
• Shape: tetrahedral, four identical Cl atoms.
• Dipoles cancel.
• Non-polar.

8. CH₂Cl₂ (dichloromethane)

• Tetrahedral, but with two H and two Cl. Even though it looks symmetrical on paper, in 3D the two C–Cl dipoles point partly in the same direction.
• Polar (about 1.6 D). Tip: flat drawings of tetrahedral molecules can mislead — think in 3D.

9. SO₂

• Bonds: S=O polar.
• Shape: bent (one lone pair on S).
• Polar. Compare CO₂, which is linear and non-polar — the lone pair on sulfur makes the difference.

10. PCl₅

• Trigonal bipyramidal, five identical Cl atoms, no lone pairs.
• Non-polar.

11. SF₄

• Seesaw (one lone pair on S).
• Polar.

12. XeF₄

• Square planar: four F atoms in a square, two lone pairs opposite each other above and below.
• Everything cancels.
• Non-polar.

Isomers can differ

Molecules with the same formula can differ in polarity. cis-1,2-Dichloroethene (both Cl on the same side of the C=C) is polar; trans-1,2-dichloroethene (opposite sides) is non-polar, because its C–Cl dipoles point in opposite directions. This affects their boiling points: cis about 60 °C, trans about 48 °C.

Quick checks

• Diatomic molecules: polar if the two atoms differ (HCl, CO); non-polar if they’re the same (O₂, N₂).
• Hydrocarbons: effectively non-polar — which is why petrol and oils don’t mix with water.
• Molecules with O–H or N–H bonds: almost always polar (and able to hydrogen-bond).
• Ionic compounds aren’t described as polar or non-polar molecules — they’re made of ions.

Why polarity changes physical properties

It’s worth connecting the yes-or-no answer to real behaviour, because that’s how polarity questions are often extended in exams.

Compare tetrachloromethane (CCl₄) and trichloromethane (CHCl₃). They’re almost the same size, and both contain polar C–Cl bonds. But CCl₄ is non-polar, because its four identical bond dipoles cancel in a tetrahedron, while CHCl₃ is polar, because the C–H bond breaks the symmetry. The polar molecules of CHCl₃ attract each other through permanent dipole–dipole forces as well as London forces.

You might expect CHCl₃ to boil higher. In fact, CCl₄ boils at about 77 °C and CHCl₃ at about 61 °C. CCl₄ has one more chlorine atom, so more electrons and stronger London forces, and in this case that outweighs the dipole–dipole forces. The lesson is that polarity is one factor among several: always consider molecular size as well (see intermolecular forces).

Polarity matters more clearly for solubility. Polar molecules such as ethanol and propanone mix with water in any proportion. Non-polar molecules such as hexane and CCl₄ form a separate layer.

A practical test

Run a thin stream of the liquid from a burette past a charged rod (rubbed with a cloth). Polar liquids such as water and ethanol are clearly deflected towards the rod; non-polar liquids such as hexane barely move. (Organic solvents are flammable and some are harmful, so this is a supervised lab demonstration, not a home experiment.)

Common mistakes

• “It has polar bonds, so it’s polar.” Symmetry can cancel them (CO₂, CCl₄, BF₃).
• Forgetting lone pairs on the central atom when finding the shape (SO₂, NH₃).
• Judging symmetry from a flat drawing of a 3D molecule (CH₂Cl₂).
• Ignoring different outer atoms (CH₃Cl vs CCl₄).

Key takeaways

1. Check whether any bonds are polar (Δχ above about 0.4).
2. Find the shape with VSEPR, including lone pairs.
3. If the shape is symmetrical and all outer atoms are identical, the dipoles cancel → non-polar.
4. Otherwise → polar.

For the underlying theory, see polar vs non-polar molecules.