A molecule can be stuffed with polar bonds and still be completely nonpolar. Whether the whole molecule is polar depends on two things: how polar the bonds are, and how they are arranged in space. This is where bond polarity and molecular shape finally meet.
📚 What you need to know
To decide if a molecule is polar you need bond polarityandmolecular shape.
Each polar bond pulls electron density in a particular direction. Those pulls are vectors — they have a size and a direction — so they can add together or cancel each other out.
Imagine four people pulling on ropes tied to a ring. If they pull equally hard in opposite directions, the ring doesn’t move. If one of them lets go, or one pulls harder, the ring shifts. Molecules work the same way.
Case 1: no polar bonds at all
If both atoms in every bond are the same element, there are no dipoles to begin with, so there is nothing to add up. H2, Cl2, O2 and N2 are all nonpolar. Simple.
Case 2: polar bonds that cancel
This is the case students get wrong. CO2 has two very polar C=O bonds — oxygen is much more electronegative than carbon. But the molecule is linear, so the two pulls point in exactly opposite directions and cancel perfectly. The result is a nonpolar molecule made of polar bonds.
Compare that with water. The O–H bonds are polar in the same way, but water is bent. The two pulls are both angled upwards towards the oxygen, so instead of cancelling they reinforce each other and give one overall dipole.
Same kind of polar bond in both, but the linear shape cancels the dipoles and the bent shape does not.
Same bond type, completely different answer — and the only thing that changed was the shape. That is why you should never try to judge polarity from the formula alone.
Symmetry is the deciding factor
The shortcut is to look for symmetry. A molecule is nonpolar when the polar bonds are spread out evenly around the central atom, so that every pull has an equal and opposite partner.
Two conditions have to be met for the dipoles to cancel:
All the outer atoms must be the same. Swap one for something different and the balance breaks.
There must be no lone pairs on the central atom, because a lone pair takes up a position but exerts no bond dipole — it leaves a gap in the symmetry.
CCl4 passes both tests: four identical Cl atoms in a perfect tetrahedron, no lone pairs. Nonpolar. CH3Cl fails the first test — three H atoms and one Cl — so the strong C–Cl pull has nothing to balance it.
Replacing just one chlorine with a hydrogen destroys the symmetry and turns a nonpolar molecule into a polar one.
And here is the lone pair effect on its own. BF3 and NH3 both have three bonds to a central atom, but BF3 is flat with no lone pair while NH3 is pyramidal because of one:
The lone pair on nitrogen means there is no bond dipole pointing upwards to balance the three pointing down.
🧩 Is this molecule polar? Four steps
Draw the Lewis formula and work out the shape using VSEPR.
Check the bonds. Are the atoms different elements? If every bond is between identical atoms, stop — it’s nonpolar.
Draw the dipole arrows, one per polar bond, pointing towards the more electronegative atom.
Add them up. Do they cancel? Symmetrical with identical outer atoms and no lone pairs → nonpolar. Anything else → polar.
Molecule
Shape
Polar bonds?
Do they cancel?
Molecule
Cl2
linear
No
—
Nonpolar
CO2
linear
Yes
Yes, symmetrical
Nonpolar
BF3
trigonal planar
Yes
Yes, symmetrical
Nonpolar
CCl4
tetrahedral
Yes
Yes, symmetrical
Nonpolar
H2O
bent
Yes
No — lone pairs
Polar
NH3
trigonal pyramidal
Yes
No — lone pair
Polar
CH3Cl
tetrahedral
Yes
No — different atoms
Polar
WORKED EXAMPLE
CCl4 and CHCl3 both contain polar C–Cl bonds. Explain why one is polar and the other is not.
Both are tetrahedral with four bonding pairs and no lone pairsCCl₄: all four outer atoms are identicalFour equal dipoles spread evenly cancel exactly.CCl₄ is nonpolarCHCl₃: one Cl has been swapped for an HThe C–H bond is barely polar, so the three C–Cl pulls are no longer balanced.CHCl₃ is polar
WORKED EXAMPLE
Deduce whether SO2 is polar.
Step 1 — shape: 2 bonding domains + 1 lone pair → bentStep 2 — bonds: O (3.4) is more electronegative than S (2.6), so S=O is polarStep 3 — the bent shape means both dipoles point the same general waySO₂ is polarCO₂ and SO₂ look similar on paper but behave completely differently — because of that one lone pair.
💡 Exam tip
Always name the shape in your answer. “The dipoles cancel” on its own won’t score — say why they cancel.
A good full answer has three parts: the bonds are polar because…, the shape is…, so the dipoles do/don’t cancel.
Molecular polarity is what drives solubility and intermolecular forces later on, so this is not a one-off topic.
⚠️ Common mix-up
Polar bonds ≠ polar molecule. CO2 and CCl4 are the classic proof.
Nonpolar molecule ≠ nonpolar bonds. The bonds in CCl4 are very polar indeed; they just cancel.
Don’t forget the lone pairs when judging symmetry. They occupy space but contribute no bond dipole.
A molecule with a net dipole is still electrically neutral overall — polar does not mean charged.
Up next: Giant Covalent Structures — what happens when covalent bonding doesn’t stop at a molecule but carries on through an entire crystal.
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