Determining the Net Dipole
A molecule can contain polar bonds but still be non-polar overall if the bond dipoles cancel
out due to symmetry.
Non-Polar Molecules
Usually occur when there are no lone pairs on the central atom and all outer atoms
are identical.
Examples: CO₂ (Linear), BF₃ (Trigonal Planar), CH₄ (Tetrahedral), CCl₄
Polar Molecules
Usually occur when there are lone pairs (breaking symmetry) or different outer
atoms.
Examples: H₂O (Bent), NH₃ (Pyramidal), CH₃Cl
Quick Decision Tree
- Any polar bonds? If no → Non-Polar.
- All outer atoms identical + no lone pairs on central? If yes → Non-Polar (dipoles
cancel).
- Otherwise? → Polar (net dipole exists).
Think About It
CO₂ has two polar bonds (C=O), yet the molecule is non-polar. CHCl₃ has polar bonds
too — but it is polar. What's the key difference?
CO₂ is linear and
symmetrical — the two C=O dipoles point in exactly opposite directions and cancel. CHCl₃ is
tetrahedral with different atoms, so the dipoles don't cancel — there is a net
dipole moment towards the Cl atoms.