Electronegativity sets the stage
Electronegativity measures how strongly an atom pulls on the electrons it shares in a bond. It’s a periodic trend: it generally increases up and to the right of the periodic table, peaking at fluorine. When two atoms bond, the one with the higher electronegativity wins more than its fair share of the electron pair. You can explore the underlying trend in the periodic trends lesson — here we use it to predict polarity. The tool above is free to embed on your own site or LMS.
From a tug-of-war to δ+ and δ−
Picture the shared pair of electrons as a rope in a tug-of-war. If both atoms pull equally — as in two oxygen atoms — the rope sits in the middle and the bond is nonpolar. If one atom pulls harder, the electron density shifts toward it. That atom takes on a small partial negative charge, written δ−, and its partner becomes δ+. We draw this imbalance as a dipole arrow that points from δ+ toward the more electronegative atom, with a small cross on its tail at the δ+ end. In an H–Cl bond, for example, chlorine is more electronegative than hydrogen, so chlorine is δ− and the arrow points its way.
Reading ΔEN: nonpolar, polar, or ionic
Subtract the two electronegativities to get ΔEN. As a rough teaching guide:
- ΔEN below ~0.4 → nonpolar covalent (electrons shared roughly equally)
- ΔEN ~0.4 to 1.7 → polar covalent (unequal sharing, real δ+/δ− charges)
- ΔEN ~1.7 or more → ionic (the electrons are essentially transferred)
These cutoffs are approximate, not sharp lines — bonding is a smooth scale from pure covalent to fully ionic, and chemists disagree on the exact boundaries. The type of bond formed connects directly to the types of chemical bonds you may have already met.
A polar bond is not the same as a polar molecule
Here’s the idea students miss most: a molecule can be packed with polar bonds and still be nonpolar overall. Each polar bond has its own little dipole arrow, and those arrows are vectors — they have direction. If the molecule’s shape is symmetric, the arrows point in opposing directions and cancel, leaving no net dipole. Carbon dioxide (CO₂) is the classic case: it’s linear, so its two C=O dipoles point exactly opposite and cancel, making CO₂ nonpolar despite very polar bonds.
Symmetry decides: CO₂ vs H₂O
Compare CO₂ with water (H₂O). Water has the same kind of polar bonds, but the two lone pairs on oxygen bend the molecule to about 104.5°. In that bent shape the two O–H dipoles no longer oppose each other — they add up to a net dipole, so water is polar. Whether dipoles cancel comes straight from molecular geometry: symmetric shapes like linear CO₂, tetrahedral CH₄, and octahedral SF₆ cancel; bent, pyramidal, or lopsided shapes usually don’t. Switch the tool to Molecule mode to test several at once.
Why polarity matters
Polarity controls how substances behave. “Like dissolves like”: polar water dissolves polar and ionic substances (salt, sugar) but not nonpolar oil. Polar molecules also attract each other more strongly, which raises boiling points — a big reason water is liquid at room temperature while nonpolar gases like CO₂ are not. The strongest of these attractions, hydrogen bonding, is just an extreme case of bond polarity at work.