The big idea: Electronegativity is how strongly an atom pulls the shared (bonding) electrons towards itself. Across a covalent bond, the atoms rarely pull equally.
If one atom pulls harder, the shared pair sits closer to it. That atom becomes slightly negative (δ−) and the other slightly positive (δ+). The bond is then a polar covalent bond — it has a small dipole.
The bigger the difference in electronegativity (Δχ) between the two atoms, the more polar the bond.
Each O–H bond is polar (O is δ−, H is δ+). The molecule is bent, so the two bond dipoles do NOT cancel — water has a net dipole and is polar.
Interactive diagram
Explore the labelled diagram, charts and maps for this topic in full study mode.
δ+ and δ− — which way?: The more electronegative atom always becomes δ− (it gets a bigger share of the electrons); the less electronegative atom becomes δ+.
Use the data booklet electronegativity values: in O–H, oxygen (3.4) > hydrogen (2.2), so O is δ− and H is δ+.
Bonding is a spectrum, set by the electronegativity difference Δχ between the two bonded atoms. There is no sharp cut-off — these are guideline ranges used to describe a bond.
| Electronegativity difference, Δχ | Bond type | Example |
|---|---|---|
| 0 (same element) | pure (non-polar) covalent | Cl–Cl in Cl2 |
| small (roughly 0 – 1.8) | polar covalent (δ+ / δ−) | H–Cl, O–H |
| large (roughly > 1.8) | ionic (electron effectively transferred) | Na⁺Cl⁻ |
From non-polar to ionic: - Δχ = 0 → the electrons are shared equally → a pure covalent (non-polar) bond. - Small Δχ → unequal sharing → a polar covalent bond with δ+ and δ− ends. - Large Δχ → the more electronegative atom takes the electrons almost completely → an ionic bond.
So the same idea — the pull on the shared electrons — runs from covalent through polar to ionic.
Stop wasting time on topics you know
Our AI identifies your weak areas and focuses your study time where it matters. No more overstudying easy topics.
The step students miss: A molecule with polar bonds is not always a polar molecule. Each polar bond is a little arrow (a bond dipole) pointing toward the δ− atom. You must add the arrows up like vectors.
- If the bond dipoles cancel (because the molecule is symmetrical) → the molecule is non-polar. - If they do not cancel → there is a net dipole → the molecule is polar.
Non-polar: dipoles CANCEL
- Symmetrical shape (linear, trigonal planar, tetrahedral).
- Equal bond dipoles point in opposite directions and sum to zero.
- Examples: CO_{2} (linear), CCl_{4} (tetrahedral), BF_{3} (trigonal planar).
Polar: dipoles DON'T cancel
- Asymmetrical shape, usually caused by lone pairs (bent, pyramidal).
- Bond dipoles add to give a net dipole.
- Examples: H_{2}O (bent), NH_{3} (pyramidal), SO_{2} (bent).
Compare CO2 and H2O. Both have polar bonds, but their shapes differ, so only one is a polar molecule:
Each C=O bond is polar (O is δ−, C is δ+), but the molecule is linear and symmetrical, so the two equal dipoles point in opposite directions and cancel.
Interactive diagram
Explore the labelled diagram, charts and maps for this topic in full study mode.
Same polar bonds, opposite result: CO_{2} is linear (O=C=O). The two equal C=O dipoles point in exactly opposite directions, so they cancel → CO2 is non-polar.
H_{2}O is bent (the two lone pairs on oxygen push the H atoms down). The two O–H dipoles point partly the same way and add up → H2O has a net dipole → polar.
The bonds are equally polar; the shape is what makes the difference.
How this is tested: This idea shows up in three reliable ways across Paper 1A and Paper 2:
- Deduce δ+ / δ− on the atoms of a bond from data-booklet electronegativities. - Identify which molecule is polar from a list (the test of shape/symmetry). - Explain why a named molecule is (or is not) polar — a 2-mark answer that needs polar bonds and the shape/symmetry point about whether the dipoles cancel.
The classic trap: Don't stop at 'it has polar bonds'. CO2 has very polar bonds yet is non-polar because they cancel. To explain polarity you must talk about the shape and whether the dipoles cancel — that is the marking point.
IB-style question — partial charges in HCl (a)
Hydrogen chloride, HCl, has a polar bond. (a) Using electronegativity, deduce which atom is partially positive (δ+) and which is partially negative (δ−), and explain your answer. [2]
How to score the marks
- Mark 1 — compare electronegativities. Chlorine (χ ≈ 3.2) is more electronegative than hydrogen (χ ≈ 2.2), so it pulls the shared pair of electrons towards itself.
- Mark 2 — assign the charges. The chlorine therefore gains a bigger share of the electrons and is δ−, leaving the hydrogen δ+.
Final answer
Cl is δ− and H is δ+, because Cl is more electronegative and pulls the bonding electrons towards itself.
IB-style question — is SO₂ polar? (b)
(b) Sulfur dioxide, SO2, is a bent molecule. Explain whether SO2 is a polar molecule. [2]
How to score the marks
- Mark 1 — the bonds are polar. Oxygen is more electronegative than sulfur, so each S–O bond is polar (O is δ−, S is δ+) and has a bond dipole.
- Mark 2 — the shape. SO2 is bent (sulfur has a lone pair), so the two bond dipoles do not cancel — they add to a net dipole, making the molecule polar.
Final answer
SO2 is polar: the S–O bonds are polar and, because the molecule is bent, the bond dipoles do not cancel, leaving a net dipole.