How to tell if a bond is polar or nonpolar
Three steps, and the whole method is one subtraction.
- Look up both electronegativities. Carbon is 2.55, oxygen is 3.44.
- Subtract the smaller from the larger. That difference is ΔEN. For a C=O bond, 3.44 − 2.55 = 0.89.
- Read the answer off the scale:
- ΔEN below about 0.4 → nonpolar covalent. The two atoms pull about equally, so the electrons sit near the middle.
- ΔEN about 0.4 to 1.7 → polar covalent. Unequal sharing, with real partial charges (δ+ and δ−) and a bond dipole.
- ΔEN about 1.7 or more, between a metal and a nonmetal → ionic. The electrons are effectively handed over rather than shared.
Two identical atoms always give ΔEN of 0, so a bond like H-H or Cl-Cl is always nonpolar. The tool above does the subtraction for any pair of atoms and draws the dipole arrow, and it is free to embed on a class site.
These cutoffs are approximate teaching guides, not sharp lines. Bonding is a smooth scale from pure covalent to fully ionic, and different textbooks put the boundaries in slightly different places.
ΔEN for 30 common bonds
Sorted from least polar to most polar, so the order itself answers “which of these is more polar?”.
| Bond | Electronegativities | ΔEN | Bond type | Partial charges |
|---|---|---|---|---|
| C-C every carbon chain | 2.55 − 2.55 | 0.00 | Nonpolar covalent | neither, they pull equally |
| Cl-Cl chlorine gas, Cl₂ | 3.16 − 3.16 | 0.00 | Nonpolar covalent | neither, they pull equally |
| H-H hydrogen gas, H₂ | 2.2 − 2.2 | 0.00 | Nonpolar covalent | neither, they pull equally |
| P-H phosphine, PH₃ | 2.2 − 2.19 | 0.01 | Nonpolar covalent | P is δ+, H is δ− |
| C-S carbon disulfide, CS₂ | 2.58 − 2.55 | 0.03 | Nonpolar covalent | C is δ+, S is δ− |
| C-I iodomethane, CH₃I | 2.66 − 2.55 | 0.11 | Nonpolar covalent | C is δ+, I is δ− |
| C-H every hydrocarbon, and fats and oils | 2.55 − 2.2 | 0.35 | Nonpolar covalent | H is δ+, C is δ− |
| S-H hydrogen sulfide, H₂S | 2.58 − 2.2 | 0.38 | Nonpolar covalent | H is δ+, S is δ− |
| N-O nitrogen dioxide, NO₂ | 3.44 − 3.04 | 0.40 | Polar covalent | N is δ+, O is δ− |
| C-Br bromoethane, C₂H₅Br | 2.96 − 2.55 | 0.41 | Polar covalent | C is δ+, Br is δ− |
| H-I hydrogen iodide, HI | 2.66 − 2.2 | 0.46 | Polar covalent | H is δ+, I is δ− |
| C-N amines, amino acids, proteins | 3.04 − 2.55 | 0.49 | Polar covalent | C is δ+, N is δ− |
| C-Cl chloroform, CHCl₃, and PVC | 3.16 − 2.55 | 0.61 | Polar covalent | C is δ+, Cl is δ− |
| H-Br hydrogen bromide, HBr | 2.96 − 2.2 | 0.76 | Polar covalent | H is δ+, Br is δ− |
| N-H ammonia, NH₃, and protein backbones | 3.04 − 2.2 | 0.84 | Polar covalent | H is δ+, N is δ− |
| S-O sulfur dioxide, SO₂ | 3.44 − 2.58 | 0.86 | Polar covalent | S is δ+, O is δ− |
| C-O carbon dioxide, alcohols, sugars | 3.44 − 2.55 | 0.89 | Polar covalent | C is δ+, O is δ− |
| H-Cl hydrochloric acid, HCl | 3.16 − 2.2 | 0.96 | Polar covalent | H is δ+, Cl is δ− |
| O-H water, alcohols, carboxylic acids | 3.44 − 2.2 | 1.24 | Polar covalent | H is δ+, O is δ− |
| C-F PTFE and refrigerants | 3.98 − 2.55 | 1.43 | Polar covalent | C is δ+, F is δ− |
| Si-O sand, quartz and glass | 3.44 − 1.9 | 1.54 | Polar covalent | Si is δ+, O is δ− |
| Al-Cl aluminium chloride, AlCl₃ | 3.16 − 1.61 | 1.55 | Polar covalent | Al is δ+, Cl is δ− |
| H-F hydrofluoric acid, HF | 3.98 − 2.2 | 1.78 | Polar covalent | H is δ+, F is δ− |
| B-F boron trifluoride, BF₃ | 3.98 − 2.04 | 1.94 | Polar covalent | B is δ+, F is δ− |
| Mg-O magnesium oxide, MgO | 3.44 − 1.31 | 2.13 | Ionic | Mg is δ+, O is δ− |
| Na-Cl table salt, NaCl | 3.16 − 0.93 | 2.23 | Ionic | Na is δ+, Cl is δ− |
| K-Cl potassium chloride, KCl | 3.16 − 0.82 | 2.34 | Ionic | K is δ+, Cl is δ− |
| Ca-O quicklime, CaO | 3.44 − 1 | 2.44 | Ionic | Ca is δ+, O is δ− |
| Li-F lithium fluoride, LiF | 3.98 − 0.98 | 3.00 | Ionic | Li is δ+, F is δ− |
| Na-F sodium fluoride, in toothpaste | 3.98 − 0.93 | 3.05 | Ionic | Na is δ+, F is δ− |
Two rows are worth pointing at. C-O is 0.89 and C-N is 0.49, which are the bonds behind most of organic chemistry. And C-H is 0.35, just under the line, which has consequences the whole of the next-but-one section is about.
What δ+ and δ− mean
The lowercase Greek letter δ (delta) means “a little bit of”. So δ− marks the atom that is slightly negative and δ+ the one that is slightly positive. They are fractions of a charge, not whole ions: in an H-Cl bond, chlorine has not taken hydrogen’s electron, it has just pulled the shared pair closer.
To mark up a bond, put δ− on the more electronegative atom and δ+ on the other:
δ+H-Clδ−
Then draw the dipole arrow along the bond, pointing from δ+ toward δ−, with a small cross on its tail at the δ+ end. The arrow points where the electrons have moved.
One thing that trips people up: the same Greek letter does two jobs here. Capital Δ in ΔEN means “the difference between”, the way it does everywhere else in science. Lowercase δ on an atom means “a partial charge”. Same letter, different size, different meaning.
Which bond is more polar?
Whichever has the bigger ΔEN. That is the whole rule, and the table above is sorted so you can read the comparison straight off it.
If you do not have the numbers to hand, use the periodic trend. Electronegativity rises going up and to the right, peaking at fluorine at 3.98 and bottoming out at the bottom left. So the further apart two elements sit along that diagonal, the more polar the bond between them. Fluorine bonded to a metal from group 1 is about as polar as chemistry gets; two neighbours in the same period barely differ at all. The periodic trends lesson covers where the values come from.
The C-H bond, and why oil and water do not mix
Carbon is 2.55 and hydrogen is 2.20, so C-H comes out at 0.35: below the 0.4 cutoff, and treated as nonpolar.
That one small number explains a lot. A hydrocarbon like petrol, candle wax or cooking oil is mostly C-C bonds (ΔEN 0) and C-H bonds (ΔEN 0.35), so it has almost no charge separation anywhere in it. Water is built from O-H bonds at ΔEN 1.24 and is strongly polar. Polar and nonpolar molecules do not attract each other enough to mix, which is “like dissolves like” in one line, and why oil floats in a separate layer.
Bond polarity is not molecular polarity
This is the distinction students lose marks on most often.
| Bond polarity | Molecular polarity | |
|---|---|---|
| What it describes | one bond | the whole molecule |
| What decides it | ΔEN between the two atoms | the bond dipoles and the shape |
| How you work it out | subtract two numbers | add the dipole arrows as vectors |
| Typical answer | nonpolar, polar, or ionic | polar or nonpolar |
A bond dipole has a direction, so bond dipoles are vectors. If a molecule’s shape points them in opposing directions, they cancel, and a molecule packed with very polar bonds comes out nonpolar overall.
Carbon dioxide is the classic case. Each C=O bond has ΔEN 0.89 and is genuinely polar, but CO₂ is linear, so the two dipoles point exactly opposite and cancel. CO₂ is nonpolar. Water has the same kind of bonds but oxygen’s two lone pairs bend it to about 104.5°, so its two O-H dipoles do not oppose each other. They add, and water is polar.
Is the whole molecule polar?
| Molecule | Shape | Largest ΔEN | Bonds | Molecule | Why |
|---|---|---|---|---|---|
| CO₂ Carbon dioxide | linear | 0.89 | polar | nonpolar | the shape is symmetric, so the bond dipoles cancel |
| H₂O Water | bent | 1.24 | polar | polar | the shape does not let the bond dipoles cancel |
| CH₄ Methane | tetrahedral | 0.35 | nonpolar | nonpolar | the bonds are barely polar to begin with |
| CCl₄ Carbon tetrachloride | tetrahedral | 0.61 | polar | nonpolar | the shape is symmetric, so the bond dipoles cancel |
| CHCl₃ Chloroform | tetrahedral | 0.61 | polar | polar | the shape does not let the bond dipoles cancel |
| NH₃ Ammonia | trigonal pyramidal | 0.84 | polar | polar | the shape does not let the bond dipoles cancel |
| NF₃ Nitrogen trifluoride | trigonal pyramidal | 0.94 | polar | polar | the shape does not let the bond dipoles cancel |
| BF₃ Boron trifluoride | trigonal planar | 1.94 | polar | nonpolar | the shape is symmetric, so the bond dipoles cancel |
| SO₂ Sulfur dioxide | bent | 0.86 | polar | polar | the shape does not let the bond dipoles cancel |
| SO₃ Sulfur trioxide | trigonal planar | 0.86 | polar | nonpolar | the shape is symmetric, so the bond dipoles cancel |
| CH₂O Formaldehyde | trigonal planar | 0.89 | polar | polar | the shape does not let the bond dipoles cancel |
| HCN Hydrogen cyanide | linear | 0.49 | polar | polar | the shape does not let the bond dipoles cancel |
| HCl Hydrogen chloride | linear | 0.96 | polar | polar | the shape does not let the bond dipoles cancel |
| HF Hydrogen fluoride | linear | 1.78 | polar | polar | the shape does not let the bond dipoles cancel |
| H₂S Hydrogen sulfide | bent | 0.38 | nonpolar | polar | the bonds are only weakly polar; the lone pairs give the dipole |
| PH₃ Phosphine | trigonal pyramidal | 0.01 | nonpolar | polar | the bonds are only weakly polar; the lone pairs give the dipole |
| SF₆ Sulfur hexafluoride | octahedral | 1.40 | polar | nonpolar | the shape is symmetric, so the bond dipoles cancel |
| XeF₄ Xenon tetrafluoride | square planar | 1.38 | polar | nonpolar | the shape is symmetric, so the bond dipoles cancel |
| SF₄ Sulfur tetrafluoride | seesaw | 1.40 | polar | polar | the shape does not let the bond dipoles cancel |
| O₂ Oxygen | linear | 0.00 | nonpolar | nonpolar | the bonds are barely polar to begin with |
| N₂ Nitrogen | linear | 0.00 | nonpolar | nonpolar | the bonds are barely polar to begin with |
| CS₂ Carbon disulfide | linear | 0.03 | nonpolar | nonpolar | the bonds are barely polar to begin with |
Reading down the “Why” column gives the four things that can happen, and two of them are worth spelling out because the usual shortcut misses them.
Methane is not the example people think it is. CH₄ is usually taught as “polar bonds cancelling by symmetry”, but its C-H bonds are at ΔEN 0.35 and are not really polar in the first place. It is nonpolar for a simpler reason than the story suggests. CCl₄ is the honest version of that example: C-Cl bonds are polar at 0.61, the tetrahedron is symmetric, and they cancel exactly.
Phosphine is polar with no polar bonds. Phosphorus is 2.19 and hydrogen is 2.20, so P-H sits at ΔEN 0.01, and yet PH₃ has a real dipole moment. It comes from the lone pair on the phosphorus, which is a region of concentrated negative charge in its own right. H₂S is the same story. Any rule that looks only at bonds and symmetry gets both of these molecules wrong.
Swap one atom and the answer flips: CHCl₃ has the same tetrahedral shape as CCl₄, but replacing one chlorine with a hydrogen breaks the symmetry, so its dipoles no longer cancel and chloroform is polar. Switch the tool above to Molecule mode to step through these, and see molecular geometry for where the shapes come from.
Why polarity matters
Polarity controls how substances behave. “Like dissolves like”: polar water dissolves polar and ionic substances such as salt and sugar, but not nonpolar oil. Polar molecules also attract each other more strongly, which raises boiling points, a large part of why water is a liquid at room temperature while nonpolar CO₂ is a gas. The strongest of those attractions, hydrogen bonding, is bond polarity taken to its extreme: an O-H or N-H bond leaves the hydrogen so exposed that it is pulled hard by a lone pair on a neighbouring molecule.