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Picture two people holding a skipping rope between them. If they are equally strong, the middle of the rope stays in the middle. If one of them is stronger, the rope drifts towards that person, even though both are still holding it. Nobody has let go, and nobody has snatched the rope away. It is simply shared unevenly.
That is almost exactly what happens in a polar covalent bond. Two atoms share a pair of electrons, but one atom attracts that pair more strongly than the other. The electrons are still shared, so the bond is still covalent, yet they spend more of their time closer to one atom. The result is a bond with a slightly negative end and a slightly positive end.
This article builds up the idea from scratch: what “sharing” means, why the sharing can be unfair, how chemists label it, and why it matters for everything from the shape of a raindrop to the way salt dissolves.
First, what is being shared?
A covalent bond forms when two atoms each contribute an electron to a pair that sits between their nuclei. Both nuclei are positively charged, and both attract that negative pair. The attraction of two nuclei to the same pair of electrons is what holds the atoms together. (If you want a refresher on this, see covalent bonding explained.)
When the two atoms are identical, as in H₂ or Cl₂, the two nuclei pull with exactly the same strength. The electron pair is shared perfectly evenly. We call this a non-polar covalent bond.
When the two atoms are different, there is no reason to expect them to pull equally. One nucleus might have more protons, or its outer electrons might be closer to it, or both. Then the shared pair is pulled towards the “stronger” atom.
Electronegativity: a number for pulling power
Chemists measure how strongly an atom attracts a shared pair with a quantity called electronegativity. The most widely used scale is Pauling’s, where values run from below 1 for metals such as caesium up to 3.98 for fluorine.
A few values worth knowing (Pauling scale):
| Element | Electronegativity |
|---|---|
| Fluorine (F) | 3.98 |
| Oxygen (O) | 3.44 |
| Chlorine (Cl) | 3.16 |
| Nitrogen (N) | 3.04 |
| Carbon (C) | 2.55 |
| Hydrogen (H) | 2.20 |
| Sodium (Na) | 0.93 |
The pattern is easy to remember. Electronegativity rises as you move right across a period, because each step adds a proton to the nucleus while the outer electrons stay in the same shell. It falls as you move down a group, because the outer electrons sit further from the nucleus and are shielded by more inner shells. That puts the champion, fluorine, in the top right corner (ignoring the noble gases). For the full story of the trend, see the electronegativity trend.
In a bond between two different atoms, the one with the higher electronegativity wins the tug-of-war.
Partial charges: δ+ and δ−
Because the shared electrons lean towards one atom, that atom carries a little more negative charge than it would on its own. The other atom is left a little short of negative charge, so it ends up slightly positive. These are partial charges: fractions of the charge on an electron, not whole units.
Chemists write them with the Greek letter delta:
- δ− (“delta minus”) on the more electronegative atom
- δ+ (“delta plus”) on the less electronegative atom
For hydrogen chloride:
δ+ H–Cl δ−
Chlorine (3.16) is more electronegative than hydrogen (2.20), so chlorine gets δ−.
For a single O–H bond in water:
δ+ H–O δ−
Oxygen (3.44) beats hydrogen (2.20) by a wide margin, so the O–H bond is strongly polar. Water has two of these bonds, so the oxygen carries a δ− charge and each hydrogen carries δ+.
Some textbooks draw an arrow along the bond instead. The arrowhead points to the negative end, and a small cross on the tail marks the positive end, so the symbol looks like a plus sign that has grown an arrow.
How big is the difference? A rough guide
The size of the electronegativity difference, written Δχ, tells you how uneven the sharing is:
| Δχ | Usual description | Examples |
|---|---|---|
| 0 to about 0.4 | Non-polar (or very slightly polar) | H–H (0), C–H (0.35) |
| about 0.4 to 1.7 | Polar covalent | H–Cl (0.96), O–H (1.24), C–O (0.89) |
| above about 1.7 | Mostly ionic | Na–Cl (2.23) |
Treat the boundaries as signposts, not walls. A bond with Δχ = 1.6 is not dramatically different from one with Δχ = 1.8. Bonding changes gradually, and the article on the bonding triangle shows how ionic, covalent and metallic bonding blend into one another.
Let us work out a few bonds to get a feel for it.
C–H: 2.55 − 2.20 = 0.35. Almost equal sharing. This is why hydrocarbons such as methane and oils behave as non-polar substances.
N–H: 3.04 − 2.20 = 0.84. Clearly polar, with δ− on nitrogen. This polarity is part of the reason ammonia dissolves so well in water.
C–Cl: 3.16 − 2.55 = 0.61. Polar, with δ− on chlorine.
H–F: 3.98 − 2.20 = 1.78. Very polar. It sits right at the fuzzy edge between “polar covalent” and “ionic”, yet hydrogen fluoride is a molecular substance that boils at about 20 °C, not a high-melting ionic solid. That is a good reminder that the numbers guide you but do not replace looking at real behaviour.
Why polar bonds matter
Partial charges might sound like a small detail, but they change how substances behave.
Molecules attract each other more strongly. The δ+ end of one molecule is drawn to the δ− end of its neighbour. These attractions, called dipole–dipole forces, raise boiling points. The most extreme case, where hydrogen is bonded to N, O or F, gives hydrogen bonding, which is why water is a liquid at room temperature while methane, a molecule of similar size, boils at around −160 °C.
Polar substances dissolve polar and ionic substances. Water’s δ− oxygen atoms cluster around positive ions, and its δ+ hydrogens cluster around negative ions. That is how water pulls sodium chloride apart. Oil, with its nearly non-polar C–H bonds, cannot do this.
Polar bonds are where reactions happen. A δ+ carbon atom attached to oxygen or a halogen is a natural target for electron-rich particles. Much of organic chemistry can be predicted by spotting which atoms are δ+ and which are δ−.
An everyday analogy that goes a step further
Think of electron density as a cloud of fog between two hills. In a non-polar bond the fog sits evenly in the valley. In a polar bond the fog drifts up one slope. The fog has not left the valley, which is the point: the electrons are still shared. If the fog blew entirely over one hill, that would be an ionic bond, where an electron has effectively moved from one atom to the other.
So a polar covalent bond is the halfway house. It has some of the character of an ionic bond (charge separation) while still being a shared pair.
Common mistakes
“Polar bond” and “polar molecule” are not the same thing. Carbon dioxide, O=C=O, has two strongly polar C=O bonds, but they point in exactly opposite directions and cancel out. The molecule as a whole is non-polar. Water has two polar O–H bonds too, but the molecule is bent, so they do not cancel and water is polar. Always look at the shape. The guide to polar and non-polar molecules walks through this.
Putting δ+ and δ− the wrong way round. The δ− always goes on the atom with the higher electronegativity, the one that attracts the electrons. A useful self-check: the δ− atom is the one closer to fluorine on the periodic table.
Thinking partial charges are whole charges. δ− on oxygen does not mean oxygen has gained an electron. It means oxygen has a larger share of the bonding pair. An O²⁻ ion is a completely different thing.
Treating the 1.7 cut-off as a law. The boundaries between non-polar, polar and ionic are conventions. Different textbooks use slightly different numbers, and some bonds on the “ionic” side behave covalently in practice (H–F is the classic example).
Assuming C–H bonds are polar because the atoms are different. Different atoms always give some polarity, but a Δχ of 0.35 is so small that C–H bonds are normally treated as non-polar.
Forgetting that electronegativity is about bonded atoms. It describes how an atom behaves in a bond. It is not the same as electron affinity, which is a measured energy change for an isolated gaseous atom gaining an electron.
Try it yourself
Use the table above to decide which atom is δ− and whether each bond is non-polar, polar or mostly ionic.
- C=O (2.55 and 3.44)
- N–Cl (3.04 and 3.16)
- Na–O (0.93 and 3.44)
Answers: (1) Δχ = 0.89, polar, δ− on oxygen. (2) Δχ = 0.12, very nearly non-polar, a tiny δ− on chlorine. (3) Δχ = 2.51, mostly ionic, oxygen carries the negative charge.
If you enjoyed that, the next step is putting a number on “how ionic” a bond is: see estimating percent ionic character, or read bond polarity for more practice with Δχ.
Key takeaways
- A polar covalent bond is a shared electron pair that sits closer to one atom than the other.
- The atom with the higher electronegativity attracts the pair more strongly and carries a partial negative charge, δ−; the other atom carries δ+.
- The electronegativity difference, Δχ, measures how uneven the sharing is: roughly 0–0.4 non-polar, 0.4–1.7 polar, above 1.7 mostly ionic, with no sharp boundaries.
- Polar bonds lead to stronger attractions between molecules, higher boiling points and the ability to dissolve ionic and polar substances.
- A molecule with polar bonds can still be non-polar overall if its shape makes the bond polarities cancel, as in CO₂.
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