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Two atoms can share one, two or three pairs of electrons. Those three possibilities — single, double and triple bonds — behave in systematically different ways. They differ in length, strength, shape, whether the atoms can rotate, and how the molecule reacts. Comparing them side by side is one of the quickest ways to understand organic chemistry and much of inorganic chemistry too.
At a glance
| Feature | Single bond | Double bond | Triple bond |
|---|---|---|---|
| Shared electron pairs | 1 | 2 | 3 |
| Electrons in the bond | 2 | 4 | 6 |
| Drawn as | A–B | A=B | A≡B |
| Made of | 1 σ | 1 σ + 1 π | 1 σ + 2 π |
| Bond order | 1 | 2 | 3 |
| Length | Longest | Shorter | Shortest |
| Bond energy | Lowest | Higher | Highest |
| Rotation around the bond | Free | Restricted | Restricted (linear, so rotation doesn’t change anything) |
| Geometry at a carbon atom | Tetrahedral (109.5°) | Trigonal planar (120°) | Linear (180°) |
| Carbon hybridisation | sp³ | sp² | sp |
| Typical reactivity (carbon compounds) | Unreactive: substitution | Reactive: addition | Reactive: addition |
| Classic examples | H–H, C–C in ethane, O–H | O=O, C=C in ethene, C=O | N≡N, C≡C in ethyne, C≡N |
Counting electrons
A single bond is one shared pair of electrons. Each atom contributes one electron. Examples: H₂, Cl₂, the C–H bonds in methane (see covalent bonding explained).
A double bond shares two pairs — four electrons. In O₂, each oxygen has six outer electrons and needs two more, so the atoms share two pairs. In carbon dioxide, O=C=O, carbon forms a double bond to each oxygen.
A triple bond shares three pairs — six electrons. In N₂, each nitrogen has five outer electrons and needs three more, so they share three pairs, leaving one lone pair on each nitrogen.
To check any structure, count electrons around each atom: for most atoms in periods 2 and 3, bonding pairs plus lone pairs should give eight (see how to draw dot-and-cross diagrams).
Length and strength: the data
More shared electrons between two nuclei pull them closer together and hold them more tightly. The pattern holds for every pair of elements:
| Bond | Length (pm) | Average bond energy (kJ mol⁻¹) |
|---|---|---|
| C–C | 154 | 347 |
| C=C | 134 | 614 |
| C≡C | 120 | 839 |
| C–O | 143 | 358 |
| C=O | 120–123 | about 745–805 (depends on the molecule) |
| C≡O (in CO) | 113 | 1,077 |
| N–N | 145 | 163 |
| N=N | 125 | 418 |
| N≡N | 110 | 945 |
| O–O | 148 | 146 |
| O=O | 121 | 498 |
Two patterns stand out:
- Shorter bonds are stronger (see bond length and bond strength).
- A double bond is stronger than a single bond but less than twice as strong (for carbon: 614 vs 2 × 347 = 694). A triple bond is less than three times as strong (839 vs 3 × 347 = 1,041).
Nitrogen is an exception to the second pattern: N≡N (945) is more than three times N–N (163). The N–N single bond is unusually weak, because the lone pairs on the two small nitrogen atoms repel each other strongly. That contrast is why nitrogen exists as N₂ and why compounds with N–N single bonds, such as hydrazine, tend to be reactive and release energy when they form N₂.
Why isn’t a double bond twice as strong? Sigma and pi
The first pair of electrons in any bond forms a sigma (σ) bond: the electron density sits directly on the line between the nuclei, where it’s most effective at holding them together.
The second and third pairs form pi (π) bonds. They come from p orbitals overlapping sideways, above and below (and in front of and behind) the line between the nuclei. Sideways overlap is less effective than head-on overlap, so each π bond is weaker than the σ bond (see sigma and pi bonds).
For carbon–carbon bonds:
- σ bond ≈ 347 kJ mol⁻¹
- first π bond ≈ 614 − 347 = 267 kJ mol⁻¹
- second π bond ≈ 839 − 614 = 225 kJ mol⁻¹
So:
- Single = σ
- Double = σ + π
- Triple = σ + 2π
Rotation and isomers
Atoms joined by a single bond can rotate freely around it, because turning doesn’t affect the head-on σ overlap. That’s why the two ends of ethane, CH₃–CH₃, can spin relative to each other.
A double bond can’t rotate without breaking the π bond, which needs about 250 kJ mol⁻¹ — far more than is available at room temperature. This restricted rotation creates cis–trans (E/Z) isomers: molecules with the same atoms joined in the same order but locked into different arrangements. But-2-ene exists as cis (both CH₃ groups on the same side) and trans (opposite sides), with different boiling points. The same effect gives unsaturated fats their kinks (see saturated vs unsaturated fats) and makes vision possible when light flips a cis double bond in retinal (see how we see).
A triple bond makes the atoms and their neighbours lie in a straight line, so there’s nothing to distinguish by rotation, and no cis–trans isomers.
Shape around the atoms
The number of electron domains around a carbon atom sets its shape (see VSEPR and molecular geometry). A double or triple bond counts as one domain:
- Ethane (C–C): each carbon has four single bonds → tetrahedral, 109.5°.
- Ethene (C=C): each carbon has three domains → trigonal planar, about 120°. All six atoms lie in one plane.
- Ethyne (C≡C): each carbon has two domains → linear, 180°. H–C≡C–H is a straight line.
Reactivity
For carbon compounds, the pattern is often surprising at first: although a C=C bond is stronger overall than a C–C bond, alkenes are more reactive than alkanes.
The reason is the π bond. Its electrons sit above and below the plane of the molecule, exposed and relatively loosely held. Electron-seeking reagents (electrophiles) such as bromine and hydrogen bromide are attracted to this electron-rich region and add across the double bond, breaking only the weaker π bond:
CH₂=CH₂ + Br₂ → CH₂Br–CH₂Br
This addition reaction is the basis of the bromine water test for unsaturation — the orange colour disappears. Alkanes, with only σ bonds, don’t react with bromine water in the dark; they need UV light to undergo substitution.
Triple bonds in alkynes can undergo addition twice. Nitrogen’s triple bond is the exception to high reactivity: N≡N is so strong, and N₂ so non-polar, that it’s famously unreactive — a large part of why air is safe to breathe and why making ammonia from nitrogen needs an iron catalyst at high temperature and pressure.
Beyond carbon
Double and triple bonds are common for period 2 elements — carbon, nitrogen and oxygen — which are small enough for their p orbitals to overlap well sideways. Heavier elements form π bonds much less readily, because their larger atoms keep p orbitals further apart. That’s why carbon dioxide is a gas of separate O=C=O molecules, while silicon dioxide is a giant network of Si–O single bonds, and why nitrogen is N₂ while phosphorus forms P₄ molecules with single bonds.
Common misconceptions
- “A double bond is twice as strong as a single bond.” It’s stronger, but less than twice as strong for carbon.
- “Stronger bonds mean less reactive molecules.” Alkenes react readily because the π bond is exposed, even though C=C is stronger than C–C.
- “A double bond is two identical bonds.” It’s one σ and one π bond.
- “All four electrons in a double bond sit between the nuclei.” The π pair sits above and below the bond axis.
Key takeaways
- Single = 1 pair (σ), double = 2 pairs (σ + π), triple = 3 pairs (σ + 2π).
- More pairs → shorter and stronger bonds, though not proportionally.
- Double bonds restrict rotation, giving cis–trans isomers; triple bonds make atoms linear.
- Carbon geometry: tetrahedral, trigonal planar, linear.
- π bonds make alkenes and alkynes reactive in addition reactions.
For bond energies in calculations, see bond enthalpy calculations.
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