A multiple covalent bond forms when two atoms share more than one pair of electrons. In a standard single bond, two atoms share one pair of electrons (two electrons total). When that isn’t enough for both atoms to reach a stable electron arrangement, they share additional pairs, forming a double bond (two shared pairs, four electrons) or a triple bond (three shared pairs, six electrons).
Why Multiple Bonds Form
Atoms form bonds to fill their outer electron shell, typically reaching eight electrons (the octet rule). Some atoms can’t get there by sharing just one pair with each neighbor. Carbon dioxide is a good example: if each oxygen shared only one electron pair with the central carbon, none of the three atoms would have a complete octet. Instead, each oxygen shares two pairs with carbon, forming two double bonds and giving every atom in the molecule a full set of eight outer electrons.
The same logic applies to nitrogen gas. A single nitrogen atom has five outer electrons and needs three more. Two nitrogen atoms solve this by sharing three pairs of electrons with each other, creating the triple bond in N₂, one of the strongest bonds found in nature.
How Double and Triple Bonds Are Built
Not all shared pairs in a multiple bond are identical. The first pair forms what chemists call a sigma bond: the electron cloud sits directly along the line between the two atomic nuclei, like a rope connecting them end to end. Every single bond is a sigma bond, and every multiple bond starts with one.
The additional pairs form pi bonds. In a pi bond, the electron cloud doesn’t sit between the nuclei. Instead, it extends above and below (or in front of and behind) the axis connecting the atoms, created by the sideways overlap of orbitals. A double bond is one sigma bond plus one pi bond. A triple bond is one sigma bond plus two pi bonds, with the two pi bonds oriented perpendicular to each other.
Strength and Length
Adding more shared electron pairs pulls two atoms closer together and makes the bond harder to break. The numbers for carbon-carbon bonds illustrate this clearly:
- Single bond (C–C): 346 kJ/mol of energy to break
- Double bond (C=C): 602 kJ/mol
- Triple bond (C≡C): 835 kJ/mol
Notice that a double bond isn’t exactly twice as strong as a single bond, and a triple bond isn’t three times as strong. That’s because pi bonds, formed by sideways orbital overlap, are inherently weaker than the head-on sigma bond. Each additional pi bond adds strength, but with diminishing returns. Bond length follows the opposite pattern: more electron pairs pull the nuclei closer, so triple bonds are shorter than double bonds, which are shorter than single bonds.
Rigidity and Restricted Rotation
One of the most important physical consequences of multiple bonds is that they lock atoms in place. A single bond (sigma only) allows the two groups on either side to rotate freely around the bond axis, like a swivel joint. Pi bonds prevent this rotation because twisting the molecule would break the sideways orbital overlap that holds the pi bond together.
This rigidity is why double bonds create distinct geometric forms in molecules. Two groups attached to the same side of a double bond stay on that side, and two groups on opposite sides stay opposite. In organic chemistry, this produces different versions of the same molecule (called cis and trans isomers) that can have very different properties. Triple bonds take this further: the two pi bonds perpendicular to each other force the molecule into a completely linear arrangement around the bond.
Effects on Molecular Shape
Multiple bonds also influence the angles between neighboring bonds. Because a double or triple bond packs more electron density into one region of space than a single bond does, it pushes nearby single bonds away more forcefully. In formaldehyde (CH₂O), the carbon sits at the center with a double bond to oxygen and single bonds to two hydrogens. The standard angle for three groups around a central atom would be 120°, but the double bond’s extra repulsion compresses the angle between the two hydrogens to about 116.5°.
The general rule is that multiple bonds act somewhat like lone pairs of electrons: they take up more room and squeeze adjacent bonds into tighter angles. The order of repulsive force goes from lone pairs (strongest) to multiple bonds to single bonds (weakest).
Common Molecules With Multiple Bonds
Multiple bonds are everywhere, both in simple molecules and in complex biological structures.
Oxygen gas (O₂) holds its two atoms together with a double bond. Nitrogen gas (N₂) uses a triple bond, which is why N₂ is so chemically inert: breaking that triple bond requires enormous energy. Carbon dioxide (CO₂) contains two carbon-oxygen double bonds, giving the molecule its linear shape. Acetylene (C₂H₂), the gas used in welding torches, features a carbon-carbon triple bond that stores a large amount of energy.
In organic chemistry, molecules with carbon-carbon double bonds are called alkenes, and those with triple bonds are called alkynes. The double bonds in alkenes are what make unsaturated fats “unsaturated,” and they’re the reactive sites where many chemical reactions occur. Benzene, the ring-shaped molecule at the heart of countless industrial chemicals, has a special arrangement where pi electrons spread out evenly across all six carbon atoms rather than staying locked between any two. This delocalization gives benzene unusual stability compared to what you’d expect from three separate double bonds.
The carbon-oxygen double bond shows up in nearly every biological molecule, from the amino acids that build proteins to the sugars that fuel your cells. The carbon-nitrogen triple bond appears in compounds called nitriles, which are widely used in manufacturing plastics and pharmaceuticals. Carbon monoxide (CO) contains a triple bond with a bond energy of 1,072 kJ/mol, making it one of the strongest covalent bonds measured.