A chemical bond forms when atoms share electrons (covalent bonding). A double bond forms when two atoms share two pairs of electrons (four electrons total). This increased electron sharing creates a stronger and shorter connection between the atoms. The presence of a double bond fundamentally alters a molecule’s structure and chemical behavior, giving it unique properties exploited in natural and manufactured compounds.
How Electron Sharing Creates the Double Bond
A double bond is composed of two distinct types of bonds. The foundation is the sigma (\(\sigma\)) bond, which forms from the direct, head-to-head overlap of electron orbitals along the axis connecting the two atomic nuclei. This sigma bond provides the primary, stable link between the two atoms.
The second part is the pi (\(\pi\)) bond, created by the side-by-side overlap of unhybridized p-orbitals. These orbitals are oriented perpendicular to the plane of the sigma bond. Because they overlap sideways, the pi bond is weaker than the sigma bond. Its electron density is distributed above and below the plane of the sigma bond.
The combination of one sigma bond and one pi bond constitutes the double bond. This structure results in a cloud of electron density surrounding the core sigma framework. This arrangement makes the double bond stronger and holds the atoms closer together than a single bond. The pi bond, while weaker individually, is the source of the double bond’s distinctive characteristics, including increased chemical reactivity.
The Physical Consequences of Double Bonding
The dual structure of the double bond directly influences the molecule’s physical characteristics. Double bonds are consistently shorter than corresponding single bonds because the increased electron attraction pulls the nuclei closer together. For example, the carbon-carbon single bond in ethane measures about 154 picometers, while the double bond in ethene is approximately 133 picometers.
The double bond is stronger than a single bond, requiring more energy to break, but it is not twice as strong because the pi component is inherently weaker than the sigma component. The pi bond’s side-by-side overlap prevents free rotation around the bond axis. It acts like a rigid clamp that locks the atoms on either side into a fixed spatial arrangement.
This restricted rotation forces the atoms involved in the double bond and the atoms attached to them to lie in a single, flat plane. This arrangement is known as trigonal planar geometry, where bond angles around the central atoms are fixed at approximately 120 degrees. This planar, rigid structure defines molecules containing double bonds, contrasting with the three-dimensional, tetrahedral geometry of single-bonded carbons.
Double Bonds in Chemistry and Biology
Double bonds are foundational to alkenes, organic compounds like ethene (ethylene), used in plastics production. The electron-rich pi bond makes the double bond a target for chemical reactions. Since the pi bond is relatively easy to break, new atoms or groups can be added across the double bond in addition reactions.
In biological systems, double bonds play an important role in fat structure. Saturated fats contain carbon chains with only single bonds, allowing the chains to remain straight and pack tightly, making them typically solid at room temperature. Unsaturated fats contain one or more carbon-carbon double bonds.
The presence of a double bond introduces a sharp bend or “kink” into the hydrocarbon chain of a fatty acid. This structural distortion prevents the chains from stacking neatly, resulting in a lower melting point. Consequently, most unsaturated fats, like vegetable oils, exist as liquids at room temperature. The difference in bonding dictates the physical state and dietary properties of these compounds.