Pi Bond
A pi bond (π bond) is a covalent bond formed by the lateral (side-by-side) overlap of unhybridised p orbitals above and below the internuclear axis. Pi bonds are always formed in addition to an existing sigma bond, making up the second bond in a double bond and the second and third bonds in a triple bond. Unlike sigma bonds, pi bonds restrict rotation around the bond axis, which is critical for cis-trans isomerism in alkenes.
Comparison of Sigma and Pi Bonds
| Property | Sigma Bond (σ) | Pi Bond (π) |
|---|---|---|
| Type of overlap | Head-on (axial) | Lateral (side-by-side) |
| Relative strength | Stronger | Weaker |
| Electron density | Along internuclear axis | Above and below axis |
| Rotation allowed? | Yes (free rotation) | No (restricted) |
| Occurrence | All covalent bonds | Double and triple bonds only |
| Orbital used | Hybridised or s/p | Unhybridised p orbital |
Interactive Tools
Wikimedia Commons, CC BY-SA
Related Terms
Sigma Bond
A sigma bond (σ bond) is the strongest type of covalent bond, formed by the direct head-on overlap of atomic orbitals along the internuclear axis. It is the first bond formed between two atoms in any covalent bond and allows free rotation around the bond axis. Sigma bonds are present in all single, double, and triple bonds and are responsible for the overall framework and shape of molecules.
Covalent Bond
A covalent bond is a type of chemical bond formed when two atoms share one or more pairs of electrons, resulting in a stable arrangement for both atoms. This sharing occurs most commonly between non-metal atoms that have similar electronegativities, allowing each atom to achieve a full valence shell without complete electron transfer. Covalent bonds are the foundation of organic chemistry and molecular biology, governing the structure of molecules ranging from water (H₂O) to complex proteins.
Resonance Structure
Resonance structures (also called canonical forms or contributing structures) are two or more valid Lewis structures for the same molecule that differ only in the placement of electrons (not atoms), used collectively to represent the actual electron distribution. No single resonance structure accurately depicts the molecule; the true structure is a hybrid (weighted average) of all contributing structures, with electron density delocalised across multiple bonds. The concept is essential for understanding the stability of molecules like benzene, ozone, and carbonate ion, where observed bond lengths are intermediate between single and double bonds.
From the Greek letter pi (π), the second letter of the Greek word "periphery," referencing the electron density existing peripherally (above and below) rather than along the axis. The term was formalised alongside sigma bond notation in molecular orbital theory.