A chemical bond is the strong attractive force that holds atoms together to form molecules and compounds. Covalent bonds involve atoms sharing electron pairs, and these bonds are categorized into two distinct types based on the geometry of orbital overlap. Understanding these bonding modes explains the vast diversity of molecular shapes and the specific chemical reactions molecules undergo. This discussion focuses on the nature and formation of the pi ($\pi$) bond.
The Foundation: Sigma Bonds vs. Pi Bonds
The distinction between sigma ($\sigma$) and pi ($\pi$) bonds is based on how atomic orbitals interact in space. A sigma bond is the first bond formed between two atoms. It results from the direct, head-on overlap of orbitals along the internuclear axis—the imaginary line connecting the two atomic nuclei. This linear overlap concentrates electron density directly between the nuclei, creating a strong, cylindrically symmetrical bond.
A pi bond forms from the side-by-side or lateral overlap of atomic orbitals. This parallel alignment means the electron density is not concentrated along the internuclear axis. Instead, it exists in two separate regions: one lobe above the plane of the sigma bond and one lobe below it. Pi bonds never exist alone; they are always the second or third bond formed between atoms already connected by a sigma bond, resulting in double or triple bonds.
The Mechanics of Pi Bond Formation
Pi bonds are created through the lateral overlap of unhybridized $p$-orbitals that lie parallel to each other. For example, in ethene, carbon atoms use $sp^2$ hybrid orbitals to form the sigma framework. This leaves one $p$-orbital on each carbon atom oriented perpendicular to the sigma bond plane. These two parallel $p$-orbitals approach sideways, allowing their electron clouds to merge above and below the sigma bond axis.
Because this sideways interaction is less efficient than head-on overlap, a pi bond is weaker than a sigma bond. The resulting pi molecular orbital consists of two separate regions of electron density. These two lobes collectively represent a single pi bond, formed by the constructive overlap of the component $p$-orbitals. The pi electron cloud shares a nodal plane, a region of zero electron density, that passes through both atomic nuclei.
Influence on Molecular Shape and Reactivity
The presence of a pi bond fundamentally alters a molecule’s structure and chemical behavior. Although individually weaker than a sigma bond, its addition creates a multiple bond that is shorter and stronger overall than a single bond. For example, the carbon-carbon triple bond, containing one sigma and two pi bonds, is the shortest and strongest carbon-carbon connection.
The most significant structural consequence is the restriction of rotation around the bond axis. The two $p$-orbitals must maintain parallel alignment for lateral overlap to persist; rotation would break the pi bond. This rigidity locks the atoms into a fixed spatial arrangement, which can lead to geometric isomers—molecules with the same chemical formula but different arrangements of atoms in space. Furthermore, the exposed pi electron cloud is highly susceptible to attack by electron-seeking chemical species, increasing the molecule’s reactivity, especially in addition reactions.
Delocalized Pi Systems and Stability
In certain molecules, the influence of pi bonds extends beyond a single pair of bonded atoms, leading to a phenomenon called delocalization. This occurs when multiple pi bonds are separated by a single bond, creating a conjugated system. The parallel $p$-orbitals overlap across three or more atoms, allowing the pi electrons to move, or become “delocalized,” over the entire length of the conjugated chain rather than being confined to a single bond.
The dispersal of electron density over a larger volume significantly stabilizes the molecule by lowering its overall potential energy. This stabilization is often described using the concept of resonance, meaning the molecule is better represented by an average of multiple hypothetical structures. The classic example is the benzene ring, where six carbon atoms form a hexagon. The six pi electrons are completely delocalized across the ring, granting the molecule exceptional stability and unique chemical properties known as aromaticity.