When atoms join to form molecules, they share electrons in a covalent bond. This sharing can involve one, two, or three pairs of electrons, creating single, double, or triple bonds, respectively. The chemical structure and behavior of a molecule depend on how these shared electron pairs are arranged. While simple chemical drawings use multiple lines for double and triple bonds, the actual arrangement of atomic orbitals is more complex. Understanding the difference between the foundational single connection and supplementary connections is necessary to explain the unique properties of multiple bonds.
Differentiating Sigma and Pi Bonds
Covalent bonds are categorized into two main types: sigma (\(\sigma\)) bonds and pi (\(\pi\)) bonds. The sigma bond is the first bond formed between any two atoms and results from the direct, head-on overlap of orbitals (e.g., two \(s\) orbitals). This overlap concentrates the electron density along the internuclear axis, the imaginary line connecting the nuclei. Due to this high degree of overlap, sigma bonds are generally the strongest type of covalent bond.
Pi bonds are formed by the side-by-side overlap of unhybridized \(p\) orbitals. This parallel overlap concentrates the electron density in two regions, one above and one below the internuclear axis. Pi bonds can only form after a sigma bond has established the connection. The lesser extent of orbital contact compared to the head-on sigma bond makes pi bonds generally weaker.
All single covalent bonds consist of just one sigma bond. When a second bond forms, creating a double bond, it must be a pi bond. The sigma bond acts as the molecular backbone, while the pi bond adds an extra layer of shared electron density.
The Composition of a Triple Bond
A triple bond between two atoms is always composed of one sigma (\(\sigma\)) bond and two pi (\(\pi\)) bonds. The single sigma bond forms first, lying directly along the axis that connects the two nuclei.
The two pi bonds are formed by the side-by-side overlap of the two pairs of unhybridized \(p\) orbitals remaining on each atom. These two pi bonds are oriented perpendicularly to one another, creating a cylindrical sheath of electron density around the central sigma bond. One pi bond exists above and below the internuclear axis, while the second exists in front of and behind the axis.
The triple bond in acetylene (\(C_2H_2\)) illustrates this structure. The two carbon atoms are connected by one sigma bond and two pi bonds, and each carbon also forms a sigma bond with a hydrogen atom. This combination results in a shorter and stronger overall connection than either a single or double bond.
How Triple Bonds Influence Molecular Shape
The structure of a triple bond dictates the geometry of the molecule through orbital mixing known as hybridization. Atoms forming a triple bond use \(sp\) hybridization, which mixes one \(s\) orbital and one \(p\) orbital to create two new \(sp\) hybrid orbitals. These two hybrid orbitals are oriented linearly, positioned 180 degrees apart.
The two \(sp\) hybrid orbitals are used to form the sigma bonds: one connects to the other atom in the triple bond, and the other connects to an adjacent atom. This mixing process leaves two \(p\) orbitals on each atom unhybridized. These unhybridized \(p\) orbitals participate in the side-by-side overlap to form the two perpendicular pi bonds.
The \(sp\) hybridization leads to a defining linear molecular geometry around the triple bond. This linearity results from the electron groups establishing a bond angle of 180 degrees. This is distinct from the trigonal planar shape (120 degrees) seen in double bonds (\(sp^2\) hybridization) and the tetrahedral shape (109.5 degrees) of single bonds (\(sp^3\) hybridization).