A covalent bond forms when atoms share electrons to achieve a stable electron configuration. The sigma (\(\sigma\)) bond represents the most basic and strongest type of covalent bond between two atoms. It is formed by the direct, head-to-head overlap of atomic orbitals, resulting in a single molecular orbital. This orbital concentrates the electron density directly along the axis connecting the two atomic nuclei. The sigma bond is the foundational connection in all single covalent bonds and is the first bond present in every multiple bond.
How Sigma Bonds Form Through Orbital Overlap
The formation of a sigma bond relies on “end-on” or “head-to-head” overlap, where the orbitals align directly along the internuclear axis. This direct alignment maximizes the shared electron density between the two nuclei. The high degree of overlap in this orientation is why sigma bonds are considered the strongest type of covalent bond.
Sigma bonds can form through the combination of various types of atomic orbitals. The simplest example is the overlap between two spherical s orbitals, such as in a hydrogen molecule (\(H_2\)). A sigma bond can also result from the overlap between one s orbital and one elongated p orbital, as seen in hydrogen fluoride (\(HF\)). Two p orbitals can also form a sigma bond, provided they overlap along the axis connecting the two atoms. In all these configurations, the shared electron cloud is centered between the atoms, which holds the positively charged nuclei together.
Defining Characteristics of Sigma Bonds
A defining feature of the sigma bond is its cylindrical symmetry, which means the electron density is uniform when viewed from any angle along the internuclear axis. If the bond were rotated around its central axis, the appearance of the electron cloud would remain unchanged. This symmetry results directly from the head-to-head overlap of the atomic orbitals.
This cylindrical symmetry allows for the property of free rotation around a single sigma bond. The atoms or groups attached to the bond can spin relative to each other without disrupting the orbital overlap. This rotation allows molecules to adopt many different spatial arrangements, known as conformations.
Sigma bonds are strong because the direct overlap maximizes the attraction between the shared electrons and the two nuclei. Their strength is a consequence of the large extent of orbital overlap possible along the bond axis. The presence of these bonds dictates the fundamental skeleton and geometry of a molecule.
Sigma Bonds vs. Pi Bonds
The primary difference between a sigma (\(\sigma\)) bond and a pi (\(\pi\)) bond lies in the way the atomic orbitals overlap. While sigma bonds are formed by head-to-head overlap along the internuclear axis, pi bonds are formed by the parallel, or side-by-side, overlap of unhybridized p orbitals. This side-by-side overlap creates two distinct regions of electron density, one located above and one below the internuclear axis.
Due to the less efficient side-by-side overlap, pi bonds are weaker than sigma bonds. Rotation is restricted in a pi bond because twisting the bond would break the parallel alignment of the p orbitals, which requires a significant amount of energy. This restriction on rotation makes molecules containing pi bonds more rigid than those with only sigma bonds.
Every single covalent bond is exclusively a sigma bond. Multiple bonds are formed by adding pi bonds to a pre-existing sigma bond. A double bond is composed of one sigma bond and one pi bond, and a triple bond consists of one sigma bond and two pi bonds. The sigma bond always forms first and acts as the anchor around which the pi bond is established.