A sigma (\(\sigma\)) bond represents the most basic form of a covalent chemical connection between two atoms. It is defined as a bond formed by the direct, head-on overlap of atomic orbitals, which results in the highest concentration of electron density located precisely between the two atomic nuclei. This direct overlap makes the sigma bond the strongest type of covalent bond, serving as the fundamental framework for all single bonds in chemistry. Every multiple bond, whether double or triple, also contains a single sigma bond as its structural base. The Greek letter sigma is used because the bond exhibits a circular symmetry when viewed along the axis connecting the two nuclei.
The Mechanism of Formation
A sigma bond forms when the atomic orbitals of two approaching atoms merge directly along the internuclear axis, which is the imaginary line connecting the centers of the two nuclei. This specific type of interaction is often described as “end-on” or “head-to-head” overlap. The efficiency of this direct alignment maximizes the orbital overlap, allowing the shared electron pair to be strongly attracted to both nuclei simultaneously.
Three primary combinations of simple atomic orbitals can participate in this head-on overlap. The first is the overlap of two spherical s orbitals, such as the bond formed in a hydrogen molecule (\(H_2\)). Another common formation involves the overlap of an s orbital with a p orbital, which occurs in molecules like hydrogen chloride (\(HCl\)). Finally, two p orbitals can overlap axially along the internuclear axis, as seen in the bond between the two chlorine atoms in \(Cl_2\).
Sigma bonds can also be formed from the overlap of hybridized orbitals, which are blends of the simple atomic orbitals created to maximize bonding efficiency. Orbitals like \(sp^3\), \(sp^2\), and \(sp\) can all engage in direct, end-on overlap with other hybrid orbitals, or with simple s or p orbitals. Regardless of whether the orbitals are simple or hybridized, the defining feature is that the region of overlap is concentrated right on the line connecting the two atomic centers.
The Resulting Shape and Symmetry
To visualize what a sigma bond looks like, one must picture the resulting electron cloud after the orbitals have merged. The shape of a sigma bond is characterized by its cylindrical symmetry around the internuclear axis. This means that if the bond were rotated around the line connecting the two nuclei, the electron density cloud would look exactly the same from every angle.
The electron density in a sigma bond is at its maximum precisely between the two nuclei, forming a continuous, cylindrical volume that encases the bond axis. This high density of negative charge between the two positively charged nuclei acts like an electrostatic glue, holding the atoms together firmly. Since the electron density is uniformly distributed around the axis, there is no restriction on the rotation of the atoms relative to each other.
This free rotation is a direct consequence of the bond’s cylindrical symmetry. Molecules like ethane (\(C_2H_6\)), which consists entirely of sigma bonds, can therefore adopt various three-dimensional shapes through the rotation of its parts.
How Sigma Bonds Differ from Pi Bonds
The distinction between a sigma bond and a pi (\(\pi\)) bond lies in the geometry of their orbital overlap. While a sigma bond forms from the head-on overlap of orbitals along the internuclear axis, a pi bond forms from the side-by-side or parallel overlap of unhybridized p orbitals. This difference in formation geometry leads to different resulting shapes for the electron clouds.
In a pi bond, the electron density is not concentrated along the internuclear axis but is instead distributed in two separate regions, one above and one below the plane of the sigma bond. This creates a region of zero electron density, called a nodal plane, that runs directly along the internuclear axis. The side-by-side overlap is less efficient than the direct head-on overlap, which is why pi bonds are generally weaker.
Pi bonds are only found in multiple bonds. A double bond consists of one sigma bond and one pi bond, while a triple bond contains one sigma bond and two pi bonds. The presence of the pi bond’s electron density above and below the axis locks the atoms into a fixed position, preventing the free rotation that is characteristic of a lone sigma bond.