Chemical bonds hold atoms together, allowing them to assemble into molecules. Understanding these linkages is key to grasping the structure and behavior of chemical substances. The covalent bond, which involves the sharing of electrons, is widespread in nature. The single bond represents the most basic and common form of this electron-sharing arrangement.
Defining the Covalent Single Bond
A single bond is a form of covalent bond created when two atoms share one pair of valence electrons between them. This shared pair helps each participating atom complete its outer electron shell and achieve a more stable electronic configuration. This process is most common between nonmetal atoms that have a relatively small difference in electronegativity. For instance, in a molecule of hydrogen gas (\(H_2\)), each hydrogen atom contributes one electron to form a single shared pair. Chemists typically represent this bond using a single line, such as in methane (\(CH_4\)).
How Single Bonds Form Through Orbital Overlap
The formation of a single bond occurs through the direct merging, or overlap, of atomic orbitals from the two participating atoms. This mechanism is known as a sigma (\(\sigma\)) bond, which is the characteristic feature of all single covalent bonds. The overlap happens head-on, meaning the orbitals combine directly along the axis that connects the two atomic nuclei. This merging results in a high concentration of electron density located symmetrically in the region directly between the two nuclei. This symmetrical arrangement of electron density defines the sigma bond.
Unique Properties of Single Bonds
The geometry of the sigma bond imparts physical characteristics to molecules. Single bonds are generally the longest and weakest type of covalent bond between any specific pair of elements. For example, a carbon-carbon single bond is longer than a carbon-carbon double or triple bond. A property arising from the single bond’s axial symmetry is the ability for free rotation around the bond axis. Because the electron density is concentrated along the central axis, the connected atoms can spin or twist relative to one another without breaking the orbital overlap. This rotational freedom is responsible for the conformational flexibility observed in many complex molecules.
Single Bonds Versus Multiple Bonds
The single bond stands in contrast to multiple bonds, such as double and triple bonds, based on the number of shared electron pairs (two pairs for a double bond, three for a triple bond). In terms of orbital overlap, a multiple bond always consists of one underlying sigma bond and one or two additional pi (\(\pi\)) bonds. These pi bonds are formed by the sideways overlap of orbitals above and below the internuclear axis, which locks the atoms into a rigid position. This pi overlap makes multiple bonds progressively shorter and stronger than single bonds between the same two atoms. The restriction of rotation in double and triple bonds fundamentally changes the three-dimensional shape and chemical reactivity of the resulting molecules.