When atoms come together, they form chemical bonds that hold them in specific arrangements, creating molecules. These bonds represent the attractive forces that stabilize atoms by sharing or transferring electrons. The way atoms connect fundamentally influences a molecule’s shape and its resulting properties.
The Foundation: Sigma Bonds
The most prevalent type of covalent bond is the sigma (σ) bond, formed by the direct, head-on overlap of atomic orbitals. This overlap concentrates electron density along the internuclear axis, the imaginary line connecting the two atomic nuclei. Sigma bonds can arise from the overlap of various orbital types. This direct overlap results in a strong bond.
The cylindrical symmetry of a sigma bond around the internuclear axis allows for free rotation of the atoms joined by this bond. This rotational freedom means parts of a molecule connected by a single sigma bond can spin relative to each other without breaking the bond. All single bonds are sigma bonds, and a sigma bond is always the first bond to form between two atoms, serving as the foundational link in multiple bonds.
The Uniqueness of Pi Bonds
In contrast to sigma bonds, pi (π) bonds form from the sideways overlap of parallel p orbitals. This lateral overlap creates electron density above and below the internuclear axis, rather than directly between the nuclei. Pi bonds are weaker than sigma bonds because the extent of orbital overlap is less effective in a sideways configuration. A pi bond cannot exist independently between two atoms; it always forms in addition to an existing sigma bond.
The presence of a pi bond restricts rotation around the bond axis, unlike the free rotation seen in sigma bonds. This is because rotating the atoms would disrupt the parallel alignment of the p orbitals, effectively breaking the pi bond.
Unpacking Double Bonds
A double bond is a combination of bonds between two atoms. It consists of one sigma (σ) bond and one pi (π) bond. The sigma bond provides the primary connection. The additional pi bond contributes further electron density, which is distributed above and below the plane of the sigma bond.
This combination makes a double bond stronger and shorter than a single (sigma) bond alone. The inability of the pi bond to rotate freely means that the atoms connected by a double bond are held rigidly in place, influencing molecular structure.
Extending the Concept: Triple Bonds
Building on the understanding of double bonds, a triple bond represents an even greater sharing of electrons between two atoms. A triple bond is comprised of one sigma (σ) bond and two pi (π) bonds. Similar to double bonds, the sigma bond forms the initial link along the internuclear axis. The two additional pi bonds form from the sideways overlap of two sets of parallel p orbitals, further increasing the electron density between the atoms.
This arrangement makes triple bonds the strongest and shortest type of covalent bond between two atoms. The presence of two pi bonds imposes even greater restrictions on rotation compared to double bonds, resulting in a linear geometry around the triple bond.