Covalent bonds, which involve the sharing of electrons between atoms, form the structural basis of nearly every molecule in the universe. Chemists classify these bonds into distinct types based on how the electron orbitals interact. This distinction is made primarily between sigma (\(\sigma\)) bonds and pi (\(\pi\)) bonds, named after the Greek letters that resemble their orbital symmetry when viewed along the bond axis. Understanding these two bonding modes is fundamental to explaining a molecule’s shape, flexibility, and chemical reactivity. The specific geometry of the orbital overlap determines the type of bond formed.
Defining the Sigma Bond
The sigma (\(\sigma\)) bond represents the most fundamental type of covalent connection between two atoms and is generally considered the strongest. This bond is formed through the direct, face-to-face overlap of atomic orbitals, often described as head-to-head overlap. The result is a concentration of electron density symmetrically centered along the imaginary line connecting the two atomic nuclei, known as the internuclear axis. This direct overlap contributes to the bond’s relative strength and stability.
Sigma bonds can form from the overlap of various orbital types, including s, p, or hybrid orbitals. Since the electron density is localized directly between the nuclei, the atoms connected by a sigma bond can rotate freely around the bond axis without disrupting the orbital overlap. Every single covalent bond found in a molecule is composed of exactly one sigma bond.
Defining the Pi Bond
In contrast to the sigma bond’s head-to-head alignment, the pi (\(\pi\)) bond is formed by the parallel or side-by-side overlap of unhybridized p orbitals. This lateral overlap creates two distinct regions of electron density, one situated above the internuclear axis and the other below it. Electron sharing in a pi bond is less effective than in a sigma bond because the overlap region is farther from the nuclei, which makes pi bonds generally weaker.
Pi bonds never exist alone; they are always created as a second or third connection between atoms already linked by a sigma bond. The side-by-side nature of the pi bond rigidly locks the two atoms into position, preventing free rotation around the bond axis. This restriction on rotation is a defining characteristic of molecules containing pi bonds.
How Bonds Combine in Molecules
The combination of sigma and pi bonds dictates the overall structure and reactivity of a molecule. A single covalent connection between two atoms consists solely of one sigma bond.
Double and Triple Bonds
The double bond, seen in molecules like ethene (\(\text{C}_2\text{H}_4\)), is a combination of one sigma bond and one pi bond. The triple bond, found in compounds such as ethyne (\(\text{C}_2\text{H}_2\)), is composed of one sigma bond and two perpendicular pi bonds.
Impact on Rotation and Reactivity
This arrangement profoundly affects the compound’s properties. Ethane has a carbon-carbon single bond (sigma bond) that allows for free rotation of the \(\text{CH}_3\) groups. Ethene, with its double bond, cannot rotate due to the presence of the pi bond, resulting in a fixed, planar structure. This rigidity makes ethene more reactive than ethane, as the weaker and more exposed pi electrons are easily targeted in chemical addition reactions. Ethyne, with two pi bonds, is forced into a linear molecular geometry.
Visualizing Orbital Overlap
The geometry of the atomic orbitals determines the type of bond that forms. Atomic orbitals are described by their shapes: s orbitals are spherical, while p orbitals are dumbbell-shaped and oriented along three perpendicular axes. When a sigma bond forms, the overlap occurs directly along the internuclear axis, which can happen between any combination of s, p, or hybrid orbitals.
To achieve the correct geometry for bonding, atoms often blend their native s and p orbitals into hybrid orbitals, such as \(sp^3\), \(sp^2\), or \(sp\). For example, the carbon atoms in ethane use \(sp^3\) hybrid orbitals, which are all equivalent and form only sigma bonds. In molecules that form pi bonds, like ethene and ethyne, the atoms use \(sp^2\) or \(sp\) hybridization. This process leaves one or two p orbitals unhybridized. These remaining unhybridized p orbitals are oriented perpendicular to the sigma bond framework, allowing their side-by-side overlap to form the characteristic pi bonds.