A Lewis structure is a two-dimensional diagram representing the sharing of valence electrons (covalent bonds) between atoms in a molecule. These diagrams are fundamental for understanding how atoms bond to achieve stability. While a single line represents a shared pair of electrons, multiple bonds (double or triple) involve a more complex connection: the pi (\(\pi\)) bond. Pi bonds differentiate single bonds from the stronger, more rigid multiple bonds that govern molecular shape and reactivity.
The Foundation of Covalent Bonds: Sigma and Pi
Covalent bonds form when atomic orbitals from two different atoms overlap in space, allowing electrons to be shared between the nuclei. This overlap can occur in two distinct geometric arrangements, which results in the two fundamental types of covalent bonds: sigma and pi. The first bond formed between any two atoms is always a sigma (\(\sigma\)) bond, which results from the head-to-head or axial overlap of orbitals.
This direct overlap positions the highest electron density directly along the internuclear axis (the imaginary line connecting the two atomic nuclei). Because this head-on alignment maximizes the overlap area, the sigma bond is the stronger component of any bond between two atoms. All single bonds are exclusively sigma bonds.
When atoms share more than one pair of electrons, additional bonds form through parallel or side-by-side overlap. This lateral interaction produces the pi bond, which is structurally and energetically distinct from the sigma bond. Pi bonds are supplementary connections that add strength and unique properties to the overall bond.
Defining the Pi Bond: Mechanism and Characteristics
The pi bond is formed exclusively through the side-by-side overlap of unhybridized p-orbitals on adjacent atoms. These p-orbitals are oriented perpendicular to the internuclear axis, meaning they are situated above and below the plane established by the sigma bond. This parallel alignment causes the electron density to be distributed into two separate regions: one above the internuclear axis and one below it.
The overlap area is less direct and less extensive compared to the head-to-head overlap of a sigma bond. Consequently, pi bonds are weaker than sigma bonds, even though their presence increases the overall strength of a multiple bond. The total energy of a double bond is less than twice the energy of a single bond, reflecting the relative weakness of the pi component.
The formation of the pi bond requires the atoms to be close enough for their p-orbitals to interact laterally. This mechanism dictates that a pi bond can never exist on its own; it must always accompany an existing sigma bond between the same two atoms. The electron cloud of the pi bond acts like a sheath around the central sigma bond, creating a region of high electron density that is exposed and chemically reactive.
Identifying Pi Bonds in Lewis Structures
Lewis structures provide a simple way to identify the number of pi bonds. The convention is that the first bond drawn between any two atoms is the sigma bond, and any subsequent bonds are pi bonds.
A single bond, represented by one line, consists of one sigma bond and zero pi bonds (e.g., the carbon-carbon bond in ethane).
A double bond, shown as two parallel lines, contains one sigma bond and one pi bond (e.g., ethene).
Finally, a triple bond, depicted by three parallel lines, is composed of one sigma bond and two pi bonds. In ethyne, the carbon-carbon triple bond consists of one central sigma bond and two pi bonds.
How Pi Bonds Influence Molecular Geometry
The side-by-side p-orbital overlap fundamentally constrains the molecule’s three-dimensional geometry. Unlike single bonds, which are free to rotate around the sigma bond axis, the presence of a pi bond imposes a rigid barrier to rotation. Rotating the atoms around a multiple bond would require breaking the parallel overlap, which is energetically costly.
This restricted rotation forces the atoms participating in the multiple bond, along with their attached groups, to lie in the same plane in space. For a double bond, this rigidity results in a planar geometry around the bonded atoms, such as the trigonal planar shape seen in ethene. This lack of free rotation is the direct cause of geometric isomerism, where two molecules have the same chemical formula but different structural arrangements.
The presence of pi bonds also contributes to the phenomenon of resonance, where the electron density of the pi bond is delocalized, or spread out, over several atoms. This delocalization is necessary for the enhanced stability observed in many cyclic and aromatic compounds. The unique geometry and rigidity imparted by pi bonds are central to determining a molecule’s shape, stability, and chemical behavior.