A chemical bond represents a lasting attraction between atoms that enables the formation of chemical compounds. Covalent bonds involve the sharing of valence electrons between two atoms, creating a stable molecular structure. A multiple bond is a type of covalent bond where two atoms share more than one pair of electrons between their nuclei. This increased sharing of electrons significantly influences a molecule’s shape, strength, and chemical behavior.
The Foundation of Covalent Bonding
Covalent bonding begins when atoms share a single pair of electrons, known as a single bond. This shared pair creates a stable link, often helping each atom achieve a full outer electron shell, a principle known as the octet rule. For many simple molecules, such as hydrogen gas (H–H) or methane, a single bond satisfies the stability requirements.
However, many atoms cannot reach a stable octet configuration with only a single shared pair. In these instances, atoms must share additional pairs of electrons to achieve the necessary stability. This drive to complete the octet is the chemical necessity that leads to the formation of multiple bonds. By sharing two or three pairs of electrons, atoms form a more electron-dense linkage that stabilizes the molecule.
The Two Forms of Multiple Bonds
Multiple bonds are categorized based on the number of electron pairs shared. The double bond involves the sharing of two pairs of electrons, totaling four electrons, between two atoms. A common example is the oxygen molecule (\(O=O\)), where each oxygen atom shares two electrons, allowing both to achieve a stable octet.
The triple bond is formed when two atoms share three pairs of electrons, or six electrons in total. Nitrogen gas (\(N\equiv N\)) is a classic example of a triple bond. In organic chemistry, a molecule like acetylene (\(H-C\equiv C-H\)) contains a carbon-carbon triple bond, which is fundamental to its structure and reactivity.
The Role of Orbital Overlap
The formation of multiple bonds is best understood by examining how atomic orbitals overlap to create molecular orbitals. Every single covalent bond, and the first bond in any multiple bond, is a sigma (\(\sigma\)) bond. A sigma bond forms from the direct, head-to-head overlap of atomic orbitals along the axis connecting the two atomic nuclei, concentrating electron density directly between them.
The second and third bonds in a multiple bond are always pi (\(\pi\)) bonds. Pi bonds are formed by the parallel, side-to-side overlap of unhybridized \(p\) orbitals, resulting in two regions of electron density located above and below the plane of the sigma bond. Consequently, a double bond consists of one sigma bond and one pi bond, with the pi bond being the secondary, weaker component.
In the case of a triple bond, the structure comprises one sigma bond and two pi bonds. The two pi bonds are oriented perpendicularly to each other, surrounding the central sigma bond axis to form a dense, cylindrical electron cloud.
Physical Properties and Molecular Structure
The presence of multiple bonds imparts specific physical characteristics to a molecule. Multiple bonds are stronger than single bonds between the same atoms because more energy is required to break the combined sigma and pi bonds. For instance, a carbon-carbon single bond requires approximately 350 kilojoules per mole (kJ/mol) to break, while a double bond requires around 600 kJ/mol, and a triple bond demands over 830 kJ/mol.
This increased electron sharing also pulls the atoms closer together, making multiple bonds shorter than their single bond counterparts. A carbon-carbon single bond measures about 154 picometers (pm), but a double bond shortens to approximately 134 pm, and a triple bond is the shortest at about 120 pm. The pi bonds, with their exposed electron density, make molecules more chemically reactive than those with only single bonds.
A final defining feature is the rigidity a multiple bond imposes on a molecule. While a single bond allows for free rotation around its axis, the side-to-side overlap of the pi bond prevents this rotation in double and triple bonds. This restriction locks the atoms into a fixed spatial arrangement, which can lead to different forms of the same molecule, known as geometric isomers.