Atoms may form multiple bonds (single, double, and/or triple) depending on he number of valence electrons they have available.
Multiple bonds consist of a sigma (σ) bond located along the axis between two atoms and one or two pi (π) bonds. The σ bonds are usually formed by the overlap of hybridized atomic orbitals, while the π bonds are formed by the side-by-side overlap of unhybridized orbitals.
For example, the Lewis structure of ethene, C2H4, shows us that each carbon atom is surrounded by one other carbon atom and two hydrogen atoms. The three bonding regions form a trigonal planar electron-pair geometry. Thus we expect the σ bonds from each carbon atom are formed using a set of sp2 hybrid orbitals that result from hybridization of two of the 2p orbitals and the 2s orbital. These orbitals form the C–H single bonds and the σ bond in the C=C double bond.
The unhybridized p orbital is perpendicular to the plane of the sp2 hybrid orbitals. Thus the unhybridized 2p orbitals overlap in a side-by-side fashion, above and below the σ bond and form a π bond. This results in a double covalent bond forming in the ethene structure between the two carbon atoms.
Multiple bonds cause rigidity in the molecular structure. A single σ bond can rotate easily because the end-to-end orbital overlap does not depend on the relative orientation of the orbitals on each atom in the bond. But with a π bond accompanying the σ bond rotation about the axis is not possible.
Double and triple covalent bonds are stronger than single covalent bonds as they are composed of σ bonds between hybridized orbitals and π bonds between unhybridized p orbitals. Double and triple bonds also have shorter bond lengths between atoms as well as higher bond energies than single bonds.
• Multiple bonds form when unhybridized p orbitals overlap on adjacent atoms which are already bonded by a σ bond. This introduces multiple bonds to the structure.
• Multiple bonds are identified as double or triple covalent bonds and have rigidity in structure due to the π bonds, there is no free rotation around the σ bond. The multiple bonds also have higher bond energy and shorter bond lengths than a single bond.
Sigma (σ) bond: A covalent bond whose electron density is concentrated in the region directly between the nuclei.
Pi (π) bond: A covalent bond formed between two neighboring atom’s unbonded p-orbitals.
Covalent bond: A type of chemical bond where two atoms are connected to each other by the sharing of two or more electrons.
Bond strength: Directly related to the amount of energy required to break the bond between two atoms. The more energy required, the stronger the bond is said to be.
Bond length: The distance between the nuclei of two bonded atoms. It can be experimentally determined.
Orbital hybridization: The concept of mixing atomic orbitals to form new hybrid orbitals suitable for the qualitative description of atomic bonding properties and geometries.
Atomic orbitals: The physical region in space around the nucleus where an electron has a probability of being.
Hybrid orbital: An orbital formed by the combination of two or more atomic orbitals.
Unhybridized orbitals: The ground state of the atom.
P orbital : The orbital of an electron shell in an atom in which the electrons have the second lowest energy.
sp2: One of a set of hybrid orbitals produced when one s orbital and two p orbitals are combined mathematically to form three new equivalent orbitals oriented toward the corners of a triangle (hence the designation trigonal planar).