Britannica Sources used in this answer

How do molecular orbitals explain covalent bonding?

Created from Britannica articles using AI. AI can make mistakes, so verify using Britannica articles.

representations of molecular structure — molecule Several methods of representing a molecule's structure. In Lewis structures, element symbols represent atoms, and dots represent electrons surrounding them. A pair of shared electrons (covalent bond) may also be shown as a single dash. The ball-and-stick model better illustrates the spatial arrangement of the atoms. For aromatic compounds, the Kekulé structure is common, in which each bond is represented by a dash, carbon atoms are implied where two or more lines meet, and hydrogen atoms are usually omitted. Bond-line formulas, similar to the Kekulé structure, are often used for complex nonaromatic organic compounds. Sugars are often drawn as Fischer projections, in which the carbon “backbone” is drawn as a straight vertical line, with carbon atoms implied where horizontal lines intersect the vertical one.

Molecular-orbital theory explains covalent bonding by describing electrons as being distributed over all the nuclei in a molecule, rather than being localized between specific atoms.12

In this model, molecular orbitals (MOs) are formed through the linear combination of atomic orbitals (LCAO), where atomic wave functions interfere with one another. When atomic orbitals combine, they create two types of molecular orbitals:2

Bonding orbitals: These result from constructive interference, which increases electron density between the nuclei. Because this region of space is at a lower energy level than the original atomic orbitals, it stabilizes the molecule.13
Antibonding orbitals: These result from destructive interference, creating a nodal plane of zero amplitude between the nuclei. These orbitals are at a higher energy level and, if occupied, destabilize the bond.13

A covalent bond forms when electrons occupy the lower-energy bonding orbital, resulting in a system with lower total energy than the separated atoms. The theory effectively accounts for bonding by showing that the energy of the paired electrons is minimized when they reside in these shared, lower-energy spatial distributions.13