Most bonding interactions involve the sharing of a pair of electrons between two atoms. Diatomic hydrogen (H2) is the simplest example of this. Each hydrogen has a single electron to contribute to the covalent bond - (H:H).

The one-electron bond is only about half as strong as an electron pair bond. For example H2+ has two hydrogens sharing only one electron (H.H). The interaction energy of the one-electron bond is approximately 61 kcal/mol, while the typical electron pair bond is closer to 102 kcal/mol in the gas phase.

A significant criterion that must be satisfied for a one-electron bond to occur is that the two participating atoms must be the same or very similar, chemically. As a result of this and the low energy of the interaction, the one-electron bond is extremely rare.

One-electron bonds are found in boron hydrides. Boron hydride is never found in the simplest form BH3 but is most often found in more complex compositions: B2H6, B4H10, B5H11, B5H9, B6H10, B10H14. The problem with most of these complexes is that if you count electrons, you will find that there are not enough valence electrons to form electron pairs between all the atoms. In B2H6, for example, there are twelve electrons available. If all these electrons were used in pairs to hold the six hydrogens to the borons, there would be none left to hold the two borons together. As a result, it appears that there are two one-electron bonds holding the complex together:

            H    H
             ..   ..
        H . B : B . H
             ..   ..
            H    H

Others have suggested that these molecules are held together by a resonance of delocalized electrons, rather than discrete one-electron bonds.

Although the idea of the one electron bond is very old, its existence is still debated in experimental and theoretical studies.

See also: three-electron bond

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