Editing Covalent bond (section)

== One- and three-electron bonds ==
[[File:Graphical comparison of bonds.svg|200px|thumb|right|[[Lewis structure|Lewis]] and [[MO diagram]]s of an individual 2e<sup>−</sup> bond and 3e<sup>−</sup> bond]]
Bonds with one or three electrons can be found in [[radical (chemistry)|radical]] species, which have an odd number of electrons. The simplest example of a 1-electron bond is found in the [[dihydrogen cation]], {{chem|H|2|+}}. One-electron bonds often have about half the bond energy of a 2-electron bond, and are therefore called "half bonds". However, there are exceptions: in the case of [[dilithium]], the bond is actually stronger for the 1-electron {{chem|Li|2|+}} than for the 2-electron Li<sub>2</sub>. This exception can be explained in terms of [[Orbital hybridisation|hybridization]] and inner-shell effects.<ref>{{cite book | title=Valency and Bonding| publisher=Cambridge | year=2005 |pages=96–100 | last1=Weinhold|first1= F. |last2= Landis|first2= C. | isbn=0-521-83128-8}}</ref>

The simplest example of three-electron bonding can be found in the [[helium dimer]] cation, {{chem|He|2|+}}. It is considered a "half bond" because it consists of only one shared electron (rather than two);<ref>{{cite book |editor-last=Harcourt |editor-first=Richard D.|title=Bonding in Electron-Rich Molecules: Qualitative Valence-Bond Approach via Increased-Valence Structures |publisher=Springer |date=2015 |chapter=Chapter 2: Pauling "3-Electron Bonds", 4-Electron 3-Centre Bonding, and the Need for an "Increased-Valence" Theory|isbn=9783319166766}}</ref> in molecular orbital terms, the third electron is in an anti-bonding orbital which cancels out half of the bond formed by the other two electrons. Another example of a molecule containing a 3-electron bond, in addition to two 2-electron bonds, is [[nitric oxide]], NO. The oxygen molecule, O<sub>2</sub> can also be regarded as having two 3-electron bonds and one 2-electron bond, which accounts for its [[paramagnetism]] and its formal bond order of 2.<ref name="pauling">{{cite book|last=Pauling|first=L.|date=1960|title=The Nature of the Chemical Bond|url=https://archive.org/details/natureofchemical00paul|url-access=registration|publisher=Cornell University Press|pages=[https://archive.org/details/natureofchemical00paul/page/340 340–354]}}</ref> [[Chlorine dioxide]] and its heavier analogues [[bromine dioxide]] and [[Iodine oxide|iodine dioxide]] also contain three-electron bonds.

Molecules with odd-electron bonds are usually highly reactive. These types of bond are only stable between atoms with similar electronegativities.<ref name="pauling" />

[[Dioxygen]] is sometimes represented as obeying the octet rule with a double bond (O=O) containing two pairs of shared electrons.<ref>For example, ''General chemistry'' by R.H.Petrucci, W.S.Harwood and F.G.Herring (8th ed., Prentice-Hall 2002, {{ISBN|0-13-014329-4}}, p.395) writes the Lewis structure with a double bond, but adds a question mark with the explanation that there is some doubt about the validity of this structure because it fails to account for the observed paramagnetism.</ref> However the ground state of this molecule is [[paramagnetic]], indicating the presence of unpaired electrons. Pauling proposed that this molecule actually contains two three-electron bonds and one normal covalent (two-electron) bond.<ref>L. Pauling ''The Nature of the Chemical Bond'' (3rd ed., Oxford University Press 1960) chapter 10.</ref> The octet on each atom then consists of two electrons from each three-electron bond, plus the two electrons of the covalent bond, plus one lone pair of non-bonding electrons. The bond order is 1+0.5+0.5=2.
{{multiple image | align = center | direction = horizontal | header = Modified Lewis structures with 3e bonds | width = 150 | image1 = Nitric oxide.svg | caption1  = Nitric oxide | image2 = Triplett-Sauerstoff.svg | caption2  = Dioxygen }}