Editing Covalent bond (section)

== Resonance ==
{{Main|Resonance (chemistry)}}
There are situations whereby a single [[Lewis structure]] is insufficient to explain the electron configuration in a molecule and its resulting experimentally-determined properties, hence a superposition of structures is needed. The same two atoms in such molecules can be bonded differently in different Lewis structures (a single bond in one, a double bond in another, or even none at all), resulting in a non-integer [[bond order]]. The [[nitrate]] ion is one such example with three equivalent structures. The bond between the [[nitrogen]] and each oxygen is a double bond in one structure and a single bond in the other two, so that the average bond order for each N–O interaction is {{sfrac|2 + 1 + 1|3}} = {{sfrac|4|3}}.<ref name=":0" />

[[File:Nitrate-ion-resonance-2D.png|400px]]

=== Aromaticity ===
{{Main|Aromaticity}}

In [[organic chemistry]], when a molecule with a planar ring obeys [[Hückel's rule]], where the number of [[pi bond|π electrons]] fit the formula 4''n''&nbsp;+&nbsp;2 (where ''n'' is an integer), it attains extra stability and symmetry. In [[benzene]], the prototypical aromatic compound, there are 6 π bonding electrons (''n''&nbsp;=&nbsp;1, 4''n''&nbsp;+&nbsp;2&nbsp;=&nbsp;6). These occupy three delocalized π molecular orbitals ([[molecular orbital theory]]) or form conjugate π bonds in two resonance structures that linearly combine ([[valence bond theory]]), creating a regular [[hexagon]] exhibiting a greater stabilization than the hypothetical 1,3,5-cyclohexatriene.<ref name=":1" />

In the case of [[heterocyclic]] aromatics and substituted [[benzene]]s, the electronegativity differences between different parts of the ring may dominate the chemical behavior of aromatic ring bonds, which otherwise are equivalent.<ref name=":1" />

=== Hypervalence ===
{{Main|Hypervalent molecule}}
Certain molecules such as [[xenon difluoride]] and [[sulfur hexafluoride]] have higher coordination numbers than would be possible due to strictly covalent bonding according to the [[octet rule]]. This is explained by the [[three-center four-electron bond]] ("3c–4e") model which interprets the molecular wavefunction in terms of non-bonding [[HOMO/LUMO|highest occupied molecular orbital]]s in [[molecular orbital theory]] and [[resonance (chemistry)|resonance]] of sigma bonds in [[valence bond theory]].<ref>{{Cite book|last1=Weinhold|first1=F.|title=Valency and Bonding|last2=Landis|first2=C.|publisher=Cambridge University Press|year=2005|isbn=0521831288|location=|pages=275–306}}</ref>

=== Electron deficiency ===
{{Main|Electron deficiency}}
In [[three-center two-electron bond]]s ("3c–2e") three atoms share two electrons in bonding. This type of bonding occurs in [[boron hydrides]] such as [[diborane]] (B<sub>2</sub>H<sub>6</sub>), which are often described as electron deficient because there are not enough valence electrons to form localized (2-centre 2-electron) bonds joining all the atoms. However, the more modern description using 3c–2e bonds does provide enough bonding orbitals to connect all the atoms so that the molecules can instead be classified as electron-precise.

Each such bond (2 per molecule in diborane) contains a pair of electrons which connect the [[boron]] atoms to each other in a banana shape, with a proton (the nucleus of a hydrogen atom) in the middle of the bond, sharing electrons with both boron atoms. In certain [[cluster chemistry|cluster compounds]], so-called [[four-center two-electron bond]]s also have been postulated.<ref>{{cite journal |title= A new 4c–2e bond in {{chem|B|6|H|7|-}} |first1= K. |last1= Hofmann |first2= M. H. |last2= Prosenc |first3= B. R. |last3= Albert |journal= Chemical Communications |date= 2007 |volume= 2007 |issue= 29 |pages= 3097–3099 |doi= 10.1039/b704944g |pmid= 17639154 }}</ref>