Editing Lone pair (section)

{{Short description|Pair of valence electrons which are not shared with another atom in a covalent bond}}
[[File:Hydroxide lone pairs-2D.svg|thumb|right|200px|Lone pairs (shown as pairs of dots) in the [[Lewis structure]] of [[hydroxide]]]]

In chemistry, a '''lone pair''' refers to a pair of [[valence electron]]s that are not shared with another atom in a [[covalent bond]]<ref name=goldbookoxstate>[[IUPAC]] ''[[Gold Book]]'' definition: [https://goldbook.iupac.org/terms/view/L03618 ''lone (electron) pair'']</ref> and is sometimes called an '''unshared pair''' or '''non-bonding pair'''. Lone pairs are found in the outermost [[electron shell]] of atoms. They can be identified by using a [[Lewis structure]]. [[Electron pair|Electron pairs]] are therefore considered lone pairs if two electrons are paired but are not used in [[chemical bonding]]. Thus, the number of [[electron]]s in lone pairs plus the number of electrons in bonds equals the number of valence electrons around an atom.

Lone pair is a concept used in [[valence shell electron pair repulsion theory]] (VSEPR theory) which explains the [[Molecular geometry|shapes of molecules]]. They are also referred to in the chemistry of [[Lewis acids and bases]]. However, not all non-bonding pairs of electrons are considered by chemists to be lone pairs. Examples are the transition metals where the non-bonding pairs do not influence molecular geometry and are said to be stereochemically inactive. In [[molecular orbital theory]] (fully [[delocalized]] canonical [[Molecular orbital|orbitals]] or localized in some form), the concept of a lone pair is less distinct, as the correspondence between an orbital and components of a Lewis structure is often not straightforward.  Nevertheless, occupied [[non-bonding orbital]]s (or orbitals of mostly nonbonding character) are frequently identified as lone pairs.

[[File:ParSolitario.png|thumb|right|300px|Lone pairs in [[ammonia]] (A), [[water]] (B), and [[hydrogen chloride]] (C)]]

A ''single'' lone pair can be found with atoms in the [[nitrogen group]], such as nitrogen in [[ammonia]]. ''Two'' lone pairs can be found with atoms in the [[chalcogen]] group, such as oxygen in water. The [[halogen]]s can carry ''three'' lone pairs, such as in [[hydrogen chloride]].

In VSEPR theory the electron pairs on the oxygen atom in water form the vertices of a tetrahedron with the lone pairs on two of the four vertices. The H–O–H [[bond angle]] is 104.5°, less than the 109° predicted for a [[tetrahedral angle]], and this can be explained by a repulsive interaction between the lone pairs.<ref name="FoxWhitesell2004">{{cite book |last1=Fox |first1=M.A. |last2=Whitesell |first2=J.K. |title=Organic Chemistry |publisher=Jones and Bartlett Publishers |year=2004 |isbn=978-0-7637-2197-8 |url=https://books.google.com/books?id=xx_uIP5LgO8C |access-date=5 May 2021 |page=}}</ref><ref name="McMurry2000">{{cite book |last=McMurry |first=J. |title=Organic Chemistry 5th Ed. |publisher=Ceneage Learning India Pvt Limited |year=2000 |isbn=978-81-315-0039-2 |url=https://books.google.com/books?id=1i84SwAACAAJ |access-date=5 May 2021}}</ref><ref name="Lee">{{cite book |last=Lee |first=J.D. |title=Concise Inorganic Chemistry |publisher=Van Nostrand |series=Student's paperback edition |year=1968 |url=https://books.google.com/books?id=THcfnwEACAAJ |access-date=5 May 2021 |page=}}</ref>

Various computational criteria for the presence of lone pairs have been proposed.  While electron density ρ('''r''') itself generally does not provide useful guidance in this regard, the [[Laplace operator|Laplacian]] of the electron density is revealing, and one criterion for the location of the lone pair is where  ''L''('''r''') ''= –''∇<sup>2</sup>ρ('''r''') is a local maximum.  The minima of the electrostatic potential ''V''('''r''') is another proposed criterion.  Yet another considers the [[electron localization function]] (ELF).<ref name=":0">{{cite journal |last1=Kumar |first1=Anmol |last2=Gadre |first2=Shridhar R. |last3=Mohan |first3=Neetha |last4=Suresh |first4=Cherumuttathu H. |date=2014-01-06 |title=Lone Pairs: An Electrostatic Viewpoint |journal=The Journal of Physical Chemistry A |language=en |volume=118 |issue=2 |pages=526–532 |doi=10.1021/jp4117003 |pmid=24372481 |issn=1089-5639 |bibcode=2014JPCA..118..526K}}</ref>