Editing Lithium (section)

== Chemistry ==
{{Main|Category:Lithium compounds}}
{{Redirect|Lithium salt|Lithium salts used in medication|Lithium (medication)}}

=== Of lithium metal ===
Lithium reacts with water easily, but with noticeably less vigor than other alkali metals. The reaction forms [[hydrogen]] gas and [[lithium hydroxide]].<ref name="krebs" /> When placed over a flame, lithium compounds give off a striking crimson color, but when the metal burns strongly, the flame becomes a brilliant silver. Lithium will ignite and burn in oxygen when exposed to water or water vapor. In moist air, lithium rapidly tarnishes to form a black coating of [[lithium hydroxide]] (LiOH and LiOH·H<sub>2</sub>O), [[lithium nitride]] (Li<sub>3</sub>N) and [[lithium carbonate]] (Li<sub>2</sub>CO<sub>3</sub>, the result of a secondary reaction between LiOH and [[carbon dioxide|CO<sub>2</sub>]]).<ref name="kamienski" /> Lithium is one of the few metals that react with [[nitrogen]] gas.<ref>{{cite book |page=47 |url={{google books |plainurl=y |id=yb9xTj72vNAC |page=47}} |title=The history and use of our earth's chemical elements: a reference guide |author=Krebs, Robert E. |publisher=Greenwood Publishing Group |date=2006 |isbn=978-0-313-33438-2 |url-status=live |archive-url=https://web.archive.org/web/20160804025424/https://books.google.com/books?id=yb9xTj72vNAC&pg=PA47 |archive-date=4 August 2016}}</ref><ref>{{Cite journal |author1=Institute, American Geological |author2=Union, American Geophysical |author3=Society, Geochemical |title=Geochemistry international |volume=31 |issue=1–4 |page=115 |date=1 January 1994 |url={{google books |plainurl=y |id=77McAQAAIAAJ}} |url-status=live |archive-url=https://web.archive.org/web/20160604195805/https://books.google.com/books?id=77McAQAAIAAJ |archive-date=4 June 2016 |website=Google Books}}</ref>

Because of its reactivity with water, and especially nitrogen, lithium metal is usually stored in a hydrocarbon sealant, often [[petroleum jelly]]. Although the heavier alkali metals can be stored under [[mineral oil]], lithium is not dense enough to fully submerge itself in these liquids.<ref name="emsley" />

Lithium has a [[diagonal relationship]] with [[magnesium]], an element of similar atomic and [[ionic radius]]. Chemical resemblances between the two metals include the formation of a [[nitride]] by reaction with N<sub>2</sub>, the formation of an [[lithium oxide|oxide]] ({{chem|Li|2|O}}) and peroxide ({{chem|Li|2|O|2}}) when burnt in O<sub>2</sub>, [[salt (chemistry)|salts]] with similar [[solubility|solubilities]], and thermal instability of the [[carbonate]]s and nitrides.<ref name="kamienski">{{Cite book |first=Conrad W. |last=Kamienski |author2=McDonald, Daniel P. |author3=Stark, Marshall W. |author4=Papcun, John R. |chapter=Lithium and lithium compounds |title=Kirk-Othmer Encyclopedia of Chemical Technology |publisher=John Wiley & Sons, Inc. |date=2004 |doi=10.1002/0471238961.1209200811011309.a01.pub2 |isbn=978-0-471-23896-6}}</ref><ref name="Greenwood">{{Greenwood&Earnshaw1st|pages=97–99}}</ref> The metal reacts with hydrogen gas at high temperatures to produce [[lithium hydride]] (LiH).<ref>{{cite web |url=http://www.lyon.edu/webdata/users/fbeckford/CHM%20120/Lecture%20Notes/Chapter-14.ppt |archive-url=https://web.archive.org/web/20051104025202/http://www.lyon.edu/webdata/users/fbeckford/CHM%20120/Lecture%20Notes/Chapter-14.ppt |archive-date=4 November 2005 |title=University of Lyon course online (powerpoint) slideshow |access-date=27 July 2008 |author=Beckford, Floyd |quote=definitions:Slides 8–10 (Chapter 14)}}</ref>

Lithium forms a variety of binary and ternary materials by direct reaction with the main group elements. These [[Zintl phase]]s, although highly covalent, can be viewed as salts of polyatomic anions such as Si<sub>4</sub><sup>4-</sup>, P<sub>7</sub><sup>3-</sup>, and Te<sub>5</sub><sup>2-</sup>.  With graphite, lithium forms a variety of [[intercalation compound]]s.<ref name="Greenwood" />

It dissolves in ammonia (and amines) to give [Li(NH<sub>3</sub>)<sub>4</sub>]<sup>+</sup> and the [[solvated electron]].<ref name="Greenwood" />

=== Inorganic compounds ===
Lithium forms salt-like derivatives with all [[halide]]s and pseudohalides. Some examples include the halides [[lithium fluoride|LiF]], [[lithium chloride|LiCl]], [[lithium bromide|LiBr]], [[Lithium iodide|LiI]], as well as the [[pseudohalide]]s and related anions. Lithium carbonate has been described as the most important compound of lithium.<ref name="Greenwood" />  This white solid is the principal product of [[beneficiation]] of lithium ores. It is a precursor to other salts including ceramics and materials for lithium batteries.

The compounds [[Lithium borohydride|{{chem|LiBH|4}}]] and [[Lithium aluminium hydride|{{chem|LiAlH|4}}]] are useful [[reagent]]s.  These salts and many other lithium salts exhibit distinctively high solubility in ethers, in contrast with salts of heavier alkali metals.

In aqueous solution, the [[coordination complex]] [Li(H<sub>2</sub>O)<sub>4</sub>]<sup>+</sup> predominates for many lithium salts. Related complexes are known with amines and ethers.

=== Organic chemistry ===
{{Main|Organolithium reagent}}
[[File:Butyllithium-hexamer-from-xtal-3D-balls-A.png|thumb|right|Hexameric structure of the [[N-Butyllithium|''n''-butyllithium]] fragment in a crystal]]
[[Organolithium compound]]s are numerous and useful.  They are defined by the presence of a [[covalent bond|bond]] between [[carbon]] and lithium. They serve as metal-stabilized [[carbanion]]s, although their solution and solid-state structures are more complex than this simplistic view.<ref>{{Cite book |url={{google books |plainurl=y |id=z76sVepirh4C |page=16}} |title=Lithium chemistry: a theoretical and experimental overview |author=Sapse, Anne-Marie |author2=von R. Schleyer, Paul |date=1995 |publisher=Wiley-IEEE |isbn=978-0-471-54930-7 |pages=3–40 |archive-url=https://web.archive.org/web/20160731221323/https://books.google.com/books?id=z76sVepirh4C&pg=PA16 |archive-date=31 July 2016 |url-status=live |name-list-style=amp}}</ref> Thus, these are extremely powerful [[base (chemistry)|bases]] and [[carbon nucleophile|nucleophiles]]. They have also been applied in asymmetric synthesis in the pharmaceutical industry. For laboratory organic synthesis, many organolithium reagents are commercially available in solution form. These reagents are highly reactive, and are sometimes [[pyrophoricity|pyrophoric]].

Like its inorganic compounds, almost all organic compounds of lithium formally follow the [[duet rule]] (e.g., [[N-Butyllithium|BuLi]], [[Methyllithium|MeLi]]). However, it is important to note that in the absence of coordinating solvents or ligands, organolithium compounds form dimeric, tetrameric, and hexameric clusters (e.g., BuLi is actually [BuLi]<sub>6</sub> and MeLi is actually [MeLi]<sub>4</sub>) which feature multi-center bonding and increase the coordination number around lithium. These clusters are broken down into smaller or monomeric units in the presence of solvents like [[dimethoxyethane]] (DME) or ligands like [[tetramethylethylenediamine]] (TMEDA).<ref>{{Cite journal |last1=Nichols |first1=Michael A. |last2=Williard |first2=Paul G. |date=1993-02-01 |title=Solid-state structures of n-butyllithium-TMEDA, -THF, and -DME complexes |journal=Journal of the American Chemical Society |volume=115 |issue=4 |pages=1568–1572 |doi=10.1021/ja00057a050 |bibcode=1993JAChS.115.1568N |issn=0002-7863}}</ref> As an exception to the duet rule, a two-coordinate lithate complex with four electrons around lithium, [Li(thf)<sub>4</sub>]<sup>+</sup>[((Me<sub>3</sub>Si)<sub>3</sub>C)<sub>2</sub>Li]<sup>–</sup>, has been characterized crystallographically.<ref>{{Cite book |title=Organometallic chemistry: a unified approach. |last=Mehrotra |first=R. C. |year=2009 |publisher=New Age International Pvt |isbn=978-81-224-1258-1 |oclc=946063142}}</ref>