View source for Alkali metal

{{short description|Group of highly reactive chemical elements}}
{{distinguish|Alkaline earth metal}}
{{Good article}}
{{Use dmy dates|date=July 2020}}
{{Infobox periodic table group/header
| title       = Alkali metals
| group number= 1
| trivial name= alkali metals
| by element  = lithium group
| CAS         = IA
| old IUPAC   = IA
| mark        = Li,Na,K,Rb,Cs,Fr
| left        = [[noble gas]]es
| right       = [[alkaline earth metal]]s}}
|-
! colspan=2 style="text-align:left;" | ↓&nbsp;<small>[[Period (periodic table)|Period]]</small>
|-
! [[Period 2 element|2]]
| {{element cell image|3|Lithium|Li| |Solid|Alkali metal|Primordial|legend=|image=Lithium paraffin.jpg|image caption=Lithium metal stored under paraffin}}
|-
! [[Period 3 element|3]]
| {{element cell image|11|Sodium|Na| |Solid|Alkali metal|Primordial|legend=|image=Na (Sodium).jpg|image caption=Sodium metal}}
|-
! [[Period 4 element|4]]
| {{element cell image|19|Potassium|K| |Solid|Alkali metal|Primordial|legend=|image=Potassium-2.jpg|image caption=Potassium metal}}
|-
! [[Period 5 element|5]]
| {{element cell image|37|Rubidium|Rb| |Solid|Alkali metal|Primordial|legend=|image=Rb5.jpg|image caption=Rubidium metal in a glass ampoule}}
|-
! [[Period 6 element|6]]
| {{element cell image|55|Caesium|Cs| |Solid|Alkali metal|Primordial|legend=|image=Cesium.jpg|image caption=Caesium metal in a glass ampoule}}
|-
! [[Period 7 element|7]]
| {{element cell image|87|Francium|Fr| |Solid|Alkali metal|from decay|legend=}}
|-
| colspan="2"|
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''Legend''
{| style="text-align:center; border:0; margin: 0 auto"
|-
| style="border:{{element color|Primordial}}; background:{{Element color|table mark}};" | [[Primordial element|primordial]]
|-
| style="border:{{element color|from decay}};  background:{{Element color|table mark}};padding:0 2px;" | [[radioactive decay|element by radioactive decay]]
|-
| Atomic number color:<br/><span style="color:{{element color|Solid}};">black=solid</span>
|}
|}

The '''alkali metals''' consist of the [[chemical element]]s [[lithium]] (Li), [[sodium]] (Na), [[potassium]] (K),<ref group=note>The symbols '''Na''' and '''K''' for sodium and potassium are derived from their Latin names, ''natrium'' and ''kalium''; these are still the origins of the names for the elements in some languages, such as German and Russian.</ref> [[rubidium]] (Rb), [[caesium]] (Cs),{{refn|''Caesium'' is the spelling recommended by the [[International Union of Pure and Applied Chemistry]] (IUPAC).<ref>{{RedBook2005|pages=248–49}}.</ref> The [[American Chemical Society]] (ACS) has used the spelling ''cesium'' since 1921,<ref>{{cite book |editor1-first= Anne M. |editor1-last= Coghill |editor2-first= Lorrin R. |editor2-last= Garson |year= 2006 |title= The ACS Style Guide: Effective Communication of Scientific Information |edition= 3rd |publisher= American Chemical Society |location= Washington, D.C. |isbn= 978-0-8412-3999-9 |page= [https://archive.org/details/acsstyleguideeff0000unse/page/127 127] |url= https://archive.org/details/acsstyleguideeff0000unse/page/127 }}</ref><ref>{{cite journal |journal=Pure Appl. Chem. |volume=70 |issue=1 |last1=Coplen |pages= 237–257 |year= 1998 |first1=T. B. |url= http://old.iupac.org/reports/1998/7001coplen/history.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://old.iupac.org/reports/1998/7001coplen/history.pdf |archive-date=2022-10-09 |url-status=live |last2=Peiser |first2=H. S. |title= History of the recommended atomic-weight values from 1882 to 1997: a comparison of differences from current values to the estimated uncertainties of earlier values |doi= 10.1351/pac199870010237|s2cid=96729044 }}</ref> following ''Webster's Third New International Dictionary''.|group=note}} and [[francium]] (Fr). Together with [[hydrogen]] they constitute [[Group (periodic table)#Group names|group 1]],{{refn|In both the old IUPAC and the [[Chemical Abstracts Service|CAS]] systems for group numbering, this group is known as '''group IA''' (pronounced as "group one A", as the "I" is a [[Roman numeral]]).<ref name = fluck>{{cite journal |last1=Fluck |first1=E. |year=1988 |title=New Notations in the Periodic Table |journal=[[Pure and Applied Chemistry|Pure Appl. Chem.]] |volume=60 |issue=3 |pages=431–436 |publisher=[[International Union of Pure and Applied Chemistry|IUPAC]] |doi=10.1351/pac198860030431 |s2cid=96704008 |url=http://www.iupac.org/publications/pac/1988/pdf/6003x0431.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://www.iupac.org/publications/pac/1988/pdf/6003x0431.pdf |archive-date=2022-10-09 |url-status=live |access-date=24 March 2012}}</ref>|name=group-numbering|group=note}} which lies in the [[s-block]] of the [[periodic table]]. All alkali metals have their outermost electron in an [[atomic orbital|s-orbital]]: this shared electron configuration results in their having very similar characteristic properties.{{refn|While hydrogen also has this electron configuration, it is not considered an alkali metal as it has very different behaviour owing to the lack of [[valence electron|valence]] p-orbitals in [[period&nbsp;1 element]]s.|group=note}} Indeed, the alkali metals provide the best example of [[periodic trends|group trends]] in properties in the periodic table, with elements exhibiting well-characterised [[Homologous series|homologous]] behaviour.<ref name=rsc>{{cite web|url=http://www.rsc.org/chemsoc/visualelements/PAGES/data/intro_groupi_data.html |title=Visual Elements: Group 1 – The Alkali Metals |author=Royal Society of Chemistry |work=Visual Elements |publisher=Royal Society of Chemistry |access-date=13 January 2012 |url-status=dead |archive-url=https://web.archive.org/web/20120805145647/http://www.rsc.org/chemsoc/visualelements/PAGES/data/intro_groupi_data.html |archive-date=5 August 2012 |author-link=Royal Society of Chemistry }}</ref> This family of elements is also known as the '''lithium family''' after its leading element.

The alkali metals are all shiny, [[hardness|soft]], highly [[reactivity (chemistry)|reactive]] [[metals]] at [[standard temperature and pressure]] and readily lose their [[valence electron|outermost electron]] to form [[cations]] with [[electric charge|charge]] +1. They can all be cut easily with a knife due to their softness, exposing a shiny surface that tarnishes rapidly in air due to [[redox|oxidation]] by atmospheric moisture and [[oxygen]] (and in the case of lithium, [[nitrogen]]). Because of their high reactivity, they must be stored under oil to prevent reaction with air, and are found naturally only in [[salt (chemistry)|salts]] and never as the free elements. Caesium, the fifth alkali metal, is the most reactive of all the metals.<!--YES, NOT FRANCIUM. SEE BELOW.--> All the alkali metals react with water, with the heavier alkali metals reacting more vigorously than the lighter ones.

All of the discovered alkali metals occur in nature as their compounds: in order of [[abundance of the chemical elements|abundance]], sodium is the most abundant, followed by potassium, lithium, rubidium, caesium, and finally francium, which is very rare due to its extremely high [[radioactivity]]; francium occurs only in minute [[trace radioisotope|traces]] in nature as an intermediate step in some obscure side branches of the natural [[decay chain]]s. Experiments have been conducted to attempt the synthesis of [[ununennium|element 119]], which is likely to be the next member of the group; none were successful. However, ununennium may not be an alkali metal due to [[relativistic quantum chemistry|relativistic effects]], which are predicted to have a large influence on the chemical properties of [[superheavy element]]s; even if it does turn out to be an alkali metal, it is predicted to have some differences in physical and chemical properties from its lighter homologues.

Most alkali metals have many different applications. One of the best-known applications of the pure elements is the use of rubidium and caesium in [[atomic clock]]s, of which caesium atomic clocks form the basis of the second. A common application of the compounds of sodium is the [[sodium-vapour lamp]], which emits light very efficiently. [[Salt|Table salt]], or sodium chloride, has been used since antiquity. [[Lithium (medication)|Lithium]] finds use as a psychiatric medication and as an [[anode]] in [[lithium batteries]]. Sodium, potassium and possibly lithium are [[essential element]]s, having major biological roles as [[electrolytes]], and although the other alkali metals are not essential, they also have various effects on the body, both beneficial and harmful.
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== History ==
[[File:Petalite.jpg|thumb|alt=A sample of petalite|[[Petalite]], the lithium mineral from which lithium was first isolated]]
Sodium compounds have been known since ancient times; salt ([[sodium chloride]]) has been an important commodity in human activities, as testified by the English word ''salary'', referring to ''salarium'', money paid to Roman soldiers for the purchase of salt.<ref>{{OEtymD|salary}}</ref>{{better source needed|date=August 2020}} While [[potash]] has been used since ancient times, it was not understood for most of its history to be a fundamentally different substance from sodium mineral salts. [[Georg Ernst Stahl]] obtained experimental evidence which led him to suggest the fundamental difference of sodium and potassium salts in 1702,<ref name="1702Suspect">{{cite book|language = de |url= https://books.google.com/books?id=b-ATAAAAQAAJ&pg=PA167 |page= 167|title= Chymische Schriften|last1= Marggraf|first1= Andreas Siegmund |year= 1761}}</ref> and [[Henri-Louis Duhamel du Monceau]] was able to prove this difference in 1736.<ref>{{cite journal |url= http://gallica.bnf.fr/ark:/12148/bpt6k3533j/f73.image.r=Memoires%20de%20l%27Academie%20royale%20des%20Sciences.langEN |journal= Mémoires de l'Académie Royale des Sciences |title= Sur la Base de Sel Marine |last= du Monceau|first= H. L. D. |year= 1736 |pages= 65–68 |language= fr}}</ref> The exact chemical composition of potassium and sodium compounds, and the status as chemical element of potassium and sodium, was not known then, and thus [[Antoine Lavoisier]] did not include either alkali in his list of chemical elements in 1789.<ref name="weeks">{{cite journal |doi= 10.1021/ed009p1035|title= The discovery of the elements. IX. Three alkali metals: Potassium, sodium, and lithium |year= 1932|last1= Weeks|first1= Mary Elvira|author-link1=Mary Elvira Weeks|journal= Journal of Chemical Education|volume= 9|issue= 6|page= 1035|bibcode= 1932JChEd...9.1035W}}</ref><ref name="disco">{{cite journal |jstor= 228541|pages= 247–258|last1= Siegfried|first1= R.|title= The Discovery of Potassium and Sodium, and the Problem of the Chemical Elements|volume= 54|issue= 2|journal= Isis|year= 1963|doi= 10.1086/349704|pmid= 14147904|s2cid= 38152048}}</ref>

Pure potassium was first isolated in 1807 in England by [[Humphry Davy]], who derived it from [[Potassium hydroxide|caustic potash]] (KOH, potassium hydroxide) by the use of electrolysis of the molten salt with the newly invented [[voltaic pile]]. Previous attempts at electrolysis of the aqueous salt were unsuccessful due to potassium's extreme reactivity.<ref name="Greenwood&Earnshaw" />{{rp|68}} Potassium was the first metal that was isolated by electrolysis.<ref name=Enghag2004>{{cite book |last=Enghag|first=P.|year=2004|title=Encyclopedia of the elements|publisher=Wiley-VCH Weinheim|isbn=978-3-527-30666-4|chapter=11. Sodium and Potassium}}</ref> Later that same year, Davy reported extraction of sodium from the similar substance [[caustic soda]] (NaOH, lye) by a similar technique, demonstrating the elements, and thus the salts, to be different.<ref name="weeks" /><ref name="disco" /><ref name=Davy1807>{{cite journal |first=Humphry|last=Davy|title=On some new phenomena of chemical changes produced by electricity, in particular the decomposition of the fixed alkalies, and the exhibition of the new substances that constitute their bases; and on the general nature of alkaline bodies|pages=1–44|year=1808|volume=98|journal=Philosophical Transactions of the Royal Society of London|url=https://books.google.com/books?id=gpwEAAAAYAAJ&pg=PA57|doi=10.1098/rstl.1808.0001|doi-access=free}}</ref><ref name="200disco">{{cite journal |doi= 10.1134/S1061934807110160|title= History of the discovery of potassium and sodium (on the 200th anniversary of the discovery of potassium and sodium)|year= 2007|last1= Shaposhnik|first1= V. A.|journal= Journal of Analytical Chemistry|volume= 62|issue= 11|pages= 1100–1102|s2cid= 96141217}}</ref>

[[File:Johann Wolfgang Döbereiner.jpg|thumb|upright|[[Johann Wolfgang Döbereiner]] was among the first to notice similarities between what are now known as the alkali metals.]]
[[Petalite]] ({{chem2|LiAlSi4O10|auto=yes}}) was discovered in 1800 by the Brazilian chemist [[José Bonifácio de Andrada]] in a mine on the island of [[Utö, Sweden]].<ref name=mindat>{{cite web |url=http://www.mindat.org/min-3171.html |title=Petalite: Petalite mineral information and data |last1=Ralph |first1=Jolyon |last2=Chau |first2=Ida |date=24 August 2011 |access-date=27 November 2011}}</ref><ref name=webelementshistory>{{cite web |url=http://www.webelements.com/lithium/history.html|title=WebElements Periodic Table of the Elements {{!}} Lithium {{!}} historical information |last=Winter |first=Mark |access-date=27 November 2011}}</ref><ref name=discovery>{{cite book |title=Discovery of the Elements |last=Weeks |first=Mary |year=2003 |page=124 |publisher=Kessinger Publishing |location=Whitefish, Montana, United States |isbn=978-0-7661-3872-8 }}</ref> However, it was not until 1817 that [[Johan August Arfwedson]], then working in the laboratory of the chemist [[Jöns Jacob Berzelius]], [[discovery of the chemical elements|detected]] the presence of a new element while analysing petalite [[ore]].<ref name=uwis>{{cite web |url=http://genchem.chem.wisc.edu/lab/PTL/PTL/BIOS/arfwdson.htm |archive-url=https://web.archive.org/web/20080605152857/http://genchem.chem.wisc.edu/lab/PTL/PTL/BIOS/arfwdson.htm |archive-date=5 June 2008 |title=Johan Arfwedson |access-date=10 August 2009}}</ref><ref name=vanderkrogt>{{cite web |publisher= Elementymology & Elements Multidict|title= Lithium |first= Peter|last= van der Krogt|url= http://elements.vanderkrogt.net/element.php?sym=Li|access-date= 5 October 2010}}</ref> This new element was noted by him to form compounds similar to those of sodium and potassium, though its [[lithium carbonate|carbonate]] and [[lithium hydroxide|hydroxide]] were less [[solubility|soluble in water]] and more [[Base (chemistry)|alkaline]] than the other alkali metals.<ref name=compounds>{{cite web |url=http://www.chemguide.co.uk/inorganic/group1/compounds.html|title=Compounds of the Group 1 Elements |access-date=10 August 2009 |last=Clark |first=Jim |year=2005 |work=chemguide}}</ref> Berzelius gave the unknown material the name ''lithion''/''lithina'', from the [[Ancient Greek|Greek]] word ''λιθoς'' (transliterated as ''lithos'', meaning "stone"), to reflect its discovery in a solid mineral, as opposed to potassium, which had been discovered in plant ashes, and sodium, which was known partly for its high abundance in animal blood. He named the metal inside the material ''lithium''.<ref name=krebs>{{cite book |last= Krebs|first= Robert E.|year= 2006|title= The History and Use of Our Earth's Chemical Elements: A Reference Guide|publisher= Greenwood Press|location= Westport, Conn.|isbn= 978-0-313-33438-2}}</ref><ref name=webelementshistory /><ref name=vanderkrogt /> Lithium, sodium, and potassium were part of the discovery of [[periodic table|periodicity]], as they are among a series of triads of elements in the same [[group (periodic table)|group]] that were noted by [[Johann Wolfgang Döbereiner]] in 1850 as having similar properties.<ref name="meta-synthesis2" />

[[File:Lepidolite-76774.jpg|thumb|upright|alt=A sample of lepidolite|[[Lepidolite]], the rubidium mineral from which rubidium was first isolated]]
Rubidium and caesium were the first elements to be discovered using the [[spectroscope]], invented in 1859 by [[Robert Bunsen]] and [[Gustav Kirchhoff]].<ref name="caesium">{{cite web |url=http://pubs.acs.org/cen/80th/print/cesium.html |title=C&EN: It's Elemental: The Periodic Table – Cesium |publisher=American Chemical Society|access-date=25 February 2010|last=Kaner|first=Richard|year= 2003}}</ref> The next year, they discovered caesium in the [[mineral water]] from [[Bad Dürkheim]], Germany. Their discovery of rubidium came the following year in [[Heidelberg]], Germany, finding it in the mineral [[lepidolite]].<ref name="BuKi1861">{{cite journal |title= Chemische Analyse durch Spectralbeobachtungen |pages= 337–381 |first1= G.|last1= Kirchhoff |first2= R.|last2= Bunsen|author-link1= Gustav Kirchhoff |author-link2 = Robert Bunsen|doi= 10.1002/andp.18611890702 |journal= Annalen der Physik und Chemie |volume= 189 |issue= 7|year= 1861 |bibcode=1861AnP...189..337K|url= http://archiv.ub.uni-heidelberg.de/volltextserver/15657/1/spektral.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://archiv.ub.uni-heidelberg.de/volltextserver/15657/1/spektral.pdf |archive-date=2022-10-09 |url-status=live }}</ref> The names of rubidium and caesium come from the most prominent lines in their [[emission spectrum|emission spectra]]: a bright red line for rubidium (from the [[Latin]] word ''rubidus'', meaning dark red or bright red), and a sky-blue line for caesium (derived from the Latin word ''caesius'', meaning sky-blue).<ref name="Weeks">{{cite journal |title= The discovery of the elements. XIII. Some spectroscopic discoveries |pages= 1413–1434|last= Weeks|first= Mary Elvira |author-link=Mary Elvira Weeks|doi=10.1021/ed009p1413|journal= [[Journal of Chemical Education]] |volume= 9 |issue= 8 |year= 1932 |bibcode=1932JChEd...9.1413W}}</ref><ref>{{Cite OED|caesium|edition = 2nd}}</ref>

Around 1865 [[John Alexander Reina Newlands|John Newlands]] produced a series of papers where he listed the elements in order of increasing atomic weight and similar physical and chemical properties that recurred at intervals of eight; he likened such periodicity to the [[octave]]s of music, where notes an octave apart have similar musical functions.<ref>{{cite journal |title= On Relations Among the Equivalents |last=Newlands|first=John A. R. |journal= Chemical News |volume= 10 |pages= 94–95 |date= 20 August 1864 |url=http://web.lemoyne.edu/~GIUNTA/EA/NEWLANDSann.HTML |url-status=live |archive-url=https://web.archive.org/web/20110101073248/http://web.lemoyne.edu/~GIUNTA/EA/NEWLANDSann.HTML |archive-date=1 January 2011 |access-date=25 November 2013}}</ref><ref>{{cite journal |title= On the Law of Octaves |last=Newlands|first=John A. R. |journal= Chemical News |volume= 12 |page= 83 |date= 18 August 1865 |url=http://web.lemoyne.edu/~GIUNTA/EA/NEWLANDSann.HTML |url-status=live |archive-url=https://web.archive.org/web/20110101073248/http://web.lemoyne.edu/~GIUNTA/EA/NEWLANDSann.HTML |archive-date=1 January 2011 |access-date=25 November 2013}}</ref> His version put all the alkali metals then known (lithium to caesium), as well as copper, silver, and [[thallium]] (which show the +1 oxidation state characteristic of the alkali metals), together into a group. His table placed hydrogen with the [[halogen]]s.<ref name="meta-synthesis2" />

[[File:Mendelejevs periodiska system 1871.png|thumb|upright=1.75|[[Dmitri Mendeleev]]'s periodic system proposed in 1871 showing hydrogen and the alkali metals as part of his group I, along with copper, silver, and gold]]
After 1869, [[Dmitri Mendeleev]] proposed his periodic table placing lithium at the top of a group with sodium, potassium, rubidium, caesium, and thallium.<ref>{{cite journal |last=Mendelejew |first=Dimitri |year=1869 |title=Über die Beziehungen der Eigenschaften zu den Atomgewichten der Elemente |journal=Zeitschrift für Chemie |pages=405–406 |language=de}}</ref> Two years later, Mendeleev revised his table, placing hydrogen in group 1 above lithium, and also moving thallium to the [[boron group]]. In this 1871 version, copper, silver, and gold were placed twice, once as part of [[group 11 element|group IB]], and once as part of a "group VIII" encompassing today's groups [[group 8 element|8]] to 11.<ref name="Jensen">{{cite journal |last1=Jensen |first1=William B.|author1-link=William B. Jensen |year=2003 |title=The Place of Zinc, Cadmium, and Mercury in the Periodic Table |journal=Journal of Chemical Education |volume=80 |issue=8 |pages=952–961 |publisher=[[American Chemical Society]] |doi=10.1021/ed080p952 |bibcode=2003JChEd..80..952J |url=http://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/091.%20Zn-Cd-Hg.pdf |access-date=2012-05-06 |url-status=dead |archive-url=https://web.archive.org/web/20100611152417/http://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/091.%20Zn-Cd-Hg.pdf |archive-date=11 June 2010}}</ref><ref group="note">In the 1869 version of Mendeleev's periodic table, copper and silver were placed in their own group, aligned with hydrogen and [[mercury (element)|mercury]], while gold was tentatively placed under [[uranium]] and the undiscovered [[gallium|eka-aluminium]] in the [[boron group]].</ref> After the introduction of the 18-column table, the group IB elements were moved to their current position in the [[d-block]], while alkali metals were left in ''group IA''. Later the group's name was changed to ''group 1'' in 1988.<ref name = fluck/> The [[trivial name]] "alkali metals" comes from the fact that the hydroxides of the group 1 elements are all strong [[alkali]]s when dissolved in water.<ref name=rsc />

There were at least four erroneous and incomplete discoveries<ref name="fontani">{{cite conference |first= Marco |last= Fontani |title= The Twilight of the Naturally-Occurring Elements: Moldavium (Ml), Sequanium (Sq) and Dor (Do) |book-title= International Conference on the History of Chemistry |pages= 1–8 |date= 10 September 2005 |location= Lisbon|url= http://5ichc-portugal.ulusofona.pt/uploads/PaperLong-MarcoFontani.doc |archive-url= https://web.archive.org/web/20060224090117/http://5ichc-portugal.ulusofona.pt/uploads/PaperLong-MarcoFontani.doc |archive-date=24 February 2006|access-date= 8 April 2007}}</ref><ref name="vanderkrogt-Fr" /><ref>{{cite news |title= Education: Alabamine & Virginium|magazine=[[Time (magazine)|Time]] |date= 15 February 1932|url= http://www.time.com/time/magazine/article/0,9171,743159,00.html |archive-url= https://web.archive.org/web/20070930015028/http://www.time.com/time/magazine/article/0,9171,743159,00.html |url-status= dead |archive-date= 30 September 2007 |access-date= 1 April 2007|url-access=subscription }}</ref><ref>{{cite journal |last= MacPherson |first= H. G. |title= An Investigation of the Magneto-Optic Method of Chemical Analysis |journal= Physical Review |volume= 47 |issue= 4 |pages= 310–315 |publisher= American Physical Society|year=1934|doi= 10.1103/PhysRev.47.310|bibcode= 1935PhRv...47..310M}}</ref> before [[Marguerite Perey]] of the [[Curie Institute (Paris)|Curie Institute]] in Paris, France discovered francium in 1939 by purifying a sample of [[actinium-227]], which had been reported to have a decay energy of 220&nbsp;[[electronvolt|keV]]. However, Perey noticed decay particles with an energy level below 80&nbsp;keV. Perey thought this decay activity might have been caused by a previously unidentified decay product, one that was separated during purification, but emerged again out of the pure [[actinium]]-227. Various tests eliminated the possibility of the unknown element being [[thorium]], [[radium]], lead, [[bismuth]], or [[thallium]]. The new product exhibited chemical properties of an alkali metal (such as coprecipitating with caesium salts), which led Perey to believe that it was element 87, caused by the [[alpha decay]] of actinium-227.<ref name="chemeducator">Adloff, Jean-Pierre; Kaufman, George B. (25 September 2005). [http://chemeducator.org/sbibs/s0010005/spapers/1050387gk.htm Francium (Atomic Number 87), the Last Discovered Natural Element] {{webarchive |url=https://web.archive.org/web/20130604212956/http://chemeducator.org/sbibs/s0010005/spapers/1050387gk.htm |date=4 June 2013 }}. ''The Chemical Educator'' '''10''' (5). Retrieved 26 March 2007.</ref> Perey then attempted to determine the proportion of [[beta decay]] to alpha decay in actinium-227. Her first test put the alpha branching at 0.6%, a figure that she later revised to 1%.<ref name="mcgraw">{{cite book |contribution= Francium |year= 2002 |title= McGraw-Hill Encyclopedia of Science & Technology |volume= 7 |pages= [https://archive.org/details/mcgrawhillencycl165newy/page/493 493–494] |publisher= McGraw-Hill Professional |isbn= 978-0-07-913665-7 |title-link= McGraw-Hill Encyclopedia of Science & Technology }}</ref>
:{{nuclide|actinium|227}} {{overunderset|→|α (1.38%)|21.77 y}} '''{{nuclide|francium|223}}''' {{overunderset|→|β<sup>−</sup>|22 min}} {{nuclide|radium|223}} {{overunderset|→|α|11.4 d}}{{nuclide|radon|219}}

The next element below francium ([[Mendeleev's predicted elements|eka]]-francium) in the periodic table would be [[ununennium]] (Uue), element 119.<ref name="Uue" />{{rp|1729–1730}} The synthesis of ununennium was first attempted in 1985 by bombarding a target of [[einsteinium]]-254 with [[calcium]]-48 ions at the superHILAC accelerator at the [[Lawrence Berkeley National Laboratory]] in Berkeley, California. No atoms were identified, leading to a limiting yield of 300 [[barn (unit)|nb]].<ref name="link">{{cite journal |title=Search for superheavy elements using <sup>48</sup>Ca + <sup>254</sup>Es<sup>g</sup> reaction|first1=R. W. |last1=Lougheed|first2=J. H.|last2=Landrum|first3=E. K.|last3=Hulet|first4=J. F.|last4=Wild|first5=R. J.|last5=Dougan|first6=A. D.|last6=Dougan|first7=H.|last7=Gäggeler|first8=M.|last8=Schädel|first9=K. J.|last9=Moody|first10=K. E.|last10=Gregorich|last11=Seaborg|journal=Physical Review C|year=1985|pages=1760–1763|volume=32|issue=5|bibcode= 1985PhRvC..32.1760L|doi=10.1103/PhysRevC.32.1760|pmid=9953034 |first11=G.}}</ref><ref name="vanderkrogt-uue">{{cite web |publisher= Elementymology & Elements Multidict|title= Ununennium |first= Peter|last= van der Krogt|url= http://elements.vanderkrogt.net/element.php?sym=Uue|access-date= 14 February 2011}}</ref>

:{{nuclide|einsteinium|254|link=y}} + {{nuclide|calcium|48|link=y}} → {{nuclide|ununennium|302}}* → ''no atoms''<ref group="note">The [[asterisk]] denotes an [[excited state]].</ref>

It is highly unlikely<ref name="link" /> that this reaction will be able to create any atoms of ununennium in the near future, given the extremely difficult task of making sufficient amounts of einsteinium-254, which is favoured for production of [[superheavy element|ultraheavy elements]] because of its large mass, relatively long half-life of 270 days, and availability in significant amounts of several micrograms,<ref>{{cite journal |last1=Schadel|first1=M.|last2=Brüchle|first2=W.|last3=Brügger|first3=M.|last4=Gäggeler|first4=H.|last5=Moody|first5=K.|last6=Schardt|first6=D.|last7=Sümmerer|first7=K.|last8=Hulet|first8=E.|last9=Dougan|first9=A.|last10=Dougan|title=Heavy isotope production by multinucleon transfer reactions with <sup>254</sup>Es|journal=Journal of the Less Common Metals|volume=122|pages=411–417|year=1986|doi=10.1016/0022-5088(86)90435-2|first10=R. J.|last11=Landrum|first11=J. H.|last12=Lougheed|first12=R. W.|last13=Wild|first13=J. F.|last14=O'Kelley|first14=G. D.|last15=Hahn|first15=R. L.|display-authors=9|url=https://zenodo.org/record/1253958}}</ref> to make a large enough target to increase the sensitivity of the experiment to the required level; einsteinium has not been found in nature and has only been produced in laboratories, and in quantities smaller than those needed for effective synthesis of superheavy elements. However, given that ununennium is only the first [[period 8 element]] on the [[extended periodic table]], it may well be discovered in the near future through other reactions, and indeed an attempt to synthesise it is currently ongoing in Japan.<ref name=Enyo>{{Cite news|url=https://www.chemistryworld.com/news/hunt-for-element-119-to-begin-this-year/3007977.article|title=Hunt for element 119 set to begin|newspaper=Chemistry World|date=12 September 2017|access-date=9 January 2018}}</ref> Currently, none of the period 8 elements has been discovered yet, and it is also possible, due to [[nucleon drip line|drip instabilities]], that only the lower period 8 elements, up to around element 128, are physically possible.<ref name=EB>{{cite encyclopedia |last=Seaborg|first=G. T.|url=https://www.britannica.com/EBchecked/topic/603220/transuranium-element|title=transuranium element (chemical element)|encyclopedia=Encyclopædia Britannica|date=c. 2006|access-date=16 March 2010}}</ref><ref name="emsley">{{cite book |last=Emsley|first=John|title=Nature's Building Blocks: An A-Z Guide to the Elements|edition=New|year=2011|publisher=Oxford University Press|location=New York, NY|isbn=978-0-19-960563-7|page=593}}</ref> No attempts at synthesis have been made for any heavier alkali metals: due to their extremely high atomic number, they would require new, more powerful methods and technology to make.<ref name="Uue" />{{rp|1737–1739}}

== Occurrence ==

=== In the Solar System ===
[[File:SolarSystemAbundances.png|thumb|upright=2.5|Estimated abundances of the chemical elements in the Solar system. Hydrogen and helium are most common, from the [[Big Bang]]. The next three elements (lithium, [[beryllium]], and [[boron]]) are rare because they are poorly synthesised in the Big Bang and also in stars. The two general trends in the remaining stellar-produced elements are: (1) an alternation of abundance in elements as they have even or odd atomic numbers, and (2) a general decrease in abundance, as elements become heavier. Iron is especially common because it represents the minimum-energy nuclide that can be made by fusion of helium in supernovae.<ref name=lodders>{{cite journal |last1= Lodders |first1= Katharina|author1-link=Katharina Lodders |year= 2003 |title= Solar System Abundances and Condensation Temperatures of the Elements |journal= The Astrophysical Journal |volume= 591 |issue= 2 |pages= 1220–1247 |bibcode= 2003ApJ...591.1220L |doi= 10.1086/375492|doi-access= free }}</ref>]]
The [[Oddo–Harkins rule]] holds that elements with even atomic numbers are more common that those with odd atomic numbers, with the exception of hydrogen. This rule argues that elements with odd atomic numbers have one unpaired proton and are more likely to capture another, thus increasing their atomic number. In elements with even atomic numbers, protons are paired, with each member of the pair offsetting the spin of the other, enhancing stability.<ref name=oddo>{{cite journal |doi= 10.1002/zaac.19140870118 |title= Die Molekularstruktur der radioaktiven Atome |year= 1914 |last1= Oddo |first1= Giuseppe |journal= Zeitschrift für Anorganische Chemie |volume= 87 |pages= 253–268|url= https://www.academia.edu/11043300 }}</ref><ref name=harkins>{{cite journal |doi= 10.1021/ja02250a002 |year= 1917 |last1= Harkins |first1= William D. |journal= Journal of the American Chemical Society |volume= 39 |issue= 5 |pages= 856–879 |title= The Evolution of the Elements and the Stability of Complex Atoms. I. A New Periodic System Which Shows a Relation Between the Abundance of the Elements and the Structure of the Nuclei of Atoms|url= https://zenodo.org/record/1429060 }}</ref><ref name=north>{{cite book |last=North|first=John|title=Cosmos an illustrated history of astronomy and cosmology|year=2008|publisher=Univ. of Chicago Press|isbn=978-0-226-59441-5|page=602|url=https://books.google.com/books?id=qq8Luhs7rTUC&q=%22william+draper+harkins%22+oddo&pg=PA602|edition=Rev. and updated}}</ref> All the alkali metals have odd atomic numbers and they are not as common as the elements with even atomic numbers adjacent to them (the [[noble gas]]es and the [[alkaline earth metal]]s) in the Solar System. The heavier alkali metals are also less abundant than the lighter ones as the alkali metals from rubidium onward can only be synthesised in [[supernova]]e and not in [[stellar nucleosynthesis]]. Lithium is also much less abundant than sodium and potassium as it is poorly synthesised in both [[Big Bang nucleosynthesis]] and in stars: the Big Bang could only produce trace quantities of lithium, [[beryllium]] and [[boron]] due to the absence of a stable nucleus with 5 or 8 [[nucleon]]s, and stellar nucleosynthesis could only pass this bottleneck by the [[triple-alpha process]], fusing three helium nuclei to form [[carbon]], and skipping over those three elements.<ref name=lodders />

=== On Earth ===
[[File:Spodumene-usa59abg.jpg|thumb|upright|[[Spodumene]], an important lithium mineral]]
The Earth formed from the same cloud of matter that formed the Sun, but the planets acquired different compositions during the [[formation and evolution of the solar system]]. In turn, the [[history of Earth|natural history of the Earth]] caused parts of this planet to have differing concentrations of the elements. The mass of the Earth is approximately 5.98{{e|24}}&nbsp;kg. It is composed mostly of iron (32.1%), [[oxygen]] (30.1%), [[silicon]] (15.1%), [[magnesium]] (13.9%), [[sulfur]] (2.9%), [[nickel]] (1.8%), [[calcium]] (1.5%), and aluminium (1.4%); with the remaining 1.2% consisting of trace amounts of other elements. Due to [[planetary differentiation]], the core region is believed to be primarily composed of iron (88.8%), with smaller amounts of nickel (5.8%), sulfur (4.5%), and less than 1% trace elements.<ref name=pnas71_12_6973>{{cite journal |last1=Morgan|first1=J. W. |last2=Anders|first2=E. |title=Chemical composition of Earth, Venus, and Mercury |journal=Proceedings of the National Academy of Sciences |year=1980 |volume=77 |issue=12 |pages=6973–6977 |doi=10.1073/pnas.77.12.6973 |pmid=16592930 |pmc=350422 |bibcode= 1980PNAS...77.6973M|doi-access=free }}</ref>

The alkali metals, due to their high reactivity, do not occur naturally in pure form in nature. They are [[Goldschmidt classification|lithophiles]] and therefore remain close to the Earth's surface because they combine readily with [[oxygen]] and so associate strongly with [[silica]], forming relatively low-density minerals that do not sink down into the Earth's core. Potassium, rubidium and caesium are also [[incompatible element]]s due to their large [[ionic radius|ionic radii]].<ref name="albarede">{{cite book |title= Geochemistry: an introduction |url= https://books.google.com/books?id=doVGzreGq14C&pg=PA17 |publisher= Cambridge University Press |year= 2003 |isbn= 978-0-521-89148-6 |first= Francis |last= Albarède}}</ref>

Sodium and potassium are very abundant on Earth, both being among the ten [[abundance of elements in Earth's crust|most common elements in Earth's crust]];<ref name="webelements-occurrence">{{cite web |url= http://www.webelements.com/webelements/properties/text/image-flash/abund-crust.html|title= Abundance in Earth's Crust|publisher= WebElements.com|access-date= 14 April 2007}}</ref><ref name="IsraelScience&Technology">{{cite web |url= https://www.science.co.il/elements/?s=Earth|title= List of Periodic Table Elements Sorted by Abundance in Earth's crust|publisher= Israel Science and Technology Directory}}</ref> sodium makes up approximately 2.6% of the Earth's crust measured by weight, making it the [[Abundance of the chemical elements|sixth most abundant element]] overall<ref name="RubberBible86th">{{RubberBible86th}}</ref> and the most abundant alkali metal. Potassium makes up approximately 1.5% of the Earth's crust and is the seventh most abundant element.<ref name="RubberBible86th" /> Sodium is found in many different minerals, of which the most common is ordinary salt (sodium chloride), which occurs in vast quantities dissolved in seawater. Other solid deposits include [[Halite (mineral)|halite]], [[amphibole]], [[cryolite]], [[nitratine]], and [[zeolite]].<ref name="RubberBible86th" /> Many of these solid deposits occur as a result of ancient seas evaporating, which still occurs now in places such as [[Utah]]'s [[Great Salt Lake]] and the [[Dead Sea]].<ref name="Greenwood&Earnshaw" />{{rp|69}} Despite their near-equal abundance in Earth's crust, sodium is far more common than potassium in the ocean, both because potassium's larger size makes its salts less soluble, and because potassium is bound by silicates in soil and what potassium leaches is absorbed far more readily by plant life than sodium.<ref name="Greenwood&Earnshaw" />{{rp|69}}

Despite its chemical similarity, lithium typically does not occur together with sodium or potassium due to its smaller size.<ref name="Greenwood&Earnshaw" />{{rp|69}} Due to its relatively low reactivity, it can be found in seawater in large amounts; it is estimated that lithium concentration in seawater is approximately 0.14 to 0.25 parts per million (ppm)<ref>{{cite web |url=http://www.ioes.saga-u.ac.jp/ioes-study/li/lithium/occurence.html |title=Lithium Occurrence |access-date=13 March 2009 |publisher=Institute of Ocean Energy, Saga University, Japan |url-status=dead |archive-url=https://web.archive.org/web/20090502142924/http://www.ioes.saga-u.ac.jp/ioes-study/li/lithium/occurence.html |archive-date=2 May 2009}}</ref><ref name=enc>{{cite web|url=http://www.enclabs.com/lithium.html|access-date=15 October 2010|title=Some Facts about Lithium|publisher=ENC Labs|archive-date=10 July 2011|archive-url=https://web.archive.org/web/20110710191644/http://www.enclabs.com/lithium.html|url-status=dead}}</ref> or 25 [[micromolar]].<ref>{{cite book |doi=10.1007/3-540-13534-0_3|chapter=Extraction of metals from sea water|volume= 124/1984|pages= 91–133|last=Schwochau|first=Klaus|year=1984|series=Topics in Current Chemistry|isbn=978-3-540-13534-0|title=Inorganic Chemistry|s2cid=93866412 }}</ref> Its diagonal relationship with magnesium often allows it to replace magnesium in [[ferromagnesium]] minerals, where its crustal concentration is about 18&nbsp;[[parts per million|ppm]], comparable to that of [[gallium]] and [[niobium]]. Commercially, the most important lithium mineral is [[spodumene]], which occurs in large deposits worldwide.<ref name="Greenwood&Earnshaw" />{{rp|69}}

Rubidium is approximately as abundant as [[zinc]] and more abundant than copper. It occurs naturally in the minerals [[leucite]], [[pollucite]], [[carnallite]], [[zinnwaldite]], and [[lepidolite]],<ref>{{cite journal |title= Trace element chemistry of lithium-rich micas from rare-element granitic pegmatites |volume= 55
|issue= 13 |year= 1995 |doi= 10.1007/BF01162588 |pages= 203–215 |journal= Mineralogy and Petrology |first= M. A. |last= Wise |bibcode= 1995MinPe..55..203W|s2cid= 140585007
}}</ref> although none of these contain only rubidium and no other alkali metals.<ref name="Greenwood&Earnshaw" />{{rp|70}} Caesium is more abundant than some commonly known elements, such as [[antimony]], [[cadmium]], [[tin]], and [[tungsten]], but is much less abundant than rubidium.<ref name="pubs.usgs" />

[[Francium-223]], the only naturally occurring isotope of francium,<ref name="atomicweights2007" /><ref name="atomicweights2009" /> is the [[decay product|product]] of the [[alpha decay]] of actinium-227 and can be found in trace amounts in [[uranium]] minerals.<ref name="CRC2006">{{cite book |year= 2006 |title= CRC Handbook of Chemistry and Physics |volume= 4|page= 12|publisher= CRC|isbn= 978-0-8493-0474-3}}</ref> In a given sample of uranium, there is estimated to be only one francium atom for every 10<sup>18</sup> uranium atoms.<ref name="nbb">{{cite book |last= Emsley|url=https://books.google.com/books?id=Yhi5X7OwuGkC&pg=PA151 |first= John |title= Nature's Building Blocks |publisher= Oxford University Press |year= 2001 |location= Oxford |pages= 151–153 |isbn= 978-0-19-850341-5}}</ref><ref name="elemental">{{cite web |last= Gagnon |first= Steve |title= Francium |publisher= Jefferson Science Associates, LLC |url= http://education.jlab.org/itselemental/ele087.html |access-date= 1 April 2007 |archive-url= https://web.archive.org/web/20070331235139/http://education.jlab.org/itselemental/ele087.html |archive-date= 31 March 2007  |url-status= live}}</ref> It has been calculated that there are at most 30&nbsp;grams of francium in the [[crust (geology)|earth's crust]] at any time, due to its extremely short [[half-life]] of 22 minutes.<ref name="Winter" /><ref name="itselemental">{{cite web |url= http://education.jlab.org/itselemental/index.html|title= It's Elemental&nbsp;— The Periodic Table of Elements|publisher= Jefferson Lab|access-date= 14 April 2007 |archive-url= https://web.archive.org/web/20070429032414/http://education.jlab.org/itselemental/index.html |archive-date= 29 April 2007  |url-status= live}}</ref>

== Properties ==

=== Physical and chemical ===
The physical and chemical properties of the alkali metals can be readily explained by their having an ns<sup>1</sup> valence [[electron configuration]], which results in weak [[metallic bonding]]. Hence, all the alkali metals are soft and have low [[density|densities]],<ref name=rsc /> [[melting point|melting]]<ref name=rsc /> and [[boiling point]]s,<ref name=rsc /> as well as [[heat of sublimation|heats of sublimation]], [[heat of vaporization|vaporisation]], and [[dissociation (chemistry)|dissociation]].<ref name="Greenwood&Earnshaw" />{{rp|74}} They all crystallise in the [[body-centered cubic]] crystal structure,<ref name="Greenwood&Earnshaw" />{{rp|73}} and have distinctive [[flame test|flame colours]] because their outer s electron is very easily excited.<ref name="Greenwood&Earnshaw" />{{rp|75}} Indeed, these flame test colours are the most common way of identifying them since all their salts with common ions are soluble.<ref name="Greenwood&Earnshaw" />{{rp|75}} The ns<sup>1</sup> configuration also results in the alkali metals having very large [[atomic radius|atomic]] and [[ionic radius|ionic radii]], as well as very high [[thermal conductivity|thermal]] and [[electrical conductivity]].<ref name="Greenwood&Earnshaw" />{{rp|75}} Their chemistry is dominated by the loss of their lone valence electron in the outermost s-orbital to form the +1 oxidation state, due to the ease of ionising this electron and the very high second ionisation energy.<ref name="Greenwood&Earnshaw" />{{rp|76}} Most of the chemistry has been observed only for the first five members of the group. The chemistry of francium is not well established due to its extreme [[radioactive decay|radioactivity]];<ref name=rsc /> thus, the presentation of its properties here is limited. What little is known about francium shows that it is very close in behaviour to caesium, as expected. The physical properties of francium are even sketchier because the bulk element has never been observed; hence any data that may be found in the literature are certainly speculative extrapolations.<ref name=RubberBible84th />
{| class="wikitable"
|+ Properties of the alkali metals<ref name="Greenwood&Earnshaw" />{{rp|75}}<ref name=generalchemistry />
! Name
! [[Lithium]]
! [[Sodium]]
! [[Potassium]]
! [[Rubidium]]
! [[Caesium]]
! [[Francium]]
|-
| style="background:lightgrey; text-align:left;"|[[Atomic number]]
| 3 || 11 || 19 || 37 || 55 || 87
|-
| style="background:lightgrey; text-align:left;"|[[Standard atomic weight]]{{refn|The number given in [[bracket|parentheses]] refers to the [[standard uncertainty|measurement uncertainty]]. This uncertainty applies to the [[significant figure|least significant figure]](s) of the number prior to the parenthesised value (ie. counting from rightmost digit to left). For instance, {{val|1.00794|(7)}} stands for {{val|1.00794|0.00007}}, while {{val|1.00794|(72)}} stands for {{val|1.00794|0.00072}}.<ref>{{cite web |url=http://physics.nist.gov/cgi-bin/cuu/Info/Constants/definitions.html|title=Standard Uncertainty and Relative Standard Uncertainty|work=[[CODATA]] reference|publisher=[[National Institute of Standards and Technology]]|access-date=26 September 2011}}</ref>|group=note}}<ref name="atomicweights2007">{{cite journal |last1=Wieser |first1=Michael E. |last2=Berglund |first2=Michael |year=2009 |title=Atomic weights of the elements 2007 (IUPAC Technical Report) |journal=[[Pure and Applied Chemistry|Pure Appl. Chem.]] |volume=81 |issue=11 |pages= 2131–2156 |publisher=[[International Union of Pure and Applied Chemistry|IUPAC]] |doi=10.1351/PAC-REP-09-08-03 |s2cid=98084907 |url=http://iupac.org/publications/pac/pdf/2009/pdf/8111x2131.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://iupac.org/publications/pac/pdf/2009/pdf/8111x2131.pdf |archive-date=2022-10-09 |url-status=live |access-date=7 February 2012}}</ref><ref name="atomicweights2009">{{cite journal |last1=Wieser |first1=Michael E. |last2=Coplen |first2=Tyler B. |year=2011 |title=Atomic weights of the elements 2009 (IUPAC Technical Report) |journal=[[Pure and Applied Chemistry|Pure Appl. Chem.]] |volume=83 |issue=2 |pages=359–396 |publisher=[[International Union of Pure and Applied Chemistry|IUPAC]] |doi=10.1351/PAC-REP-10-09-14 |s2cid=95898322 |url=http://iupac.org/publications/pac/pdf/2011/pdf/8302x0359.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://iupac.org/publications/pac/pdf/2011/pdf/8302x0359.pdf |archive-date=2022-10-09 |url-status=live |access-date=11 February 2012}}</ref>
| 6.94(1){{refn|The value listed is the conventional value suitable for trade and commerce; the actual value may range from 6.938 to 6.997 depending on the isotopic composition of the sample.<ref name="atomicweights2009" />|group=note}} || 22.98976928(2) || 39.0983(1) || 85.4678(3) || 132.9054519(2) || [223]{{refn|The element does not have any stable [[nuclide]]s, and a value in brackets indicates the [[mass number]] of the longest-lived [[isotope]] of the element.<ref name="atomicweights2007" /><ref name="atomicweights2009" />|group=note}}
|-
| style="background:lightgrey; text-align:left;"|[[Electron configuration]]
| &#91;[[Helium|He]]&#93; 2s<sup>1</sup> || &#91;[[Neon|Ne]]&#93; 3s<sup>1</sup> || &#91;[[Argon|Ar]]&#93; 4s<sup>1</sup> || &#91;[[Krypton|Kr]]&#93; 5s<sup>1</sup> || &#91;[[Xenon|Xe]]&#93; 6s<sup>1</sup> || &#91;[[Radon|Rn]]&#93; 7s<sup>1</sup>
|-
| style="background:lightgrey; text-align:left;"|[[Melting point]] (°C)
| 180.54 || 97.72|| 63.38 || 39.31 || 28.44 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Boiling point]] (°C)
| 1342 || 883 || 759 || 688 || 671 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Density]]&nbsp;(g·cm<sup>−3</sup>)
| 0.534 || 0.968 || 0.89 || 1.532 || 1.93 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Heat of fusion]]&nbsp;(kJ·mol<sup>−1</sup>)
| 3.00 || 2.60 || 2.321 || 2.19 || 2.09 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Heat of vaporization|Heat of vaporisation]]&nbsp;(kJ·mol<sup>−1</sup>)
| 136 || 97.42 || 79.1 || 69 || 66.1 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Heat of formation]]&nbsp;of&nbsp;monatomic&nbsp;gas&nbsp;(kJ·mol<sup>−1</sup>)
| 162 || 108 || 89.6 || 82.0 || 78.2 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Electrical resistivity]]&nbsp;at&nbsp;25&nbsp;°C&nbsp;(n[[ohm|Ω]]·cm)
| 94.7 || 48.8 || 73.9 || 131 || 208 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Atomic radius]]&nbsp;([[picometer|pm]])
| 152 || 186 || 227 || 248 || 265 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Ionic radius]] of hexacoordinate M<sup>+</sup> ion&nbsp;(pm)
| 76 || 102 || 138 || 152 || 167 ||?
|-
| style="background:lightgrey; text-align:left;"|First [[ionization energy|ionisation energy]]&nbsp;([[kilojoule per mole|kJ·mol<sup>−1</sup>]])
| 520.2 || 495.8 || 418.8 || 403.0 || 375.7 || 392.8<ref name="andreev">{{cite journal |last1=Andreev|first1=S.V. |last2=Letokhov|first2=V.S. |last3=Mishin|first3=V.I. |title= Laser resonance photoionization spectroscopy of Rydberg levels in Fr |journal= [[Physical Review Letters|Phys. Rev. Lett.]] |year= 1987 |volume= 59 |pages= 1274–76 |doi= 10.1103/PhysRevLett.59.1274 |pmid=10035190 |bibcode=1987PhRvL..59.1274A |issue= 12}}</ref>
|-
| style="background:lightgrey; text-align:left;"|[[Electron affinity]]&nbsp;(kJ·mol<sup>−1</sup>)
| 59.62 || 52.87 || 48.38 || 46.89 || 45.51 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Enthalpy of dissociation]]&nbsp;of&nbsp;M<sub>2</sub>&nbsp;(kJ·mol<sup>−1</sup>)
| 106.5 || 73.6 || 57.3 || 45.6 || 44.77 ||?
|-
| style="background:lightgrey; text-align:left;"|Pauling [[electronegativity]]
| 0.98 || 0.93 || 0.82 || 0.82 || 0.79 ||?{{refn|[[Linus Pauling]] estimated the electronegativity of francium at 0.7 on the [[Pauling scale]], the same as caesium;<ref>{{cite book |last= Pauling |first= Linus |title= The Nature of the Chemical Bond|url= https://archive.org/details/natureofchemical00paul |url-access= registration |edition= Third |author-link= Linus Pauling |publisher= Cornell University Press |year= 1960 |isbn= 978-0-8014-0333-0 |page= [https://archive.org/details/natureofchemical00paul/page/93 93]}}</ref> the value for caesium has since been refined to 0.79, although there are no experimental data to allow a refinement of the value for francium.<ref>{{cite journal |last=Allred|first=A. L. |year= 1961 |journal= J. Inorg. Nucl. Chem.|volume= 17 |issue= 3–4 |pages= 215–221 |title= Electronegativity values from thermochemical data |doi= 10.1016/0022-1902(61)80142-5}}</ref> Francium has a slightly higher ionisation energy than caesium,<ref name="andreev" /> 392.811(4)&nbsp;kJ/mol as opposed to 375.7041(2)&nbsp;kJ/mol for caesium, as would be expected from [[relativistic quantum chemistry|relativistic effects]], and this would imply that caesium is the less electronegative of the two.|name=Fr-electronegativity|group=note}}
|-
| style="background:lightgrey; text-align:left;"|Allen [[electronegativity]]
|0.91
|0.87
|0.73
|0.71
|0.66
|0.67
|-
| style="background:lightgrey; text-align:left;"|[[Standard electrode potential]] (''E''°(M<sup>+</sup>→M<sup>0</sup>); [[volt|V]])<ref name=van92>Vanýsek, Petr (2011). [http://www.hbcpnetbase.com/articles/05_22_92.pdf “Electrochemical Series”], in [http://www.hbcpnetbase.com/ ''Handbook of Chemistry and Physics: 92nd Edition''] {{Webarchive|url=https://web.archive.org/web/20170724011402/http://www.hbcpnetbase.com/ |date=24 July 2017 }} (Chemical Rubber Company).</ref>
| −3.04 || −2.71 || −2.93 || −2.98 || −3.03 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Flame test]] colour<br />Principal emission/absorption wavelength&nbsp;([[nanometer|nm]])
| Crimson<br />670.8 || Yellow<br />589.2 || Violet<br />766.5 || Red-violet<br />780.0 || Blue<br />455.5 ||?
|}
The alkali metals are more similar to each other than the elements in any other [[group (periodic table)|group]] are to each other.<ref name=rsc /> Indeed, the similarity is so great that it is quite difficult to separate potassium, rubidium, and caesium, due to their similar [[ionic radius|ionic radii]]; lithium and sodium are more distinct. For instance, when moving down the table, all known alkali metals show increasing [[atomic radius]],<ref name=chemguide /> decreasing [[electronegativity]],<ref name=chemguide>{{cite web |url=http://www.chemguide.co.uk/inorganic/group1/properties.html |title=Atomic and Physical Properties of the Group 1 Elements |last=Clark|first=Jim |year=2005 |work=chemguide |access-date=30 January 2012}}</ref> increasing [[Reactivity (chemistry)|reactivity]],<ref name=rsc /> and decreasing melting and boiling points<ref name=chemguide /> as well as heats of fusion and vaporisation.<ref name="Greenwood&Earnshaw" />{{rp|75}} In general, their [[density|densities]] increase when moving down the table, with the exception that potassium is less dense than sodium.<ref name=chemguide /> One of the very few properties of the alkali metals that does not display a very smooth trend is their [[reduction potential]]s: lithium's value is anomalous, being more negative than the others.<ref name="Greenwood&Earnshaw" />{{rp|75}} This is because the Li<sup>+</sup> ion has a very high [[hydration energy]] in the gas phase: though the lithium ion disrupts the structure of water significantly, causing a higher change in entropy, this high hydration energy is enough to make the reduction potentials indicate it as being the most electropositive alkali metal, despite the difficulty of ionising it in the gas phase.<ref name="Greenwood&Earnshaw" />{{rp|75}}

The stable alkali metals are all silver-coloured metals except for caesium, which has a pale golden tint:<ref name="theodoregray-caesium">{{cite web |url=http://www.theodoregray.com/periodictable/Elements/055/index.s7.html |title=Facts, pictures, stories about the element Cesium in the Periodic Table |last=Gray|first=Theodore|author-link=Theodore Gray|work=The Wooden Periodic Table Table |access-date=13 January 2012}}</ref> it is one of only three metals that are clearly coloured (the other two being copper and gold).<ref name="Greenwood&Earnshaw" />{{rp|74}} Additionally, the heavy [[alkaline earth metal]]s [[calcium]], [[strontium]], and [[barium]], as well as the divalent [[lanthanide]]s [[europium]] and [[ytterbium]], are pale yellow, though the colour is much less prominent than it is for caesium.<ref name="Greenwood&Earnshaw" />{{rp|74}} Their lustre tarnishes rapidly in air due to oxidation.<ref name=rsc />

[[File:Potassium water 20.theora.ogv|thumb|right|Potassium reacts violently with water at room temperature]]
[[File:Cesium water.theora.ogv|thumb|right|Caesium reacts explosively with water even at low temperatures]]
All the alkali metals are highly reactive and are never found in elemental forms in nature.<ref name="krebs" /> Because of this, they are usually stored in [[mineral oil]] or [[kerosene]] (paraffin oil).<ref name="OU">{{cite web |url=http://www.open.edu/openlearn/science-maths-technology/science/chemistry/alkali-metals |title=Alkali metals |author=The OpenLearn team |year=2012 |work=OpenLearn |publisher=The Open University |access-date=9 July 2012}}</ref> They react aggressively with the [[halogen]]s to form the [[alkali metal halide]]s, which are white [[ionic crystal]]line compounds that are all [[solubility|soluble]] in water except [[lithium fluoride]] (LiF).<ref name=rsc /> The alkali metals also react with water to form strongly [[alkali]]ne [[hydroxide]]s and thus should be handled with great care. The heavier alkali metals react more vigorously than the lighter ones; for example, when dropped into water, caesium produces a larger explosion than potassium if the same number of moles of each metal is used.<ref name=rsc /><ref name="alkalibangs">{{cite web |last=Gray|first=Theodore|title=Alkali Metal Bangs|url=http://www.theodoregray.com/periodictable/AlkaliBangs/index.html|publisher=[[Theodore Gray]]|access-date=13 May 2012}}</ref><ref name="pubs.usgs" /> The alkali metals have the lowest first [[ionization energy|ionisation energies]] in their respective periods of the [[periodic table]]<ref name="RubberBible84th">{{cite book |editor= Lide, D. R. |title= CRC Handbook of Chemistry and Physics |edition= 84th |location= Boca Raton, FL |publisher= CRC Press |year= 2003}}</ref> because of their low [[effective nuclear charge]]<ref name=rsc /> and the ability to attain a [[noble gas]] configuration by losing just one [[electron]].<ref name=rsc /> Not only do the alkali metals react with water, but also with proton donors like [[Alcohol (chemistry)|alcohol]]s and [[phenols]], gaseous [[ammonia]], and [[alkyne]]s, the last demonstrating the phenomenal degree of their reactivity. Their great power as reducing agents makes them very useful in liberating other metals from their oxides or halides.<ref name="Greenwood&Earnshaw" />{{rp|76}}

The second ionisation energy of all of the alkali metals is very high<ref name=rsc /><ref name="RubberBible84th" /><!--the second ionisation energy for francium is not given in [[ionization energies of the elements (data page)]]--> as it is in a full shell that is also closer to the nucleus;<ref name=rsc /> thus, they almost always lose a single electron, forming cations.<ref name="Greenwood&Earnshaw" />{{rp|28}} The [[alkalide]]s are an exception: they are unstable compounds which contain alkali metals in a −1 oxidation state, which is very unusual as before the discovery of the alkalides, the alkali metals were not expected to be able to form [[anion]]s and were thought to be able to appear in [[salt (chemistry)|salts]] only as cations. The alkalide anions have filled [[s-orbital|s-subshells]], which gives them enough stability to exist. All the stable alkali metals except lithium are known to be able to form alkalides,<ref>{{cite journal |journal= [[J. Am. Chem. Soc.]] |title= Crystalline salt of the sodium anion (Na<sup>−</sup>) |year= 1974 |volume= 96 |issue= 2 |pages= 608–609 |doi= 10.1021/ja00809a060|last1= Dye |first1= James L. |last2= Ceraso |first2= Joseph M. |last3= Lok |first3= Mei |last4= Barnett |first4= B. L. |last5= Tehan |first5= Frederick J. }}</ref><ref>{{cite journal |title= Alkali anions. Preparation and crystal structure of a compound which contains the cryptated sodium cation and the sodium anion |journal= [[J. Am. Chem. Soc.]] |year= 1974 |volume= 96 |issue= 23 |pages= 7203–7208 |doi= 10.1021/ja00830a005|last1= Tehan |first1= Frederick J. |last2= Barnett |first2= B. L. |last3= Dye |first3= James L. }}</ref><ref>{{cite journal |journal= [[Angew. Chem. Int. Ed. Engl.]] |year= 1979 |last=Dye|first=J. L. |title= Compounds of Alkali Metal Anions |volume= 18 |issue= 8 |pages= 587–598 |doi= 10.1002/anie.197905871}}</ref> and the alkalides have much theoretical interest due to their unusual [[stoichiometry]] and low [[ionization potential|ionisation potentials]]. Alkalides are chemically similar to the [[electride]]s, which are salts with trapped [[electron]]s acting as anions.<ref name="Redko">{{cite journal |year= 2003 |title= Barium azacryptand sodide, the first alkalide with an alkaline Earth cation, also contains a novel dimer, (Na<sub>2</sub>)<sup>2−</sup> |journal= [[Journal of the American Chemical Society|J. Am. Chem. Soc.]] |volume= 125 |issue= 8 |pages= 2259–2263 |doi= 10.1021/ja027241m |pmid= 12590555|url=https://www.researchgate.net/publication/10896204 |last1= Redko |first1= M. Y. |last2= Huang |first2= R. H. |last3= Jackson |first3= J. E. |last4= Harrison |first4= J. F. |last5= Dye |first5= J. L. }}</ref> A particularly striking example of an alkalide is "inverse [[sodium hydride]]", H<sup>+</sup>Na<sup>−</sup> (both ions being [[coordination complex|complexed]]), as opposed to the usual sodium hydride, Na<sup>+</sup>H<sup>−</sup>:<ref name="HNa">{{cite journal |year= 2002 |title="Inverse sodium hydride": a crystalline salt that contains H<sup>+</sup> and Na<sup>−</sup> |journal= [[Journal of the American Chemical Society|J. Am. Chem. Soc.]] |volume= 124 |issue= 21 |pages= 5928–5929 |doi= 10.1021/ja025655+|pmid=12022811 |last1=Redko |first1=M. Y. |last2=Vlassa |first2=M. |last3=Jackson |first3=J. E. |last4=Misiolek |first4=A. W. |last5=Huang |first5=R. H. |last6=Dye |first6=J. L. }}</ref> it is unstable in isolation, due to its high energy resulting from the displacement of two electrons from hydrogen to sodium, although several derivatives are predicted to be [[metastability|metastable]] or stable.<ref name="HNa" /><ref name="HNa-theory">{{cite journal |url=http://simons.hec.utah.edu/papers/266.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://simons.hec.utah.edu/papers/266.pdf |archive-date=2022-10-09 |url-status=live|title=Inverse Sodium Hydride: A Theoretical Study|journal=J. Am. Chem. Soc.|year=2003|volume=125|pages=3954–3958|doi=10.1021/ja021136v|pmid=12656631|issue=13|last1=Sawicka|first1=A.|last2=Skurski|first2=P.|last3=Simons|first3=J.}}</ref>

In aqueous solution, the alkali metal ions form [[metal ions in aqueous solution|aqua ions]] of the formula [M(H<sub>2</sub>O)<sub>''n''</sub>]<sup>+</sup>, where ''n'' is the solvation number. Their [[coordination number]]s and shapes agree well with those expected from their ionic radii. In aqueous solution the water molecules directly attached to the metal ion are said to belong to the [[first coordination sphere]], also known as the first, or primary, solvation shell. The bond between a water molecule and the metal ion is a [[dative covalent bond]], with the oxygen atom donating both electrons to the bond. Each coordinated water molecule may be attached by [[hydrogen bond]]s to other water molecules. The latter are said to reside in the second coordination sphere. However, for the alkali metal cations, the second coordination sphere is not well-defined as the +1 charge on the cation is not high enough to [[Polarizability|polarise]] the water molecules in the primary solvation shell enough for them to form strong hydrogen bonds with those in the second coordination sphere, producing a more stable entity.<ref>{{cite book |last=Burgess |first=John |title=Metal Ions in Solution |year=1978 |publisher=Ellis Horwood |location=Chichester |page=20 |isbn=978-0-85312-027-8}}</ref><ref name=Richens />{{rp|25}} The solvation number for Li<sup>+</sup> has been experimentally determined to be 4, forming the [[tetrahedron|tetrahedral]] [Li(H<sub>2</sub>O)<sub>4</sub>]<sup>+</sup>: while solvation numbers of 3 to 6 have been found for lithium aqua ions, solvation numbers less than 4 may be the result of the formation of contact [[ion pair]]s, and the higher solvation numbers may be interpreted in terms of water molecules that approach [Li(H<sub>2</sub>O)<sub>4</sub>]<sup>+</sup> through a face of the tetrahedron, though molecular dynamic simulations may indicate the existence of an [[octahedron|octahedral]] hexaaqua ion. There are also probably six water molecules in the primary solvation sphere of the sodium ion, forming the octahedral [Na(H<sub>2</sub>O)<sub>6</sub>]<sup>+</sup> ion.<ref name=generalchemistry /><ref name=Richens>{{cite book |last=Richens |first=David. T. |title=The Chemistry of Aqua Ions |year=1997 |publisher=Wiley |isbn=978-0-471-97058-3}}</ref>{{rp|126–127}} While it was previously thought that the heavier alkali metals also formed octahedral hexaaqua ions, it has since been found that potassium and rubidium probably form the [K(H<sub>2</sub>O)<sub>8</sub>]<sup>+</sup> and [Rb(H<sub>2</sub>O)<sub>8</sub>]<sup>+</sup> ions, which have the [[square antiprism]]atic structure, and that caesium forms the 12-coordinate [Cs(H<sub>2</sub>O)<sub>12</sub>]<sup>+</sup> ion.<ref>{{cite journal |last=Persson |first=Ingmar |date=2010 |title=Hydrated metal ions in aqueous solution: How regular are their structures? |url=http://pac.iupac.org/publications/pac/pdf/2010/pdf/8210x1901.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://pac.iupac.org/publications/pac/pdf/2010/pdf/8210x1901.pdf |archive-date=2022-10-09 |url-status=live |journal=Pure Appl. Chem. |volume=82 |issue=10 |pages=1901–1917 |doi=10.1351/PAC-CON-09-10-22 |s2cid=98411500 |access-date=23 August 2014}}</ref>
{{clear left}}

==== Lithium ====
The chemistry of lithium shows several differences from that of the rest of the group as the small Li<sup>+</sup> cation [[chemical polarity|polarises]] [[anion]]s and gives its compounds a more [[covalent bond|covalent]] character.<ref name=rsc /> Lithium and [[magnesium]] have a [[diagonal relationship]] due to their similar atomic radii,<ref name=rsc /> so that they show some similarities. For example, lithium forms a stable [[nitride]], a property common among all the [[alkaline earth metal]]s (magnesium's group) but unique among the alkali metals.<ref name="alkalireact" /> In addition, among their respective groups, only lithium and magnesium form [[organometallic compound]]s with significant covalent character (e.g. Li[[methyl group|Me]] and MgMe<sub>2</sub>).<ref name="Shriver&Atkins">{{cite book |title=Inorganic Chemistry |first1=Duward |last1=Shriver |first2=Peter |last2=Atkins |publisher=W. H. Freeman |year=2006 |isbn=978-0-7167-4878-6 |page=259 |access-date=10 November 2012 |url=https://books.google.com/books?id=NwOTQAAACAAJ }}{{Dead link|date=August 2023 |bot=InternetArchiveBot |fix-attempted=yes }}</ref>

Lithium fluoride is the only alkali metal halide that is poorly soluble in water,<ref name=rsc /> and [[lithium hydroxide]] is the only alkali metal hydroxide that is not [[deliquescent]].<ref name=rsc /> Conversely, [[lithium perchlorate]] and other lithium salts with large anions that cannot be polarised are much more stable than the analogous compounds of the other alkali metals, probably because Li<sup>+</sup> has a high [[solvation energy]].<ref name="Greenwood&Earnshaw" />{{rp|76}} This effect also means that most simple lithium salts are commonly encountered in hydrated form, because the anhydrous forms are extremely [[hygroscopic]]: this allows salts like [[lithium chloride]] and [[lithium bromide]] to be used in [[dehumidifier]]s and [[air-conditioner]]s.<ref name="Greenwood&Earnshaw" />{{rp|76}}

==== Francium ====
Francium is also predicted to show some differences due to its high [[atomic weight]], causing its electrons to travel at considerable fractions of the speed of light and thus making [[relativistic quantum chemistry|relativistic effects]] more prominent. In contrast to the trend of decreasing [[electronegativity|electronegativities]] and [[ionisation energy|ionisation energies]] of the alkali metals, francium's electronegativity and ionisation energy are predicted to be higher than caesium's due to the relativistic stabilisation of the 7s electrons; also, its [[atomic radius]] is expected to be abnormally low.<!--Haire says this happens for Uue because of the analogous effect for 8s – seems likely for Fr too--> Thus, contrary to expectation, caesium is the most reactive of the alkali metals, not francium.<ref name="andreev" /><ref name="Uue">{{cite book |title= The Chemistry of the Actinide and Transactinide Elements |editor1-last= Morss|editor2-first= Norman M. |editor2-last= Edelstein |editor3-last= Fuger|editor3-first= Jean |last1= Hoffman|first1= Darleane C. |last2=Lee|first2=Diana M. |last3=Pershina|first3=Valeria |chapter= Transactinides and the future elements |publisher= Springer |year= 2006 |isbn= 978-1-4020-3555-5 |location= Dordrecht, The Netherlands |edition= 3rd }}</ref>{{rp|1729}}<ref name=Thayer /> All known physical properties of francium also deviate from the clear trends going from lithium to caesium, such as the first ionisation energy, electron affinity, and anion polarisability, though due to the paucity of known data about francium many sources give extrapolated values, ignoring that relativistic effects make the trend from lithium to caesium become inapplicable at francium.<ref name=Thayer /> Some of the few properties of francium that have been predicted taking relativity into account are the electron affinity (47.2 kJ/mol)<ref name=Landaualkalis>{{cite journal |last1= Landau |first1= A. |last2= Eliav |first2= E. |last3= Ishikawa |first3= Y. |last4= Kaldor |first4= U. |year= 2001 |title= Benchmark calculations of electron affinities of the alkali atoms sodium to eka-francium (element 119) |url= https://www.academia.edu/20466410|journal= J. Chem. Phys. |volume= 115 |issue= 6 |page= 2389 |doi= 10.1063/1.1386413|bibcode= 2001JChPh.115.2389L }}</ref> and the enthalpy of dissociation of the Fr<sub>2</sub> molecule (42.1 kJ/mol).<ref name=Liddle>{{cite book |last1=Jones |first1=Cameron |last2=Mountford |first2=Philip |last3=Stasch |first3=Andreas |last4=Blake |first4=Matthew P. |editor-last=Liddle |editor-first=Stephen T. |title=Molecular Metal-Metal Bonds: Compounds, Synthesis, Properties |publisher=John Wiley and Sons |date=22 June 2015 |pages=23–24 |chapter=s-block Metal-Metal Bonds |isbn=978-3-527-33541-1}}</ref> The CsFr molecule is polarised as Cs<sup>+</sup>Fr<sup>−</sup>, showing that the 7s subshell of francium is much more strongly affected by relativistic effects than the 6s subshell of caesium.<ref name=Thayer /> Additionally, francium superoxide (FrO<sub>2</sub>) is expected to have significant covalent character, unlike the other alkali metal superoxides, because of bonding contributions from the 6p electrons of francium.<ref name=Thayer />

=== Nuclear ===
<div style="float: right; margin: 5px;">
{| class="sortable wikitable" style="text-align:center;"
|+Primordial isotopes of the alkali metals
|-
! Z<br />
! Alkali metal<br />
! <small>[[stable isotope|Stable]]</small><br />
! <small>''[[primordial element|Decays]]''</small><br />
! class="unsortable" colspan="3"|<small>''unstable: italics''<div style="background:pink">odd–odd isotopes coloured pink</div></small>
|-
|  3 ||[[lithium]]      || [[isotopes of lithium|2]]    || — || {{SimpleNuclide|lithium|7}}|| style="background:pink;"|{{SimpleNuclide|lithium|6}}||&nbsp;
|-
| 11 ||[[sodium]]       || [[isotopes of sodium|1]]     || — ||{{SimpleNuclide|sodium|23}}||&nbsp;||&nbsp;
|-
| 19 ||[[potassium]]    || [[isotopes of potassium|2]]  || 1 ||{{SimpleNuclide|potassium|39}}||{{SimpleNuclide|potassium|41}}|| style="background:pink;"|''{{SimpleNuclide|potassium|40}}''
|-
| 37 ||[[rubidium]]     || [[isotopes of rubidium|1]]   || 1 ||{{SimpleNuclide|rubidium|85}}|||''{{SimpleNuclide|rubidium|87}}''||&nbsp;
|-
| 55 ||[[caesium]]      || [[isotopes of caesium|1]]    || — ||{{SimpleNuclide|caesium|133}}||&nbsp;||&nbsp;
|-
| 87 ||[[francium]]     || [[isotopes of francium|—]]   || — ||colspan="3"|''No primordial isotopes''<br />(''{{SimpleNuclide|francium|223}}'' is a [[radiogenic nuclide]])
|-
| colspan="7"|<small>Radioactive: {{nowrap|<sup>40</sup>K, [[half-life|t<sub>1/2</sub>]] 1.25 × 10<sup>9</sup> years;}} {{nowrap|<sup>87</sup>Rb, t<sub>1/2</sub> 4.9 × 10<sup>10</sup> years;}} {{nowrap|<sup>223</sup>Fr, t<sub>1/2</sub> 22.0 min.}}</small>
|}</div>
All the alkali metals have odd atomic numbers; hence, their isotopes must be either [[odd–odd nuclei|odd–odd]] (both proton and [[neutron number]] are odd) or [[odd–even nuclei|odd–even]] ([[proton number]] is odd, but neutron number is even). Odd–odd nuclei have even [[mass number]]s, whereas odd–even nuclei have odd mass numbers. Odd–odd [[primordial nuclide]]s are rare because most odd–odd nuclei are highly unstable with respect to [[beta decay]], because the decay products are even–even, and are therefore more strongly bound, due to [[Semi-empirical mass formula#Pairing term|nuclear pairing effects]].<ref name="Lide02">{{cite book |editor-last=Lide |editor-first=David R. |year=2002 |title=Handbook of Chemistry & Physics |edition=88th |publisher=CRC |url=http://www.hbcpnetbase.com/ |access-date=2008-05-23 |isbn=978-0-8493-0486-6 |oclc=179976746 |archive-date=24 July 2017 |archive-url=https://web.archive.org/web/20170724011402/http://www.hbcpnetbase.com/ |url-status=dead }}</ref>

Due to the great rarity of odd–odd nuclei, almost all the primordial isotopes of the alkali metals are odd–even (the exceptions being the light stable isotope lithium-6 and the long-lived [[radioisotope]] potassium-40). For a given odd mass number, there can be only a single [[beta-decay stable isobars|beta-stable nuclide]], since there is not a difference in binding energy between even–odd and odd–even comparable to that between even–even and odd–odd, leaving other nuclides of the same mass number ([[isobar (nuclide)|isobars]]) free to [[beta decay]] toward the lowest-mass nuclide. An effect of the instability of an odd number of either type of nucleons is that odd-numbered elements, such as the alkali metals, tend to have fewer stable isotopes than even-numbered elements. Of the 26 [[monoisotopic element]]s that have only a single stable isotope, all but one have an odd atomic number and all but one also have an even number of neutrons. [[Beryllium]] is the single exception to both rules, due to its low atomic number.<ref name="Lide02" />

All of the alkali metals except lithium and caesium have at least one naturally occurring [[radioisotope]]: [[sodium-22]] and [[sodium-24]] are [[trace radioisotope]]s produced [[cosmogenic]]ally,<ref>{{cite web |url=http://www.nucleonica.net/unc.aspx |title=Universal Nuclide Chart |date=2007–2012 |work=Nucleonica |publisher=Institute for Transuranium Elements |access-date=2011-04-17}}</ref> potassium-40 and [[rubidium-87]] have very long [[half-life|half-lives]] and thus occur naturally,<ref name="nuclideschart" /> and all [[isotopes of francium]] are [[radioactive decay|radioactive]].<ref name="nuclideschart" /> Caesium was also thought to be radioactive in the early 20th century,<ref name="Patt1926">{{cite journal |doi= 10.1021/cr60009a003 |title= The Radioactivity of the Alkali Metals |year= 1926 |last1= Patton |first1= I. Jocelyn |last2= Waldbauer |first2= L. J. |journal= Chemical Reviews |volume= 3 |pages= 81–93}}</ref><ref name="Kenn1908">{{cite journal |doi= 10.1080/14786440908636519 |title= On the radioactivity of potassium and other alkali metals |year= 1908 |last1= McLennan |first1= J. C. |last2= Kennedy |first2= W. T. |journal= [[Philosophical Magazine]] |series= 6 |volume= 16 |issue= 93 |pages= 377–395|url= https://zenodo.org/record/1430860 }}</ref> although it has no naturally occurring radioisotopes.<ref name="nuclideschart">{{cite web|url=http://www.nndc.bnl.gov/chart/|title=Interactive Chart of Nuclides|publisher=Brookhaven National Laboratory|last=Sonzogni|first=Alejandro|location=National Nuclear Data Center|access-date=4 October 2012|archive-date=21 July 2011|archive-url=https://web.archive.org/web/20110721051025/http://www.nndc.bnl.gov/chart/|url-status=dead}}</ref> (Francium had not been discovered yet at that time.) The natural long-lived radioisotope of potassium, potassium-40, makes up about 0.012% of natural potassium,<ref>{{cite web |url=http://www.ead.anl.gov/pub/doc/potassium.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://www.ead.anl.gov/pub/doc/potassium.pdf |archive-date=2022-10-09 |url-status=live |title=Potassium-40 |date=August 2005 |work=Human Health Fact Sheet |publisher=[[Argonne National Laboratory]], Environmental Science Division |access-date=7 February 2012}}</ref> and thus natural potassium is weakly radioactive. This natural radioactivity became a basis for a mistaken claim of the discovery for element 87 (the next alkali metal after caesium) in 1925.<ref name="fontani" /><ref name="vanderkrogt-Fr">{{cite web |last= Van der Krogt |first= Peter |title= Francium |work= Elementymology & Elements Multidict |date= 10 January 2006 |url= http://elements.vanderkrogt.net/element.php?sym=Fr |access-date= 8 April 2007}}</ref> Natural rubidium is similarly slightly radioactive, with 27.83% being the long-lived radioisotope rubidium-87.<ref name="Greenwood&Earnshaw" />{{rp|74}}

[[Caesium-137]], with a half-life of 30.17&nbsp;years, is one of the two principal [[medium-lived fission product]]s, along with [[strontium-90]], which are responsible for most of the [[radioactivity]] of [[spent nuclear fuel]] after several years of cooling, up to several hundred years after use. It constitutes most of the radioactivity still left from the [[Chernobyl accident]]. Caesium-137 undergoes high-energy beta decay and eventually becomes stable [[barium-137]]. It is a strong emitter of gamma radiation. Caesium-137 has a very low rate of neutron capture and cannot be feasibly disposed of in this way, but must be allowed to decay.<ref name="Cs-137">{{cite web|title=Radionuclide Half-Life Measurements|url=https://www.nist.gov/pml/data/halflife-html.cfm|author=National Institute of Standards and Technology|date=6 September 2009|access-date=2011-11-07|archive-date=12 August 2016|archive-url=https://web.archive.org/web/20160812133216/http://nist.gov/pml/data/halflife-html.cfm|url-status=dead}}</ref> Caesium-137 has been used as a [[Flow tracer|tracer]] in hydrologic studies, analogous to the use of [[tritium]].<ref>[https://web.archive.org/web/20160329120038/http://www.bt.cdc.gov/radiation/isotopes/cesium.asp Radioisotope Brief: Cesium-137 (Cs-137)]. U.S. National Center for Environmental Health</ref> Small amounts of [[caesium-134]] and caesium-137 were released into the environment during nearly all [[nuclear weapon test]]s and some [[nuclear accident]]s, most notably the [[Goiânia accident]] and the [[Chernobyl disaster]]. As of 2005, caesium-137 is the principal source of radiation in the [[zone of alienation]] around the [[Chernobyl nuclear power plant]].<ref name="IAEA">{{cite book |title=The Radiological Accident in Goiânia |publisher=[[IAEA]] |year=1988 |url=http://www-pub.iaea.org/MTCD/publications/PubDetAR.asp?pubId=3684}}</ref> Its chemical properties as one of the alkali metals make it one of the most problematic of the short-to-medium-lifetime fission products because it easily moves and spreads in nature due to the high water solubility of its salts, and is taken up by the body, which mistakes it for its essential congeners sodium and potassium.<ref name="RPD">{{cite book |title=Radionuclide and Radiation Protection Data Handbook 2002 |isbn=978-1-870965-87-3 |year=2002 |first1=D. |last1=Delacroix |first2=J. P. |last2=Guerre |first3=P. |last3=Leblanc |first4=C. |last4=Hickman |publisher=Nuclear Technology Publishing |edition=2nd}}</ref>{{rp|114}}

== Periodic trends ==
The alkali metals are more similar to each other than the elements in any other [[group (periodic table)|group]] are to each other.<ref name=rsc /> For instance, when moving down the table, all known alkali metals show increasing [[atomic radius]],<ref name=chemguide /> decreasing [[electronegativity]],<ref name=chemguide /> increasing [[reactivity (chemistry)|reactivity]],<ref name=rsc /> and decreasing melting and boiling points<ref name=chemguide /> as well as heats of fusion and vaporisation.<ref name="Greenwood&Earnshaw">{{Greenwood&Earnshaw2nd}}</ref>{{rp|75}} In general, their [[density|densities]] increase when moving down the table, with the exception that potassium is less dense than sodium.<ref name=chemguide />

=== Atomic and ionic radii ===
[[File:Effective Nuclear Charge.svg|thumb|250px|[[Effective nuclear charge]] on an atomic electron]]
The [[Atomic radius|atomic radii]] of the alkali metals increase going down the group.<ref name=chemguide /> Because of the [[shielding effect]], when an atom has more than one [[electron shell]], each electron feels electric repulsion from the other electrons as well as electric attraction from the nucleus.<ref name=shielding>{{cite book |first1=Theodore |last1=L. Brown | first2=H. Eugene Jr. | last2=LeMay |first3=Bruce E. |last3=Bursten |first4=Julia R. |last4=Burdge |year=2003 |title=Chemistry: The Central Science |edition=8th |publisher=Pearson Education |location=US |isbn=978-0-13-061142-0 }}</ref> In the alkali metals, the [[valence electron|outermost electron]] only feels a net charge of +1, as some of the [[nuclear charge]] (which is equal to the [[atomic number]]) is cancelled by the inner electrons; the number of inner electrons of an alkali metal is always one less than the nuclear charge. Therefore, the only factor which affects the atomic radius of the alkali metals is the number of electron shells. Since this number increases down the group, the atomic radius must also increase down the group.<ref name=chemguide />

The [[ionic radius|ionic radii]] of the alkali metals are much smaller than their atomic radii. This is because the outermost electron of the alkali metals is in a different [[electron shell]] than the inner electrons, and thus when it is removed the resulting atom has one fewer electron shell and is smaller. Additionally, the [[effective nuclear charge]] has increased, and thus the electrons are attracted more strongly towards the nucleus and the ionic radius decreases.<ref name=rsc />

=== First ionisation energy ===
[[File:First Ionization Energy blocks.svg|thumb|upright=2.7|Periodic trend for ionisation energy: each period begins at a minimum for the alkali metals, and ends at a maximum for the [[noble gas]]es. Predicted values are used for elements beyond 104.]]
The first [[ionisation energy]] of an [[chemical element|element]] or [[molecule]] is the energy required to move the most loosely held electron from one [[mole (unit)|mole]] of gaseous atoms of the element or molecules to form one mole of gaseous ions with [[electric charge]] +1. The factors affecting the first ionisation energy are the [[nuclear charge]], the amount of [[shielding effect|shielding]] by the inner electrons and the distance from the most loosely held electron from the nucleus, which is always an outer electron in [[main group element]]s. The first two factors change the effective nuclear charge the most loosely held electron feels. Since the outermost electron of alkali metals always feels the same effective nuclear charge (+1), the only factor which affects the first ionisation energy is the distance from the outermost electron to the nucleus. Since this distance increases down the group, the outermost electron feels less attraction from the nucleus and thus the first ionisation energy decreases.<ref name=chemguide /> This trend is broken in francium due to the [[relativistic quantum chemistry|relativistic]] stabilisation and contraction of the 7s orbital, bringing francium's valence electron closer to the nucleus than would be expected from non-relativistic calculations. This makes francium's outermost electron feel more attraction from the nucleus, increasing its first ionisation energy slightly beyond that of caesium.<ref name="Uue" />{{Rp|1729}}<!--Also explain why the alkali metals have the lowest ionization energies in their period.-->

The second ionisation energy of the alkali metals is much higher than the first as the second-most loosely held electron is part of a fully filled [[electron shell]] and is thus difficult to remove.<ref name=rsc />

=== Reactivity ===
The [[Reactivity (chemistry)|reactivities]] of the alkali metals increase going down the group. This is the result of a combination of two factors: the first ionisation energies and [[atomisation energy|atomisation energies]] of the alkali metals. Because the first ionisation energy of the alkali metals decreases down the group, it is easier for the outermost electron to be removed from the atom and participate in [[chemical reaction]]s, thus increasing reactivity down the group. The atomisation energy measures the strength of the [[metallic bond]] of an element, which falls down the group as the atoms increase in [[atomic radius|radius]] and thus the metallic bond must increase in length, making the [[delocalized electron|delocalised electrons]] further away from the attraction of the nuclei of the heavier alkali metals. Adding the atomisation and first ionisation energies gives a quantity closely related to (but not equal to) the [[activation energy]] of the reaction of an alkali metal with another substance. This quantity decreases going down the group, and so does the activation energy; thus, chemical reactions can occur faster and the reactivity increases down the group.<ref name="alkaliwater">{{cite web |url=http://www.chemguide.co.uk/inorganic/group1/reacth2o.html#top |title=Reaction of the Group 1 Elements with Water |last=Clark|first=Jim |year=2005 |work=chemguide |access-date=18 June 2012}}</ref>

=== Electronegativity ===
[[File:Periodic variation of Pauling electronegativities.svg|thumb|upright=1.25|Periodic variation of Pauling electronegativities as one descends the [[main group element|main groups]] of the periodic table from the [[period 2 element|second]] to the [[period 6 element|sixth period]].]]

[[Electronegativity]] is a [[chemical property]] that describes the tendency of an [[atom]] or a [[functional group]] to attract [[electron]]s (or [[electron density]]) towards itself.<ref name="definition">{{GoldBookRef|file=E01990|title=Electronegativity}}</ref> If the bond between [[sodium]] and [[chlorine]] in [[sodium chloride]] were [[covalent bond|covalent]], the pair of shared electrons would be attracted to the chlorine because the effective nuclear charge on the outer electrons is +7 in chlorine but is only +1 in sodium. The electron pair is attracted so close to the chlorine atom that they are practically transferred to the chlorine atom (an [[ionic bond]]). However, if the sodium atom was replaced by a lithium atom, the electrons will not be attracted as close to the chlorine atom as before because the lithium atom is smaller, making the electron pair more strongly attracted to the closer effective nuclear charge from lithium. Hence, the larger alkali metal atoms (further down the group) will be less electronegative as the bonding pair is less strongly attracted towards them. As mentioned previously, francium is expected to be an exception.<ref name=chemguide />

Because of the higher electronegativity of lithium, some of its compounds have a more covalent character. For example, [[lithium iodide]] (LiI) will dissolve in [[organic solvent]]s, a property of most covalent compounds.<ref name=chemguide /> [[Lithium fluoride]] (LiF) is the only [[alkali halide]] that is not soluble in water,<ref name=rsc /> and [[lithium hydroxide]] (LiOH) is the only [[alkali hydroxide|alkali metal hydroxide]] that is not [[deliquescent]].<ref name=rsc />

=== Melting and boiling points ===
The [[melting point]] of a substance is the point where it changes [[state of matter|state]] from solid to liquid while the [[boiling point]] of a substance (in liquid state) is the point where the [[vapor pressure|vapour pressure]] of the liquid equals the environmental pressure surrounding the liquid<ref>{{cite book |last=Goldberg|first=David E. |title=3,000 Solved Problems in Chemistry|edition=1st|publisher=McGraw-Hill|year=1988|isbn=978-0-07-023684-4}} Section 17.43, page 321</ref><ref>{{cite book |editor1=Theodore, Louis |editor2=Dupont, R. Ryan |editor3=Ganesan, Kumar |title=Pollution Prevention: The Waste Management Approach to the 21st Century|publisher=CRC Press|year=1999|isbn=978-1-56670-495-3|page=15 Section 27}}</ref> and all the liquid changes state to gas. As a metal is heated to its melting point, the [[metallic bond]]s keeping the atoms in place weaken so that the atoms can move around, and the metallic bonds eventually break completely at the metal's boiling point.<ref name=chemguide /><ref name="metallic-bonding">{{cite web |url=http://www.chemguide.co.uk/atoms/bonding/metallic.html |title=Metallic Bonding |last=Clark|first=Jim |year=2000 |work=chemguide |access-date=23 March 2012}}</ref> Therefore, the falling melting and boiling points of the alkali metals indicate that the strength of the metallic bonds of the alkali metals decreases down the group.<ref name=chemguide /> This is because metal atoms are held together by the electromagnetic attraction from the positive ions to the delocalised electrons.<ref name=chemguide /><ref name="metallic-bonding" /> As the atoms increase in size going down the group (because their atomic radius increases), the nuclei of the ions move further away from the delocalised electrons and hence the metallic bond becomes weaker so that the metal can more easily melt and boil, thus lowering the melting and boiling points.<ref name=chemguide /> The increased nuclear charge is not a relevant factor due to the shielding effect.<ref name=chemguide />

=== Density ===
The alkali metals all have the same [[crystal structure]] ([[body-centered cubic|body-centred cubic]])<ref name="Greenwood&Earnshaw" /> and thus the only relevant factors are the number of atoms that can fit into a certain volume and the mass of one of the atoms, since density is defined as mass per unit volume. The first factor depends on the volume of the atom and thus the atomic radius, which increases going down the group; thus, the volume of an alkali metal atom increases going down the group. The mass of an alkali metal atom also increases going down the group. Thus, the trend for the densities of the alkali metals depends on their atomic weights and atomic radii; if figures for these two factors are known, the ratios between the densities of the alkali metals can then be calculated. The resultant trend is that the densities of the alkali metals increase down the table, with an exception at potassium. Due to having the lowest atomic weight and the largest atomic radius of all the elements in their periods, the alkali metals are the least dense metals in the periodic table.<ref name=chemguide /> Lithium, sodium, and potassium are the only three metals in the periodic table that are less dense than water:<ref name=rsc /> in fact, lithium is the least dense known solid at [[room temperature]].<ref name="Greenwood&Earnshaw" />{{rp|75}}

== Compounds ==
The alkali metals form complete series of compounds with all usually encountered anions, which well illustrate group trends. These compounds can be described as involving the alkali metals losing electrons to acceptor species and forming monopositive ions.<ref name="Greenwood&Earnshaw" />{{rp|79}} This description is most accurate for alkali halides and becomes less and less accurate as cationic and anionic charge increase, and as the anion becomes larger and more polarisable. For instance, [[ionic bond]]ing gives way to [[metallic bond]]ing along the series NaCl, Na<sub>2</sub>O, Na<sub>2</sub>S, Na<sub>3</sub>P, Na<sub>3</sub>As, Na<sub>3</sub>Sb, Na<sub>3</sub>Bi, Na.<ref name="Greenwood&Earnshaw" />{{rp|81}}

=== Hydroxides ===
{{External media
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| video1= [https://www.youtube.com/watch?v=QSZ-3wScePM Reactions of the alkali metals with water], conducted by [[The Open University]]
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[[File:Large Sodium Explosion.jpg|thumb|right|alt=A large orange-yellow explosion|A reaction of 3 [[pound (mass)|pounds]] (≈ 1.4&nbsp;kg) of sodium with water]]
All the alkali metals react vigorously or explosively with cold water, producing an [[aqueous solution]] of a strongly [[base (chemistry)|basic]] alkali metal [[hydroxide]] and releasing hydrogen gas.<ref name="alkaliwater" /> This reaction becomes more vigorous going down the group: lithium reacts steadily with [[effervescence]], but sodium and potassium can ignite, and rubidium and caesium sink in water and generate hydrogen gas so rapidly that shock waves form in the water that may shatter glass containers.<ref name=rsc /> When an alkali metal is dropped into water, it produces an explosion, of which there are two separate stages. The metal reacts with the water first, breaking the hydrogen bonds in the water and producing [[hydrogen]] gas; this takes place faster for the more reactive heavier alkali metals. Second, the heat generated by the first part of the reaction often ignites the hydrogen gas, causing it to burn explosively into the surrounding air. This secondary hydrogen gas explosion produces the visible flame above the bowl of water, lake or other body of water, not the initial reaction of the metal with water (which tends to happen mostly under water).<ref name="alkalibangs" /> The alkali metal hydroxides are the most basic known hydroxides.<ref name="Greenwood&Earnshaw" />{{rp|87}}

Recent research has suggested that the explosive behavior of alkali metals in water is driven by a [[Coulomb explosion]] rather than solely by rapid generation of hydrogen itself.<ref name=coulomb>{{cite journal |title=Coulomb explosion during the early stages of the reaction of alkali metals with water|journal=Nature Chemistry|doi=10.1038/nchem.2161|pmid = 25698335|date=26 January 2015|volume=7|issue=3|pages=250–254|bibcode=2015NatCh...7..250M | last1 = Mason | first1 = Philip E.}}</ref> All alkali metals melt as a part of the reaction with water. Water molecules ionise the bare metallic surface of the liquid metal, leaving a positively charged metal surface and negatively charged water ions. The attraction between the charged metal and water ions will rapidly increase the surface area, causing an exponential increase of ionisation. When the repulsive forces within the liquid metal surface exceeds the forces of the surface tension, it vigorously explodes.<ref name=coulomb />

The hydroxides themselves are the most basic hydroxides known, reacting with acids to give salts and with alcohols to give [[oligomer]]ic [[alkoxide]]s. They easily react with [[carbon dioxide]] to form [[carbonate]]s or [[bicarbonate]]s, or with [[hydrogen sulfide]] to form [[sulfide]]s or [[bisulfide]]s, and may be used to separate [[thiol]]s from petroleum. They react with amphoteric oxides: for example, the oxides of [[aluminium oxide|aluminium]], [[zinc oxide|zinc]], [[tin(IV) oxide|tin]], and [[lead dioxide|lead]] react with the alkali metal hydroxides to give aluminates, zincates, stannates, and plumbates. [[Silicon dioxide]] is acidic, and thus the alkali metal hydroxides can also attack [[silicate glass]].<ref name="Greenwood&Earnshaw" />{{rp|87}}

=== Intermetallic compounds ===
[[File:NaK alloy.jpg|thumb|right|Liquid NaK alloy at room temperature]]
The alkali metals form many [[intermetallic compound]]s with each other and the elements from groups [[alkaline earth metal|2]] to [[boron group|13]] in the periodic table of varying stoichiometries,<ref name="Greenwood&Earnshaw" />{{rp|81}} such as the [[sodium amalgam]]s with [[mercury (element)|mercury]], including Na<sub>5</sub>Hg<sub>8</sub> and Na<sub>3</sub>Hg.<ref>{{cite book | last=Buszek | first=Keith R. | title=Encyclopedia of Reagents for Organic Synthesis | chapter=Sodium Amalgam | publisher=John Wiley & Sons, Ltd | publication-place=Chichester, UK | date=2001-04-15 | isbn=978-0-471-93623-7 | doi=10.1002/047084289x.rs040 | page=}}</ref> Some of these have ionic characteristics: taking the alloys with gold, the most electronegative of metals, as an example, NaAu and KAu are metallic, but RbAu and [[caesium auride|CsAu]] are semiconductors.<ref name="Greenwood&Earnshaw" />{{rp|81}} [[NaK]] is an alloy of sodium and potassium that is very useful because it is liquid at room temperature, although precautions must be taken due to its extreme reactivity towards water and air. The [[eutectic mixture]] melts at −12.6&nbsp;°C.<ref name="basf-ds-NaK">{{cite web |publisher= [[BASF]] |title= Sodium-Potassium Alloy (NaK) |url= http://www.basf.com/inorganics/pdfs/tech_datasheet/NaK.pdf |date= December 2004 |archive-url= https://web.archive.org/web/20070927210800/https://www.basf.com/inorganics/pdfs/tech_datasheet/NaK.pdf |archive-date= 27 September 2007}}</ref> An alloy of 41% caesium, 47% sodium, and 12% potassium has the lowest known melting point of any metal or alloy, −78&nbsp;°C.<ref name="caesium" />

=== Compounds with the group 13 elements ===
The intermetallic compounds of the alkali metals with the heavier group 13 elements (aluminium, [[gallium]], [[indium]], and [[thallium]]), such as NaTl, are poor [[Electrical conductor|conductors]] or [[semiconductor]]s, unlike the normal alloys with the preceding elements, implying that the alkali metal involved has lost an electron to the [[Zintl phase|Zintl anions]] involved.<ref name="Sevov">{{cite book | title=Intermetallic Compounds - Principles and Practice: Progress | publisher=Wiley | date=2002-05-04 | isbn=978-0-471-49315-0 | doi=10.1002/0470845856 | pages=113–132 |chapter=Zintl Phases |chapter-url=http://www3.nd.edu/~sevovlab/articles/SlaviChapter.pdf | editor-last1=Westbrook | editor-last2=Fleischer | editor-first1=J. H. | editor-first2=R. L. }}</ref> Nevertheless, while the elements in group 14 and beyond tend to form discrete anionic clusters, group 13 elements tend to form polymeric ions with the alkali metal cations located between the giant ionic lattice. For example, NaTl consists of a polymeric anion (—Tl<sup>−</sup>—)<sub>n</sub> with a covalent [[diamond cubic]] structure with Na<sup>+</sup> ions located between the anionic lattice. The larger alkali metals cannot fit similarly into an anionic lattice and tend to force the heavier group 13 elements to form anionic clusters.<ref name="Kauzlarich">S.M. Kauzlarich, Encyclopedia of Inorganic chemistry, 1994, John Wiley & Sons, {{ISBN|0-471-93620-0}}</ref>

[[Boron]] is a special case, being the only nonmetal in group 13. The alkali metal [[boride]]s tend to be boron-rich, involving appreciable boron–boron bonding involving [[deltahedron|deltahedral]] structures,<ref name="Greenwood&Earnshaw" />{{rp|147–8}} and are thermally unstable due to the alkali metals having a very high [[vapour pressure]] at elevated temperatures. This makes direct synthesis problematic because the alkali metals do not react with boron below 700&nbsp;°C, and thus this must be accomplished in sealed containers with the alkali metal in excess. Furthermore, exceptionally in this group, reactivity with boron decreases down the group: lithium reacts completely at 700&nbsp;°C, but sodium at 900&nbsp;°C and potassium not until 1200&nbsp;°C, and the reaction is instantaneous for lithium but takes hours for potassium. Rubidium and caesium borides have not even been characterised. Various phases are known, such as LiB<sub>10</sub>, NaB<sub>6</sub>, NaB<sub>15</sub>, and KB<sub>6</sub>.<ref>{{cite book |last=Hagen |first=A. P. |date=17 September 2009 |title=Inorganic Reactions and Methods, The Formation of Bonds to Group-I, -II, and -IIIB Elements |publisher=John Wiley & Sons |pages=204–5 |isbn=978-0-470-14549-4}}</ref><ref>{{cite book |last=Matkovich |first=V. I. |date=6 December 2012 |title=Boron and Refractory Borides |publisher=Springer |pages=262–92 |isbn=978-3-642-66620-9}}</ref> Under high pressure the boron–boron bonding in the lithium borides changes from following [[polyhedral skeletal electron pair theory|Wade's rules]] to forming Zintl anions like the rest of group 13.<ref>{{cite journal |last1=Hermann |first1=Andreas |last2=McSorley |first2=Alexandra |first3=Ashcroft |last3=N. W. |first4=Roald |last4=Hoffmann |date=2012 |title=From Wade–Mingos to Zintl–Klemm at 100 GPa: Binary Compounds of Boron and Lithium |url=http://www2.ph.ed.ac.uk/~aherman2/Andreas_Hermann,_Edinburgh/Research_files/JAmChemSoc_134_18606_2012.pdf |journal=[[Journal of the American Chemical Society]] |volume=2012 |issue=134 |pages=18606–18 |doi=10.1021/ja308492g |pmid=23066852 |access-date=21 August 2016 |archive-date=27 September 2020 |archive-url=https://web.archive.org/web/20200927024724/https://www2.ph.ed.ac.uk/~aherman2/Andreas_Hermann,_Edinburgh/Research_files/JAmChemSoc_134_18606_2012.pdf |url-status=dead }}</ref>

=== Compounds with the group 14 elements ===
{{multiple image
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| image2 = Potassium-graphite-xtal-3D-SF-B.png
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| footer = Side ''(left)'' and top ''(right)'' views of the [[graphite intercalation compound]] KC<sub>8</sub>
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Lithium and sodium react with [[carbon]] to form [[acetylide]]s, Li<sub>2</sub>C<sub>2</sub> and Na<sub>2</sub>C<sub>2</sub>, which can also be obtained by reaction of the metal with [[acetylene]]. Potassium, rubidium, and caesium react with [[graphite]]; their atoms are [[intercalation (chemistry)|intercalated]] between the hexagonal graphite layers, forming [[graphite intercalation compound]]s of formulae MC<sub>60</sub> (dark grey, almost black), MC<sub>48</sub> (dark grey, almost black), MC<sub>36</sub> (blue), MC<sub>24</sub> (steel blue), and MC<sub>8</sub> (bronze) (M = K, Rb, or Cs). These compounds are over 200 times more electrically conductive than pure graphite, suggesting that the valence electron of the alkali metal is transferred to the graphite layers (e.g. {{chem2|M+C8-}}).<ref name=generalchemistry /> Upon heating of KC<sub>8</sub>, the elimination of potassium atoms results in the conversion in sequence to KC<sub>24</sub>, KC<sub>36</sub>, KC<sub>48</sub> and finally KC<sub>60</sub>. KC<sub>8</sub> is a very strong [[reducing agent]] and is pyrophoric and explodes on contact with water.<ref name="InorgChem">{{cite book |title= Inorganic Chemistry, 3rd Edition |chapter= Chapter 14: The group 14 elements |last1=Housecroft|first1=Catherine E.|last2=Sharpe|first2=Alan G.|publisher= Pearson |year= 2008 |isbn= 978-0-13-175553-6 |page= 386}}</ref><ref>[https://web.archive.org/web/20141129035434/http://physics.nist.gov/TechAct.2001/Div846/div846h.html NIST Ionizing Radiation Division 2001 – Technical Highlights]. physics.nist.gov</ref> While the larger alkali metals (K, Rb, and Cs) initially form MC<sub>8</sub>, the smaller ones initially form MC<sub>6</sub>, and indeed they require reaction of the metals with graphite at high temperatures around 500&nbsp;°C to form.<ref name=cac6>{{cite journal |last1=Emery|first1=N. |display-authors=1|title=Review: Synthesis and superconducting properties of CaC6|journal=Sci. Technol. Adv. Mater.|volume=9|year=2008|page=044102|doi=10.1088/1468-6996/9/4/044102|bibcode=2008STAdM...9d4102E|issue=4|first2=Claire|last2=Hérold|first3=Jean-François|last3=Marêché|first4=Philippe|last4=Lagrange|pmc=5099629|pmid=27878015}}</ref> Apart from this, the alkali metals are such strong reducing agents that they can even reduce [[buckminsterfullerene]] to produce solid [[fulleride]]s M<sub>''n''</sub>C<sub>60</sub>; sodium, potassium, rubidium, and caesium can form fullerides where ''n'' = 2, 3, 4, or 6, and rubidium and caesium additionally can achieve ''n'' = 1.<ref name="Greenwood&Earnshaw" />{{rp|285}}

When the alkali metals react with the heavier elements in the [[carbon group]] ([[silicon]], [[germanium]], [[tin]], and lead), ionic substances with cage-like structures are formed, such as the [[silicide]]s M<sub>4</sub>[[silicon|Si]]<sub>4</sub> (M = K, Rb, or Cs), which contains M<sup>+</sup> and tetrahedral {{chem2|Si4(4-)}} ions.<ref name=generalchemistry /> The chemistry of alkali metal [[germanide]]s, involving the germanide ion [[germanium|Ge]]<sup>4−</sup> and other cluster ([[Zintl ion|Zintl]]) ions such as {{chem2|Ge4(2-)}}, {{chem2|Ge9(4-)}}, {{chem2|Ge9(2-)}}, and [(Ge<sub>9</sub>)<sub>2</sub>]<sup>6−</sup>, is largely analogous to that of the corresponding silicides.<ref name="Greenwood&Earnshaw" />{{rp|393}} Alkali metal [[stannide]]s are mostly ionic, sometimes with the stannide ion ([[tin|Sn]]<sup>4−</sup>),<ref name="Kauzlarich" /> and sometimes with more complex Zintl ions such as {{chem2|Sn9(4-)}}, which appears in tetrapotassium nonastannide (K<sub>4</sub>Sn<sub>9</sub>).<ref name="Hoch">{{cite journal |doi=10.1107/S0108270102002032|pmid=11932511|title=Tetrapotassium nonastannide, K4Sn9|year=2002|last1=Hoch|first1=Constantin|last2=Wendorff|first2=Marco|last3=Röhr|first3=Caroline|journal=Acta Crystallographica Section C|volume=58|issue=4|pages=I45–I46}}</ref> The monatomic [[plumbide]] ion ([[lead|Pb]]<sup>4−</sup>) is unknown, and indeed its formation is predicted to be energetically unfavourable; alkali metal plumbides have complex Zintl ions, such as {{chem2|Pb9(4-)}}. These alkali metal germanides, stannides, and plumbides may be produced by reducing germanium, tin, and lead with sodium metal in liquid ammonia.<ref name="Greenwood&Earnshaw" />{{rp|394}}

=== Nitrides and pnictides ===
[[File:Lithium-nitride-xtal-CM-3D-polyhedra.png|thumb|[[Unit cell]] [[ball-and-stick model]] of [[lithium nitride]].<ref>{{cite journal |title=Structure of Lithium Nitride and Transition-Metal-Doped Derivatives, Li<sub>3−''x''−''y''</sub>M<sub>''x''</sub>N (M= Ni, Cu): A Powder Neutron Diffraction Study|last1=Gregory|first1=Duncan H.|last2=O'Meara|first2=Paul M.|last3=Gordon|first3=Alexandra G.|last4=Hodges|first4=Jason P.|last5=Short|first5=Simine|last6=Jorgensen|first6=James D.|journal=Chem. Mater.|year=2002|volume=14|issue=5|pages=2063–2070|doi=10.1021/cm010718t}}</ref> On the basis of size a [[tetrahedron|tetrahedral]] structure would be expected, but that would be geometrically impossible: thus lithium nitride takes on this unique crystal structure.<ref name="Greenwood&Earnshaw" />{{rp|76}}]]
Lithium, the lightest of the alkali metals, is the only alkali metal which reacts with [[nitrogen]] at [[standard conditions]], and its [[nitride]] is the only stable alkali metal nitride. Nitrogen is an [[reactivity (chemistry)|unreactive]] gas because breaking the strong [[triple bond]] in the [[dinitrogen]] molecule (N<sub>2</sub>) requires a lot of energy. The formation of an alkali metal nitride would consume the ionisation energy of the alkali metal (forming M<sup>+</sup> ions), the energy required to break the triple bond in N<sub>2</sub> and the formation of N<sup>3−</sup> ions, and all the energy released from the formation of an alkali metal nitride is from the [[lattice energy]] of the alkali metal nitride. The lattice energy is maximised with small, highly charged ions; the alkali metals do not form highly charged ions, only forming ions with a charge of +1, so only lithium, the smallest alkali metal, can release enough lattice energy to make the reaction with nitrogen [[exothermic]], forming [[lithium nitride]]. The reactions of the other alkali metals with nitrogen would not release enough lattice energy and would thus be [[endothermic]], so they do not form nitrides at standard conditions.<ref name="alkalireact">{{cite web |url=http://www.chemguide.co.uk/inorganic/group1/reacto2.html#top |title=Reaction of the Group 1 Elements with Oxygen and Chlorine |last=Clark|first=Jim |year=2005 |work=chemguide |access-date=27 June 2012}}</ref> [[Sodium nitride]] (Na<sub>3</sub>N) and [[potassium nitride]] (K<sub>3</sub>N), while existing, are extremely unstable, being prone to decomposing back into their constituent elements, and cannot be produced by reacting the elements with each other at standard conditions.<ref name=Jansen1>{{cite journal |title=Synthesis and structure of Na<sub>3</sub>N|last1=Fischer|first1=D. |last2=Jansen|first2=M.|journal= Angew Chem|volume=41|issue=10|pages=1755–1756|year=2002|doi=10.1002/1521-3773(20020517)41:10<1755::AID-ANIE1755>3.0.CO;2-C|pmid=19750706 }}</ref><ref name="Jansen2">{{cite journal |title=Synthesis and structure of K<sub>3</sub>N|last1=Fischer|first1=D. |last2=Cancarevic|first2=Z. |last3=Schön|first3=J. C. |last4=Jansen|first4=M. Z. |journal=Z. Anorg. Allg. Chem.|volume= 630|issue=1|pages=156–160|doi=10.1002/zaac.200300280|year=2004}}. [http://pubs.acs.org/cen/topstory/8020/8020notw9.html 'Elusive Binary Compound Prepared'] ''Chemical & Engineering News'' '''80''' No. 20 (20 May 2002)</ref> Steric hindrance forbids the existence of rubidium or caesium nitride.<ref name="Greenwood&Earnshaw" />{{rp|417}} However, sodium and potassium form colourless [[azide]] salts involving the linear {{chem2|N3-}} anion; due to the large size of the alkali metal cations, they are thermally stable enough to be able to melt before decomposing.<ref name="Greenwood&Earnshaw" />{{rp|417}}

All the alkali metals react readily with [[phosphorus]] and [[arsenic]] to form [[phosphide]]s and [[arsenide]]s with the formula M<sub>3</sub>Pn (where M represents an alkali metal and Pn represents a [[pnictogen]] – phosphorus, arsenic, [[antimony]], or [[bismuth]]). This is due to the greater size of the P<sup>3−</sup> and As<sup>3−</sup> ions, so that less lattice energy needs to be released for the salts to form.<ref name=generalchemistry /> These are not the only phosphides and arsenides of the alkali metals: for example, potassium has nine different known phosphides, with formulae K<sub>3</sub>P, K<sub>4</sub>P<sub>3</sub>, K<sub>5</sub>P<sub>4</sub>, KP, K<sub>4</sub>P<sub>6</sub>, K<sub>3</sub>P<sub>7</sub>, K<sub>3</sub>P<sub>11</sub>, KP<sub>10.3</sub>, and KP<sub>15</sub>.<ref name="Schnering">H.G. Von Schnering, W. Hönle ''Phosphides – Solid-state Chemistry'' Encyclopedia of Inorganic Chemistry Ed. R. Bruce King (1994) John Wiley & Sons {{ISBN|0-471-93620-0}}</ref> While most metals form arsenides, only the alkali and alkaline earth metals form mostly ionic arsenides. The structure of Na<sub>3</sub>As is complex with unusually short Na–Na distances of 328–330 pm which are shorter than in sodium metal, and this indicates that even with these electropositive metals the bonding cannot be straightforwardly ionic.<ref name="Greenwood&Earnshaw" /> Other alkali metal arsenides not conforming to the formula M<sub>3</sub>As are known, such as LiAs, which has a metallic lustre and electrical conductivity indicating the presence of some [[metallic bond]]ing.<ref name="Greenwood&Earnshaw" /> The [[antimonide]]s are unstable and reactive as the Sb<sup>3−</sup> ion is a strong reducing agent; reaction of them with acids form the toxic and unstable gas [[stibine]] (SbH<sub>3</sub>).<ref>{{cite book |title=Outlines of Chemistry&nbsp;– A Textbook for College Students|last=Kahlenberg|first=Louis|publisher=READ BOOKS|year=2008|isbn=978-1-4097-6995-8|pages=324–325}}</ref> Indeed, they have some metallic properties, and the alkali metal antimonides of stoichiometry MSb involve antimony atoms bonded in a spiral Zintl structure.<ref name=King>{{cite book |last=King |first=R. Bruce |date=1995 |title=Inorganic Chemistry of Main Group Elements |publisher=Wiley-VCH |isbn=978-0-471-18602-1}}</ref> [[Bismuthide]]s are not even wholly ionic; they are [[intermetallic compound]]s containing partially metallic and partially ionic bonds.<ref>{{cite web |url=http://xray.chem.ualberta.ca/mar/ |title=Welcome to Arthur Mar's Research Group |date=1999–2013 |work=University of Alberta |access-date=24 June 2013 |archive-date=4 December 2012 |archive-url=https://web.archive.org/web/20121204000908/http://xray.chem.ualberta.ca/mar/ |url-status=dead }}</ref>

=== Oxides and chalcogenides ===
{{See also|Alkali metal oxide}}
{{multiple image
| align = right
| image1 = Rb9O2 cluster.png
| width1 = 150
| alt1 = The ball-and-stick diagram shows two regular octahedra which are connected to each other by one face. All nine vertices of the structure are purple spheres representing rubidium, and at the centre of each octahedron is a small red sphere representing oxygen.
| caption1 = {{chem2|Rb9O2}} cluster, composed of two regular [[octahedron|octahedra]] connected to each other by one face
| image2 = Cs11O3 cluster.png
| width2 = 150
| alt2 = The ball-and-stick diagram shows three regular octahedra where each octahedron is connected to both of the others by one face each. All three octahedra have one edge in common. All eleven vertices of the structure are violet spheres representing caesium, and at the centre of each octahedron is a small red sphere representing oxygen.
| caption2 = {{chem2|Cs11O3}} cluster, composed of three regular octahedra where each octahedron is connected to both of the others by one face each. All three octahedra have one edge in common.
| footer = 
}}
All the alkali metals react vigorously with [[oxygen]] at standard conditions. They form various types of oxides, such as simple [[oxide]]s (containing the O<sup>2−</sup> ion), [[peroxide]]s (containing the {{chem2|O2(2-)}} ion, where there is a [[single bond]] between the two oxygen atoms), [[superoxide]]s (containing the {{chem2|O2-}} ion), and many others. Lithium burns in air to form [[lithium oxide]], but sodium reacts with oxygen to form a mixture of [[sodium oxide]] and [[sodium peroxide]]. Potassium forms a mixture of [[potassium peroxide]] and [[potassium superoxide]], while rubidium and caesium form the superoxide exclusively. Their reactivity increases going down the group: while lithium, sodium and potassium merely burn in air, rubidium and caesium are [[pyrophoric]] (spontaneously catch fire in air).<ref name="alkalireact" />

The smaller alkali metals tend to polarise the larger anions (the peroxide and superoxide) due to their small size. This attracts the electrons in the more complex anions towards one of its constituent oxygen atoms, forming an oxide ion and an oxygen atom. This causes lithium to form the oxide exclusively on reaction with oxygen at room temperature. This effect becomes drastically weaker for the larger sodium and potassium, allowing them to form the less stable peroxides. Rubidium and caesium, at the bottom of the group, are so large that even the least stable superoxides can form. Because the superoxide releases the most energy when formed, the superoxide is preferentially formed for the larger alkali metals where the more complex anions are not polarised. The oxides and peroxides for these alkali metals do exist, but do not form upon direct reaction of the metal with oxygen at standard conditions.<ref name="alkalireact" /> In addition, the small size of the Li<sup>+</sup> and O<sup>2−</sup> ions contributes to their forming a stable ionic lattice structure. Under controlled conditions, however, all the alkali metals, with the exception of francium, are known to form their oxides, peroxides, and superoxides. The alkali metal peroxides and superoxides are powerful [[oxidising agent]]s. [[Sodium peroxide]] and [[potassium superoxide]] react with [[carbon dioxide]] to form the alkali metal carbonate and oxygen gas, which allows them to be used in [[submarine]] air purifiers; the presence of [[water vapour]], naturally present in breath, makes the removal of carbon dioxide by potassium superoxide even more efficient.<ref name=generalchemistry /><ref>{{cite journal |last1=Lindsay |first1=D. M. |last2=Garland |first2=D. A. |year=1987 |title=ESR spectra of matrix-isolated lithium superoxide |journal=The Journal of Physical Chemistry |volume=91 |issue=24 |pages=6158–61 |doi=10.1021/j100308a020}}</ref> All the stable alkali metals except lithium can form red [[ozonide]]s (MO<sub>3</sub>) through low-temperature reaction of the powdered anhydrous hydroxide with [[ozone]]: the ozonides may be then extracted using liquid [[ammonia]]. They slowly decompose at standard conditions to the superoxides and oxygen, and hydrolyse immediately to the hydroxides when in contact with water.<ref name="Greenwood&Earnshaw" />{{rp|85}} Potassium, rubidium, and caesium also form sesquioxides M<sub>2</sub>O<sub>3</sub>, which may be better considered peroxide disuperoxides, {{chem2|[(M+)4(O2(2-))(O2-)2]}}.<ref name="Greenwood&Earnshaw" />{{rp|85}}

Rubidium and caesium can form a great variety of suboxides with the metals in formal oxidation states below +1.<ref name="Greenwood&Earnshaw" />{{rp|85}} Rubidium can form Rb<sub>6</sub>O and Rb<sub>9</sub>O<sub>2</sub> (copper-coloured) upon oxidation in air, while caesium forms an immense variety of oxides, such as the ozonide CsO<sub>3</sub><ref>{{cite journal |doi= 10.1007/BF00845494|title= Synthesis of cesium ozonide through cesium superoxide|year= 1963|last1= Vol'nov|first1= I. I.|last2= Matveev|first2= V. V.|journal= Bulletin of the Academy of Sciences, USSR Division of Chemical Science|volume= 12|pages= 1040–1043|issue= 6}}</ref><ref>{{cite journal |doi= 10.1070/RC1971v040n02ABEH001903|title= Alkali and Alkaline Earth Metal Ozonides|year= 1971|last1= Tokareva|first1= S. A.|journal= Russian Chemical Reviews|volume= 40|pages= 165–174|bibcode= 1971RuCRv..40..165T|issue= 2|s2cid= 250883291}}</ref> and several brightly coloured [[suboxide]]s,<ref name=Simon>{{cite journal |last= Simon|first= A.|title= Group 1 and 2 Suboxides and Subnitrides – Metals with Atomic Size Holes and Tunnels|journal= Coordination Chemistry Reviews |year= 1997|volume= 163|pages= 253–270|doi= 10.1016/S0010-8545(97)00013-1}}</ref> such as Cs<sub>7</sub>O (bronze), Cs<sub>4</sub>O (red-violet), Cs<sub>11</sub>O<sub>3</sub> (violet), Cs<sub>3</sub>O (dark green),<ref>{{cite journal |doi= 10.1021/j150537a023|year= 1956|last1= Tsai|first1= Khi-Ruey|last2= Harris|first2= P. M.|last3= Lassettre |first3= E. N.|journal= Journal of Physical Chemistry|volume= 60|pages= 345–347|title=The Crystal Structure of Tricesium Monoxide|issue= 3}}</ref> CsO, Cs<sub>3</sub>O<sub>2</sub>,<ref>{{cite journal |doi= 10.1007/s11669-009-9636-5|title= Cs-O (Cesium-Oxygen)|year= 2009 |last1= Okamoto|first1= H.|journal= Journal of Phase Equilibria and Diffusion|volume= 31|pages= 86–87|s2cid= 96084147}}</ref> as well as Cs<sub>7</sub>O<sub>2</sub>.<ref>{{cite journal |doi= 10.1021/jp036432o|title= Characterization of Oxides of Cesium|year= 2004|last1= Band|first1= A.|last2= Albu-Yaron|first2= A.|last3= Livneh|first3= T.|last4= Cohen|first4= H.|last5= Feldman|first5= Y.|last6= Shimon |first6= L.|last7= Popovitz-Biro|first7= R.|last8= Lyahovitskaya|first8= V.|last9= Tenne|first9= R.|journal= The Journal of Physical Chemistry B|volume= 108|pages= 12360–12367|issue= 33}}</ref><ref>{{cite journal |doi= 10.1002/zaac.19472550110|title= Untersuchungen über das System Cäsium-Sauerstoff|year= 1947|last1= Brauer|first1= G.|journal= Zeitschrift für Anorganische Chemie|volume= 255|issue= 1–3|pages= 101–124}}</ref> The last of these may be heated under vacuum to generate Cs<sub>2</sub>O.<ref name="pubs.usgs" />

The alkali metals can also react analogously with the heavier chalcogens ([[sulfur]], [[selenium]], [[tellurium]], and [[polonium]]), and all the alkali metal chalcogenides are known (with the exception of francium's). Reaction with an excess of the chalcogen can similarly result in lower chalcogenides, with chalcogen ions containing chains of the chalcogen atoms in question. For example, sodium can react with sulfur to form the [[sulfide]] ([[sodium sulfide|Na<sub>2</sub>S]]) and various [[polysulfide]]s with the formula Na<sub>2</sub>S<sub>''x''</sub> (''x'' from 2 to 6), containing the {{chem|S|''x''|2-}} ions.<ref name=generalchemistry /> Due to the basicity of the Se<sup>2−</sup> and Te<sup>2−</sup> ions, the alkali metal [[selenide]]s and [[telluride (chemistry)|tellurides]] are alkaline in solution; when reacted directly with selenium and tellurium, alkali metal polyselenides and polytellurides are formed along with the selenides and tellurides with the {{chem|Se|''x''|2-}} and {{chem|Te|''x''|2-}} ions.<ref name="house2008">{{cite book |title= Inorganic chemistry |first= James E.|last= House |publisher= Academic Press |year= 2008 |isbn= 978-0-12-356786-4 |page= 524}}</ref> They may be obtained directly from the elements in liquid ammonia or when air is not present, and are colourless, water-soluble compounds that air oxidises quickly back to selenium or tellurium.<ref name="Greenwood&Earnshaw" />{{rp|766}} The alkali metal [[polonide]]s are all ionic compounds containing the Po<sup>2−</sup> ion; they are very chemically stable and can be produced by direct reaction of the elements at around 300–400&nbsp;°C.<ref name="Greenwood&Earnshaw" />{{rp|766}}<ref name="AEC-chem">{{cite book |last= Moyer |first= Harvey V. |contribution= Chemical Properties of Polonium |pages= 33–96 |title= Polonium |url= http://www.osti.gov/bridge/servlets/purl/4367751-nEJIbm/ |editor-last= Moyer |editor-first= Harvey V. |id= TID-5221 |doi= 10.2172/4367751 |year= 1956 |location= Oak Ridge, Tenn. |publisher= United States Atomic Energy Commission }}</ref><ref name="Bagnall">{{cite journal |first= K. W. |last= Bagnall |title= The Chemistry of Polonium |journal= Adv. Inorg. Chem. Radiochem. |year= 1962 |volume= 4 |pages= 197–229 |url= https://books.google.com/books?id=8qePsa3V8GQC&pg=PA197 |isbn= 978-0-12-023604-6 |doi= 10.1016/S0065-2792(08)60268-X |series= Advances in Inorganic Chemistry and Radiochemistry }}</ref>

=== Halides, hydrides, and pseudohalides ===
{{Main|Alkali metal halide}}
The alkali metals are among the most [[electropositive]] elements on the periodic table and thus tend to [[ionic bond|bond ionically]] to the most [[electronegative]] elements on the periodic table, the [[halogen]]s ([[fluorine]], [[chlorine]], [[bromine]], [[iodine]], and [[astatine]]), forming [[salt (chemistry)|salts]] known as the alkali metal halides. The reaction is very vigorous and can sometimes result in explosions.<ref name="Greenwood&Earnshaw" />{{rp|76}} All twenty stable alkali metal halides are known; the unstable ones are not known, with the exception of sodium astatide, because of the great instability and rarity of astatine and francium. The most well-known of the twenty is certainly [[sodium chloride]], otherwise known as common salt. All of the stable alkali metal halides have the formula MX where M is an alkali metal and X is a halogen. They are all white ionic crystalline solids that have high melting points.<ref name=rsc /><ref name="alkalireact" /> All the alkali metal halides are [[solubility|soluble]] in water except for [[lithium fluoride]] (LiF), which is insoluble in water due to its very high [[lattice enthalpy]]. The high lattice enthalpy of lithium fluoride is due to the small sizes of the Li<sup>+</sup> and F<sup>−</sup> ions, causing the [[electrostatic interaction]]s between them to be strong:<ref name=rsc /> a similar effect occurs for [[magnesium fluoride]], consistent with the diagonal relationship between lithium and magnesium.<ref name="Greenwood&Earnshaw" />{{rp|76}}

The alkali metals also react similarly with hydrogen to form ionic alkali metal hydrides, where the [[hydride]] anion acts as a [[pseudohalogen|pseudohalide]]: these are often used as reducing agents, producing hydrides, complex metal hydrides, or hydrogen gas.<ref name="Greenwood&Earnshaw" />{{rp|83}}<ref name="generalchemistry">{{cite book |last1=Averill |first1=Bruce A. |last2=Eldredge |first2=Patricia |title=Chemistry: Principles, Patterns, and Applications with Student Access Kit for Mastering General Chemistry |chapter-url=http://2012books.lardbucket.org/books/general-chemistry-principles-patterns-and-applications-v1.0/section_25_03.html |access-date=24 June 2013 |year=2007 |publisher=Prentice Hall |edition=1st |isbn=978-0-8053-3799-0 |chapter=21.3: The Alkali Metals}}</ref> Other pseudohalides are also known, notably the [[cyanide]]s. These are isostructural to the respective halides except for [[lithium cyanide]], indicating that the cyanide ions may rotate freely.<ref name="Greenwood&Earnshaw" />{{rp|322}} Ternary alkali metal halide oxides, such as Na<sub>3</sub>ClO, K<sub>3</sub>BrO (yellow), Na<sub>4</sub>Br<sub>2</sub>O, Na<sub>4</sub>I<sub>2</sub>O, and K<sub>4</sub>Br<sub>2</sub>O, are also known.<ref name="Greenwood&Earnshaw" />{{rp|83}} The polyhalides are rather unstable, although those of rubidium and caesium are greatly stabilised by the feeble polarising power of these extremely large cations.<ref name="Greenwood&Earnshaw" />{{rp|835}}

=== Coordination complexes ===
{{multiple image
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| image1 = 18-crown-6-potassium-3D-balls-A.png
| width1 = 150
| alt1 = 
| caption1 = [[18-crown-6]] coordinating a potassium ion
| image2 = Cryptate of potassium cation.jpg
| width2 = 150
| alt2 = 
| caption2 = Structure of [[2.2.2-Cryptand]] encapsulating a potassium cation (purple). At crystalline state, obtained with an X-ray diffraction.<ref>{{cite journal |last1=Alberto|first1=R. |last2=Ortner|first2=K. |last3=Wheatley|first3=N. |last4=Schibli|first4=R. |last5=Schubiger|first5=A. P. |title= Synthesis and properties of boranocarbonate: a convenient in situ CO source for the aqueous preparation of [<sup>99m</sup>Tc(OH<sub>2</sub>)<sub>3</sub>(CO)<sub>3</sub>]<sup>+</sup> |journal= [[J. Am. Chem. Soc.]] |year= 2001 |volume= 121 |pages= 3135–3136 |doi= 10.1021/ja003932b |pmid=11457025 |issue= 13}}</ref>
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}}
Alkali metal cations do not usually form [[coordination complex]]es with simple [[Lewis base]]s due to their low charge of just +1 and their relatively large size; thus the Li<sup>+</sup> ion forms most complexes and the heavier alkali metal ions form less and less (though exceptions occur for weak complexes).<ref name="Greenwood&Earnshaw" />{{rp|90}} Lithium in particular has a very rich coordination chemistry in which it exhibits [[coordination number]]s from 1 to 12, although octahedral hexacoordination is its preferred mode.<ref name="Greenwood&Earnshaw" />{{rp|90–1}} In [[aqueous solution]], the alkali metal ions exist as octahedral hexahydrate complexes [M(H<sub>2</sub>O)<sub>6</sub>]<sup>+</sup>, with the exception of the lithium ion, which due to its small size forms tetrahedral tetrahydrate complexes [Li(H<sub>2</sub>O)<sub>4</sub>]<sup>+</sup>; the alkali metals form these complexes because their ions are attracted by electrostatic forces of attraction to the polar water molecules. Because of this, [[anhydrous]] salts containing alkali metal cations are often used as [[desiccant]]s.<ref name=generalchemistry /> Alkali metals also readily form complexes with [[crown ether]]s (e.g. [[12-crown-4]] for Li<sup>+</sup>, [[15-crown-5]] for Na<sup>+</sup>, [[18-crown-6]] for K<sup>+</sup>, and [[21-crown-7]] for Rb<sup>+</sup>) and [[cryptand]]s due to electrostatic attraction.<ref name=generalchemistry />

=== Ammonia solutions ===
The alkali metals dissolve slowly in liquid [[ammonia]], forming ammoniacal solutions of solvated metal cation M<sup>+</sup> and [[solvated electron]] e<sup>−</sup>, which react to form hydrogen gas and the [[metal amide#Alkali metal amides|alkali metal amide]] (MNH<sub>2</sub>, where M represents an alkali metal): this was first noted by [[Humphry Davy]] in 1809 and rediscovered by W. Weyl in 1864. The process may be speeded up by a [[catalyst]]. Similar solutions are formed by the heavy divalent [[alkaline earth metal]]s [[calcium]], [[strontium]], [[barium]], as well as the divalent [[lanthanide]]s, [[europium]] and [[ytterbium]]. The amide salt is quite insoluble and readily precipitates out of solution, leaving intensely coloured ammonia solutions of the alkali metals. In 1907, [[Charles A. Kraus]] identified the colour as being due to the presence of [[solvated electron]]s, which contribute to the high electrical conductivity of these solutions. At low concentrations (below 3 M), the solution is dark blue and has ten times the conductivity of aqueous [[sodium chloride]]; at higher concentrations (above 3 M), the solution is copper-coloured and has approximately the conductivity of liquid metals like [[mercury (element)|mercury]].<ref name="Greenwood&Earnshaw" /><ref name=generalchemistry /><ref name="c&w">{{cite book |last1=Cotton |first1=F. A. |first2=G.|last2=Wilkinson |title=Advanced Inorganic Chemistry |year=1972 |publisher=John Wiley and Sons Inc |isbn=978-0-471-17560-5}}</ref> In addition to the alkali metal amide salt and solvated electrons, such ammonia solutions also contain the alkali metal cation (M<sup>+</sup>), the neutral alkali metal atom (M), [[diatomic molecule|diatomic]] alkali metal molecules (M<sub>2</sub>) and alkali metal anions (M<sup>−</sup>). These are unstable and eventually become the more thermodynamically stable alkali metal amide and hydrogen gas. Solvated electrons are powerful [[reducing agent]]s and are often used in chemical synthesis.<ref name=generalchemistry />

=== Organometallic ===

==== Organolithium ====
{{Main|Organolithium reagent}}
[[File:Butyllithium-hexamer-from-xtal-3D-balls-A.png|thumb|upright=1.15|Structure of the octahedral [[n-butyllithium|''n''-butyllithium]] hexamer, (C<sub>4</sub>H<sub>9</sub>Li)<sub>6</sub>.<ref>{{cite journal
|title= Structures of Classical Reagents in Chemical Synthesis: (nBuLi)<sub>6</sub>, (tBuLi)<sub>4</sub>, and the Metastable (tBuLi · Et<sub>2</sub>O)<sub>2</sub>
|last=T. Kottke|first=D. Stalke
|journal= Angew. Chem. Int. Ed. Engl.
|date= September 1993
|volume= 32
|issue= 4
|pages= 580–582
|doi= 10.1002/anie.199305801
|url=http://resolver.sub.uni-goettingen.de/purl?goescholar/3373}}</ref> The aggregates are held together by delocalised covalent bonds between lithium and the terminal carbon of the butyl chain.<ref>Elschenbroich, C. "Organometallics" (2006) Wiley-VCH: Weinheim. {{ISBN|3-527-29390-6}}.</ref> There is no direct lithium–lithium bonding in any organolithium compound.<ref name=King />{{rp|264}}]]
[[File:Phenyllithium-chain-from-xtal-Mercury-3D-balls.png|thumb|upright=1.15|Solid [[phenyllithium]] forms monoclinic crystals that can be described as consisting of dimeric Li<sub>2</sub>([[phenyl group|C<sub>6</sub>H<sub>5</sub>]])<sub>2</sub> subunits. The lithium atoms and the ''[[arene substitution pattern|ipso]]'' carbons of the phenyl rings form a planar four-membered ring. The plane of the phenyl groups is perpendicular to the plane of this Li<sub>2</sub>C<sub>2</sub> ring. Additional strong intermolecular bonding occurs between these phenyllithium dimers and the π electrons of the phenyl groups in the adjacent dimers, resulting in an infinite polymeric ladder structure.<ref>{{Cite journal |last1= Dinnebier |first1= R. E. |last2= Behrens |first2= U. |last3= Olbrich |first3= F. |title= Lewis Base-Free Phenyllithium: Determination of the Solid-State Structure by Synchrotron Powder Diffraction |journal= [[Journal of the American Chemical Society]] |year= 1998 |volume= 120 |issue= 7 |pages= 1430–1433 |doi= 10.1021/ja972816e}}</ref>]]

Being the smallest alkali metal, lithium forms the widest variety of and most stable [[organometallic compound]]s, which are bonded covalently. [[Organolithium reagent|Organolithium]] compounds are electrically non-conducting volatile solids or liquids that melt at low temperatures, and tend to form [[oligomer]]s with the structure (RLi)<sub>''x''</sub> where R is the organic group. As the electropositive nature of lithium puts most of the [[charge density]] of the bond on the carbon atom, effectively creating a [[carbanion]], organolithium compounds are extremely powerful [[base (chemistry)|bases]] and [[carbon nucleophile|nucleophiles]]. For use as bases, [[butyllithium]]s are often used and are commercially available. An example of an organolithium compound is [[methyllithium]] ((CH<sub>3</sub>Li)<sub>''x''</sub>), which exists in tetrameric (''x'' = 4, tetrahedral) and hexameric (''x'' = 6, octahedral) forms.<ref name=generalchemistry /><ref name=Brown1957>{{cite journal |last1=Brown|first1=T. L. |last2=Rogers|first2=M. T. |title= The Preparation and Properties of Crystalline Lithium Alkyls |journal= Journal of the American Chemical Society |year= 1957 |volume= 79 |issue= 8 |pages= 1859–1861 |doi= 10.1021/ja01565a024}}</ref> Organolithium compounds, especially ''n''-butyllithium, are useful reagents in organic synthesis, as might be expected given lithium's diagonal relationship with magnesium, which plays an important role in the [[Grignard reaction]].<ref name="Greenwood&Earnshaw" />{{rp|102}} For example, alkyllithiums and aryllithiums may be used to synthesise [[aldehyde]]s and [[ketone]]s by reaction with metal [[carbonyl]]s. The reaction with [[nickel tetracarbonyl]], for example, proceeds through an unstable acyl nickel carbonyl complex which then undergoes [[electrophilic substitution]] to give the desired aldehyde (using H<sup>+</sup> as the electrophile) or ketone (using an alkyl halide) product.<ref name="Greenwood&Earnshaw" />{{rp|105}}

:<chem>LiR \ + \ Ni(CO)4 \ \longrightarrow Li^{+}[RCONi(CO)3]^{-}</chem>
:<chem>Li^{+}[RCONi(CO)3]^{-}->[\ce{H^{+}}][\ce{solvent}] \ Li^{+} \ + \ RCHO \ + \ [(solvent)Ni(CO)3]</chem>
:<chem>Li^{+}[RCONi(CO)3]^{-}->[\ce{R^{'}Br}][\ce{solvent}] \ Li^{+} \ + \ RR^{'}CO \ + \ [(solvent)Ni(CO)3]</chem>

Alkyllithiums and aryllithiums may also react with ''N'',''N''-disubstituted [[amide]]s to give aldehydes and ketones, and symmetrical ketones by reacting with [[carbon monoxide]]. They thermally decompose to eliminate a β-hydrogen, producing [[alkene]]s and [[lithium hydride]]: another route is the reaction of [[ether]]s with alkyl- and aryllithiums that act as strong bases.<ref name="Greenwood&Earnshaw" />{{rp|105}} In non-polar solvents, aryllithiums react as the carbanions they effectively are, turning carbon dioxide to aromatic [[carboxylic acid]]s (ArCO<sub>2</sub>H) and aryl ketones to tertiary carbinols (Ar'<sub>2</sub>C(Ar)OH). Finally, they may be used to synthesise other organometallic compounds through metal-halogen exchange.<ref name="Greenwood&Earnshaw" />{{rp|106}}

==== Heavier alkali metals ====
Unlike the organolithium compounds, the organometallic compounds of the heavier alkali metals are predominantly ionic. The application of [[organosodium chemistry|organosodium]] compounds in chemistry is limited in part due to competition from [[organolithium compound]]s, which are commercially available and exhibit more convenient reactivity. The principal organosodium compound of commercial importance is [[sodium cyclopentadienide]]. [[Sodium tetraphenylborate]] can also be classified as an organosodium compound since in the solid state sodium is bound to the aryl groups. Organometallic compounds of the higher alkali metals are even more reactive than organosodium compounds and of limited utility. A notable reagent is [[Schlosser's base]], a mixture of [[n-Butyllithium|''n''-butyllithium]] and [[potassium tert-butoxide|potassium ''tert''-butoxide]]. This reagent reacts with [[propene]] to form the compound [[allylpotassium]] (KCH<sub>2</sub>CHCH<sub>2</sub>). [[cis-2-butene|''cis''-2-Butene]] and [[trans-2-butene|''trans''-2-butene]] equilibrate when in contact with alkali metals. Whereas [[isomerisation]] is fast with lithium and sodium, it is slow with the heavier alkali metals. The heavier alkali metals also favour the [[steric hindrance|sterically]] congested conformation.<ref>{{cite journal |title= Superbases for organic synthesis |last=Schlosser|first=Manfred|journal= Pure Appl. Chem. |volume= 60 |issue= 11 |pages= 1627–1634 |year= 1988 |doi= 10.1351/pac198860111627|s2cid=39746336|url= http://old.iupac.org/publications/pac/1988/pdf/6011x1627.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://old.iupac.org/publications/pac/1988/pdf/6011x1627.pdf |archive-date=2022-10-09 |url-status=live}}</ref> Several crystal structures of organopotassium compounds have been reported, establishing that they, like the sodium compounds, are polymeric.<ref name=Klett>{{cite journal |doi=10.1002/ejic.201000983|title=Synthesis and Structures of \(Trimethylsilyl)methyl]sodium and -potassium with Bi- and Tridentate N-Donor Ligands|year=2011|last1=Clegg|first1=William|last2=Conway|first2=Ben|last3=Kennedy|first3=Alan R.|last4=Klett|first4=Jan|last5=Mulvey|first5=Robert E.|last6=Russo|first6=Luca|journal=European Journal of Inorganic Chemistry|volume=2011|issue=5|pages=721–726|url=https://www.researchgate.net/publication/210286264|doi-access=}}</ref> Organosodium, organopotassium, organorubidium and organocaesium compounds are all mostly ionic and are insoluble (or nearly so) in nonpolar solvents.<ref name=generalchemistry />

Alkyl and aryl derivatives of sodium and potassium tend to react with air. They cause the cleavage of [[ether]]s, generating alkoxides. Unlike alkyllithium compounds, alkylsodiums and alkylpotassiums cannot be made by reacting the metals with alkyl halides because [[Wurtz coupling]] occurs:<ref name=King />{{rp|265}}
:RM + R'X → R–R' + MX

As such, they have to be made by reacting [[organomercury compound|alkylmercury]] compounds with sodium or potassium metal in inert hydrocarbon solvents. While methylsodium forms tetramers like methyllithium, methylpotassium is more ionic and has the [[nickel arsenide]] structure with discrete methyl anions and potassium cations.<ref name=King />{{rp|265}}

The alkali metals and their hydrides react with acidic hydrocarbons, for example [[cyclopentadiene]]s and terminal alkynes, to give salts. Liquid ammonia, ether, or hydrocarbon solvents are used, the most common of which being [[tetrahydrofuran]]. The most important of these compounds is [[sodium cyclopentadienide]], NaC<sub>5</sub>H<sub>5</sub>, an important precursor to many transition metal cyclopentadienyl derivatives.<ref name=King />{{rp|265}} Similarly, the alkali metals react with [[cyclooctatetraene]] in tetrahydrofuran to give alkali metal [[cyclooctatetraenide]]s; for example, [[dipotassium cyclooctatetraenide]] (K<sub>2</sub>C<sub>8</sub>H<sub>8</sub>) is an important precursor to many metal cyclooctatetraenyl derivatives, such as [[uranocene]].<ref name=King />{{rp|266}} The large and very weakly polarising alkali metal cations can stabilise large, aromatic, polarisable radical anions, such as the dark-green [[sodium naphthalenide]], Na<sup>+</sup>[C<sub>10</sub>H<sub>8</sub>•]<sup>−</sup>, a strong reducing agent.<ref name=King />{{rp|266}}

==Representative reactions of alkali metals==
===Reaction with oxygen===

Upon reacting with oxygen, alkali metals form [[oxide]]s, [[peroxide]]s, [[superoxide]]s and [[suboxide]]s. However, the first three are more common. The table below<ref name="miessler">"Inorganic Chemistry" by Gary L. Miessler and Donald A. Tar, 6th edition, Pearson</ref> shows the types of compounds formed in reaction with oxygen. The compound in brackets represents the minor product of combustion.
{| class="wikitable"
|-
|'''Alkali metal'''||'''Oxide'''||'''Peroxide'''||'''Superoxide'''
|-
|[[Lithium|Li]]||Li<sub>2</sub>O||(Li<sub>2</sub>O<sub>2</sub>)||
|-
|[[Sodium|Na]]||(Na<sub>2</sub>O)||Na<sub>2</sub>O<sub>2</sub>||
|-
|[[Potassium|K]]|| || ||KO<sub>2</sub>
|-
|[[Rubidium|Rb]]|| || ||RbO<sub>2</sub>
|-
|[[Caesium|Cs]]|| || ||CsO<sub>2</sub>
|}
The alkali metal peroxides are ionic compounds that are unstable in water. The peroxide anion is weakly bound to the cation, and it is hydrolysed, forming stronger covalent bonds.
:Na<sub>2</sub>O<sub>2</sub> + 2H<sub>2</sub>O → 2NaOH + H<sub>2</sub>O<sub>2</sub>
The other oxygen compounds are also unstable in water.
:2KO<sub>2</sub> + 2H<sub>2</sub>O → 2KOH + H<sub>2</sub>O<sub>2</sub> + O<sub>2</sub><ref>Kumar De, Anil (2007). A Text Book of Inorganic Chemistry. New Age International. p. 247. {{ISBN|978-8122413847}}.</ref>
:Li<sub>2</sub>O + H<sub>2</sub>O → 2LiOH

===Reaction with sulfur===

With sulfur, they form [[sulfide]]s and [[polysulfide]]s.<ref>"The chemistry of the Elements" by Greenwood and Earnshaw, 2nd edition, Elsevier</ref>
:2Na + 1/8S<sub>8</sub> → Na<sub>2</sub>S + 1/8S<sub>8</sub> → Na<sub>2</sub>S<sub>2</sub>...Na<sub>2</sub>S<sub>7</sub>
Because alkali metal sulfides are essentially salts of a weak acid and a strong base, they form basic solutions.
:S<sup>2-</sup> + H<sub>2</sub>O → HS<sup>−</sup> + HO<sup>−</sup>
:HS<sup>−</sup> + H<sub>2</sub>O → H<sub>2</sub>S + HO<sup>−</sup>

===Reaction with nitrogen===

Lithium is the only metal that combines directly with nitrogen at room temperature.
:3Li + 1/2N<sub>2</sub> → Li<sub>3</sub>N
Li<sub>3</sub>N can react with water to liberate ammonia.
:Li<sub>3</sub>N + 3H<sub>2</sub>O → 3LiOH + NH<sub>3</sub>

===Reaction with hydrogen===

With hydrogen, alkali metals form saline [[hydride]]s that hydrolyse in water. 
:<chem>2 Na \ + H2 \ ->[\ce{\Delta}] \ 2 NaH</chem>
:<chem>2 NaH \ + \ 2 H2O \ \longrightarrow \ 2 NaOH \ + \ H2 \uparrow</chem>

===Reaction with carbon===

Lithium is the only metal that reacts directly with carbon to give [[dilithium acetylide]]. Na and K can react with [[acetylene]] to give [[acetylide]]s.<ref>"Inorganic Chemistry" by Cotton and Wilkinson</ref>
:<chem>2 Li \ + \ 2 C \ \longrightarrow \ Li2C2</chem>
:<p style="line-height: 1.6;vertical-align: text-bottom; " ><chem>2 Na \ + \ 2 C2H2 \ ->[\ce{150 \ ^{o}C}] \ 2 NaC2H \ + \ H2</chem></p>
:<p style="line-height: 1.6;vertical-align: text-bottom; " ><chem>2 Na \ + \ 2 NaC2H \ ->[\ce{220 \ ^{o}C}] \ 2 Na2C2 \ + \ H2</chem></p>

===Reaction with water===

On reaction with water, they generate [[hydroxide]] ions and [[hydrogen]] gas. This reaction is vigorous and highly exothermic and the hydrogen resulted may ignite in air or even explode in the case of Rb and Cs.<ref name="miessler"/>
:Na + H<sub>2</sub>O → NaOH + 1/2H<sub>2</sub>

===Reaction with other salts===

The alkali metals are very good reducing agents. They can reduce metal cations that are less electropositive. [[Titanium]] is produced industrially by the reduction of [[titanium tetrachloride]] with Na at 400&nbsp;°C ([[van Arkel–de Boer process]]).
:TiCl<sub>4</sub> + 4Na → 4NaCl + Ti

===Reaction with organohalide compounds===

Alkali metals react with halogen derivatives to generate hydrocarbon via the [[Wurtz reaction]].
:2CH<sub>3</sub>-Cl + 2Na → H<sub>3</sub>C-CH<sub>3</sub> + 2NaCl

===Alkali metals in liquid ammonia===

Alkali metals dissolve in liquid [[ammonia]] or other donor solvents like aliphatic [[amine]]s or [[hexamethylphosphoramide]] to give blue solutions. These solutions are believed to contain free electrons.<ref name="miessler"/>
:Na + xNH<sub>3</sub> → Na<sup>+</sup> + e(NH<sub>3</sub>)<sub>x</sub><sup>−</sup>
Due to the presence of [[solvated electron]]s, these solutions are very powerful reducing agents used in organic synthesis.

[[File:Na in lie. am.jpg|thumb|upright=1.25|centre|Reduction reactions using sodium in liquid ammonia]]

Reaction 1) is known as [[Birch reduction]].
Other reductions<ref name="miessler"/> that can be carried by these solutions are:
:S<sub>8</sub> + 2e<sup>−</sup> → S<sub>8</sub><sup>2-</sup> 
:Fe(CO)<sub>5</sub> + 2e<sup>−</sup> → Fe(CO)<sub>4</sub><sup>2-</sup> + CO

== Extensions ==
[[File:Atomic radius of alkali metals and alkaline earth metals.svg|thumb|upright=1.12|[[Empirical evidence|Empirical]] (Na–Cs, Mg–Ra) and predicted (Fr–Uhp, Ubn–Uhh) atomic radius of the alkali and alkaline earth metals from the [[period 3 element|third]] to the [[period 9 element|ninth period]], measured in [[angstrom]]s<ref name="Uue" />{{rp|1730}}<ref name="pyykko" />]]
Although francium is the heaviest alkali metal that has been discovered, there has been some theoretical work predicting the physical and chemical characteristics of hypothetical heavier alkali metals. Being the first [[period 8 element]], the undiscovered element [[ununennium]] (element 119) is predicted to be the next alkali metal after francium and behave much like their lighter [[Congener (chemistry)|congeners]]; however, it is also predicted to differ from the lighter alkali metals in some properties.<ref name="Uue" />{{rp|1729–1730}} Its chemistry is predicted to be closer to that of potassium<ref name=EB /> or rubidium<ref name="Uue" />{{rp|1729–1730}} instead of caesium or francium. This is unusual as [[periodic trends]], ignoring relativistic effects would predict ununennium to be even more reactive than caesium and francium. This lowered [[reactivity (chemistry)|reactivity]] is due to the relativistic stabilisation of ununennium's valence electron, increasing ununennium's first ionisation energy and decreasing the [[metallic radius|metallic]] and [[ionic radius|ionic radii]];<ref name="EB" /> this effect is already seen for francium.<ref name="Uue" />{{rp|1729–1730}} This assumes that ununennium will behave chemically as an alkali metal, which, although likely, may not be true due to relativistic effects.<ref name="tanm">{{cite web|url=http://lch.web.psi.ch/files/lectures/TexasA&M/TexasA&M.pdf |title=Gas Phase Chemistry of Superheavy Elements |last=Gäggeler |first=Heinz W. |date=5–7 November 2007 |work=Lecture Course Texas A&M |access-date=26 February 2012 |url-status=dead |archive-url=https://web.archive.org/web/20120220090755/http://lch.web.psi.ch/files/lectures/TexasA%26M/TexasA%26M.pdf |archive-date=20 February 2012 }}</ref> The relativistic stabilisation of the 8s orbital also increases ununennium's [[electron affinity]] far beyond that of caesium and francium; indeed, ununennium is expected to have an electron affinity higher than all the alkali metals lighter than it. Relativistic effects also cause a very large drop in the [[polarisability]] of ununennium.<ref name="Uue" />{{rp|1729–1730}} On the other hand, ununennium is predicted to continue the trend of melting points decreasing going down the group, being expected to have a melting point between 0&nbsp;°C and 30&nbsp;°C.<ref name="Uue" />{{rp|1724}}

[[File:Electron affinity of alkali metals.svg|thumb|left|Empirical (Na–Fr) and predicted (Uue) electron affinity of the alkali metals from the third to the [[period 8 element|eighth period]], measured in [[electron volt]]s<ref name="Uue" />{{rp|1730}}<ref name="pyykko" />]]
The stabilisation of ununennium's valence electron and thus the contraction of the 8s orbital cause its atomic radius to be lowered to 240&nbsp;[[picometer|pm]],<ref name="Uue" />{{rp|1729–1730}} very close to that of rubidium (247&nbsp;pm),<ref name=rsc /> so that the chemistry of ununennium in the +1 oxidation state should be more similar to the chemistry of rubidium than to that of francium. On the other hand, the ionic radius of the Uue<sup>+</sup> ion is predicted to be larger than that of Rb<sup>+</sup>, because the 7p orbitals are destabilised and are thus larger than the p-orbitals of the lower shells. Ununennium may also show the +3<ref name="Uue" />{{rp|1729–1730}} and +5 [[oxidation state]]s,<ref>{{cite journal |last1=Cao |first1=Chang-Su |last2=Hu |first2=Han-Shi |last3=Schwarz |first3=W. H. Eugen |last4=Li |first4=Jun |date=2022 |title=Periodic Law of Chemistry Overturns for Superheavy Elements |type=preprint |url=https://chemrxiv.org/engage/chemrxiv/article-details/63730be974b7b6d84cfdda35 |journal=[[ChemRxiv]] |volume= |issue= |pages= |doi=10.26434/chemrxiv-2022-l798p |access-date=16 November 2022}}</ref> which are not seen in any other alkali metal,<ref name="Greenwood&Earnshaw" />{{rp|28}} in addition to the +1 oxidation state that is characteristic of the other alkali metals and is also the main oxidation state of all the known alkali metals: this is because of the destabilisation and expansion of the 7p<sub>3/2</sub> spinor, causing its outermost electrons to have a lower ionisation energy than what would otherwise be expected.<ref name="Greenwood&Earnshaw" />{{rp|28}}<ref name="Uue" />{{rp|1729–1730}} Indeed, many ununennium compounds are expected to have a large [[covalent]] character, due to the involvement of the 7p<sub>3/2</sub> electrons in the bonding.<ref name="Thayer">{{cite book |last1=Thayer |first1=John S. |chapter=Relativistic Effects and the Chemistry of the Heavier Main Group Elements |title=Relativistic Methods for Chemists |year=2010 |pages=81, 84 |doi=10.1007/978-1-4020-9975-5_2 |volume=10 |isbn=978-1-4020-9974-8 |series=Challenges and Advances in Computational Chemistry and Physics }}</ref>

[[File:Ionization energy of alkali metals and alkaline earth metals.svg|thumb|Empirical (Na–Fr, Mg–Ra) and predicted (Uue–Uhp, Ubn–Uhh) ionisation energy of the alkali and alkaline earth metals from the third to the ninth period, measured in electron volts<ref name="Uue" />{{rp|1730}}<ref name="pyykko" />]]
Not as much work has been done predicting the properties of the alkali metals beyond ununennium. Although a simple extrapolation of the periodic table (by the [[Aufbau principle]]) would put element 169, unhexennium, under ununennium, Dirac-Fock calculations predict that the next element after ununennium with alkali-metal-like properties may be element 165, unhexpentium, which is predicted to have the electron configuration [Og] 5g<sup>18</sup> 6f<sup>14</sup> 7d<sup>10</sup> 8s<sup>2</sup> 8p<sub>1/2</sub><sup>2</sup> 9s<sup>1</sup>.<ref name="Uue" />{{rp|1729–1730}}<ref name="pyykko">{{cite journal |last1=Pyykkö|first1=Pekka|title=A suggested periodic table up to Z ≤ 172, based on Dirac–Fock calculations on atoms and ions|journal=Physical Chemistry Chemical Physics|volume=13|issue=1|pages=161–8|year=2011|pmid=20967377|doi=10.1039/c0cp01575j|bibcode= 2011PCCP...13..161P|url=https://www.researchgate.net/publication/47521056}}</ref> This element would be intermediate in properties between an alkali metal and a [[group 11 element]], and while its physical and atomic properties would be closer to the former, its chemistry may be closer to that of the latter. Further calculations show that unhexpentium would follow the trend of increasing ionisation energy beyond caesium, having an ionisation energy comparable to that of sodium, and that it should also continue the trend of decreasing atomic radii beyond caesium, having an atomic radius comparable to that of potassium.<ref name="Uue" />{{rp|1729–1730}} However, the 7d electrons of unhexpentium may also be able to participate in chemical reactions along with the 9s electron, possibly allowing oxidation states beyond +1, whence the likely transition metal behaviour of unhexpentium.<ref name="Uue" />{{rp|1732–1733}}<ref name=BFricke>{{cite journal |last1=Fricke |first1=Burkhard |year=1975 |title=Superheavy elements: a prediction of their chemical and physical properties |journal=Recent Impact of Physics on Inorganic Chemistry |volume=21 |pages=[https://archive.org/details/recentimpactofph0000unse/page/89 89–144] |doi=10.1007/BFb0116498 |url=https://archive.org/details/recentimpactofph0000unse/page/89 |access-date=4 October 2013 |series=Structure and Bonding |isbn=978-3-540-07109-9 }}</ref> Due to the alkali and [[alkaline earth metal]]s both being [[s-block]] elements, these predictions for the trends and properties of ununennium and unhexpentium also mostly hold quite similarly for the corresponding alkaline earth metals [[unbinilium]] (Ubn) and unhexhexium (Uhh).<ref name="Uue" />{{rp|1729–1733}} Unsepttrium, element 173, may be an even better heavier homologue of ununennium; with a predicted electron configuration of [Usb] 6g<sup>1</sup>, it returns to the alkali-metal-like situation of having one easily removed electron far above a closed p-shell in energy, and is expected to be even more reactive than caesium.<ref name="BFricke1977">{{cite journal |last1=Fricke |first1=Burkhard |year=1977 |title=Dirac-Fock-Slater calculations for the elements Z = 100, fermium, to Z = 173 |journal=Recent Impact of Physics on Inorganic Chemistry |volume=19 |pages=83–192 |doi=10.1016/0092-640X(77)90010-9 |url=http://kobra.bibliothek.uni-kassel.de/bitstream/urn:nbn:de:hebis:34-2008071622807/1/Fricke_Dirac_1977.pdf |access-date=25 February 2016 |bibcode=1977ADNDT..19...83F |archive-url=https://web.archive.org/web/20160322072636/http://kobra.bibliothek.uni-kassel.de/bitstream/urn:nbn:de:hebis:34-2008071622807/1/Fricke_Dirac_1977.pdf |archive-date=22 March 2016 |url-status=dead }}</ref><ref name=primefan>{{cite web |url=http://www.primefan.ru/stuff/chem/ptable/ptable.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://www.primefan.ru/stuff/chem/ptable/ptable.pdf |archive-date=2022-10-09 |url-status=live |title=Есть ли граница у таблицы Менделеева? |trans-title=Is there a boundary to the Mendeleev table? |last=Kul'sha |first=A. V. |website=www.primefan.ru |access-date=8 September 2018 |language=ru}}</ref>

The probable properties of further alkali metals beyond unsepttrium have not been explored yet as of 2019, and they may or may not be able to exist.<ref name=pyykko /> In periods 8 and above of the periodic table, relativistic and shell-structure effects become so strong that extrapolations from lighter congeners become completely inaccurate. In addition, the relativistic and shell-structure effects (which stabilise the s-orbitals and destabilise and expand the d-, f-, and g-orbitals of higher shells) have opposite effects, causing even larger difference between relativistic and non-relativistic calculations of the properties of elements with such high atomic numbers.<ref name="Uue" />{{rp|1732–1733}} Interest in the chemical properties of ununennium, unhexpentium, and unsepttrium stems from the fact that they are located close to the expected locations of [[island of stability|islands of stability]], centered at elements 122 (<sup>306</sup>Ubb) and 164 (<sup>482</sup>Uhq).<ref name=Kratz>{{cite conference |last1=Kratz |first1=J. V. |date=5 September 2011 |title=The Impact of Superheavy Elements on the Chemical and Physical Sciences |url=http://tan11.jinr.ru/pdf/06_Sep/S_1/02_Kratz.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://tan11.jinr.ru/pdf/06_Sep/S_1/02_Kratz.pdf |archive-date=2022-10-09 |url-status=live |conference=4th International Conference on the Chemistry and Physics of the Transactinide Elements |access-date=27 August 2013}}</ref><ref>[http://www.eurekalert.org/pub_releases/2008-04/acs-nse031108.php Nuclear scientists eye future landfall on a second 'island of stability']. EurekAlert! (2008-04-06). Retrieved on 2016-11-25.</ref><ref>{{cite journal |doi=10.1007/BF01406719 |volume=228 |issue=5 |title=Investigation of the stability of superheavy nuclei around Z=114 and Z=164 |journal=Zeitschrift für Physik |pages=371–386|year=1969 |last1=Grumann |first1=Jens |last2=Mosel |first2=Ulrich |last3=Fink |first3=Bernd |last4=Greiner |first4=Walter |bibcode=1969ZPhy..228..371G |s2cid=120251297 }}</ref>

== Pseudo-alkali metals ==
Many other substances are similar to the alkali metals in their tendency to form monopositive cations. Analogously to the [[pseudohalogen]]s, they have sometimes been called "pseudo-alkali metals". These substances include some elements and many more [[polyatomic ion]]s; the polyatomic ions are especially similar to the alkali metals in their large size and weak polarising power.<ref name=pseudo />

=== Hydrogen ===
The element [[hydrogen]], with one electron per neutral atom, is usually placed at the top of Group 1 of the periodic table because of its electron configuration. But hydrogen is not normally considered to be an alkali metal.<ref name="iupac" />  [[Metallic hydrogen]], which only exists at very high pressures, is known for its electrical and magnetic properties, not its chemical properties.<ref name="Folden">{{cite web |url=http://cyclotron.tamu.edu/smp/The%20Heaviest%20Elements%20in%20the%20Universe.pdf |title=The Heaviest Elements in the Universe |last=Folden|first=Cody |date=31 January 2009 |work=Saturday Morning Physics at Texas A&M |access-date=9 March 2012 |url-status=dead |archive-url=https://web.archive.org/web/20140810213232/http://cyclotron.tamu.edu/smp/The%20Heaviest%20Elements%20in%20the%20Universe.pdf |archive-date=10 August 2014}}</ref> Under typical conditions, pure hydrogen exists as a [[diatomic]] gas consisting of two atoms per molecule (H<sub>2</sub>);<ref>{{cite book |last=Emsley|first=J. |title= The Elements |publisher= Oxford: Clarendon Press |year= 1989 |pages= 22–23 }}</ref><!--It is uncertain if this reference (from [[diatomic molecule]]) refers to astatine usually not being considered with the other halogens or the list of elements that form diatomic molecules.--> however, the alkali metals form diatomic molecules (such as [[dilithium]], Li<sub>2</sub>) only at high temperatures, when they are in the gaseous state.<ref>Winter, Mark J. (1994) ''Chemical Bonding'', Oxford University Press, {{ISBN|0-19-855694-2}}</ref>

Hydrogen, like the alkali metals, has one [[valence electron]]<ref name=King/> and reacts easily with the [[halogen]]s,<ref name=King/> but the similarities mostly end there because of the small size of a bare proton H<sup>+</sup> compared to the alkali metal cations.<ref name=King/> Its placement above lithium is primarily due to its [[electron configuration]].<ref name="iupac">{{cite web |url=http://old.iupac.org/reports/periodic_table/ |title=International Union of Pure and Applied Chemistry > Periodic Table of the Elements |publisher=IUPAC |access-date=1 May 2011}}</ref> It is sometimes placed above [[fluorine]] due to their similar chemical properties, though the resemblance is likewise not absolute.<ref name="hydrogen">{{cite journal |last=Cronyn |first=Marshall W. |title=The Proper Place for Hydrogen in the Periodic Table |journal=Journal of Chemical Education |volume=80 |issue=8 |date=August 2003 |pages=947–951 |doi=10.1021/ed080p947 |url=http://www.reed.edu/reed_magazine/summer2009/columns/noaa/downloads/CronynHydrogen.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://www.reed.edu/reed_magazine/summer2009/columns/noaa/downloads/CronynHydrogen.pdf |archive-date=2022-10-09 |url-status=live |bibcode=2003JChEd..80..947C}}</ref>

The first ionisation energy of hydrogen (1312.0 [[kilojoule per mole|kJ/mol]]) is much higher than that of the alkali metals.<ref name="huheey">Huheey, J.E.; Keiter, E.A. and Keiter, R.L. (1993) ''Inorganic Chemistry: Principles of Structure and Reactivity'', 4th edition, HarperCollins, New York, USA.</ref><ref name="macmillan">James, A.M. and Lord, M.P. (1992) ''Macmillan's Chemical and Physical Data'', Macmillan, London, UK.</ref> As only one additional electron is required to fill in the outermost shell of the hydrogen atom, hydrogen often behaves like a halogen, forming the negative [[hydride]] ion, and is very occasionally considered to be a halogen on that basis. (The alkali metals can also form negative ions, known as [[alkalide]]s, but these are little more than laboratory curiosities, being unstable.)<ref name="HNa" /><ref name="HNa-theory" /> An argument against this placement is that formation of hydride from hydrogen is endothermic, unlike the exothermic formation of halides from halogens. The radius of the H<sup>−</sup> anion also does not fit the trend of increasing size going down the halogens: indeed, H<sup>−</sup> is very diffuse because its single proton cannot easily control both electrons.<ref name=King />{{rp|15–6}} It was expected for some time that liquid hydrogen would show metallic properties;<ref name=hydrogen /> while this has been shown to not be the case, under extremely high [[pressure]]s, such as those found at the cores of [[Jupiter]] and [[Saturn]], hydrogen does become metallic and behaves like an alkali metal; in this phase, it is known as [[metallic hydrogen]].<ref>{{cite journal |last1=Wigner|first1=E.|last2=Huntington|first2=H. B.|year=1935|title=On the possibility of a metallic modification of hydrogen|journal=[[Journal of Chemical Physics]]|volume=3 |page=764|doi=10.1063/1.1749590|bibcode= 1935JChPh...3..764W|issue=12}}</ref> The [[resistivity|electrical resistivity]] of liquid [[metallic hydrogen]] at 3000 K is approximately equal to that of liquid [[rubidium]] and [[caesium]] at 2000 K at the respective pressures when they undergo a nonmetal-to-metal transition.<ref>{{cite journal |last1=Nellis |first1=W. J. |last2=Weir |first2=S. T. |last3=Mitchell |first3=A. C. |year=1999 |title=Metallization of fluid hydrogen at 140 GPa (1.4 Mbar) by shock compression |journal=Shock Waves |volume=9 |issue=5 |pages=301–305 |doi= 10.1007/s001930050189|bibcode= 1999ShWav...9..301N|s2cid=97261131 |url=https://zenodo.org/record/1232611 }}</ref>

The 1s<sup>1</sup> electron configuration of hydrogen, while analogous to that of the alkali metals (ns<sup>1</sup>), is unique because there is no 1p subshell. Hence it can lose an electron to form the [[hydron (chemistry)|hydron]] H<sup>+</sup>, or gain one to form the [[hydride]] ion H<sup>−</sup>.<ref name="Greenwood&Earnshaw" />{{rp|43}} In the former case it resembles superficially the alkali metals; in the latter case, the halogens, but the differences due to the lack of a 1p subshell are important enough that neither group fits the properties of hydrogen well.<ref name="Greenwood&Earnshaw" />{{rp|43}} Group 14 is also a good fit in terms of thermodynamic properties such as [[ionisation energy]] and [[electron affinity]], but hydrogen cannot be tetravalent. Thus none of the three placements are entirely satisfactory, although group 1 is the most common placement (if one is chosen) because of the electron configuration and the fact that the hydron is by far the most important of all monatomic hydrogen species, being the foundation of acid-base chemistry.<ref name=hydrogen /> As an example of hydrogen's unorthodox properties stemming from its unusual electron configuration and small size, the hydrogen ion is very small (radius around 150&nbsp;[[femtometre|fm]] compared to the 50–220&nbsp;pm size of most other atoms and ions) and so is nonexistent in condensed systems other than in association with other atoms or molecules. Indeed, transferring of protons between chemicals is the basis of [[acid-base chemistry]].<ref name="Greenwood&Earnshaw" />{{rp|43}} Also unique is hydrogen's ability to form [[hydrogen bond]]s, which are an effect of charge-transfer, [[electrostatic]], and electron correlative contributing phenomena.<ref name=hydrogen /> While analogous lithium bonds are also known, they are mostly electrostatic.<ref name=hydrogen /> Nevertheless, hydrogen can take on the same structural role as the alkali metals in some molecular crystals, and has a close relationship with the lightest alkali metals (especially lithium).<ref>{{cite journal |last1=Cousins |first1=David M. |last2=Davidson |first2=Matthew G. |last3=García-Vivó |first3=Daniel |date=2013 |title=Unprecedented participation of a four-coordinate hydrogen atom in the cubane core of lithium and sodium phenolates |url=http://pubs.rsc.org/en/content/articlepdf/2013/cc/c3cc47393g |journal=Chem. Commun. |volume=49 |issue=100 |pages=11809–11811 |doi=10.1039/c3cc47393g |pmid=24217230 |access-date=7 August 2014|doi-access=free }}</ref>

=== Ammonium and derivatives ===
[[File:Hydrochloric acid ammonia.jpg|thumb|right|Similarly to the alkali metals, [[ammonia]] reacts with [[hydrochloric acid]] to form the salt [[ammonium chloride]].]]
The [[ammonium]] ion ({{chem2|NH4+}}) has very similar properties to the heavier alkali metals, acting as an alkali metal intermediate between potassium and rubidium,<ref name=pseudo>{{cite journal |last1=Dietzel |first1=P. D. |last2=Kremer |first2=R. K. |last3=Jansen |first3=M. |date=8 January 2007 |title=Superoxide compounds of the large pseudo-alkali-metal ions tetramethylammonium, -phosphonium, and -arsonium. |journal=Chemistry: An Asian Journal |volume=2 |issue=1 |pages=66–75 |doi=10.1002/asia.200600306 |pmid=17441140}}</ref><ref>{{cite web |url=http://www.meta-synthesis.com/webbook/35_pt/pt_database.php?PT_id=429|title=2002 Inorganic Chemist's Periodic Table|last=Leach|first=Mark R. |access-date=16 October 2012}}</ref> and is often considered a close relative.<ref name="Holleman&Wiberg">{{Holleman&Wiberg}}</ref><ref name="Stevenson" /><ref name="Bernal&Massey" /> For example, most alkali metal [[salt (chemistry)|salts]] are [[solubility|soluble]] in water, a property which ammonium salts share.<ref>{{cite web |title=Solubility Rules!|url=http://www.chem.sc.edu/faculty/morgan/resources/solubility/|work=chem.sc.edu}}</ref> Ammonium is expected to behave stably as a metal ({{chem2|NH4+}} ions in a sea of delocalised electrons) at very high pressures (though less than the typical pressure where transitions from insulating to metallic behaviour occur around, 100&nbsp;[[pascal (unit)|GPa]]), and could possibly occur inside the [[Gas giant#Uranus and Neptune|ice giants]] [[Uranus]] and [[Neptune]], which may have significant impacts on their interior magnetic fields.<ref name="Stevenson">{{cite journal |last1=Stevenson |first1=D. J. |date=20 November 1975 |title=Does metallic ammonium exist? |journal=[[Nature (journal)|Nature]] |volume=258 |issue= 5532 |pages=222–223 |publisher=[[Nature Publishing Group]] |doi=10.1038/258222a0 |bibcode= 1975Natur.258..222S|s2cid=4199721 }}</ref><ref name="Bernal&Massey">{{cite journal |last1=Bernal |first1=M. J. M. |last2=Massey |first2=H. S. W. |date=3 February 1954 |title=Metallic Ammonium |journal=[[Monthly Notices of the Royal Astronomical Society]] |volume=114 |issue=2 |pages=172–179 |publisher=[[Wiley-Blackwell]] for the [[Royal Astronomical Society]] |bibcode=1954MNRAS.114..172B  |doi=10.1093/mnras/114.2.172|doi-access=free }}</ref> It has been estimated that the transition from a mixture of [[ammonia]] and dihydrogen molecules to metallic ammonium may occur at pressures just below 25&nbsp;GPa.<ref name="Stevenson" /> Under standard conditions, ammonium can form a metallic amalgam with mercury.<ref>{{cite journal |last1=Reedy |first1=J. H.|date=1 October 1929|title=Lecture demonstration of ammonium amalgam |journal=Journal of Chemical Education |volume=6 |issue=10 |page=1767|doi=10.1021/ed006p1767|bibcode=1929JChEd...6.1767R}}</ref>

Other "pseudo-alkali metals" include the [[alkylammonium]] cations, in which some of the hydrogen atoms in the ammonium cation are replaced by alkyl or aryl groups. In particular, the [[quaternary ammonium cation]]s ({{chem2|NR4+}}) are very useful since they are permanently charged, and they are often used as an alternative to the expensive Cs<sup>+</sup> to stabilise very large and very easily polarisable anions such as {{chem2|HI2-}}.<ref name="Greenwood&Earnshaw" />{{rp|812–9}} Tetraalkylammonium hydroxides, like alkali metal hydroxides, are very strong bases that react with atmospheric carbon dioxide to form carbonates.<ref name=King />{{rp|256}} Furthermore, the nitrogen atom may be replaced by a phosphorus, arsenic, or antimony atom (the heavier nonmetallic [[pnictogen]]s), creating a [[phosphonium]] ({{chem2|PH4+}}) or [[arsonium]] ({{chem2|AsH4+}}) cation that can itself be substituted similarly; while [[stibonium]] ({{chem2|SbH4+}}) itself is not known, some of its organic derivatives are characterised.<ref name=pseudo />

=== Cobaltocene and derivatives ===
[[Cobaltocene]], Co(C<sub>5</sub>H<sub>5</sub>)<sub>2</sub>, is a [[metallocene]], the [[cobalt]] analogue of [[ferrocene]]. It is a dark purple solid. Cobaltocene has 19 valence electrons, one more than usually found in organotransition metal complexes, such as its very stable relative, ferrocene, in accordance with the [[18-electron rule]]. This additional electron occupies an orbital that is antibonding with respect to the Co–C bonds. Consequently, many chemical reactions of Co(C<sub>5</sub>H<sub>5</sub>)<sub>2</sub> are characterized by its tendency to lose this "extra" electron, yielding a very stable 18-electron cation known as cobaltocenium. Many cobaltocenium salts coprecipitate with caesium salts, and cobaltocenium hydroxide is a strong base that absorbs atmospheric carbon dioxide to form cobaltocenium carbonate.<ref name=King />{{rp|256}} Like the alkali metals, cobaltocene is a strong reducing agent, and [[decamethylcobaltocene]] is stronger still due to the combined [[inductive effect]] of the ten methyl groups.<ref>{{cite journal|doi=10.1021/cr940053x|pmid=11848774|title=Chemical Redox Agents for Organometallic Chemistry|journal=Chemical Reviews|volume=96|issue=2|pages=877–910|year=1996|last1=Connelly|first1=Neil G.|last2=Geiger|first2=William E.}}</ref> Cobalt may be substituted by its heavier congener [[rhodium]] to give [[rhodocene]], an even stronger reducing agent.<ref>{{cite journal |last1= El Murr |first1= N. |last2= Sheats |first2= J. E. |last3= Geiger |first3= W. E. |last4= Holloway |first4= J. D. L. |year= 1979 |title= Electrochemical Reduction Pathways of the Rhodocenium Ion. Dimerization and Reduction of Rhodocene |journal= [[Inorganic Chemistry (journal)|Inorg. Chem.]] |volume= 18 |issue= 6 |pages= 1443–1446 |doi= 10.1021/ic50196a007}}</ref> [[Iridocene]] (involving [[iridium]]) would presumably be still more potent, but is not very well-studied due to its instability.<ref name="Keller_1967">{{cite journal |last1= Keller |first1= H. J. |last2= Wawersik |first2= H. |year= 1967 |title= Spektroskopische Untersuchungen an Komplexverbindungen. VI. EPR-spektren von (C<sub>5</sub>H<sub>5</sub>)<sub>2</sub>Rh und (C<sub>5</sub>H<sub>5</sub>)<sub>2</sub>Ir |journal= [[Journal of Organometallic Chemistry|J. Organomet. Chem.]] |volume= 8 |issue= 1 |pages= 185–188 |language= de |doi= 10.1016/S0022-328X(00)84718-X}}</ref>

=== Thallium ===
[[File:Thallium pieces in ampoule.jpg|thumb|right|Very pure thallium pieces in a glass [[ampoule]], stored under [[argon]] gas]]

[[Thallium]] is the heaviest stable element in group 13 of the periodic table. At the bottom of the periodic table, the [[inert-pair effect]] is quite strong, because of the [[relativistic effects|relativistic]] stabilisation of the 6s orbital and the decreasing bond energy as the atoms increase in size so that the amount of energy released in forming two more bonds is not worth the high ionisation energies of the 6s electrons.<ref name="Greenwood&Earnshaw" />{{rp|226–7}} It displays the +1 [[oxidation state]]<ref name="Greenwood&Earnshaw" />{{rp|28}} that all the known alkali metals display,<ref name="Greenwood&Earnshaw" />{{rp|28}} and thallium compounds with thallium in its +1 [[oxidation state]] closely resemble the corresponding potassium or silver compounds stoichiometrically due to the similar ionic radii of the Tl<sup>+</sup> (164&nbsp;[[picometer|pm]]), K<sup>+</sup> (152&nbsp;pm) and Ag<sup>+</sup> (129&nbsp;pm) ions.<ref name=Shannon>{{cite journal|doi=10.1107/S0567739476001551|title=Revised effective ionic radii and systematic studies of interatomic distances in halides and chalcogenides|last=Shannon|first=R. D.|journal=Acta Crystallogr A|volume=32|year=1976|pages=751–767|bibcode=1976AcCrA..32..751S|issue=5|url=http://journals.iucr.org/a/issues/1976/05/00/a12967/a12967.pdf|access-date=4 September 2019|archive-date=17 March 2020|archive-url=https://web.archive.org/web/20200317120028/http://journals.iucr.org/a/issues/1976/05/00/a12967/a12967.pdf|url-status=dead}}</ref><ref name=Crookes /> It was sometimes considered an alkali metal in [[continental Europe]] (but not in England) in the years immediately following its discovery,<ref name=Crookes>{{cite journal |last1=Crookes |first1=William |author-link=William Crookes |year=1864 |title=On Thallium |journal=Journal of the Chemical Society |volume=17 |pages=112–152 |publisher=Harrison & Sons |url=https://zenodo.org/records/1429753/files/article.pdf|doi=10.1039/js8641700112|url-access= |url-status= |archive-url= |archive-date= }}</ref>{{rp|126}} and was placed just after caesium as the sixth alkali metal in [[Dmitri Mendeleev]]'s 1869 [[periodic table]] and [[Julius Lothar Meyer]]'s 1868 periodic table.<ref name="meta-synthesis2">{{cite web |url=http://www.meta-synthesis.com/webbook/35_pt/pt_database.php?Button=pre-1900+Formulations |title=The Internet Database of Periodic Tables |last=Leach|first=Mark R. |date=1999–2012 |work=meta-synthesis.com |access-date=6 April 2012}}</ref> Mendeleev's 1871 periodic table and Meyer's 1870 periodic table put thallium in its current position in the [[boron group]] and left the space below caesium blank.<ref name="meta-synthesis2" /> However, thallium also displays the oxidation state +3,<ref name="Greenwood&Earnshaw" />{{rp|28}} which no known alkali metal displays<ref name="Greenwood&Earnshaw" />{{rp|28}} (although ununennium, the undiscovered seventh alkali metal, is predicted to possibly display the +3 oxidation state).<ref name="Uue" />{{rp|1729–1730}} The sixth alkali metal is now considered to be francium.<ref name="redbook">{{RedBook2005|pages=51}}.</ref> While Tl<sup>+</sup> is stabilised by the inert-pair effect, this inert pair of 6s electrons is still able to participate chemically, so that these electrons are [[stereochemistry|stereochemically]] active in aqueous solution. Additionally, the thallium halides (except [[thallium(I) fluoride|TlF]]) are quite insoluble in water, and [[thallium(I) iodide|TlI]] has an unusual structure because of the presence of the stereochemically active inert pair in thallium.<ref>{{cite journal |title= Thallium Halides – New Aspects of the Stereochemical Activity of Electron Lone Pairs of Heavier Main-Group Elements |last=Mudring|first=Anja-Verena |journal= [[European Journal of Inorganic Chemistry]] |volume= 2007 |issue= 6 |pages= 882–890 |doi= 10.1002/ejic.200600975 |year= 2007}}</ref>

=== Copper, silver, and gold ===
{{multiple image
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|image1= NatCopper.jpg
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|alt1= A crystal of a coppery-colored metal mineral of standing on a white surface
|caption1= Copper
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|alt2= A crystal of a silvery metal crystal lying on a grey surface
|caption2= Silver
|image3= Gold-crystals.jpg
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|alt3= A crystal of a yellow metal lying on a white surface
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The [[group 11 element|group 11 metals]] (or coinage metals), copper, silver, and gold, are typically categorised as transition metals given they can form ions with incomplete d-shells. Physically, they have the relatively low melting points and high electronegativity values associated with [[post-transition metal]]s. "The filled ''d'' subshell and free ''s'' electron of Cu, Ag, and Au contribute to their high electrical and thermal conductivity. Transition metals to the left of group 11 experience interactions between ''s'' electrons and the partially filled ''d'' subshell that lower electron mobility."<ref>Russell AM & Lee KL (2005) [https://books.google.com/books?id=fIu58uZTE-gC ''Structure-property relations in nonferrous metals'']. Wiley-Interscience, New York. p.&nbsp;302. {{ISBN|0-471-64952-X}}</ref> Chemically, the group 11 metals behave like main-group metals in their +1 valence states, and are hence somewhat related to the alkali metals: this is one reason for their previously being labelled as "group IB", paralleling the alkali metals' "group IA". They are occasionally classified as post-transition metals.<ref>Deming HG (1940) ''Fundamental Chemistry,'' John Wiley & Sons, New York, pp.&nbsp;705–7</ref> Their spectra are analogous to those of the alkali metals.<ref name=Jensen /> Their monopositive ions are [[paramagnetic]] and contribute no colour to their salts, like those of the alkali metals.<ref>Bailar, J. C. (1973) ''Comprehensive inorganic chemistry'', vol. 3, p. 16. {{ISBN|1-57215-291-5}}</ref>

In Mendeleev's 1871 periodic table, copper, silver, and gold are listed twice, once under group VIII (with the [[iron triad]] and [[platinum group metal]]s), and once under group IB. Group IB was nonetheless parenthesised to note that it was tentative. Mendeleev's main criterion for group assignment was the maximum oxidation state of an element: on that basis, the group 11 elements could not be classified in group IB, due to the existence of copper(II) and gold(III) compounds being known at that time.<ref name=Jensen /> However, eliminating group IB would make group I the only main group (group VIII was labelled a transition group) to lack an A–B bifurcation.<ref name=Jensen /> Soon afterward, a majority of chemists chose to classify these elements in group IB and remove them from group VIII for the resulting symmetry: this was the predominant classification until the rise of the modern medium-long 18-column periodic table, which separated the alkali metals and group 11 metals.<ref name=Jensen />

The coinage metals were traditionally regarded as a subdivision of the alkali metal group, due to them sharing the characteristic s<sup>1</sup> electron configuration of the alkali metals (group 1: p<sup>6</sup>s<sup>1</sup>; group 11: d<sup>10</sup>s<sup>1</sup>). However, the similarities are largely confined to the [[stochiometry|stoichiometries]] of the +1 compounds of both groups, and not their chemical properties.<ref name="Greenwood&Earnshaw" />{{rp|1177}} This stems from the filled d subshell providing a much weaker shielding effect on the outermost s electron than the filled p subshell, so that the coinage metals have much higher first ionisation energies and smaller ionic radii than do the corresponding alkali metals.<ref name="Greenwood&Earnshaw" />{{rp|1177}} Furthermore, they have higher melting points, hardnesses, and densities, and lower reactivities and solubilities in liquid [[ammonia]], as well as having more covalent character in their compounds.<ref name="Greenwood&Earnshaw" />{{rp|1177}} Finally, the alkali metals are at the top of the [[electrochemical series]], whereas the coinage metals are almost at the very bottom.<ref name="Greenwood&Earnshaw" />{{rp|1177}} The coinage metals' filled d shell is much more easily disrupted than the alkali metals' filled p shell, so that the second and third ionisation energies are lower, enabling higher oxidation states than +1 and a richer coordination chemistry, thus giving the group 11 metals clear [[transition metal]] character.<ref name="Greenwood&Earnshaw" />{{rp|1177}} Particularly noteworthy is gold forming ionic compounds with rubidium and caesium, in which it forms the auride ion (Au<sup>−</sup>) which also occurs in solvated form in liquid ammonia solution: here gold behaves as a [[pseudohalogen]] because its 5d<sup>10</sup>6s<sup>1</sup> configuration has one electron less than the quasi-closed shell 5d<sup>10</sup>6s<sup>2</sup> configuration of [[mercury (element)|mercury]].<ref name="Greenwood&Earnshaw" />{{rp|1177}}

== Production and isolation ==
{{multiple image
|footer= [[Salt pan (geology)|Salt flats]] are rich in lithium, such as these in Salar del Hombre Muerto, Argentina (left) and [[Salar de Uyuni|Uyuni]], Bolivia (right). The lithium-rich brine is concentrated by pumping it into [[salt evaporation pond|solar evaporation ponds]] (visible in Argentina image).
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|image1= Lithium mine, Salar del Hombre Muerto, Argentina.jpg
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The production of pure alkali metals is somewhat complicated due to their extreme reactivity with commonly used substances, such as water.<ref name=rsc /><ref name=generalchemistry /> From their [[silicate]] ores, all the stable alkali metals may be obtained the same way: [[sulfuric acid]] is first used to dissolve the desired alkali metal ion and aluminium(III) ions from the ore (leaching), whereupon basic precipitation removes aluminium ions from the mixture by precipitating it as the [[aluminium hydroxide|hydroxide]]. The remaining insoluble alkali metal [[carbonate]] is then precipitated selectively; the salt is then dissolved in [[hydrochloric acid]] to produce the chloride. The result is then left to evaporate and the alkali metal can then be isolated.<ref name=generalchemistry /> Lithium and sodium are typically isolated through electrolysis from their liquid chlorides, with [[calcium chloride]] typically added to lower the melting point of the mixture. The heavier alkali metals, however, are more typically isolated in a different way, where a reducing agent (typically sodium for potassium and [[magnesium]] or [[calcium]] for the heaviest alkali metals) is used to reduce the alkali metal chloride. The liquid or gaseous product (the alkali metal) then undergoes [[fractional distillation]] for purification.<ref name=generalchemistry /> Most routes to the pure alkali metals require the use of electrolysis due to their high reactivity; one of the few which does not is the [[pyrolysis]] of the corresponding alkali metal [[azide]], which yields the metal for sodium, potassium, rubidium, and caesium and the nitride for lithium.<ref name=King/>{{rp|77}}

Lithium salts have to be extracted from the water of [[mineral spring]]s, [[brine]] pools, and brine deposits. The metal is produced electrolytically from a mixture of fused [[lithium chloride]] and [[potassium chloride]].<ref name="ober">{{cite web |url=http://minerals.usgs.gov/minerals/pubs/commodity/lithium/450798.pdf |title=Lithium|access-date= 19 August 2007|last=Ober |first=Joyce A. |pages= 77–78 |publisher=[[United States Geological Survey]] |archive-url= https://web.archive.org/web/20070711062102/http://minerals.usgs.gov/minerals/pubs/commodity/lithium/450798.pdf |archive-date= 11 July 2007  |url-status= live}}</ref>

Sodium occurs mostly in seawater and dried [[seabed]],<ref name=rsc /> but is now produced through [[electrolysis]] of [[sodium chloride]] by lowering the melting point of the substance to below 700&nbsp;°C through the use of a [[Downs cell]].<ref name="pauling">{{cite book |last=Pauling |first=Linus |title= General Chemistry |edition=1970 |publisher=Dover Publications}}</ref><ref name="losal">{{cite web |url=http://periodic.lanl.gov/11.shtml|title=Los Alamos National Laboratory – Sodium|access-date=8 June 2007}}</ref> Extremely pure sodium can be produced through the thermal decomposition of [[sodium azide]].<ref>Merck Index, 9th ed., monograph 8325</ref> Potassium occurs in many minerals, such as [[sylvite]] ([[potassium chloride]]).<ref name=rsc /> Previously, potassium was generally made from the electrolysis of [[potassium chloride]] or [[potassium hydroxide]],<ref>{{cite web |publisher=Webelements|title=WebElements Periodic Table of the Elements {{pipe}} Potassium {{pipe}} Essential information |url=http://www.webelements.com/potassium/ |last=Winter|first=Mark|access-date=27 November 2011}}</ref> found extensively in places such as Canada, Russia, Belarus, Germany, Israel, United States, and Jordan, in a method similar to how sodium was produced in the late 1800s and early 1900s.<ref name=kirk>{{cite encyclopedia|doi=10.1002/0471238961.1915040912051311.a01.pub2|chapter=Sodium and Sodium Alloys|title=Kirk-Othmer Encyclopedia of Chemical Technology|year=2001|last1=Lemke|first1=Charles H.|last2=Markant|first2=Vernon H.|isbn=978-0-471-23896-6}}</ref> It can also be produced from [[seawater]].<ref name=rsc /> However, these methods are problematic because the potassium metal tends to dissolve in its molten chloride and vaporises significantly at the operating temperatures, potentially forming the explosive superoxide. As a result, pure potassium metal is now produced by reducing molten potassium chloride with sodium metal at 850&nbsp;°C.<ref name="Greenwood&Earnshaw" />{{rp|74}}
:Na (g) + KCl (l) {{eqm}} NaCl (l) + K (g)
Although sodium is less reactive than potassium, this process works because at such high temperatures potassium is more volatile than sodium and can easily be distilled off, so that the equilibrium shifts towards the right to produce more potassium gas and proceeds almost to completion.<ref name="Greenwood&Earnshaw" />{{rp|74}}

Metals like sodium are obtained by electrolysis of molten salts. Rb & Cs obtained mainly as by products of Li processing. To make pure caesium, ores of caesium and rubidium are crushed and heated to 650&nbsp;°C with sodium metal, generating an alloy that can then be separated via a [[fractional distillation]] technique. Because metallic caesium is too reactive to handle, it is normally offered as [[caesium azide]] (CsN3). [[Caesium hydroxide]] is formed when caesium interacts aggressively with water and ice (CsOH).<ref>{{Cite web|title=Cesium {{!}} Cs (Element) – PubChem|url=https://pubchem.ncbi.nlm.nih.gov/element/Cesium#section=History|access-date=2021-12-18|website=pubchem.ncbi.nlm.nih.gov}}</ref>

[[Rubidium]] is the 16th most abundant element in the earth's crust; however, it is quite rare. Some minerals found in North America, South Africa, Russia, and Canada contain rubidium. Some potassium minerals ([[lepidolite]]s, [[biotite]]s, [[feldspar]], [[carnallite]]) contain it, together with caesium. [[Pollucite]], [[carnallite]], [[leucite]], and [[lepidolite]] are all minerals that contain rubidium. As a by-product of lithium extraction, it is commercially obtained from [[lepidolite]]. Rubidium is also found in potassium rocks and [[brine]]s, which is a commercial supply. The majority of rubidium is now obtained as a byproduct of refining lithium. Rubidium is used in [[vacuum tube]]s as a [[getter]], a material that combines with and removes trace gases from vacuum tubes.<ref>{{Cite web|title=WebElements Periodic Table » Rubidium » geological information|url=https://www.webelements.com/rubidium/geology.html|access-date=2021-12-18|website=www.webelements.com}}</ref><ref>Liu, Jinlian & Yin, Zhoulan & Li, Xinhai & Hu, Qiyang & Liu, Wei. (2019). A novel process for the selective precipitation of valuable metals from lepidolite. Minerals Engineering. 135. 29–36. 10.1016/j.mineng.2018.11.046.</ref>[[File:Pichblende.jpg|thumb|This sample of [[uraninite]] contains about 100,000 atoms (3.3{{e|-20}}&nbsp;g) of francium-223 at any given time.<ref name="nbb" />|alt=A shiny gray 5-centimeter piece of matter with a rough surface.]]
For several years in the 1950s and 1960s, a by-product of the potassium production called Alkarb was a main source for rubidium. Alkarb contained 21% rubidium while the rest was potassium and a small fraction of caesium.<ref>{{cite journal |title= Cesium and Rubidium Hit Market|journal= Chemical & Engineering News |volume= 37|issue= 22|pages= 50–56|year= 1959|doi= 10.1021/cen-v037n022.p050}}</ref> Today the largest producers of caesium, for example the [[Tanco Mine]] in Manitoba, Canada, produce rubidium as by-product from [[pollucite]].<ref name=USGS /> Today, a common method for separating rubidium from potassium and caesium is the [[fractional crystallization (chemistry)|fractional crystallisation]] of a rubidium and caesium [[alum]] ([[Caesium|Cs]], [[Rubidium|Rb]])[[Aluminium|Al]]([[Sulfate|SO<sub>4</sub>]])<sub>2</sub>·12[[Water|H<sub>2</sub>O]], which yields pure rubidium alum after approximately 30 recrystallisations.<ref name=USGS /><ref>{{cite book |url= https://books.google.com/books?id=1ikjAQAAIAAJ&q=ferrocyanide+rubidium|publisher= United States. Bureau of Mines|title= bulletin 585|year= 1995}}</ref> The limited applications and the lack of a mineral rich in rubidium limit the production of rubidium compounds to 2 to 4 [[tonne]]s per year.<ref name=USGS>{{cite web |url= http://pubs.usgs.gov/of/2003/of03-045/of03-045.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://pubs.usgs.gov/of/2003/of03-045/of03-045.pdf |archive-date=2022-10-09 |url-status=live |publisher= United States Geological Survey|access-date= 4 December 2010|title= Mineral Commodity Profile: Rubidium|first1= William C.|last1= Butterman|first2= William E.|last2= Brooks|first3= Robert G. Jr.|last3= Reese|year=2003}}</ref> Caesium, however, is not produced from the above reaction. Instead, the mining of [[pollucite]] ore is the main method of obtaining pure caesium, extracted from the ore mainly by three methods: acid digestion, alkaline decomposition, and direct reduction.<ref name=USGS /><ref name=Burt>{{cite book |last= Burt|first= R. O.|year= 1993|chapter= Caesium and cesium compounds|title= Kirk-Othmer encyclopedia of chemical technology|edition= 4th|place= New York|publisher= John Wiley & Sons, Inc.|volume= 5|pages= 749–764|isbn= 978-0-471-48494-3}}</ref> Both metals are produced as by-products of lithium production: after 1958, when interest in lithium's thermonuclear properties increased sharply, the production of rubidium and caesium also increased correspondingly.<ref name="Greenwood&Earnshaw" />{{rp|71}} Pure rubidium and caesium metals are produced by reducing their chlorides with [[calcium]] metal at 750&nbsp;°C and low pressure.<ref name="Greenwood&Earnshaw" />{{rp|74}}

As a result of its extreme rarity in nature,<ref name="Winter">{{cite web |last= Winter|first= Mark |title= Geological information |work= Francium|publisher= The University of Sheffield |url= http://www.webelements.com/webelements/elements/text/Fr/geol.html |access-date= 26 March 2007}}</ref> most francium is synthesised in the nuclear reaction <sup>197</sup>[[Gold|Au]] + <sup>18</sup>[[Oxygen|O]] → <sup>210</sup>[[Francium|Fr]] + 5 [[neutron|n]], yielding [[francium-209]], [[francium-210]], and [[francium-211]].<ref>{{cite journal |last1=Stancari |first1=G. |last2=Veronesi |first2=S. |last3=Corradi |first3=L. |last4=Atutov |first4=S. N. |last5=Calabrese |first5=R. |last6=Dainelli |first6=A. |last7=Mariotti |first7=E. |last8=Moi |first8=L. |last9=Sanguinetti |first9=S. |first10=L.|last10=Tomassetti|year=2006 |title=Production of Radioactive Beams of Francium |journal=Nuclear Instruments and Methods in Physics Research Section A: Accelerators, Spectrometers, Detectors and Associated Equipment |volume=557 |issue=2 |pages=390–396 |doi=10.1016/j.nima.2005.11.193 |bibcode= 2006NIMPA.557..390S}}</ref> The greatest quantity of francium ever assembled to date is about 300,000 neutral atoms,<ref name="chemnews">{{cite journal |url=http://pubs.acs.org/cen/80th/francium.html|title=Francium|journal=Chemical and Engineering News|volume=81|issue=36|page=159|year=2003|last=Orozco|first=Luis A. |doi=10.1021/cen-v081n036.p159}}</ref> which were synthesised using the nuclear reaction given above.<ref name="chemnews" /> When the only natural isotope francium-223 is specifically required, it is produced as the alpha daughter of actinium-227, itself produced synthetically from the neutron irradiation of natural radium-226, one of the daughters of natural uranium-238.<ref name="andyscouse">{{cite web |last= Price |first= Andy |title= Francium |date= 20 December 2004 |url= http://www.andyscouse.com/pages/francium.htm |access-date= 19 February 2012}}</ref>

== Applications ==
Lithium, sodium, and potassium have many applications, while rubidium and caesium are very useful in academic contexts but do not have many applications yet.<ref name="Greenwood&Earnshaw" />{{rp|68}} Lithium is often used in [[lithium-ion battery|lithium-ion batteries]], and [[lithium oxide]] can help process silica. [[Lithium stearate]] is a thickener and can be used to make lubricating greases; it is produced from lithium hydroxide, which is also used to absorb [[carbon dioxide]] in space capsules and submarines.<ref name="Greenwood&Earnshaw" />{{rp|70}} [[Lithium chloride]] is used as a brazing alloy for aluminium parts.<ref>{{cite news|author=USGS |year=2011|title=Lithium|url= http://minerals.usgs.gov/minerals/pubs/commodity/lithium/mcs-2011-lithi.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://minerals.usgs.gov/minerals/pubs/commodity/lithium/mcs-2011-lithi.pdf |archive-date=2022-10-09 |url-status=live|access-date=4 December 2011}}</ref> Metallic lithium is used in alloys with magnesium and aluminium to give very tough and light alloys.<ref name="Greenwood&Earnshaw" />{{rp|70}}

Sodium compounds have many applications, the most well-known being sodium chloride as [[table salt]]. Sodium salts of [[fatty acid]]s are used as soap.<ref>{{cite web |title=Soaps & Detergents: Chemistry|url=http://www.cleaninginstitute.org/clean_living/soaps__detergents_chemistry.aspx|access-date=20 July 2015}}</ref> Pure sodium metal also has many applications, including use in [[sodium-vapor lamp|sodium-vapour lamps]], which produce very efficient light compared to other types of lighting,<ref name="lamp1">{{cite book |url= https://books.google.com/books?id=0d7u9Nr33zIC&pg=PA112 |page= 112 |title= Applied illumination engineering |isbn= 978-0-88173-212-2 |last1=Lindsey|first1=Jack L |year= 1997|publisher= The Fairmont Press }}</ref><ref name="lamp2">{{cite book |url= https://books.google.com/books?id=AFNwNAFYtCAC&pg=PA241 |page= 241 |title= Revolution in lamps: A chronicle of 50 years of progress |isbn= 978-0-88173-351-8 |last1=Kane|first1=Raymond |last2= Sell |first2= Heinz |year= 2001|publisher= Fairmont Press }}</ref> and can help smooth the surface of other metals.<ref>{{cite book |url= https://books.google.com/books?id=kyVWAAAAYAAJ&q=METALLIC+SODIUM+++DESCALING+SEVERAL |title= Metal treatment and drop forging |last1=Stampers|first1=National Association of Drop Forgers and|year= 1957}}</ref><ref>{{cite book |url= https://books.google.com/books?id=LI4KmKqca78C&pg=PA76 |page= 76 |title= Metal cleaning bibliographical abstracts |last1=Harris|first1=Jay C |year= 1949}}</ref> Being a strong reducing agent, it is often used to reduce many other metals, such as [[titanium]] and [[zirconium]], from their chlorides. Furthermore, it is very useful as a heat-exchange liquid in [[fast breeder nuclear reactor]]s due to its low melting point, viscosity, and [[cross-section (physics)|cross-section]] towards neutron absorption.<ref name="Greenwood&Earnshaw" />{{rp|74}}

Potassium compounds are often used as [[fertiliser]]s<ref name="Greenwood&Earnshaw" />{{rp|73}}<ref>{{cite book |last=Cordel|first=Oskar |title=Die Stassfurter Kalisalze in der Landwirthschalt: Eine Besprechung ...|url=https://books.google.com/books?id=EYpIAAAAYAAJ|year=1868|publisher=L. Schnock |language= de}}</ref> as potassium is an important element for plant nutrition. [[Potassium hydroxide]] is a very strong base, and is used to control the [[pH]] of various substances.<ref>{{cite book|publisher=Greenwood Publishing Group|chapter-url=https://books.google.com/books?id=UnjD4aBm9ZcC&pg=PA4|chapter=Personal Cleansing Products: Bar Soap|title=Chemical composition of everyday products|isbn=978-0-313-32579-3|last1=Toedt|first1=John|last2=Koza|first2=Darrell|last3=Cleef-Toedt|first3=Kathleen Van|year=2005|url-access=registration|url=https://archive.org/details/chemicalcomposit0000toed}}</ref><ref>{{cite book |doi= 10.1002/14356007.a22_031.pub2|title= Ullmann's Encyclopedia of Industrial Chemistry|year= 2006|last=Schultz|first=H.|chapter= Potassium compounds|isbn= 978-3-527-30673-2|volume=A22|page=95|display-authors=etal}}</ref> [[Potassium nitrate]] and [[potassium permanganate]] are often used as powerful oxidising agents.<ref name="Greenwood&Earnshaw" />{{rp|73}} [[Potassium superoxide]] is used in breathing masks, as it reacts with carbon dioxide to give potassium carbonate and oxygen gas. Pure potassium metal is not often used, but its alloys with sodium may substitute for pure sodium in fast breeder nuclear reactors.<ref name="Greenwood&Earnshaw" />{{rp|74}}

Rubidium and caesium are often used in [[atomic clock]]s.<ref name="atomic-clocks">{{cite web |title=Cesium Atoms at Work |publisher=Time Service Department—U.S. Naval Observatory—Department of the Navy |url=http://tycho.usno.navy.mil/cesium.html |access-date=20 December 2009 |url-status=dead |archive-url=https://web.archive.org/web/20150223231150/http://tycho.usno.navy.mil/cesium.html |archive-date=23 February 2015}}</ref> Caesium atomic clocks are extraordinarily accurate; if a clock had been made at the time of the dinosaurs, it would be off by less than four seconds (after 80 million years).<ref name="pubs.usgs">{{cite web |url=http://pubs.usgs.gov/of/2004/1432/2004-1432.pdf |publisher=United States Geological Survey |access-date=27 December 2009 |title=Mineral Commodity Profile: Cesium |first1=William C. |last1=Butterman |first2=William E. |last2=Brooks |first3=Robert G. Jr. |last3=Reese |year=2004 |archive-url=https://web.archive.org/web/20091122210358/http://pubs.usgs.gov/of/2004/1432/2004-1432.pdf |archive-date=22 November 2009 |url-status=dead}}</ref> For that reason, caesium atoms are used as the definition of the second.<ref name="nist-second">{{cite web |title= The NIST reference on Constants, Units, and Uncertainty|date=5 February 2015 |publisher=National Institute of Standards and Technology|url=http://physics.nist.gov/cuu/Units/second.html}}</ref> Rubidium ions are often used in purple [[firework]]s,<ref>{{cite journal |first= E.-C. |last= Koch |title= Special Materials in Pyrotechnics, Part II: Application of Caesium and Rubidium Compounds in Pyrotechnics |journal= Journal Pyrotechnics |year= 2002 |volume= 15 |pages= 9–24 |url= http://www.jpyro.com/wp/?p=179 |access-date= 3 November 2011 |archive-url= https://web.archive.org/web/20110713122322/http://www.jpyro.com/wp/?p=179 |archive-date= 13 July 2011 |url-status= dead}}</ref> and caesium is often used in drilling fluids in the petroleum industry.<ref name="pubs.usgs" /><ref>{{cite book|title= Exploring Chemical Elements and their Compounds|last= Heiserman|first= David L.|publisher= McGraw-Hill|year= 1992|isbn= 978-0-8306-3015-8|pages= [https://archive.org/details/exploringchemica00heis/page/201 201]–203|url= https://archive.org/details/exploringchemica00heis|url-access= registration}}</ref>

Francium has no commercial applications,<ref name="nbb" /><ref name="elemental" /><ref>{{cite web |last= Winter |first= Mark |title= Uses |work= Francium |publisher= The University of Sheffield|url= http://www.webelements.com/webelements/elements/text/Fr/uses.html |access-date= 25 March 2007 |archive-url= https://web.archive.org/web/20070331031655/http://www.webelements.com/webelements/elements/text/Fr/uses.html |archive-date= 31 March 2007  |url-status= live}}</ref> but because of francium's relatively simple [[atomic structure]], among other things, it has been used in [[spectroscopy]] experiments, leading to more information regarding [[energy level]]s and the [[coupling constant]]s between [[subatomic particle]]s.<ref>{{cite journal |last1= Gomez |first1= E. |last2= Orozco|first2=L. A.|last3=Sprouse|first3=G. D. |title= Spectroscopy with trapped francium: advances and perspectives for weak interaction studies |journal= Rep. Prog. Phys. |volume= 69 |issue= 1 |pages= 79–118 |date= 7 November 2005|doi= 10.1088/0034-4885/69/1/R02|bibcode= 2006RPPh...69...79G|s2cid= 15917603 |url= https://www.researchgate.net/publication/228616649}}</ref> Studies on the light emitted by laser-trapped francium-210 ions have provided accurate data on transitions between atomic energy levels, similar to those predicted by [[quantum mechanics|quantum theory]].<ref>{{cite journal |last= Peterson|first= I.|title= Creating, cooling, trapping francium atoms|page= 294|journal= Science News|date= 11 May 1996|url= http://www.sciencenews.org/pages/pdfs/data/1996/149-19/14919-06.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://www.sciencenews.org/pages/pdfs/data/1996/149-19/14919-06.pdf |archive-date=2022-10-09 |url-status=live|access-date= 11 September 2009|volume=149|issue=19|doi= 10.2307/3979560|jstor= 3979560}}</ref>

== Biological role and precautions ==

=== Metals ===
Pure alkali metals are dangerously reactive with air and water and must be kept away from heat, fire, oxidising agents, acids, most organic compounds, [[halocarbon]]s, plastics, and moisture. They also react with carbon dioxide and carbon tetrachloride, so that normal fire extinguishers are counterproductive when used on alkali metal fires.<ref name=osu /> Some Class D dry powder [[fire extinguisher|extinguishers]] designed for metal fires are effective, depriving the fire of oxygen and cooling the alkali metal.<ref>{{cite book |url= https://books.google.com/books?id=2fHsoobsCNwC&pg=PA459|page= 459 |title= Fire and Life Safety Inspection Manual |isbn= 978-0-87765-472-8|publisher=Jones & Bartlett Learning |last= Solomon |first=Robert E. |date= 2002}}</ref>

Experiments are usually conducted using only small quantities of a few grams in a [[fume hood]]. Small quantities of lithium may be disposed of by reaction with cool water, but the heavier alkali metals should be dissolved in the less reactive [[isopropanol]].<ref name=osu /><ref>{{cite book |last=Angelici|first=R. J.|title= Synthesis and Technique in Inorganic Chemistry |publisher= University Science Books |place= Mill Valley, CA |date= 1999 |isbn= 978-0-935702-48-4}}</ref> The alkali metals must be stored under [[mineral oil]] or an inert atmosphere. The inert atmosphere used may be [[argon]] or nitrogen gas, except for lithium, which reacts with nitrogen.<ref name=osu>{{cite web |url=http://chemsafety.chem.oregonstate.edu/content/sop-alkali-metals |title=Standard Operating Procedure: Storage and Handling of Alkali Metals |last1=Lerner |first1=Michael M. |date=2013 |publisher=[[Oregon State University]] |access-date=26 August 2016}}</ref> Rubidium and caesium must be kept away from air, even under oil, because even a small amount of air diffused into the oil may trigger formation of the dangerously explosive peroxide; for the same reason, potassium should not be stored under oil in an oxygen-containing atmosphere for longer than 6 months.<ref>{{cite book |chapter-url= https://books.google.com/books?id=vKBqqiCTB7MC&pg=PA215 |page= 215 |chapter= Rubidium |title= Chemical risk analysis: a practical handbook |isbn= 978-1-903996-65-2 |last1=Martel|first1=Bernard |last2=Cassidy|first2=Keith |date= 2004-07-01|publisher= Butterworth-Heinemann }}</ref><ref>{{cite web|url=http://www.ncsu.edu/ehs/www99/right/handsMan/lab/Peroxide.pdf|title=Danger: peroxidazable chemicals|last=Wray|first=Thomas K.|publisher=Environmental Health & Public Safety ([[North Carolina State University]])|url-status=dead|archive-url=https://web.archive.org/web/20110608024458/http://www.ncsu.edu/ehs/www99/right/handsMan/lab/Peroxide.pdf|archive-date=8 June 2011}}</ref>

=== Ions ===
[[File:Lithium carbonate.jpg|thumb|right|[[Lithium carbonate]]]]
The bioinorganic chemistry of the alkali metal ions has been extensively reviewed.<ref>{{cite book |publisher= Springer|date= 2016|series= Metal Ions in Life Sciences|volume=16|title= The Alkali Metal Ions: Their Role in Life|editor1-last=Astrid|editor1-first= Sigel|editor2-last=Helmut|editor2-first=Sigel|editor3-last=Roland K.O. |editor3-first= Sigel|doi=10.1007/978-3-319-21756-7|isbn= 978-3-319-21755-0|s2cid= 5983458}}</ref>
Solid state crystal structures have been determined for many complexes of alkali metal ions in small peptides, nucleic acid constituents, carbohydrates and ionophore complexes.<ref>{{cite book |last1= Katsuyuki |first1= Aoki|last2= Kazutaka |first2= Murayama |last3=Hu|first3= Ning-Hai|publisher= Springer|date= 2016 |series= Metal Ions in Life Sciences|volume=16|title= The Alkali Metal Ions: Their Role in Life|editor1-last=Astrid|editor1-first= Sigel|editor2-last=Helmut|editor2-first=Sigel|editor3-last=Roland K.O.|editor3-first= Sigel|chapter= Chapter 3. Solid State Structures of Alkali Metal Ion Complexes Formed by Low-Molecular-Weight Ligands of Biological Relevance|pages= 27–101|doi=10.1007/978-3-319-21756-7_3|pmid= 26860299|isbn= 978-3-319-21755-0}}</ref>

Lithium naturally only occurs in traces in biological systems and has no known biological role, but does have effects on the body when ingested.<ref name="webelements-lithium" /> [[Lithium carbonate]] is used as a [[mood stabiliser]] in [[psychiatry]] to treat [[bipolar disorder]] ([[manic-depression]]) in daily doses of about 0.5 to 2&nbsp;grams, although there are side-effects.<ref name="webelements-lithium" /> Excessive ingestion of lithium causes drowsiness, slurred speech and vomiting, among other symptoms,<ref name="webelements-lithium" /> and [[poison]]s the [[central nervous system]],<ref name="webelements-lithium" /> which is dangerous as the required dosage of lithium to treat bipolar disorder is only slightly lower than the toxic dosage.<ref name="webelements-lithium">{{cite web |publisher=Webelements|title=WebElements Periodic Table of the Elements {{pipe}} Lithium {{pipe}} biological information |url=http://www.webelements.com/lithium/biology.html |last=Winter|first=Mark |access-date=15 February 2011}}</ref><ref name="theodoregray-lithium">{{cite web |url=http://www.theodoregray.com/periodictable/Elements/003/index.s7.html |title=Facts, pictures, stories about the element Lithium in the Periodic Table |last=Gray|first=Theodore |publisher=theodoregray.com |access-date=9 January 2012}}</ref> Its biochemistry, the way it is handled by the human body and studies using rats and goats suggest that it is an [[essential element|essential]] [[trace element]], although the natural biological function of lithium in humans has yet to be identified.<ref>{{cite journal |last1=Howland |first1=Robert H. |date=September 2007 |title=Lithium: Underappreciated and Underused? |journal=Psychiatric Annals |volume=37 |issue=9 |pages=13–17 |doi=10.3928/00485713-20070901-06 |pmid=17848039 }}</ref><ref>{{cite journal |last1=Zarse |first1=Kim |last2=Terao |first2=Takeshi |last3=Tian |first3=Jing |last4=Iwata |first4=Noboru |last5=Ishii |first5=Nobuyoshi |last6=Ristow |first6=Michael |date=August 2011 |title=Low-dose lithium uptake promotes longevity in humans and metazoans |journal=European Journal of Nutrition |volume=50 |issue=5 |pages=387–9 |doi=10.1007/s00394-011-0171-x |pmc=3151375 |pmid=21301855}}</ref>

Sodium and potassium occur in all known biological systems, generally functioning as [[electrolytes]] inside and outside [[cell (biology)|cells]].<ref name="webelements-potassium" /><ref name="webelements-sodium" /> Sodium is an essential nutrient that regulates blood volume, blood pressure, osmotic equilibrium and [[pH]]; the minimum physiological requirement for sodium is 500&nbsp;milligrams per day.<ref name=r31>{{cite web |url=http://nuinfo-proto4.northwestern.edu/nutrition/factsheets/sodium.pdf |title=Sodium |publisher=Northwestern University |access-date=21 November 2011 |url-status=dead |archive-url=https://web.archive.org/web/20110823114818/http://nuinfo-proto4.northwestern.edu/nutrition/factsheets/sodium.pdf |archive-date=23 August 2011}}</ref> [[Sodium chloride]] (also known as common salt) is the principal source of sodium in the diet, and is used as seasoning and preservative, such as for [[pickling]] and [[jerky (food)|jerky]]; most of it comes from processed foods.<ref>{{cite web |url=http://health.ltgovernors.com/sodium-and-potassium-health-facts.html|title=Sodium and Potassium Quick Health Facts|access-date=7 November 2011}}</ref> The [[Dietary Reference Intake]] for sodium is 1.5&nbsp;grams per day,<ref>{{cite web|title=Dietary Reference Intakes: Water, Potassium, Sodium, Chloride, and Sulfate|url=http://www.iom.edu/Reports/2004/Dietary-Reference-Intakes-Water-Potassium-Sodium-Chloride-and-Sulfate.aspx|publisher=Food and Nutrition Board, [[Institute of Medicine]], [[United States National Academies]]|date=11 February 2004|access-date=23 November 2011|url-status=dead|archive-url=https://web.archive.org/web/20111006174858/http://www.iom.edu/Reports/2004/Dietary-Reference-Intakes-Water-Potassium-Sodium-Chloride-and-Sulfate.aspx|archive-date=6 October 2011}}</ref> but most people in the United States consume more than 2.3&nbsp;grams per day,<ref>{{cite book|author1=U.S. Department of Agriculture |author2=U.S. Department of Health and Human Services |author-link1=United States Department of Agriculture |author-link2=United States Department of Health and Human Services |title=Dietary Guidelines for Americans, 2010 |page=22 |edition=7th |date=December 2010 |publisher=U.S. Government Printing Office |url=http://www.cnpp.usda.gov/Publications/DietaryGuidelines/2010/PolicyDoc/PolicyDoc.pdf |access-date=23 November 2011 |isbn=978-0-16-087941-8 |oclc=738512922 |url-status=dead |archive-url=https://web.archive.org/web/20111027053444/http://www.cnpp.usda.gov/Publications/DietaryGuidelines/2010/PolicyDoc/PolicyDoc.pdf |archive-date=27 October 2011 }}</ref> the minimum amount that promotes hypertension;<ref>{{cite journal|pmid=15369026|year=2004|last1=Geleijnse|first1=J. M.|last2=Kok|first2=F. J.|last3=Grobbee|first3=D. E.|title=Impact of dietary and lifestyle factors on the prevalence of hypertension in Western populations|volume=14|issue=3|pages=235–239|journal=European Journal of Public Health|doi=10.1093/eurpub/14.3.235|doi-access=free|url=http://dspace.library.uu.nl:8080/handle/1874/12616|access-date=30 August 2017|archive-url=https://web.archive.org/web/20180801155836/http://dspace.library.uu.nl:8080/handle/1874/12616|archive-date=1 August 2018|url-status=dead}}</ref> this in turn causes 7.6&nbsp;million premature deaths worldwide.<ref>{{cite journal |pmid=18456100 |url=http://www.worldactiononsalt.com/evidence/docs/thelancet_hypertension_05.08.pdf |year=2008 |last1=Lawes |first1=C. M. |last2=Vander Hoorn |first2=S. |last3=Rodgers |first3=A. |author4=International Society of Hypertension |title=Global burden of blood-pressure-related disease, 2001 |volume=371 |issue=9623 |pages=1513–1518 |doi=10.1016/S0140-6736(08)60655-8 |journal=Lancet |s2cid=19315480 |url-status=dead |archive-url=https://web.archive.org/web/20120128072727/http://www.worldactiononsalt.com/evidence/docs/thelancet_hypertension_05.08.pdf |archive-date=28 January 2012}}</ref>

Potassium is the major [[cation]] (positive ion) inside [[cell (biology)|animal cells]],<ref name="webelements-potassium">{{cite web |url=http://www.webelements.com/potassium/biology.html|title=WebElements Periodic Table of the Elements {{pipe}} Potassium {{pipe}} biological information |publisher=WebElements |last=Winter|first=Mark |access-date=13 January 2012}}</ref> while sodium is the major cation outside animal cells.<ref name="webelements-potassium" /><ref name="webelements-sodium">{{cite web |url=http://www.webelements.com/sodium/biology.html|title=WebElements Periodic Table of the Elements {{pipe}} Sodium {{pipe}} biological information |publisher=WebElements |last=Winter|first=Mark |access-date=13 January 2012}}</ref> The [[concentration]] differences of these charged particles causes a difference in [[electric potential]] between the inside and outside of cells, known as the [[membrane potential]]. The balance between potassium and sodium is maintained by [[ion transporter]] proteins in the [[cell membrane]].<ref name="pmid16253415">{{cite journal |last1=Hellgren|first1=Mikko|last2=Sandberg|first2=Lars|last3=Edholm|first3=Olle|title=A comparison between two prokaryotic potassium channels (K<sub>ir</sub>Bac1.1 and KcsA) in a molecular dynamics (MD) simulation study |journal=Biophys. Chem. |volume=120 |issue=1 |pages=1–9 |year=2006 |pmid=16253415 |doi=10.1016/j.bpc.2005.10.002}}</ref> The cell membrane potential created by potassium and sodium ions allows the cell to generate an [[action potential]]—a "spike" of electrical discharge. The ability of cells to produce electrical discharge is critical for body functions such as [[neurotransmission]], muscle contraction, and heart function.<ref name="pmid16253415" /> Disruption of this balance may thus be fatal: for example, ingestion of large amounts of potassium compounds can lead to [[hyperkalemia]] strongly influencing the cardiovascular system.<ref name="hyper">{{cite book |publisher=Lippincott Williams & Wilkins|chapter-url= https://books.google.com/books?id=BfdighlyGiwC&pg=PA903 |chapter= Potassium Chloride and Potassium Permanganate|pages= 903–5|title= Medical toxicology|isbn= 978-0-7817-2845-4|last= Schonwald|first= Seth|date= 2004}}</ref><ref>{{cite book |url= https://books.google.com/books?id=l8RkPU1-M5wC&pg=PA223|publisher=Elsevier Health Sciences|page= 223|title= Emergency medicine secrets|isbn= 978-1-56053-503-4|last1= Markovchick |first1=Vincent J.|last2= Pons |first2=Peter T.|name-list-style= amp|date= 2003}}</ref> Potassium chloride is used in the United States for [[lethal injection]] executions.<ref name="hyper" />

[[File:GoiâniaRadiationsource.gif|thumb|400px|right|A wheel type radiotherapy device which has a long [[collimator]] to focus the radiation into a narrow beam. The caesium-137 chloride radioactive source is the blue square, and gamma rays are represented by the beam emerging from the aperture. This was the radiation source involved in the Goiânia accident, containing about 93&nbsp;grams of caesium-137 chloride.]]
Due to their similar atomic radii, rubidium and caesium in the body mimic potassium and are taken up similarly. Rubidium has no known biological role, but may help stimulate [[metabolism]],<ref name="webelements-rubidium">{{cite web |publisher=Webelements|title= WebElements Periodic Table of the Elements {{pipe}} Rubidium {{pipe}} biological information |url=http://www.webelements.com/rubidium/biology.html |last=Winter|first=Mark |access-date=15 February 2011}}</ref><ref name=yale>{{cite journal |last1= Relman |first1= A. S. |title= The Physiological Behavior of Rubidium and Cesium in Relation to That of Potassium |journal= The Yale Journal of Biology and Medicine |volume= 29 |issue= 3 |pages= 248–62 |year= 1956 |pmid= 13409924|pmc= 2603856}}</ref><ref name="jcp.sagepub.com">{{cite journal |last1= Meltzer |first1= H. L. |title= A pharmacokinetic analysis of long-term administration of rubidium chloride |url= http://jcp.sagepub.com/content/31/2/179 |journal= Journal of Clinical Pharmacology |volume= 31 |issue= 2 |pages= 179–84 |year= 1991 |pmid= 2010564 |doi= 10.1002/j.1552-4604.1991.tb03704.x |s2cid= 2574742 |url-status= dead |archive-url= https://archive.today/20120709223213/http://jcp.sagepub.com/content/31/2/179 |archive-date= 9 July 2012}}</ref> and, similarly to caesium,<ref name="webelements-rubidium" /><ref name="webelements-caesium" /> replace potassium in the body causing [[hypokalemia|potassium deficiency]].<ref name="webelements-rubidium" /><ref name="jcp.sagepub.com" /> Partial substitution is quite possible and rather non-toxic: a 70&nbsp;kg person contains on average 0.36&nbsp;g of rubidium, and an increase in this value by 50 to 100 times did not show negative effects in test persons.<ref>{{cite journal |last1= Fieve |first1= Ronald R. |last2= Meltzer |first2= Herbert L. |last3= Taylor |first3= Reginald M. |title= Rubidium chloride ingestion by volunteer subjects: Initial experience |journal= Psychopharmacologia |volume= 20 |issue= 4 |pages= 307–14 |date= 1971 |pmid= 5561654 |doi= 10.1007/BF00403562|s2cid= 33738527 }}</ref> Rats can survive up to 50% substitution of potassium by rubidium.<ref name="jcp.sagepub.com" /><ref>{{cite journal |author= <nowiki>Follis, Richard H., Jr.</nowiki> |title= Histological Effects in rats resulting from adding Rubidium or Cesium to a diet deficient in potassium |journal= American Journal of Physiology. Legacy Content|volume= 138 |issue= 2 |page= 246 |date= 1943 |doi= 10.1152/ajplegacy.1943.138.2.246 |doi-access=  }}</ref> Rubidium (and to a much lesser extent caesium) can function as temporary cures for hypokalemia; while rubidium can adequately physiologically substitute potassium in some systems, caesium is never able to do so.<ref name=yale /> There is only very limited evidence in the form of deficiency symptoms for rubidium being possibly essential in goats; even if this is true, the trace amounts usually present in food are more than enough.<ref name="Gottschlich2001">{{cite book |last=Gottschlich |first= Michele M.|title=The Science and Practice of Nutrition Support: A Case-based Core Curriculum|url=https://books.google.com/books?id=a5LjQ4POQswC&pg=PA98|year=2001|publisher=Kendall Hunt|isbn=978-0-7872-7680-5|page=98}}</ref><ref name="InselTurner2004">{{cite book |last1=Insel|first1=Paul M.|last2=Turner|first2=R. Elaine|last3=Ross|first3=Don|title= Nutrition |url=https://books.google.com/books?id=46o0PzPI07YC&pg=PA499|year=2004|publisher=Jones & Bartlett Learning |isbn=978-0-7637-0765-1|page=499}}</ref>

Caesium compounds are rarely encountered by most people, but most caesium compounds are mildly toxic. Like rubidium, caesium tends to substitute potassium in the body, but is significantly larger and is therefore a poorer substitute.<ref name="webelements-caesium">{{cite web |url=http://www.webelements.com/caesium/biology.html|title=WebElements Periodic Table of the Elements {{pipe}} Caesium {{pipe}} biological information |publisher=WebElements |last=Winter|first=Mark |access-date=13 January 2012}}</ref> Excess caesium can lead to [[hypokalemia]], [[arrythmia]], and acute cardiac arrest,<ref>{{cite journal |last1=Melnikov|first1=P.|last2=Zanoni|first2=L. Z. |title=Clinical effects of cesium intake|journal=Biological Trace Element Research|date=June 2010|volume=135|issue=1–3|pages=1–9 |pmid= 19655100|doi=10.1007/s12011-009-8486-7|s2cid=19186683|url=https://www.researchgate.net/publication/26717303}}</ref> but such amounts would not ordinarily be encountered in natural sources.<ref name= pinsky>{{cite journal |doi= 10.1080/10934528109375003|title= Cesium in mammals: Acute toxicity, organ changes and tissue accumulation|date= 1981|last1= Pinsky|first1= Carl|first2= Ranjan|first3= J. R.|first4= Jasper|first5= Claude|first6= James|journal= Journal of Environmental Science and Health, Part A|volume= 16|pages= 549–567 |last2= Bose|last3= Taylor |last4= McKee|last5= Lapointe|last6= Birchall|issue= 5}}</ref> As such, caesium is not a major chemical environmental pollutant.<ref name = pinsky/> The [[median lethal dose]] (LD<sub>50</sub>) value for [[caesium chloride]] in mice is 2.3&nbsp;g per kilogram, which is comparable to the LD<sub>50</sub> values of [[potassium chloride]] and [[sodium chloride]].<ref>{{cite journal |doi= 10.1016/0041-008X(75)90216-1 |title= Acute toxicity of cesium and rubidium compounds|year= 1975|last1= Johnson|first1= Garland T.|journal= [[Toxicology and Applied Pharmacology]]|volume= 32|pages= 239–245|pmid= 1154391|first2= Trent R.|first3= D. Wagner|issue= 2|last2= Lewis |last3= Wagner}}</ref> Caesium chloride has been promoted as an alternative cancer therapy,<ref>{{cite journal |last= Sartori|first=H. E. |year= 1984 |title= Cesium therapy in cancer patients |journal= Pharmacol Biochem Behav |volume= 21 |issue= Suppl 1 |pages= 11–13 |pmid= 6522427 |doi= 10.1016/0091-3057(84)90154-0|s2cid=11947121 }}</ref> but has been linked to the deaths of over 50 patients, on whom it was used as part of a scientifically unvalidated cancer treatment.<ref>Wood, Leonie. {{cite web |url=http://www.smh.com.au/lifestyle/lifematters/cured-cancer-patients-died-court-told-20101119-180z9.html |title='Cured' cancer patients died, court told |work=The Sydney Morning Herald |date=20 November 2010}}</ref>

[[Radioisotope]]s of caesium require special precautions: the improper handling of caesium-137 [[gamma ray]] sources can lead to release of this radioisotope and radiation injuries. Perhaps the best-known case is the Goiânia accident of 1987, in which an improperly-disposed-of radiation therapy system from an abandoned clinic in the city of [[Goiânia]], Brazil, was scavenged from a junkyard, and the glowing [[caesium chloride|caesium salt]] sold to curious, uneducated buyers. This led to four deaths and serious injuries from radiation exposure. Together with [[caesium-134]], [[iodine-131]], and [[strontium-90]], caesium-137 was among the isotopes distributed by the [[Chernobyl disaster]] which constitute the greatest risk to health.<ref name="IAEA" /> Radioisotopes of francium would presumably be dangerous as well due to their high decay energy and short half-life, but none have been produced in large enough amounts to pose any serious risk.<ref name=andyscouse />

== Notes ==
{{reflist|group="note"|30em}}

== References ==
{{reflist|30em}}

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