Editing Alkali metal (section)

=== Oxides and chalcogenides ===
{{See also|Alkali metal oxide}}
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| 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 [[octahedra]] connected to each other by one face
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| 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.
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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 [[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 |archive-date= 1 July 2019 |access-date= 24 June 2013 |archive-url= https://web.archive.org/web/20190701105103/https://www.osti.gov/biblio/4367751 |url-status= live }}</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>