Array ( [0] => {{Distinguish|cerium}} [1] => {{Featured article}} [2] => {{Use Oxford spelling|date=January 2018}} [3] => {{use dmy dates|date=November 2022}} [4] => {{Infobox caesium|engvar=en-OED}} [5] => '''Caesium''' ([[IUPAC]] spelling;{{cite web |url=https://iupac.org/what-we-do/periodic-table-of-elements/ |title=IUPAC Periodic Table of Elements |publisher=International Union of Pure and Applied Chemistry}} '''cesium''' in [[American English]]){{refn|''Caesium'' is the spelling recommended by the [[International Union of Pure and Applied Chemistry]] (IUPAC).{{RedBook2005|pages=248–49}}. The [[American Chemical Society]] (ACS) has used the spelling ''cesium'' since 1921,{{cite book |editor1-first=Anne M. |editor1-last=Coghill |editor2-first=Lorrin R. |editor2-last=Garson |date=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}}{{cite journal |journal=Pure Appl. Chem. |volume=70 |issue=1 |last1=Coplen |pages=237–257 |date=1998 |first1=T. B. |url=http://old.iupac.org/reports/1998/7001coplen/history.pdf |archive-url=https://web.archive.org/web/20110521151650/http://old.iupac.org/reports/1998/7001coplen/history.pdf |archive-date=2011-05-21 |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}} following ''Webster's New International Dictionary''. The element was named after the Latin word ''[[wikt:caesius#Adjective|caesius]]'', meaning "bluish grey".[http://www.oed.com/view/Entry/26023 OED entry for "caesium"]. Second edition, 1989; online version June 2012. Retrieved 7 September 2012. Earlier version first published in ''New English Dictionary'', 1888. In medieval and early modern writings ''caesius'' was spelled with the [[typographic ligature|ligature]] ''[[Æ#Latin and Greek|æ]]'' as ''cæsius''; hence, an alternative but now old-fashioned orthography is ''cæsium''. More spelling explanation at [[American and British English spelling differences#ae and oe|ae/oe vs e]].|group=note}} is a [[chemical element]]; it has [[Symbol (chemistry)|symbol]] '''Cs''' and [[atomic number]] 55. It is a soft, silvery-golden [[alkali metal]] with a melting point of {{convert|28.5|C|F K}}, which makes it one of only five elemental [[metal]]s that are [[liquid]] at or near [[room temperature]].{{refn|Along with [[rubidium]] ({{convert|39|C|F|disp=sqbr}}), [[francium]] (estimated at {{convert|27|C|F|disp=sqbr}}), [[mercury (element)|mercury]] ({{convert|−39|C|F|disp=sqbr}}), and [[gallium]] ({{convert|30|C|F|disp=sqbr}}); bromine is also liquid at room temperature (melting at {{convert|−7.2|C|F|disp=sqbr}}), but it is a [[halogen]] and not a metal. Preliminary work with [[copernicium]] and [[flerovium]] suggests that they are gaseous metals at room temperature.|group=note}} Caesium has physical and chemical properties similar to those of [[rubidium]] and [[potassium]]. It is [[pyrophoricity|pyrophoric]] and reacts with [[water]] even at {{convert|−116|C}}. It is the least [[electronegativity|electronegative]] element, with a value of 0.79 on the [[Pauling scale]]. It has only one stable [[isotope]], [[caesium-133]]. Caesium is mined mostly from [[pollucite]]. [[Caesium-137]], a [[fission product]], is extracted from waste produced by [[nuclear reactor technology|nuclear reactors]]. It has the largest [[atomic radius]] of all elements whose radii have been measured or calculated, at about 260 [[picometer]]s. [6] => [7] => The German chemist [[Robert Bunsen]] and physicist [[Gustav Kirchhoff]] discovered caesium in 1860 by the newly developed method of [[atomic emission spectroscopy#Flame emission spectroscopy|flame spectroscopy]]. The first small-scale applications for caesium were as a "[[getter]]" in [[vacuum tube]]s and in [[solar cell|photoelectric cells]]. Caesium is widely used in highly accurate [[atomic clock]]s. In 1967, the [[International System of Units]] began using a specific [[Hyperfine structure|hyperfine]] transition of neutral caesium-133 atoms to define the [[SI base unit|basic unit]] of time, the [[second]]. [8] => [9] => Since the 1990s, the largest [[#Applications|application of the element]] has been as [[caesium formate]] for [[drilling fluid]]s, but it has a range of applications in the production of electricity, in electronics, and in chemistry. The radioactive isotope caesium-137 has a [[half-life]] of about 30 years and is used in medical applications, industrial gauges, and hydrology. Nonradioactive caesium compounds are only mildly [[toxicity|toxic]], but the pure metal's tendency to react explosively with water means that caesium is considered a hazardous material, and the [[radioisotopes]] present a significant health and environmental hazard. [10] => [11] => ==Characteristics== [12] => [13] => ===Physical properties=== [14] => [[File:CsCrystals.JPG|left|thumb|High-purity caesium-133 stored in [[argon]].|alt=Y-shaped yellowish crystal in glass ampoule, looking like the branch of a pine tree]] [15] => Of all elements that are solid at room temperature, caesium is the softest: it has a hardness of 0.2 Mohs. It is a very [[ductility|ductile]], pale metal, which darkens in the presence of trace amounts of [[oxygen]].{{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 |date=2004 |url-status=dead |archive-url=https://web.archive.org/web/20070207015229/http://pubs.usgs.gov/of/2004/1432/2004-1432.pdf |archive-date=7 February 2007}}{{cite book |title=Exploring Chemical Elements and their Compounds |author=Heiserman, David L. |publisher=McGraw-Hill |date=1992 |isbn=978-0-8306-3015-8 |pages=[https://archive.org/details/exploringchemica00heis/page/201 201]–203 |url-access=registration |url=https://archive.org/details/exploringchemica00heis}}{{cite book |title=The Chemistry of the Liquid Alkali Metals |last=Addison |first=C. C. |date=1984 |publisher=Wiley |isbn=978-0-471-90508-0 |access-date=28 September 2012 |url=http://www.cs.rochester.edu/users/faculty/nelson/cesium/cesium_color.html}} When in the presence of [[mineral oil]] (where it is best kept during transport), it loses its metallic [[lustre (mineralogy)|lustre]] and takes on a duller, grey appearance. It has a [[melting point]] of {{convert|28.5|C}}, making it one of the few elemental metals that are liquid near [[room temperature]]. [[Mercury (element)|Mercury]] is the only stable elemental metal with a known melting point lower than caesium.{{refn|The radioactive element [[francium]] may also have a lower melting point, but its radioactivity prevents enough of it from being isolated for direct testing.{{cite web |url=http://periodic.lanl.gov/87.shtml |title=Francium |publisher=Periodic.lanl.gov |access-date=23 February 2010}} [[Copernicium]] and [[flerovium]] may also have lower melting points.|group=note}}{{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 |author=Kaner, Richard |date=2003}} In addition, the metal has a rather low [[boiling point]], {{convert|641|C}}, the [[list of elements by boiling point|lowest]] of all metals other than mercury. Its compounds burn with a blue{{cite book |url=https://books.google.com/books?id=QdU-lRMjOsgC&pg=PA13 |page=13 |first=Charles T. |last=Lynch |publisher=CRC Press |date=1974 |title=CRC Handbook of Materials Science |isbn=978-0-8493-2321-8}} or violet{{cite web |url=http://www.chemguide.co.uk/inorganic/group1/flametests.html |title=Flame Tests |author=Clark, Jim |date=2005 |work=chemguide |access-date=29 January 2012}} colour. [16] => [17] => [[File:Rb&Cs crystals.jpg|left|thumb|Caesium crystals (golden) compared to [[rubidium]] crystals (silvery)]] [18] => Caesium forms [[alloy]]s with the other alkali metals, [[gold]], and mercury ([[amalgam (chemistry)|amalgams]]). At temperatures below {{convert|650|°C}}, it does not alloy with [[cobalt]], [[iron]], [[molybdenum]], [[nickel]], [[platinum]], [[tantalum]], or [[tungsten]]. It forms well-defined [[intermetallics|intermetallic compounds]] with [[antimony]], [[gallium]], [[indium]], and [[thorium]], which are [[photosensitive]]. It mixes with all the other alkali metals (except lithium); the alloy with a [[molar concentration|molar]] distribution of 41% caesium, 47% [[potassium]], and 12% [[sodium]] has the lowest melting point of any known metal alloy, at {{convert|−78|C}}.{{cite conference |url=http://symp15.nist.gov/pdf/p564.pdf |title=Density of melts of alkali metals and their Na-K-Cs and Na-K-Rb ternary systems |author=Taova, T. M. |conference=Fifteenth symposium on thermophysical properties, Boulder, Colorado, United States |date=22 June 2003 |access-date=26 September 2010 |display-authors=etal |url-status=dead |archive-url=https://web.archive.org/web/20061009133313/http://symp15.nist.gov/pdf/p564.pdf |archive-date=9 October 2006}} A few amalgams have been studied: {{chem|CsHg|2}} is black with a purple metallic [[lustre (mineralogy)|lustre]], while CsHg is golden-coloured, also with a metallic lustre.{{cite journal |doi=10.1016/S0079-6786(97)81004-7 |journal=Progress in Solid State Chemistry |volume=25 |date=1997 |pages=73–123 |title=Alkali metal amalgams, a group of unusual alloys |first=H. J. |last=Deiseroth |issue=1–2}} [19] => [20] => The golden colour of caesium comes from the decreasing frequency of light required to excite electrons of the alkali metals as the group is descended. For lithium through rubidium this frequency is in the ultraviolet, but for caesium it enters the blue–violet end of the spectrum; in other words, the [[plasma oscillation|plasmonic frequency]] of the alkali metals becomes lower from lithium to caesium. Thus caesium transmits and partially absorbs violet light preferentially while other colours (having lower frequency) are reflected; hence it appears yellowish.{{cite book |last=Addison |first=C. C. |date=1984 |title=The chemistry of the liquid alkali metals |publisher=Wiley |page=7 |isbn=9780471905080}} [21] => [22] => === Allotropes === [23] => Caesium exists in the form of different [[Allotropy|allotropes]], One of which is a dimer called dicaesium{{Cite journal |last=C. A. |first=Onate |date=18 March 2021 |title=Ro-vibrational energies of cesium dimer and lithium dimer with molecular attractive potential |journal=Scientific Reports |volume=11 |issue=1 |page=6198 |doi=10.1038/s41598-021-85761-x |pmid=33737625 |pmc=7973739 }}. [24] => [25] => ===Chemical properties=== [26] => [[File:Cesium water.theora.ogv|left|thumb|Addition of a small amount of caesium to cold water is explosive.|alt=A person adds a small amount of metal to a petri dish with cold water which produces a small explosion.]] [27] => Caesium metal is highly reactive and [[pyrophoricity|pyrophoric]]. It ignites spontaneously in air, and reacts explosively with water even at low temperatures, more so than the other [[alkali metal]]s. It reacts with ice at temperatures as low as {{convert|−116|C}}. Because of this high reactivity, caesium metal is classified as a [[hazardous material]]. It is stored and shipped in dry, saturated hydrocarbons such as [[mineral oil]]. It can be handled only under [[inert gas]], such as [[argon]]. However, a caesium-water explosion is often less powerful than a [[sodium]]-water explosion with a similar amount of sodium. This is because caesium explodes instantly upon contact with water, leaving little time for [[hydrogen]] to accumulate.Gray, Theodore (2012) ''The Elements'', Black Dog & Leventhal Publishers, p. 131, {{ISBN|1-57912-895-5}}. Caesium can be stored in vacuum-sealed [[borosilicate glass]] [[ampoule]]s. In quantities of more than about {{convert|100|g|oz}}, caesium is shipped in hermetically sealed, stainless steel containers. [28] => [29] => The chemistry of caesium is similar to that of other alkali metals, in particular [[rubidium]], the element above caesium in the periodic table. As expected for an alkali metal, the only common oxidation state is +1.{{refn|It differs from this value in caesides, which contain the Cs− anion and thus have caesium in the −1 oxidation state. Additionally, 2013 calculations by Mao-sheng Miao indicate that under conditions of extreme pressure (greater than 30 [[pascal (unit)|GPa]]), the inner 5p electrons could form chemical bonds, where caesium would behave as the seventh 5p element. This discovery indicates that higher caesium fluorides with caesium in oxidation states from +2 to +6 could exist under such conditions.{{cite web |last=Moskowitz |first=Clara |title=A Basic Rule of Chemistry Can Be Broken, Calculations Show |url=http://www.scientificamerican.com/article.cfm?id=chemical-bonds-inner-shell-electrons |work=Scientific American |access-date=22 November 2013}}|name=oxistates|group=note}} Some slight differences arise from the fact that it has a higher [[atomic mass]] and is more [[electronegativity|electropositive]] than other (nonradioactive) alkali metals.{{cite book |publisher=Walter de Gruyter |date=1985 |edition=91–100 |pages=953–955 |isbn=978-3-11-007511-3 |title=Lehrbuch der Anorganischen Chemie |first1=Arnold F. |last1=Holleman |last2=Wiberg |first2=Egon |last3=Wiberg |first3=Nils |chapter=Vergleichende Übersicht über die Gruppe der Alkalimetalle |language=de}} Caesium is the most electropositive chemical element.{{refn|[[Francium]]'s electropositivity has not been experimentally measured due to its high radioactivity. Measurements of the first [[ionization energy]] of francium suggest that its [[relativistic quantum chemistry|relativistic effects]] may lower its reactivity and raise its electronegativity above that expected from [[periodic trend]]s.{{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]] |date=1987 |volume=59 |pages=1274–76 |doi=10.1103/PhysRevLett.59.1274 |pmid=10035190 |bibcode=1987PhRvL..59.1274A |issue=12}}|group=note}} The caesium ion is also larger and [[HSAB theory|less "hard"]] than those of the lighter [[alkali metals]]. [30] => [31] => ===Compounds=== [32] => [[File:CsCl polyhedra.png|thumb|left|upright|Ball-and-stick model of the cubic coordination of Cs and Cl in CsCl| alt=27 small grey spheres in 3 evenly spaced layers of nine. 8 spheres form a regular cube and 8 of those cubes form a larger cube. The grey spheres represent the caesium atoms. The center of each small cube is occupied by a small green sphere representing a chlorine atom. Thus, every chlorine is in the middle of a cube formed by caesium atoms and every caesium is in the middle of a cube formed by chlorine.]] [33] => [34] => Most caesium compounds contain the element as the [[cation]] {{chem|Cs|+}}, which [[ionic bond|binds ionically]] to a wide variety of [[anion]]s. One noteworthy exception is the [[alkalide|caeside]] anion ({{chem|Cs|−}}),{{cite journal |journal=[[Angewandte Chemie |Angewandte Chemie International Edition]] |year=1979 |first=J. L. |last=Dye |title=Compounds of Alkali Metal Anions |volume=18 |issue=8 |pages=587–598 |doi=10.1002/anie.197905871}} and others are the several suboxides (see section on oxides below). More recently, caesium is predicted to behave as a [[p-block]] element and capable of forming higher fluorides with higher [[oxidation state|oxidation states]] (i.e., CsFn with n > 1) under high pressure.{{cite journal |last=Miao |first=Mao-sheng |date=2013 |title=Caesium in high oxidation states and as a p-block element |url=https://www.nature.com/articles/nchem.1754 |journal=Nature Chemistry |language=en |volume=5 |issue=10 |pages=846–852 |doi=10.1038/nchem.1754 |pmid=24056341 |arxiv=1212.6290 |bibcode=2013NatCh...5..846M |s2cid=38839337 |issn=1755-4349}} This prediction needs to be validated by further experiments.{{cite journal |last1=Sneed |first1=D. |last2=Pravica |first2=M. |last3=Kim |first3=E. |last4=Chen |first4=N. |last5=Park |first5=C. |last6=White |first6=M. |date=1 October 2017 |title=Forcing Cesium into Higher Oxidation States Using Useful hard x-ray Induced Chemistry under High Pressure |journal=Journal of Physics: Conference Series |language=ENGLISH |volume=950 |issue=11, 2017 |page=042055 |doi=10.1088/1742-6596/950/4/042055 |bibcode=2017JPhCS.950d2055S |osti=1409108 |s2cid=102912809 |issn=1742-6588|doi-access=free }} [35] => [36] => Salts of Cs+ are usually colourless unless the anion itself is coloured. Many of the simple salts are [[hygroscopic]], but less so than the corresponding salts of lighter alkali metals. The [[phosphate]],Hogan, C. M. (2011).{{cite web |url=http://www.eoearth.org/article/Phosphate?topic=49557 |title=Phosphate |access-date=17 June 2012 |archive-url=https://web.archive.org/web/20121025180158/http://www.eoearth.org/article/Phosphate?topic=49557 |archive-date=25 October 2012}} in ''Encyclopedia of Earth''. Jorgensen, A. and Cleveland, C.J. (eds.). National Council for Science and the Environment. Washington DC [[acetate]], [[carbonate]], [[halide]]s, [[oxide]], [[nitrate]], and [[sulfate]] salts are water-soluble. Its [[double salt]]s are often less soluble, and the low solubility of caesium aluminium sulfate is exploited in refining Cs from ores. The double salts with antimony (such as {{chem|CsSbCl|4}}), [[bismuth]], [[cadmium]], [[copper]], [[iron]], and [[lead]] are also poorly [[dissolution (chemistry)|soluble]]. [37] => [38] => [[Caesium hydroxide]] (CsOH) is [[hygroscopic]] and strongly [[base (chemistry)|basic]]. It rapidly [[etching|etches]] the surface of [[semiconductor]]s such as [[silicon]].{{cite book |url=https://books.google.com/books?id=F-8SltAKSF8C&pg=PA90 |title=Etching in microsystem technology |author=Köhler, Michael J. |page=90 |publisher=Wiley-VCH |isbn=978-3-527-29561-6 |date=1999 }}{{Dead link|date=November 2023 |bot=InternetArchiveBot |fix-attempted=yes }} CsOH has been previously regarded by chemists as the "strongest base", reflecting the relatively weak attraction between the large Cs+ ion and OH−; it is indeed the strongest [[Arrhenius base]]; however, a number of compounds such as [[n-butyllithium|''n''-butyllithium]], [[sodium amide]], [[sodium hydride]], [[caesium hydride]], etc., which cannot be dissolved in water as reacting violently with it but rather only used in some [[anhydrous]] [[polar aprotic solvents]], are far more basic on the basis of the [[Brønsted–Lowry acid–base theory]]. [39] => [40] => A [[stoichiometry|stoichiometric]] mixture of caesium and [[gold]] will react to form yellow [[caesium auride]] (Cs+Au−) upon heating. The auride anion here behaves as a [[pseudohalogen]]. The compound reacts violently with water, yielding [[caesium hydroxide]], metallic gold, and hydrogen gas; in liquid ammonia it can be reacted with a caesium-specific ion exchange resin to produce [[tetramethylammonium auride]]. The analogous [[platinum]] compound, red caesium platinide ({{chem|Cs2Pt}}), contains the platinide ion that behaves as a {{chem name|pseudo[[chalcogen]]}}.{{cite journal |title=Effects of relativistic motion of electrons on the chemistry of gold and platinum |journal=Solid State Sciences |date=30 November 2005 |volume=7 |issue=12 |pages=1464–1474 |doi=10.1016/j.solidstatesciences.2005.06.015 |last=Jansen |first=Martin |bibcode=2005SSSci...7.1464J |doi-access=free}} [41] => [42] => ====Complexes==== [43] => Like all metal cations, Cs+ forms complexes with [[Lewis base]]s in solution. Because of its large size, Cs+ usually adopts [[coordination number]]s greater than 6, the number typical for the smaller alkali metal cations. This difference is apparent in the 8-coordination of CsCl. This high coordination number and [[HSAB|softness]] (tendency to form covalent bonds) are properties exploited in separating Cs+ from other cations in the remediation of nuclear wastes, where 137Cs+ must be separated from large amounts of nonradioactive K+.{{cite book |last1=Moyer |first1=Bruce A. |last2=Birdwell |first2=Joseph F. |last3=Bonnesen |first3=Peter V. |last4=Delmau |first4=Laetitia H. |journal=Macrocyclic Chemistry |pages=383–405 |date=2005 |doi=10.1007/1-4020-3687-6_24 |isbn=978-1-4020-3364-3 |title=Use of Macrocycles in Nuclear-Waste Cleanup: A Realworld Application of a Calixcrown in Cesium Separation Technology}}. [44] => [45] => ====Halides==== [46] => [[File:CsX@DWNT.jpg|thumb|upright|Monatomic caesium halide wires grown inside double-wall [[carbon nanotube]]s ([[transmission electron microscopy|TEM image]]).{{cite journal |doi=10.1038/ncomms8943 |pmid=26228378 |pmc=4532884 |title=Single-atom electron energy loss spectroscopy of light elements |journal=Nature Communications |volume=6 |pages=7943 |year=2015 |last1=Senga |first1=Ryosuke |last2=Suenaga |first2=Kazu |bibcode=2015NatCo...6.7943S}}]] [47] => [[Caesium fluoride]] (CsF) is a [[hygroscopic]] white solid that is widely used in [[organofluorine chemistry]] as a source of [[fluoride]] anions.{{cite journal |author=Evans, F. W. |author2=Litt, M. H. |author3=Weidler-Kubanek, A. M. |author4=Avonda, F. P. |title=Reactions Catalyzed by Potassium Fluoride. 111. The Knoevenagel Reaction |date=1968 |journal=Journal of Organic Chemistry |volume=33 |pages=1837–1839 |doi=10.1021/jo01269a028 |issue=5}} Caesium fluoride has the halite structure, which means that the Cs+ and F− pack in a [[cubic closest packed]] array as do Na+ and Cl− in [[sodium chloride]]. Notably, caesium and fluorine have the lowest and highest [[electronegativity|electronegativities]], respectively, among all the known elements. [48] => [49] => [[Caesium chloride]] (CsCl) crystallizes in the simple [[cubic crystal system]]. Also called the "caesium chloride structure", this structural motif is composed of a [[primitive cell|primitive]] cubic lattice with a two-atom basis, each with an eightfold [[coordination number|coordination]]; the chloride atoms lie upon the lattice points at the edges of the cube, while the caesium atoms lie in the holes in the centre of the cubes. This structure is shared with [[caesium bromide|CsBr]] and [[caesium iodide|CsI]], and many other compounds that do not contain Cs. In contrast, most other alkaline halides have the [[sodium chloride]] (NaCl) structure. The CsCl structure is preferred because Cs+ has an [[ionic radius]] of 174 [[picometer|pm]] and {{chem|Cl|−}} 181 pm.{{cite book |last=Wells |first=A. F. |date=1984 |title=Structural Inorganic Chemistry |edition=5th |publisher=Oxford Science Publications |isbn=978-0-19-855370-0}} [50] => [51] => ====Oxides==== [52] => [[File:Cs11O3 cluster.png|thumb|left|upright=0.7|{{chem|Cs|11|O|3}} cluster|alt=The stick and ball diagram shows three regular octahedra, which are connected to the next one by one surface and the last one shares one surface with the first. All three have one edge in common. All eleven vertices are purple spheres representing caesium, and at the center of each octahedron is a small red sphere representing oxygen.]] [53] => [54] => More so than the other alkali metals, caesium forms numerous binary compounds with [[oxygen]]. When caesium burns in air, the [[superoxide]] {{chem|CsO|2}} is the main product.{{cite book |last=Cotton |first=F. Albert |author2=Wilkinson, G. |title=Advanced Inorganic Chemistry |date=1962 |publisher=John Wiley & Sons, Inc. |page=318 |isbn=978-0-471-84997-1}} The "normal" [[caesium oxide]] ({{chem|Cs|2|O}}) forms yellow-orange [[hexagonal crystal system|hexagonal]] crystals,{{RubberBible87th|pages=451, 514}} and is the only oxide of the anti-[[cadmium chloride|{{chem|CdCl|2}}]] type.{{cite journal |doi=10.1021/j150537a022 |date=1956 |last1=Tsai |first1=Khi-Ruey |last2=Harris |first2=P. M. |last3=Lassettre |first3=E. N. |journal=Journal of Physical Chemistry |volume=60 |pages=338–344 |title=The Crystal Structure of Cesium Monoxide |issue=3 |url=http://www.dtic.mil/get-tr-doc/pdf?AD=AD0026963 |url-status=dead |archive-url=https://web.archive.org/web/20170924131429/http://www.dtic.mil/get-tr-doc/pdf?AD=AD0026963 |archive-date=24 September 2017}} It vaporizes at {{convert|250|°C}}, and decomposes to caesium metal and the [[peroxide]] [[caesium peroxide|{{chem|Cs|2|O|2}}]] at temperatures above {{convert|400|°C}}. In addition to the superoxide and the [[ozonide]] [[caesium ozonide|{{chem|CsO|3}}]],{{cite journal |doi=10.1007/BF00845494 |title=Synthesis of cesium ozonide through cesium superoxide |date=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}}{{cite journal |doi=10.1070/RC1971v040n02ABEH001903 |title=Alkali and Alkaline Earth Metal Ozonides |date=1971 |last1=Tokareva |first1=S. A. |journal=Russian Chemical Reviews |volume=40 |pages=165–174 |bibcode=1971RuCRv..40..165T |issue=2 |s2cid=250883291}} several brightly coloured [[suboxide]]s have also been studied.{{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 |date=1997 |volume=163 |pages=253–270 |doi=10.1016/S0010-8545(97)00013-1}} These include {{chem|Cs|7|O}}, {{chem|Cs|4|O}}, {{chem|Cs|11|O|3}}, {{chem|Cs|3|O}} (dark-green{{cite journal |doi=10.1021/j150537a023 |date=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}}), CsO, {{chem|Cs|3|O|2}},{{cite journal |doi=10.1007/s11669-009-9636-5 |title=Cs-O (Cesium-Oxygen) |date=2009 |last1=Okamoto |first1=H. |journal=Journal of Phase Equilibria and Diffusion |volume=31 |pages=86–87 |s2cid=96084147}} as well as {{chem|Cs|7|O|2}}.{{cite journal |doi=10.1021/jp036432o |title=Characterization of Oxides of Cesium |date=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}}{{cite journal |doi=10.1002/zaac.19472550110 |title=Untersuchungen ber das System Csium-Sauerstoff |date=1947 |last1=Brauer |first1=G. |journal=Zeitschrift für Anorganische Chemie |volume=255 |issue=1–3 |pages=101–124}} The latter may be heated in a vacuum to generate {{chem|Cs|2|O}}. Binary compounds with [[sulfur]], [[selenium]], and [[tellurium]] also exist. [55] => [56] => ===Isotopes=== [57] => {{Main|Isotopes of caesium}} [58] => Caesium has 41 known [[isotope]]s, ranging in [[mass number]] (i.e. number of [[nucleon]]s in the nucleus) from 112 to 152. Several of these are synthesized from lighter elements by the slow neutron capture process ([[S-process]]) inside old stars{{cite journal |doi=10.1146/annurev.astro.37.1.239 |author=Busso, M. |author2=Gallino, R. |author3=Wasserburg, G. J. |title=Nucleosynthesis in Asymptotic Giant Branch Stars: Relevance for Galactic Enrichment and Solar System Formation |journal=Annual Review of Astronomy and Astrophysics |volume=37 |date=1999 |pages=239–309 |url=http://authors.library.caltech.edu/1194/1/BUSaraa99.pdf |archive-url=https://ghostarchive.org/archive/20221010/http://authors.library.caltech.edu/1194/1/BUSaraa99.pdf |archive-date=10 October 2022 |url-status=live |access-date=20 February 2010 |bibcode=1999ARA&A..37..239B}} and by the [[R-process]] in [[supernova]] explosions.{{cite book |first=David |last=Arnett |date=1996 |title=Supernovae and Nucleosynthesis: An Investigation of the History of Matter, from the Big Bang to the Present |publisher=Princeton University Press |page=527 |isbn=978-0-691-01147-9}} The only [[stable nuclide|stable]] caesium isotope is 133Cs, with 78 [[neutron]]s. Although it has a large [[nuclear spin]] ({{sfrac|7|2}}+), [[nuclear magnetic resonance]] studies can use this isotope at a resonating frequency of 11.7 [[hertz|MHz]].{{cite journal |doi=10.1016/0277-5387(96)00018-6 |title=Complexation of caesium and rubidium cations with crown ethers in N,N-dimethylformamide |date=1996 |last1=Goff |first1=C. |journal=Polyhedron |volume=15 |pages=3897–3903 |last2=Matchette |first2=Michael A. |last3=Shabestary |first3=Nahid |last4=Khazaeli |first4=Sadegh |issue=21}} [59] => [[File:Cs-137-decay.svg|thumb|Decay of caesium-137|alt=A graph showing the energetics of caesium-137 (nuclear spin: I={{sfrac|7|2}}+, half-life of about 30 years) decay. With a 94.6% probability, it decays by a 512 keV beta emission into barium-137m (I=11/2-, t=2.55min); this further decays by a 662 keV gamma emission with an 85.1% probability into barium-137 (I={{sfrac|3|2}}+). Alternatively, caesium-137 may decay directly into barium-137 by a 0.4% probability beta emission.]] [60] => The radioactive [[caesium-135|135Cs]] has a very long half-life of about 2.3 million years, the longest of all radioactive isotopes of caesium. [[caesium-137|137Cs]] and [[caesium-134|134Cs]] have half-lives of 30 and two years, respectively. 137Cs decomposes to a short-lived [[barium-137m|137mBa]] by [[beta decay]], and then to nonradioactive barium, while 134Cs transforms into 134Ba directly. The isotopes with mass numbers of 129, 131, 132 and 136, have half-lives between a day and two weeks, while most of the other isotopes have half-lives from a few seconds to fractions of a second. At least 21 metastable [[nuclear isomer]]s exist. Other than 134mCs (with a half-life of just under 3 hours), all are very unstable and decay with half-lives of a few minutes or less.{{cite journal |doi=10.1016/0022-1902(55)80027-9 |title=The half-life of Cs137 |date=1955 |last1=Brown |first1=F. |last2=Hall |first2=G. R. |last3=Walter |first3=A. J. |journal=Journal of Inorganic and Nuclear Chemistry |volume=1 |pages=241–247 |issue=4–5 |bibcode=1955PhRv...99..188W}}{{cite web |url=http://www.nndc.bnl.gov/chart/ |title=Interactive Chart of Nuclides |publisher=Brookhaven National Laboratory |author=Sonzogni, Alejandro |location=National Nuclear Data Center |access-date=6 June 2008 |archive-date=22 May 2008 |archive-url=https://web.archive.org/web/20080522125027/http://www.nndc.bnl.gov/chart |url-status=dead}} [61] => [62] => The isotope 135Cs is one of the [[long-lived fission product]]s of [[uranium]] produced in [[nuclear reactor technology|nuclear reactors]].{{cite conference |conference=Seventh Information Exchange Meeting on Actinide and Fission Product Partitioning and Transmutation |date=14–16 October 2002 |place=Jeju, Korea |first1=Shigeo |last1=Ohki |first2=Naoyuki |last2=Takaki |title=Transmutation of Cesium-135 with Fast Reactors |url=http://www.oecd-nea.org/pt/docs/iem/jeju02/session6/SessionVI-08.pdf |access-date=26 September 2010 |archive-date=28 September 2011 |archive-url=https://web.archive.org/web/20110928005357/http://www.oecd-nea.org/pt/docs/iem/jeju02/session6/SessionVI-08.pdf |url-status=dead}} However, this [[fission product yield]] is reduced in most reactors because the predecessor, [[xenon-135|135Xe]], is a potent [[neutron poison]] and frequently transmutes to stable [[xenon-136|136Xe]] before it can decay to 135Cs.{{cite report |chapter-url=http://canteach.candu.org/library/20040720.pdf |title=CANDU Fundamentals |publisher=[[CANDU Owners Group]] Inc. |chapter=20 Xenon: A Fission Product Poison |access-date=15 September 2010 |url-status=dead |archive-url=https://web.archive.org/web/20110723231319/http://canteach.candu.org/library/20040720.pdf |archive-date=23 July 2011}}{{cite journal |journal=Journal of Environmental Radioactivity |title=Preliminary evaluation of 135Cs/137Cs as a forensic tool for identifying source of radioactive contamination |first1=V. F. |last1=Taylor |first2=R. D. |last2=Evans |first3=R. J. |last3=Cornett |doi=10.1016/j.jenvrad.2007.07.006 |volume=99 |issue=1 |date=2008 |pages=109–118 |pmid=17869392}} [63] => [64] => The [[beta decay]] from 137Cs to 137mBa results in [[gamma ray|gamma radiation]] as the 137mBa relaxes to ground state 137Ba, with the emitted photons having an energy of 0.6617 MeV.{{cite web |url=http://www.epa.gov/rpdweb00/radionuclides/cesium.html |title=Cesium {{pipe}} Radiation Protection |publisher=U.S. Environmental Protection Agency |date=28 June 2006 |access-date=15 February 2010 |url-status=dead |archive-url=https://web.archive.org/web/20110315034747/http://www.epa.gov/rpdweb00/radionuclides/cesium.html |archive-date=15 March 2011}} 137Cs and [[strontium-90|90Sr]] are the principal [[medium-lived fission product|medium-lived]] products of [[nuclear fission]], and the prime sources of [[radioactive decay|radioactivity]] from [[spent nuclear fuel]] after several years of cooling, lasting several hundred years.{{cite report |url=http://www.ieer.org/reports/transm/hisham.html |title=IEER Report: Transmutation – Nuclear Alchemy Gamble |publisher=Institute for Energy and Environmental Research |date=24 May 2000 |access-date=15 February 2010 |first=Hisham |last=Zerriffi}} Those two isotopes are the largest source of residual radioactivity in the area of the [[Chernobyl disaster]].{{cite report |url=http://www.iaea.org/Publications/Booklets/Chernobyl/chernobyl.pdf |title=Chernobyl's Legacy: Health, Environmental and Socia-Economic Impacts and Recommendations to the Governments of Belarus, Russian Federation and Ukraine |publisher=International Atomic Energy Agency |access-date=18 February 2010 |url-status=dead |archive-url=https://web.archive.org/web/20100215212227/http://www.iaea.org/Publications/Booklets/Chernobyl/chernobyl.pdf |archive-date=15 February 2010}} Because of the low capture rate, disposing of 137Cs through [[neutron capture]] is not feasible and the only current solution is to allow it to decay over time.{{cite journal |doi=10.3327/jnst.30.911 |title=Transmutation of Cesium-137 Using Proton Accelerator |first1=Takeshi |last1=Kase |first2=Kenji |last2=Konashi |first3=Hiroshi |last3=Takahashi |first4=Yasuo |last4=Hirao |volume=30 |issue=9 |date=1993 |pages=911–918 |journal=Journal of Nuclear Science and Technology |doi-access=free}} [65] => [66] => Almost all caesium produced from nuclear fission comes from the [[beta decay]] of originally more neutron-rich fission products, passing through various [[isotopes of iodine]] and [[isotopes of xenon|xenon]].{{cite book |isbn=978-1-56032-088-3 |publisher=Taylor & Francis |date=1992 |first=Ronald Allen |last=Knief |chapter-url=https://books.google.com/books?id=EpuaUEQaeoUC&pg=PA43 |page=42 |chapter=Fission Fragments |title=Nuclear engineering: theory and technology of commercial nuclear power}} Because iodine and xenon are volatile and can diffuse through nuclear fuel or air, radioactive caesium is often created far from the original site of fission.{{cite journal |title=Release of xenon-137 and iodine-137 from UO2 pellet by pulse neutron irradiation at NSRR |last1=Ishiwatari |first1=N. |last2=Nagai |first2=H. |pages=843–850 |volume=23 |issue=11 |journal=Nippon Genshiryoku Gakkaishi |osti=5714707}} With [[nuclear weapons testing]] in the 1950s through the 1980s, 137Cs was released into the [[atmosphere of Earth|atmosphere]] and returned to the surface of the earth as a component of [[nuclear fallout|radioactive fallout]]. It is a ready marker of the movement of soil and sediment from those times. [67] => [68] => ===Occurrence=== [69] => [[File:Pollucite-RoyalOntarioMuseum-Jan18-09.jpg|thumb|Pollucite, a caesium mineral|alt=A white mineral, from which white and pale pink crystals protrude]] [70] => {{See also|:Category:Caesium minerals|l1=Caesium minerals}} [71] => Caesium is a relatively rare element, estimated to average 3 [[parts per million]] in the [[abundance of elements in Earth's crust|Earth's crust]].{{cite journal |last1=Turekian |first1=K. K. |last2=Wedepohl |first2=K. H. |title=Distribution of the elements in some major units of the Earth's crust |journal=Geological Society of America Bulletin |volume=72 |issue=2 |pages=175–192 |doi=10.1130/0016-7606(1961)72[175:DOTEIS]2.0.CO;2 |issn=0016-7606 |bibcode=1961GSAB...72..175T |year=1961 |doi-access=free}} Nevertheless, it is more abundant than such elements as [[antimony]], [[cadmium]], [[tin]], and [[tungsten]], and two orders of magnitude more abundant than mercury and [[silver]]; it is 3.3% as abundant as [[rubidium]], with which it is closely associated, chemically. [72] => [73] => Due to its large [[ionic radius]], caesium is one of the "[[incompatible element]]s".{{cite web |url=http://www.asi.org/adb/02/13/02/cesium-occurrence-uses.html |title=Cesium as a Raw Material: Occurrence and Uses |first=Simon |last=Rowland |publisher=Artemis Society International |date=4 July 1998 |access-date=15 February 2010 |archive-date=8 July 2021 |archive-url=https://web.archive.org/web/20210708104437/http://www.asi.org/adb/02/13/02/cesium-occurrence-uses.html |url-status=dead }} During [[fractional crystallization (geology)|magma crystallization]], caesium is concentrated in the liquid phase and crystallizes last. Therefore, the largest deposits of caesium are zone [[pegmatite]] ore bodies formed by this enrichment process. Because caesium does not substitute for [[potassium]] as readily as rubidium does, the alkali evaporite minerals [[sylvite]] (KCl) and [[carnallite]] ({{chem|KMgCl|3|·6H|2|O}}) may contain only 0.002% caesium. Consequently, caesium is found in few minerals. Percentage amounts of caesium may be found in [[beryl]] ({{chem|Be|3|Al|2|(SiO|3|)|6}}) and [[avogadrite]] ({{chem|(K,Cs)BF|4}}), up to 15 wt% Cs2O in the closely related mineral [[pezzottaite]] ({{chem|Cs|(Be|2|Li)|Al|2|Si|6|O|18}}), up to 8.4 wt% Cs2O in the rare mineral [[londonite]] ({{chem|(Cs,K)Al|4|Be|4|(B,Be)|12|O|28}}), and less in the more widespread [[rhodizite]]. The only economically important ore for caesium is [[pollucite]] {{chem|Cs(AlSi|2|O|6|)}}, which is found in a few places around the world in zoned pegmatites, associated with the more commercially important [[lithium]] minerals, [[lepidolite]] and [[petalite]]. Within the pegmatites, the large grain size and the strong separation of the minerals results in high-grade ore for mining.{{cite journal |title=The Tanco Pegmatite at Bernic Lake, Manitoba: X. Pollucite |first1=Petr |last1=Černý |author-link1=Petr Černý |first2=F. M. |last2=Simpson |journal=Canadian Mineralogist |volume=16 |pages=325–333 |date=1978 |url=http://rruff.geo.arizona.edu/doclib/cm/vol38/CM38_877.pdf |archive-url=https://ghostarchive.org/archive/20221010/http://rruff.geo.arizona.edu/doclib/cm/vol38/CM38_877.pdf |archive-date=10 October 2022 |url-status=live |access-date=26 September 2010}} [74] => [75] => The world's most significant and richest known source of caesium is the [[Tanco Mine]] at [[Bernic Lake]] in [[Manitoba]], Canada, estimated to contain 350,000 [[tonne|metric tons]] of pollucite ore, representing more than two-thirds of the world's reserve base. Although the stoichiometric content of caesium in pollucite is 42.6%, pure pollucite samples from this deposit contain only about 34% caesium, while the average content is 24 wt%.{{cite web |title=Cesium |last=Polyak |first=Désirée E. |url=http://minerals.usgs.gov/minerals/pubs/commodity/cesium/mcs-2009-cesiu.pdf |archive-url=https://web.archive.org/web/20090508233834/http://minerals.usgs.gov/minerals/pubs/commodity/cesium/mcs-2009-cesiu.pdf |archive-date=8 May 2009 |url-status=live |publisher=U.S. Geological Survey |access-date=17 October 2009}} Commercial pollucite contains more than 19% caesium.{{cite book |last=Norton |first=J. J. |date=1973 |chapter=Lithium, cesium, and rubidium—The rare alkali metals |editor=Brobst, D. A. |editor2=Pratt, W. P. |title=United States mineral resources |publisher=U.S. Geological Survey Professional |volume=Paper 820 |pages=365–378 |chapter-url=https://pubs.er.usgs.gov/usgspubs/pp/pp820 |access-date=26 September 2010 |archive-date=21 July 2010 |archive-url=https://web.archive.org/web/20100721060544/http://pubs.er.usgs.gov/usgspubs/pp/pp820 |url-status=dead}} The [[Bikita District|Bikita]] pegmatite deposit in [[Zimbabwe]] is mined for its petalite, but it also contains a significant amount of pollucite. Another notable source of pollucite is in the [[Erongo Region|Karibib Desert]], [[Namibia]]. At the present rate of world mine production of 5 to 10 metric tons per year, reserves will last for thousands of years. [76] => [77] => ==Production== [78] => Mining and refining pollucite ore is a selective process and is conducted on a smaller scale than for most other metals. The ore is crushed, hand-sorted, but not usually concentrated, and then ground. Caesium is then extracted from pollucite primarily by three methods: acid digestion, alkaline decomposition, and direct reduction.{{cite book |last=Burt |first=R. O. |date=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}} [79] => [80] => In the acid digestion, the [[silicate]] pollucite rock is dissolved with strong acids, such as [[hydrochloric acid|hydrochloric]] (HCl), [[sulfuric acid|sulfuric]] ({{chem|H|2|SO|4}}), [[hydrobromic acid|hydrobromic]] (HBr), or [[hydrofluoric acid|hydrofluoric]] (HF) acids. With hydrochloric acid, a mixture of soluble chlorides is produced, and the insoluble chloride double salts of caesium are precipitated as caesium antimony chloride ({{chem|Cs|4|SbCl|7}}), caesium iodine chloride ({{chem|Cs|2|ICl}}), or caesium hexachlorocerate ({{chem|Cs|2|(CeCl|6|)}}). After separation, the pure precipitated double salt is decomposed, and pure CsCl is precipitated by evaporating the water. [81] => [82] => The sulfuric acid method yields the insoluble double salt directly as caesium [[alum]] ({{chem|CsAl(SO|4|)|2|·12H|2|O}}). The [[aluminium sulfate]] component is converted to insoluble [[aluminium oxide]] by roasting the alum with [[carbon]], and the resulting product is [[leaching (metallurgy)|leached]] with water to yield a {{chem|Cs|2|SO|4}} solution. [83] => [84] => Roasting pollucite with [[calcium carbonate]] and [[calcium chloride]] yields insoluble calcium silicates and soluble caesium chloride. Leaching with water or dilute [[ammonia]] ({{chem|NH|4|OH}}) yields a dilute chloride (CsCl) solution. This solution can be evaporated to produce caesium chloride or transformed into caesium alum or caesium carbonate. Though not commercially feasible, the ore can be directly reduced with potassium, sodium, or calcium in vacuum to produce caesium metal directly. [85] => [86] => Most of the mined caesium (as salts) is directly converted into [[formate|caesium formate]] (HCOO−Cs+) for applications such as [[oil drilling]]. To supply the developing market, [[Cabot Corporation]] built a production plant in 1997 at the [[Tanco mine]] near [[Bernic Lake]] in [[Manitoba]], with a capacity of {{convert|12000|oilbbl|m3}} per year of caesium formate solution.{{cite journal |last1=Benton |first1=William |last2=Turner |first2=Jim |date=2000 |title=Cesium formate fluid succeeds in North Sea HPHT field trials |journal=Drilling Contractor |issue=May/June |pages=38–41 |url=http://www.iadc.org/dcpi/dc-mayjun00/m-cabot.pdf |archive-url=https://web.archive.org/web/20010706223144/http://www.iadc.org/dcpi/dc-mayjun00/m-cabot.pdf |archive-date=2001-07-06 |url-status=live |access-date=26 September 2010}} The primary smaller-scale commercial compounds of caesium are [[caesium chloride]] and [[caesium nitrate|nitrate]]. [87] => [88] => Alternatively, caesium metal may be obtained from the purified compounds derived from the ore. [[Caesium chloride]] and the other caesium halides can be reduced at {{convert|700|to|800|°C|°F}} with calcium or [[barium]], and caesium metal distilled from the result. In the same way, the aluminate, carbonate, or hydroxide may be reduced by [[magnesium]]. [89] => [90] => The metal can also be isolated by [[electrolysis]] of fused caesium [[cyanide]] (CsCN). Exceptionally pure and gas-free caesium can be produced by {{convert|390|°C}} thermal decomposition of caesium [[azide]] {{chem|CsN|3}}, which can be produced from aqueous [[caesium sulfate]] and [[barium azide]]. In vacuum applications, caesium [[dichromate]] can be reacted with [[zirconium]] to produce pure caesium metal without other gaseous products.{{cite book |isbn=978-3-11-011451-5 |url=https://books.google.com/books?id=Owuv-c9L_IMC&pg=PA198 |page=198 |others=Eagleson, Mary |editor=Eagleson, Mary |year=1994 |publisher=de Gruyter |location=Berlin |title=Concise encyclopedia chemistry}} [91] => :{{chem|Cs|2|Cr|2|O|7}} + 2 {{chem|Zr}} → 2 {{chem|Cs}} + 2 {{chem|ZrO|2}}+ {{chem|Cr|2|O|3}} [92] => [93] => The price of 99.8% pure caesium (metal basis) in 2009 was about {{convert|10|$/g|$/oz|-1}}, but the compounds are significantly cheaper. [94] => [95] => ==History== [96] => [[File:Kirchhoff Bunsen Roscoe.jpg|thumb|[[Gustav Kirchhoff]] (left) and [[Robert Bunsen]] (centre) discovered caesium with their newly invented spectroscope.| alt=Three middle-aged men, with the one in the middle sitting down. All wear long jackets, and the shorter man on the left has a beard.]] [97] => In 1860, [[Robert Bunsen]] and [[Gustav Kirchhoff]] discovered caesium in the [[mineral water]] from [[Bad Dürkheim|Dürkheim]], Germany. Because of the bright blue lines in the [[emission spectrum]], they derived the name from the [[Latin]] word {{lang|la|caesius}}, meaning {{gloss|bluish grey}}.Bunsen quotes [[Aulus Gellius|Aulus Gellius Noctes Atticae]] II, 26 by [[Nigidius Figulus]]: ''Nostris autem veteribus caesia dicts est quae Graecis, ut Nigidus ait, de colore coeli quasi coelia.''[[Oxford English Dictionary]], 2nd Edition{{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 |date=1932 |bibcode=1932JChEd...9.1413W}} Caesium was the first element to be discovered with a [[spectroscopy|spectroscope]], which had been invented by Bunsen and Kirchhoff only a year previously. [98] => [99] => To obtain a pure sample of caesium, {{convert|44,000|litre}} of mineral water had to be evaporated to yield {{convert|240|kg}} of concentrated salt solution. The [[alkaline earth metal]]s were precipitated either as sulfates or [[oxalate]]s, leaving the alkali metal in the solution. After conversion to the [[nitrate]]s and extraction with [[ethanol]], a sodium-free mixture was obtained. From this mixture, the lithium was precipitated by [[ammonium carbonate]]. Potassium, rubidium, and caesium form insoluble salts with [[chloroplatinic acid]], but these salts show a slight difference in solubility in hot water, and the less-soluble caesium and rubidium hexachloroplatinate ({{chem2|(Cs,Rb)2PtCl6}}) were obtained by [[fractional crystallization (chemistry)|fractional crystallization]]. After reduction of the hexachloroplatinate with [[hydrogen]], caesium and rubidium were separated by the difference in solubility of their carbonates in alcohol. The process yielded {{convert|9.2|g}} of [[rubidium chloride]] and {{convert|7.3|g}} of caesium chloride from the initial 44,000 litres of mineral water.{{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 |Annalen der Physik und Chemie]] |volume=189 |issue=7 |date=1861 |bibcode=1861AnP...189..337K |url=http://archiv.ub.uni-heidelberg.de/volltextserver/15657/1/spektral.pdf |archive-url=https://web.archive.org/web/20160302113524/http://archiv.ub.uni-heidelberg.de/volltextserver/15657/1/spektral.pdf |archive-date=2016-03-02 |url-status=live |hdl=2027/hvd.32044080591324}} [100] => [101] => From the caesium chloride, the two scientists estimated the [[atomic weight]] of the new element at 123.35 (compared to the currently accepted one of 132.9). They tried to generate elemental caesium by electrolysis of molten caesium chloride, but instead of a metal, they obtained a blue homogeneous substance which "neither under the naked eye nor under the microscope showed the slightest trace of metallic substance"; as a result, they assigned it as a [[non-stoichiometric compound|subchloride]] ({{chem|Cs|2|Cl}}). In reality, the product was probably a [[colloid]]al mixture of the metal and caesium chloride.{{cite book |last=Zsigmondy |first=Richard |title=Colloids and the Ultra Microscope |publisher=Read books |date=2007 |isbn=978-1-4067-5938-9 |page=69 |url=https://books.google.com/books?id=Ac2mGhqjgUkC&pg=PAPA69}} The electrolysis of the aqueous solution of chloride with a mercury cathode produced a caesium amalgam which readily decomposed under the aqueous conditions. The pure metal was eventually isolated by the Swedish chemist [[Carl Setterberg]] while working on his doctorate with [[Friedrich August Kekulé von Stradonitz|Kekulé]] and Bunsen. In 1882, he produced caesium metal by electrolysing [[caesium cyanide]], avoiding the problems with the chloride.{{cite journal |title=Ueber die Darstellung von Rubidium- und Cäsiumverbindungen und über die Gewinnung der Metalle selbst |doi=10.1002/jlac.18822110105 |date=1882 |last1=Setterberg |first1=Carl |journal=Justus Liebig's Annalen der Chemie |volume=211 |pages=100–116 |url=https://zenodo.org/record/1447367}} [102] => [103] => Historically, the most important use for caesium has been in research and development, primarily in chemical and electrical fields. Very few applications existed for caesium until the 1920s, when it came into use in radio [[vacuum tube]]s, where it had two functions; as a [[getter]], it removed excess oxygen after manufacture, and as a coating on the heated [[cathode]], it increased the [[electrical conductivity]]. Caesium was not recognized as a high-performance industrial metal until the 1950s.{{cite journal |last=Strod |first=A. J. |date=1957 |title=Cesium—A new industrial metal |journal=American Ceramic Bulletin |volume=36 |issue=6 |pages=212–213}} Applications for nonradioactive caesium included [[solar cell|photoelectric cells]], [[photomultiplier]] tubes, optical components of [[infrared spectroscopy|infrared spectrophotometers]], catalysts for several organic reactions, crystals for [[scintillation counter]]s, and in [[MHD generator|magnetohydrodynamic power generators]]. Caesium is also used as a source of positive ions in [[secondary ion mass spectrometry]] (SIMS). [104] => [105] => Since 1967, the [[International System of Units|International System of Measurements]] has based the primary unit of time, the second, on the properties of caesium. The International System of Units (SI) defines the second as the duration of 9,192,631,770 cycles at the [[microwave]] [[frequency]] of the [[spectral line]] corresponding to the transition between two [[hyperfine structure|hyperfine]] [[energy level]]s of the [[ground state]] of [[caesium-133]].{{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}} The 13th [[General Conference on Weights and Measures]] of 1967 defined a second as: "the duration of 9,192,631,770 cycles of microwave light absorbed or emitted by the hyperfine transition of caesium-133 atoms in their ground state undisturbed by external fields". [106] => [107] => ==Applications== [108] => [109] => ===Petroleum exploration=== [110] => The largest present-day use of nonradioactive caesium is in [[formate|caesium formate]] [[drilling fluid]]s for the [[extractive oil industry]]. Aqueous solutions of caesium formate (HCOO−Cs+)—made by reacting caesium hydroxide with [[formic acid]]—were developed in the mid-1990s for use as oil well drilling and [[completion (oil and gas wells)|completion fluids]]. The function of a drilling fluid is to lubricate drill bits, to bring rock cuttings to the surface, and to maintain pressure on the formation during drilling of the well. Completion fluids assist the emplacement of control hardware after drilling but prior to production by maintaining the pressure. [111] => [112] => The high density of the caesium formate brine (up to 2.3 g/cm3, or 19.2 pounds per gallon),{{cite conference |conference=IADC/SPE Drilling Conference |date=February 2006 |location=Miami, Florida, USASociety of Petroleum Engineers |first1=J. D. |last1=Downs |first2=M. |last2=Blaszczynski |first3=J. |last3=Turner |first4=M. |last4=Harris |doi=10.2118/99068-MS |url=http://www.spe.org/elibinfo/eLibrary_Papers/spe/2006/06DC/SPE-99068-MS/SPE-99068-MS.htm |archive-url=https://web.archive.org/web/20071012122901/http://spe.org/elibinfo/eLibrary_Papers/spe/2006/06DC/SPE-99068-MS/SPE-99068-MS.htm |archive-date=12 October 2007 |title=Drilling and Completing Difficult HP/HT Wells With the Aid of Cesium Formate Brines-A Performance Review}} coupled with the relatively benign nature of most caesium compounds, reduces the requirement for toxic high-density suspended solids in the drilling fluid—a significant technological, engineering and environmental advantage. Unlike the components of many other heavy liquids, caesium formate is relatively environment-friendly. Caesium formate brine can be blended with potassium and sodium formates to decrease the density of the fluids to that of water (1.0 g/cm3, or 8.3 pounds per gallon). Furthermore, it is biodegradable and may be recycled, which is important in view of its high cost (about $4,000 per [[barrel (volume)#Oil barrel|barrel]] in 2001).{{cite journal |last=Flatern |first=Rick |date=2001 |title=Keeping cool in the HPHT environment |journal=Offshore Engineer |issue=February |pages=33–37}} Alkali formates are safe to handle and do not damage the producing formation or downhole metals as corrosive alternative, high-density brines (such as [[zinc bromide]] {{Chem|ZnBr|2}} solutions) sometimes do; they also require less cleanup and reduce disposal costs. [113] => [114] => ===Atomic clocks=== [115] => [[File:Usno-mc.jpg|thumb|Atomic clock ensemble at the U.S. Naval Observatory|alt=A room with a black box in the foreground and six control cabinets with space for five to six racks each. Most, but not all, of the cabinets are filled with white boxes.]] [116] => Caesium-based [[atomic clock]]s use the [[electromagnetic radiation|electromagnetic transitions]] in the [[hyperfine structure]] of caesium-133 atoms as a reference point. The first accurate caesium clock was built by [[Louis Essen]] in 1955 at the [[National Physical Laboratory, UK|National Physical Laboratory]] in the UK.{{cite journal |first1=L. |last1=Essen |first2=J. V. L. |last2=Parry |date=1955 |title=An Atomic Standard of Frequency and Time Interval: A Caesium Resonator |journal=[[Nature (journal) |Nature]] |volume=176 |pages=280–282 |doi=10.1038/176280a0 |bibcode=1955Natur.176..280E |issue=4476 |s2cid=4191481}} Caesium clocks have improved over the past half-century and are regarded as "the most accurate realization of a unit that mankind has yet achieved." These clocks measure frequency with an error of 2 to 3 parts in 1014, which corresponds to an accuracy of 2 [[nanosecond]]s per day, or one second in 1.4 million years. The latest versions are more accurate than 1 part in 1015, about 1 second in 20 million years. The [[caesium standard]] is the primary standard for standards-compliant time and frequency measurements.{{cite journal |last1=Markowitz |first1=W. |last2=Hall |first2=R. |last3=Essen |first3=L. |last4=Parry |first4=J. |title=Frequency of Cesium in Terms of Ephemeris Time |doi=10.1103/PhysRevLett.1.105 |journal=Physical Review Letters |volume=1 |issue=3 |pages=105–107 |year=1958 |bibcode=1958PhRvL...1..105M}} Caesium clocks regulate the timing of cell phone networks and the Internet.{{cite news |first=Monte |last=Reel |date=22 July 2003 |title=Where timing truly is everything |newspaper=The Washington Post |page=B1 |url=http://www.highbeam.com/doc/1P2-284155.html |access-date=26 January 2010 |archive-url=https://web.archive.org/web/20130429044454/http://www.highbeam.com/doc/1P2-284155.html |archive-date=29 April 2013 |url-status=dead}} [117] => [118] => ====Definition of the second==== [119] => The second, symbol ''s'', is the SI unit of time. The [[BIPM]] restated its definition at its 26th conference in 2018: "[The second] is defined by taking the fixed numerical value of the caesium frequency {{math|Δ''ν''Cs}}, the unperturbed ground-state hyperfine transition frequency of the caesium-133 atom, to be {{val|9192631770}} when expressed in the unit [[Hz]], which is equal to s−1."{{cite web |title=Resolution 1 of the 26th CGPM |url=https://www.bipm.org/en/CGPM/db/26/1/ |publisher=Bureau International des Poids et Mesures |location=Paris |pages=472 of the official French publication |language=FR,EN |date=2018 |access-date=2019-12-29 |archive-date=2021-02-04 |archive-url=https://web.archive.org/web/20210204120336/https://www.bipm.org/en/CGPM/db/26/1/ |url-status=dead }} [120] => [121] => ===Electric power and electronics=== [122] => Caesium vapour [[thermionic converter|thermionic generators]] are low-power devices that convert heat energy to electrical energy. In the two-electrode [[vacuum tube]] converter, caesium neutralizes the space charge near the cathode and enhances the current flow.{{cite journal |last1=Rasor |first1=Ned S. |first2=Charles |last2=Warner |title=Correlation of Emission Processes for Adsorbed Alkali Films on Metal Surfaces |journal=Journal of Applied Physics |volume=35 |issue=9 |pages=2589–2600 |date=September 1964 |doi=10.1063/1.1713806 |bibcode=1964JAP....35.2589R}} [123] => [124] => Caesium is also important for its [[photoelectric effect|photoemissive]] properties, converting light to electron flow. It is used in [[solar cell|photoelectric cells]] because caesium-based cathodes, such as the intermetallic compound {{chem|K|2|CsSb}}, have a low threshold voltage for emission of [[electron]]s.{{cite web |url=https://www.americanelements.com/cs.html |title=Cesium Supplier & Technical Information |publisher=American Elements |access-date=25 January 2010}} The range of photoemissive devices using caesium include [[optical character recognition]] devices, [[photomultiplier|photomultiplier tubes]], and [[video camera tube]]s.{{cite journal |doi=10.1063/1.3215593 |title=K2CsSb Cathode Development |journal=AIP Conference Proceedings |date=2009 |volume=1149 |issue=1 |pages=1062–1066 |first1=John |last1=Smedley |first2=Triveni |last2=Rao |first3=Erdong |last3=Wang |bibcode=2009AIPC.1149.1062S}}{{cite journal |first=P. |last=Görlich |title=Über zusammengesetzte, durchsichtige Photokathoden |journal=Zeitschrift für Physik |volume=101 |pages=335–342 |date=1936 |doi=10.1007/BF01342330 |bibcode=1936ZPhy..101..335G |issue=5–6 |s2cid=121613539}} Nevertheless, [[germanium]], rubidium, selenium, silicon, tellurium, and several other elements can be substituted for caesium in photosensitive materials. [125] => [126] => [[Caesium iodide]] (CsI), [[caesium bromide|bromide]] (CsBr) and [[caesium fluoride|fluoride]] (CsF) crystals are employed for [[scintillator]]s in [[scintillation counter]]s widely used in mineral exploration and particle physics research to detect [[gamma ray|gamma]] and [[X-ray]] radiation. Being a heavy element, caesium provides good stopping power with better detection. Caesium compounds may provide a faster response (CsF) and be less hygroscopic (CsI). [127] => [128] => Caesium vapour is used in many common [[magnetometer]]s.{{cite journal |doi=10.1007/s00340-005-1773-x |title=Comparison of discharge lamp and laser pumped cesium magnetometers |date=2005 |last1=Groeger |first1=S. |first2=A. S. |first3=A. |journal=Applied Physics B |volume=80 |pages=645–654 |last2=Pazgalev |last3=Weis |arxiv=physics/0412011 |bibcode=2005ApPhB..80..645G |issue=6 |s2cid=36065775}} [129] => [130] => The element is used as an [[internal standard]] in [[spectrophotometry]].{{cite book |chapter-url=https://books.google.com/books?id=z9SzvsSCHv4C&pg=PA108 |page=108 |isbn=978-0-471-28572-4 |chapter=Internal Standards |date=1994 |first1=Mary C. |last1=Haven |first2=Gregory A. |last2=Tetrault |first3=Jerald R. |last3=Schenken |publisher=John Wiley and Sons |location=New York |title=Laboratory instrumentation}} Like other [[alkali metal]]s, caesium has a great affinity for [[oxygen]] and is used as a "[[getter]]" in [[vacuum tube]]s.{{cite book |url=https://books.google.com/books?id=1o1WECNJkscC&pg=PA391 |title=Photo-electronic image devices: proceedings of the fourth symposium held at Imperial College, London, 16–20 September 1968 |volume=1 |publisher=Academic Press |date=1969 |first=James D. |last=McGee |page=391 |isbn=978-0-12-014528-7}} Other uses of the metal include high-energy [[laser]]s, [[fluorescent lamp|vapour glow lamps]], and vapour [[rectifier]]s. [131] => [132] => ===Centrifugation fluids=== [133] => The high density of the caesium ion makes solutions of caesium chloride, caesium sulfate, and caesium [[trifluoroacetic acid|trifluoroacetate]] ({{chem|Cs(O|2|CCF|3|)}}) useful in molecular biology for density gradient [[differential centrifugation|ultracentrifugation]].Manfred Bick, Horst Prinz, "Cesium and Cesium Compounds" in Ullmann's Encyclopedia of Industrial Chemistry 2005, Wiley-VCH, Weinheim. {{doi|10.1002/14356007.a06_153}}. This technology is used primarily in the isolation of [[virus|viral particles]], subcellular [[organelle]]s and fractions, and [[nucleic acid]]s from biological samples.{{cite book |chapter-url=https://books.google.com/books?id=1kn89nI2gUsC&pg=PA61 |pages=61–62 |isbn=978-0-89603-564-5 |chapter=Gradient Materials |editor=Desai, Mohamed A. |date=2000 |publisher=Humana Press |location=Totowa, N.J. |title=Downstream processing methods}} [134] => [135] => ===Chemical and medical use=== [136] => [[File:Caesium chloride.jpg|thumb|alt=Some fine white powder on a laboratory watch glass|Caesium chloride powder]] [137] => Relatively few chemical applications use caesium.{{cite book |last=Burt |first=R. O. |date=1993 |chapter=Cesium and cesium compounds |title=Kirk-Othmer encyclopedia of chemical technology |edition=4th |place=New York |publisher=John Wiley & Sons |volume=5 |page=759 |isbn=978-0-471-15158-6}} Doping with caesium compounds enhances the effectiveness of several metal-ion catalysts for chemical synthesis, such as [[acrylic acid]], [[anthraquinone]], [[ethylene oxide]], [[methanol]], [[phthalic anhydride]], [[styrene]], [[methyl methacrylate]] monomers, and various [[alkene|olefins]]. It is also used in the catalytic conversion of [[sulfur dioxide]] into [[sulfur trioxide]] in the production of [[sulfuric acid]]. [138] => [139] => [[Caesium fluoride]] enjoys a niche use in [[organic chemistry]] as a [[base (chemistry)|base]]{{cite book |last1=Greenwood |first1=N. N. |last2=Earnshaw |first2=A. |title=Chemistry of the Elements |publisher=Pergamon Press |place=Oxford, UK |date=1984 |isbn=978-0-08-022057-4}} and as an [[anhydrous]] source of [[fluoride]] ion. [140] => Friestad, Gregory K.; Branchaud, Bruce P.; Navarrini, Walter and Sansotera, Maurizio (2007) "Cesium Fluoride" in ''Encyclopedia of Reagents for Organic Synthesis'', John Wiley & Sons. {{doi|10.1002/047084289X.rc050.pub2}} Caesium salts sometimes replace potassium or sodium salts in [[organic synthesis]], such as [[cyclic compound|cyclization]], [[esterification]], and [[polymerization]]. Caesium has also been used in thermoluminescent radiation [[dosimetry]] (TLD): When exposed to radiation, it acquires crystal defects that, when heated, revert with emission of light proportionate to the received dose. Thus, measuring the light pulse with a [[photomultiplier tube]] can allow the accumulated radiation dose to be quantified. [141] => [142] => ===Nuclear and isotope applications=== [143] => [[Caesium-137]] is a [[radionuclide|radioisotope]] commonly used as a [[gamma ray|gamma]]-emitter in industrial applications. Its advantages include a half-life of roughly 30 years, its availability from the [[nuclear fuel cycle]], and having [[Isotopes of barium|137Ba]] as a stable end product. The high water solubility is a disadvantage which makes it incompatible with large pool irradiators for food and medical supplies.{{cite web |url=http://earth1.epa.gov/radiation/docs/source-management/csfinallongtakeshi.pdf |title=The material flow of radioactive cesium-137 in the U.S. 2000 |first=Takeshi |last=Okumura |date=21 October 2003 |access-date=20 December 2009 |publisher=United States Environmental Protection Agency |url-status=dead |archive-url=https://web.archive.org/web/20110720163223/http://earth1.epa.gov/radiation/docs/source-management/csfinallongtakeshi.pdf |archive-date=20 July 2011}} It has been used in agriculture, cancer treatment, and the [[sterilization (microbiology)|sterilization]] of food, sewage sludge, and surgical equipment.{{cite book |last=Jensen |first=N. L. |date=1985 |chapter=Cesium |title=Mineral facts and problems |publisher=U.S. Bureau of Mines |volume=Bulletin 675 |pages=133–138}} Radioactive [[isotopes of caesium]] in [[radiation therapy|radiation devices]] were used in the medical field to treat certain types of cancer,{{cite web |url=http://www.medicalnewstoday.com/releases/91994.php |title=IsoRay's Cesium-131 Medical Isotope Used In Milestone Procedure Treating Eye Cancers At Tufts-New England Medical Center |date=17 December 2007 |work=Medical News Today |access-date=15 February 2010}} but emergence of better alternatives and the use of water-soluble caesium chloride in the sources, which could create wide-ranging contamination, gradually put some of these caesium sources out of use.{{cite book |chapter-url=https://books.google.com/books?id=bk0go_-FO5QC&pg=PA22 |isbn=978-0-07-005115-7 |chapter=Caesium-137 Machines |title=Radiation therapy planning |first=Gunilla Carleson |last=Bentel |publisher=McGraw-Hill Professional |date=1996 |access-date=26 September 2010 |pages=22–23}}{{cite book |isbn=978-0-309-11014-3 |url=https://books.google.com/books?id=3cT2REdXJ98C |title=Radiation source use and replacement: abbreviated version |author=National Research Council (U.S.). Committee on Radiation Source Use and Replacement |publisher=National Academies Press |date=2008}} Caesium-137 has been employed in a variety of industrial measurement gauges, including moisture, density, levelling, and thickness gauges.{{cite book |chapter=Level and density measurement using non-contact nuclear gauges |isbn=978-0-412-53400-3 |chapter-url=https://books.google.com/books?id=RwsoQbHYjvwC&pg=PA82 |pages=82–85 |editor=Loxton, R. |editor2=Pope, P. |date=1995 |publisher=Chapman & Hall |location=London |title=Instrumentation : A Reader}} It has also been used in [[well logging]] devices for measuring the [[electron density]] of the rock formations, which is analogous to the bulk density of the formations.{{cite journal |doi=10.1146/annurev.ea.13.050185.001531 |title=Downhole Geophysical Logging |date=1985 |last1=Timur |first1=A. |last2=Toksoz |first2=M. N. |journal=Annual Review of Earth and Planetary Sciences |volume=13 |pages=315–344 |bibcode=1985AREPS..13..315T}} [144] => [145] => Caesium-137 has been used in [[hydrology|hydrologic]] studies analogous to those with [[tritium]]. As a daughter product of fission bomb testing from the 1950s through the mid-1980s, caesium-137 was released into the atmosphere, where it was absorbed readily into solution. Known year-to-year variation within that period allows correlation with soil and sediment layers. Caesium-134, and to a lesser extent caesium-135, have also been used in hydrology to measure the caesium output by the nuclear power industry. While they are less prevalent than either caesium-133 or caesium-137, these bellwether isotopes are produced solely from anthropogenic sources.{{cite web |first=Carol |last=Kendall |author-link=Carol Kendall (scientist) |url=http://wwwrcamnl.wr.usgs.gov/isoig/period/cs_iig.html |title=Isotope Tracers Project – Resources on Isotopes – Cesium |publisher=National Research Program – U.S. Geological Survey |access-date=25 January 2010}} [146] => [147] => ===Other uses=== [148] => [[File:Electrostatic ion thruster-en.svg|thumb|upright=1.4|Schematics of an electrostatic ion thruster developed for use with caesium or mercury fuel|alt=Electrons beamed from an electron gun hit and ionize neutral fuel atoms; in a chamber surrounded by magnets, the positive ions are directed toward a negative grid that accelerates them. The force of the engine is created by expelling the ions from the rear at high velocity. On exiting, the positive ions are neutralized from another electron gun, ensuring that neither the ship nor the exhaust is electrically charged and are not attracted.]] [149] => Caesium and mercury were used as a propellant in early [[ion thruster|ion engines]] designed for [[spacecraft propulsion]] on very long interplanetary or extraplanetary missions. The fuel was ionized by contact with a charged [[tungsten]] electrode. But corrosion by caesium on spacecraft components has pushed development in the direction of inert gas propellants, such as [[xenon]], which are easier to handle in ground-based tests and do less potential damage to the spacecraft. Xenon was used in the experimental spacecraft [[Deep Space 1]] launched in 1998.{{cite journal |doi=10.1063/1.1150468 |title=NSTAR Xenon Ion Thruster on Deep Space 1: Ground and flight tests (invited) |date=2000 |last1=Marcucci |first1=M. G. |last2=Polk |first2=J. E. |journal=Review of Scientific Instruments |volume=71 |pages=1389–1400 |bibcode=2000RScI...71.1389M |issue=3}}{{cite web |url=http://gltrs.grc.nasa.gov/reports/1999/TM-1999-209439.pdf |title=A Synopsis of Ion Propulsion Development Projects in the United States: SERT I to Deep Space I |first1=James S. |last1=Sovey |first2=Vincent K. |last2=Rawlin |first3=Michael J. |last3=Patterson |publisher=NASA |access-date=12 December 2009 |url-status=dead |archive-url=https://web.archive.org/web/20090629225625/http://gltrs.grc.nasa.gov/reports/1999/TM-1999-209439.pdf |archive-date=29 June 2009}} Nevertheless, [[field-emission electric propulsion]] thrusters that accelerate liquid metal ions such as caesium have been built.{{cite conference |url=http://trs-new.jpl.nasa.gov/dspace/handle/2014/11649 |title=In-FEEP Thruster Ion Beam Neutralization with Thermionic and Field Emission Cathodes |format=PDF |access-date=25 January 2010 |conference=27th International Electric Propulsion Conference |place=Pasadena, California |date=October 2001 |pages=1–15 |author=Marrese, C. |author2=Polk, J. |author3=Mueller, J. |author4=Owens, A. |author5=Tajmar, M. |author6=Fink, R. |author7=Spindt, C. |name-list-style=amp |url-status=dead |archive-url=https://web.archive.org/web/20100527071653/http://trs-new.jpl.nasa.gov/dspace/handle/2014/11649 |archive-date=27 May 2010}} [150] => [151] => [[Caesium nitrate]] is used as an [[oxidizing agent|oxidizer]] and [[pyrotechnic colorant]] to burn [[silicon]] in [[infrared]] [[flare (pyrotechnic)|flares]],{{cite web |url=http://www.freepatentsonline.com/6230628.html |work=United States Patent 6230628 |title=Infrared illumination compositions and articles containing the same |publisher=Freepatentsonline.com |access-date=25 January 2010}} such as the LUU-19 flare,{{cite web |url=https://fas.org/man/dod-101/sys/dumb/luu19.htm |title=LUU-19 Flare |publisher=Federation of American Scientists |date=23 April 2000 |access-date=12 December 2009 |url-status=dead |archive-url=https://web.archive.org/web/20100806093502/http://www.fas.org/man/dod-101/sys/dumb/luu19.htm |archive-date=6 August 2010}} because it emits much of its light in the [[infrared|near infrared]] spectrum.{{cite journal |doi=10.1016/j.tca.2006.04.002 |title=Determination of the temperature and enthalpy of the solid–solid phase transition of caesium nitrate by differential scanning calorimetry |date=2006 |last1=Charrier |first1=E. |first2=E. L. |first3=P. G. |first4=H. M. |first5=B. |first6=T. T. |journal=Thermochimica Acta |volume=445 |pages=36–39 |last2=Charsley |last3=Laye |last4=Markham |last5=Berger |last6=Griffiths}} Caesium compounds may have been used as fuel additives to reduce the [[radar cross-section|radar signature]] of [[exhaust gas|exhaust plumes]] in the [[Lockheed A-12]] [[CIA]] reconnaissance aircraft.{{cite book |isbn=978-1-84176-098-8 |page=47 |title=Lockheed SR-71: the secret missions exposed |last=Crickmore |first=Paul F. |publisher=Osprey |date=2000}} Caesium and rubidium have been added as a [[carbonate]] to glass because they reduce electrical conductivity and improve stability and durability of [[optical fiber|fibre optics]] and [[night vision]] devices. Caesium fluoride or caesium aluminium fluoride are used in fluxes formulated for brazing [[aluminium]] alloys that contain [[magnesium]]. [152] => [153] => [[MHD generator|Magnetohydrodynamic (MHD) power]]-generating systems were researched, but failed to gain widespread acceptance.{{cite book |author=National Research Council (U.S.) |publisher=National Academy Press |date=2001 |title=Energy research at DOE—Was it worth it? |access-date=26 September 2010 |url=http://books.nap.edu/openbook.php?isbn=0309074487&page=52 |isbn=978-0-309-07448-3 |pages=190–194 |doi=10.17226/10165}} Caesium metal has also been considered as the working fluid in high-temperature [[Rankine cycle]] turboelectric generators.{{cite book |title=Economics of Caesium and Rubidium (Reports on Metals & Minerals) |publisher=Roskill Information Services |date=1984 |place=London, United Kingdom |author=Roskill Information Services |page=51 |isbn=978-0-86214-250-6}} [154] => [155] => Caesium salts have been evaluated as antishock reagents following the administration of [[arsenic toxicity|arsenical drugs]]. Because of their effect on heart rhythms, however, they are less likely to be used than potassium or rubidium salts. They have also been used to treat [[epilepsy]]. [156] => [157] => Caesium-133 can be [[laser cooling|laser cooled]] and used to probe fundamental and [[Quantum technology|technological]] problems in [[quantum mechanics|quantum physics]]. It has a particularly convenient [[Feshbach resonance|Feshbach]] spectrum to enable studies of [[ultracold atom]]s requiring tunable interactions.{{cite journal |last1=Chin |first1=Cheng |last2=Grimm |first2=Rudolf |last3=Julienne |first3=Paul |last4=Tiesinga |first4=Eite |date=29 April 2010 |title=Feshbach resonances in ultracold gases |journal=Reviews of Modern Physics |volume=82 |issue=2 |pages=1225–1286 |doi=10.1103/RevModPhys.82.1225 |arxiv=0812.1496 |bibcode=2010RvMP...82.1225C |s2cid=118340314}} [158] => [159] => ==Health and safety hazards== [160] => {{Chembox [161] => |container_only = yes [162] => |Section7 = {{Chembox Hazards [163] => | ExternalSDS = [164] => | GHSPictograms = {{GHS flame}} {{GHS corrosion}} [165] => | GHSSignalWord = Danger [166] => | HPhrases = {{H-phrases|H260|H314}} [167] => | PPhrases = {{P-phrases|P223|P231 + P232|P280|P305 + P351 + P338|P370 + P378 |P422}} [168] => | GHS_ref = {{cite web |url=https://www.sigmaaldrich.com/catalog/product/aldrich/239240?lang=en®ion=US |title=Cesium 239240 |publisher=Sigma-Aldrich |date=26 September 2021 |access-date=21 December 2021}} [169] => | NFPA-H = 3 [170] => | NFPA-F = 4 [171] => | NFPA-R = 3 [172] => | NFPA-S = w [173] => | NFPA_ref = [174] => }} [175] => }} [176] => [[File:AirDoseChernobylVector.svg|thumb|right|upright=1.4|alt=Graph of percentage of the radioactive output by each nuclide that form after a nuclear fallout vs. logarithm of time after the incident. In curves of various colours, the predominant source of radiation are depicted in order: Te-132/I-132 for the first five or so days; I-131 for the next five; Ba-140/La-140 briefly; Zr-95/Nb-95 from day 10 until about day 200; and finally Cs-137. Other nuclides producing radioactivity, but not peaking as a major component are Ru, peaking at about 50 days, and Cs-134 at around 600 days.|The portion of the total radiation dose (in air) contributed by each isotope plotted against time after the [[Chernobyl disaster]]. Caesium-137 became the primary source of radiation about 200 days after the accident.Data from [https://archive.org/details/TheRadiochemicalManual The Radiochemical Manual] and Wilson, B. J. (1966) ''The Radiochemical Manual'' (2nd ed.).]] [177] => Nonradioactive caesium compounds are only mildly toxic, and nonradioactive caesium is not a significant environmental hazard. Because biochemical processes can confuse and substitute caesium with [[potassium]], excess caesium can lead to [[hypokalemia]], [[heart arrhythmia|arrhythmia]], and acute [[cardiac arrest]], but such amounts would not ordinarily be encountered in natural sources.{{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}}{{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|bibcode=1981JESHA..16..549P }} [178] => [179] => The [[median lethal dose]] (LD50) for [[caesium chloride]] in mice is 2.3 g per kilogram, which is comparable to the LD50 values of [[potassium chloride]] and [[sodium chloride]].{{cite journal |doi=10.1016/0041-008X(75)90216-1 |title=Acute toxicity of cesium and rubidium compounds |date=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}} The principal use of nonradioactive caesium is as caesium formate in petroleum [[drilling fluid]]s because it is much less toxic than alternatives, though it is more costly. [180] => [181] => Caesium metal is one of the most reactive elements and is highly [[explosive material|explosive]] in the presence of water. The hydrogen gas produced by the reaction is heated by the thermal energy released at the same time, causing ignition and a violent explosion. This can occur with other alkali metals, but caesium is so potent that this explosive reaction can be triggered even by cold water. [182] => [183] => It is highly [[pyrophoricity|pyrophoric]]: the [[autoignition temperature]] of caesium is {{convert|−116|C}}, and it ignites explosively in air to form [[caesium hydroxide]] and various oxides. Caesium hydroxide is a very strong [[base (chemistry)|base]], and will rapidly corrode glass.{{cite web |url=http://www.rsc.org/periodic-table/element/55/caesium |access-date=27 September 2010 |publisher=Royal Society of Chemistry |title=Chemical Data – Caesium – Cs}} [184] => [185] => The [[isotope]]s [[Caesium-134|134]] and 137 are present in the [[biosphere]] in small amounts from human activities, differing by location. Radiocaesium does not accumulate in the body as readily as other fission products (such as radioiodine and radiostrontium). About 10% of absorbed radiocaesium washes out of the body relatively quickly in sweat and urine. The remaining 90% has a biological half-life between 50 and 150 days.{{cite journal |journal=British Journal of Radiology |title=A Survey of the Metabolism of Caesium in Man |date=1964 |last1=Rundo |issue=434 |pages=108–114 |doi=10.1259/0007-1285-37-434-108 |pmid=14120787 |first1=J. |volume=37}} Radiocaesium follows potassium and tends to accumulate in plant tissues, including fruits and vegetables.{{cite journal |doi=10.1007/BF01376226 |title=Accumulation of Cs and K and growth of bean plants in nutrient solution and soils |date=1962 |last1=Nishita |first1=H. |last2=Dixon |first2=D. |last3=Larson |first3=K. H. |journal=Plant and Soil |volume=17 |pages=221–242 |issue=2 |bibcode=1962PlSoi..17..221N |s2cid=10293954}}{{cite journal |doi=10.1016/0265-931X(96)89276-9 |title=Fate of caesium in the environment: Distribution between the abiotic and biotic components of aquatic and terrestrial ecosystems |date=1996 |last1=Avery |first1=S. |journal=Journal of Environmental Radioactivity |volume=30 |pages=139–171 |issue=2}}{{cite journal |doi=10.1039/AN9921700487 |title=Availability of caesium isotopes in vegetation estimated from incubation and extraction experiments |journal=Analyst |date=1992 |volume=117 |pages=487–491 |first1=Brit |last1=Salbu |first2=Georg |last2=Østby |first3=Torstein H. |last3=Garmo |first4=Knut |last4=Hove |pmid=1580386 |issue=3 |bibcode=1992Ana...117..487S}} Plants vary widely in the absorption of caesium, sometimes displaying great resistance to it. It is also well-documented that mushrooms from contaminated forests accumulate radiocaesium (caesium-137) in the fungal [[sporocarp (fungi)|sporocarps]].{{cite journal |url=https://www.researchgate.net/publication/42541094 |doi=10.1016/j.scitotenv.2010.02.024 |pmid=20334900 |title=Accumulation of potassium, rubidium and caesium (133Cs and 137Cs) in various fractions of soil and fungi in a Swedish forest |journal=Science of the Total Environment |volume=408 |issue=12 |year=2010 |pages=2543–2548 |last1=Vinichuk |first1=M. |bibcode=2010ScTEn.408.2543V}} Accumulation of caesium-137 in lakes has been a great concern after the [[Chernobyl disaster]].{{cite book |first1=Jim T. |last1=Smith |first2=Nicholas A. |last2=Beresford |title=Chernobyl: Catastrophe and Consequences |date=2005 |publisher=Springer |place=Berlin |isbn=978-3-540-23866-9}}{{cite journal |doi=10.1007/BF02197418 |title=Radioactive isotopes of caesium in the waters and near-water atmospheric layer of the Black Sea |first1=V. N. |last1=Eremeev |first2=T. V. |last2=Chudinovskikh |first3=G. F. |last3=Batrakov |first4=T. M. |last4=Ivanova |volume=2 |issue=1 |date=1991 |journal=Physical Oceanography |pages=57–64 |s2cid=127482742}} Experiments with dogs showed that a single dose of 3.8 [[curie (unit)|millicuries]] (140 [[Becquerel|MBq]], 4.1 μg of caesium-137) per kilogram is lethal within three weeks;{{cite journal |title=Toxicity of 137-CsCl in the Beagle. Early Biological Effects |first1=H. C. |last1=Redman |first2=R. O. |last2=McClellan |first3=R. K. |last3=Jones |first4=B. B. |last4=Boecker |first5=T. L. |last5=Chiffelle |first6=J. A. |last6=Pickrell |first7=E. W. |last7=Rypka |volume=50 |issue=3 |date=1972 |journal=Radiation Research |pages=629–648 |doi=10.2307/3573559 |pmid=5030090 |jstor=3573559 |bibcode=1972RadR...50..629R}} smaller amounts may cause infertility and cancer.{{cite news |url=http://news.bbc.co.uk/2/hi/asia-pacific/7967285.stm |title=Chinese 'find' radioactive ball |publisher=BBC News |date=27 March 2009 |access-date=25 January 2010}} The [[International Atomic Energy Agency]] and other sources have warned that radioactive materials, such as caesium-137, could be used in radiological dispersion devices, or "[[dirty bomb]]s".{{cite news |last=Charbonneau |first=Louis |title=IAEA director warns of 'dirty bomb' risk |newspaper=The Washington Post |page=A15 |url=http://www.highbeam.com/doc/1P2-250680.html |agency=Reuters |date=12 March 2003 |access-date=28 April 2010 |archive-url=https://web.archive.org/web/20081205004052/http://www.highbeam.com/doc/1P2-250680.html |archive-date=5 December 2008 |url-status=dead}} [186] => [187] => ==See also== [188] => * {{Section link|Caesium-137#Incidents and accidents}} [189] => * [[Acerinox accident]], a caesium-137 contamination accident in 1998 [190] => * [[Goiânia accident]], a major radioactive contamination incident in 1987 involving caesium-137 [191] => * [[Kramatorsk radiological accident]], a 137Cs lost-source incident between 1980 and 1989 [192] => [193] => ==Notes== [194] => {{Reflist |group="note" |30em}} [195] => [196] => ==References== [197] => {{Reflist |30em}} [198] => [199] => ==External links== [200] => {{Spoken Wikipedia|Caesium.ogg|date=2010-11-29}} [201] => * [https://www.periodicvideos.com/videos/055.htm Caesium or Cesium] at ''[[The Periodic Table of Videos]]'' (University of Nottingham) [202] => * [https://web.archive.org/web/20171104215850/http://richannel.org/the-modern-alchemist-reacting-fluorine-with-caesium View the reaction of Caesium (most reactive metal in the periodic table) with Fluorine (most reactive non-metal)] courtesy of The Royal Institution. [203] => * {{cite journal |title=Molecular CsF5and CsF2+ |journal=Angewandte Chemie |volume=127 |issue=28 |pages=8393–8396 |doi=10.1002/ange.201500402 |year=2015 |last1=Rogachev |first1=Andrey Yu. |last2=Miao |first2=Mao-Sheng |last3=Merino |first3=Gabriel |last4=Hoffmann |first4=Roald |bibcode=2015AngCh.127.8393R}} [204] => [205] => {{Subject bar [206] => |portal=Chemistry [207] => |book1=Caesium [208] => |book2=Period 6 elements [209] => |book3=Alkali metals [210] => |book4=Chemical elements (sorted alphabetically) [211] => |book5=Chemical elements (sorted by number) [212] => |commons=y [213] => |wikt=y [214] => |wikt-search=caesium [215] => |v=y [216] => |v-search=Caesium atom [217] => |s=y [218] => |s-search=1911 Encyclopædia Britannica/Caesium}} [219] => [220] => {{Periodic table (navbox)}} [221] => {{Caesium compounds}} [222] => {{Authority control}} [223] => [224] => [[Category:Caesium| ]] [225] => [[Category:Alkali metals]] [226] => [[Category:Chemical elements with body-centered cubic structure]] [227] => [[Category:Chemical elements]] [228] => [[Category:Glycine receptor agonists]] [229] => [[Category:Reducing agents]] [230] => [[Category:Articles containing video clips]] [231] => [[Category:Pyrophoric materials]] [] => )
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Caesium

Caesium is a chemical element with the symbol Cs and atomic number 55. It is a soft, silvery-gold alkali metal that is highly reactive and easily oxidizes in air.

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It is a soft, silvery-gold alkali metal that is highly reactive and easily oxidizes in air. Caesium is the most electropositive element and has the lowest ionization energy of all known elements, making it highly reactive and able to form a wide variety of compounds. Discovered in 1860 by Robert Bunsen and Gustav Kirchhoff, caesium was named after the Latin word "caesius," meaning "sky blue," due to the blue lines observed in its spectrum. It is primarily sourced from pollucite minerals and is often found in small amounts in the Earth's crust. Caesium has several important applications, including its use in atomic clocks, where its high levels of stability and accuracy make it crucial for precise timekeeping. It is also utilized in photoelectric cells, catalysis, and research in fundamental physics. Additionally, caesium has been used in the past as a component of cesium formate brines, which can be used as drilling fluids in oil and gas exploration. However, caesium must be handled with caution as it is highly toxic and can cause severe burns. Its extreme reactivity also poses a risk of explosions when in contact with water or certain other substances. Due to its potential hazards, strict safety measures are followed when handling and storing caesium. Overall, caesium is a fascinating element with a unique set of properties. Its low melting point, high electrical conductivity, and ability to form alloys make it an important element for various scientific and industrial applications.

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