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Caesium or cesium (pronounced /ˈsiːziəm/) is the chemical element with the symbol Cs and atomic number 55. It is a soft silvery-gold alkali metal with a melting point of 28 °C (83 °F), which makes it one of the metals that are liquid at or near room temperature, along with rubidium (39 °C [102 °F]), francium (27 °C [81 °F]), mercury (−39 °C [−38 °F]), and gallium (30 °C [86 °F]). This element is most notably used in atomic clocks.
The variant spelling cesium is used especially in North American English, and the IUPAC has recognized it as a variant spelling since 1993, but caesium is the spelling used by the IUPAC.
Additional recommended knowledge
The emission spectrum of caesium has two bright lines in the blue part of the spectrum along with several other lines in the red, yellow, and green. This metal is silvery gold in color and is both soft and ductile. Caesium is the second most electropositive and alkaline of the chemical elements and has the second lowest ionization potential (after francium). Caesium is the least abundant of the five non-radioactive alkali metals. (Technically, francium is the least common alkali metal, but since it is highly radioactive with an estimated 30 grams in the entire Earth's crust at one time, its abundance can be considered zero in practical terms.)
Along with gallium, francium, and mercury, caesium is among the only metals that are liquid at or near room temperature. Caesium reacts explosively in cold water and also reacts with ice at temperatures above −116 °C (−177 °F, 157 K).
Caesium hydroxide (CsOH) is a very strong base and will rapidly etch the surface of glass. CsOH is often stated to be the "strongest base", but in fact many compounds such as n-butyllithium and sodium amide are stronger.
There is an account that caesium, reacting with fluorine, takes up more fluorine than it stoichiometrically should. It is possible that, after the salt Cs+F− has formed, the Cs+ ion, which has the same electronic structure as elemental xenon, can, like xenon, be oxidised further by fluorine and form traces of a higher fluoride such as CsF3, analogous to XeF2.
Probably the most widespread use of caesium today is in caesium formate-based drilling fluids for the oil industry. The high density of the caesium formate brine (up to 2.3 sg), coupled with the relative benignity of 133Cs , reduces the requirement for toxic high-density suspended solids in the drilling fluid, which is a significant technological, engineering and environmental advantage.  
Caesium is also notably used in atomic clocks, which are accurate to seconds in many thousands of years. Since 1967, the International System of Measurements bases its unit of time, the second, on the properties of caesium. SI defines the second as 9,192,631,770 cycles of the radiation which corresponds to the transition between two hyperfine energy levels of the ground state of the 133Cs atom.
Caesium (Latin caesius meaning "blueish grey" ) was spectroscopically discovered by Robert Bunsen and Gustav Kirchhoff in 1860 in mineral water from Dürkheim, Germany. Its identification was based upon the bright blue lines in its spectrum and it was the first element discovered by spectrum analysis. The first caesium metal was produced in 1882 by Carl Setterberg. Historically, the most important use for caesium has been in research and development, primarily in chemical and electrical applications.
An alkali metal, caesium occurs in lepidolite, pollucite (hydrated silicate of aluminium and caesium) and within other sources. One of the world's most significant and rich sources of this metal is at Bernic Lake in Manitoba. The deposits there are estimated to contain 300,000 metric tons of pollucite at an average of 20% caesium.
It can be isolated by electrolysis of fused caesium cyanide and in a number of other ways. Exceptionally pure and gas-free caesium can be made by the thermal decomposition of caesium azide. The primary compounds of caesium are caesium chloride and its nitrate. The price of caesium metal in 1997 was about US$30 per gram, but its compounds are much cheaper.
Caesium has at least 39 known isotopes, which is more than any other element except francium. The atomic masses of these isotopes range from 112 to 151. Even though this element has a large number of isotopes, it has only one naturally occurring stable isotope, 133Cs. Most of the other isotopes have half-lives from a few days to fractions of a second. The radiogenic isotope 137Cs has been used in hydrologic studies, analogous to the use of 3H. 137Cs is produced from the detonation of nuclear weapons and is produced in nuclear power plants, and was released to the atmosphere most notably from the 1986 Chernobyl meltdown. It's because this isotope (137Cs) is one of the numerous products of fission, directly issued from the fission of uranium.
All alkali metals are highly reactive. Caesium, being one of the heavier alkali metals, is also one of the most reactive and is highly explosive when it comes in contact with water, as the hydrogen gas produced by the reaction is heated by the thermal energy released at the same time, causing ignition, and a violent explosion (the same as all alkali metals) - but caesium is so reactive, this explosive reaction can even be triggered by cold water or ice. Caesium hydroxide is an extremely strong base, and can etch glass.
Caesium compounds are encountered rarely by most persons. All caesium compounds should be regarded as mildly toxic because of its chemical similarity to potassium. Large amounts cause hyperirritability and spasms, but such amounts would not ordinarily be encountered in natural sources, so Cs is not a major chemical environmental pollutant. Rats fed caesium in place of potassium in their diet die, so this element cannot replace potassium in function.
The isotopes 134Cs and 137Cs (present in the biosphere in small amounts as a result of radiation leaks) represent a radioactivity burden which varies depending on location. Radiocaesium does not accumulate in the body as effectively as many other fission products (such as radioiodine and radiostrontium), which are actively accumulated by the body.
Wikisource has the text of the 1911 Encyclopædia Britannica article Caesium.
|This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Caesium". A list of authors is available in Wikipedia.|