Caesium

Caesium (also spelled cesium, IPA: /ˈsiːziəm/) is a chemical element in the periodic table that has 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), francium (27 °C), mercury (-39 °C), and gallium (30 °C). This element is most notably used in atomic clocks.

The variant spelling cesium is sometimes used, especially in North American English, but caesium is the spelling used by the IUPAC, although since 1993 it has recognized cesium as a variant as well.

Notable characteristics

The electromagnetic 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 550 grams in the entire Earth's crust at one time,^[1] 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 (157K).

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" (after FrOH), 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.^{[citation needed]} 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.

Applications

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. ^[2]

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 electron spin energy levels of the ground state of the ¹³³Cs atom.

• ¹³⁴Cs has been used in hydrology as a measure of caesium output by the nuclear power industry. This isotope is used because, while it is less prevalent than either ¹³³Cs or ¹³⁷Cs, ¹³⁴Cs can be produced solely by nuclear reactions. ¹³⁵Cs has also been used in this function.
• Like other elements of group 1, caesium has a great affinity for oxygen and is used as a "getter" in vacuum tubes.
• This metal is also used in photoelectric cells due to its ready emission of electrons.
• Caesium is used as a catalyst in the hydrogenation of certain organic compounds.
• Radioactive isotopes of caesium are used in the medical field to treat certain types of cancer.
• Caesium fluoride is widely used in organic chemistry as a base and as a source of anhydrous fluoride ion.
• Caesium vapor is used in many common magnetometers.
• Because of their high density, caesium chloride solutions are commonly used in molecular biology for density gradient ultracentrifugation, primarily for the isolation of viral particles, subcellular organelles and fractions, and nucleic acids from biological samples.
• Cesium nitrate is used as oxidiser to burn silicon in infrared flares^[3] like the LUU-19 flare^[4], because it emits much of its light in the near infrared spectrum.
• More recently this metal has been used in ion propulsion systems.^{[citation needed]}
• Caesium-137 is an extremely common radioisotope used as a gamma-emitter in industrial applications such as:
  • moisture density gauges
  • leveling gauges
  • thickness gauges
  • well-logging devices (used to measure the thickness of rock-strata)
• also used as an internal standard in spectrophotometry

History

Caesium (Latin caesius meaning "sky blue" or "light blue") 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.

Occurrence

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.

See also Caesium minerals.

Isotopes

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 issue from the fission of an uranium core.

Beginning in 1945 with the commencement of nuclear testing, ¹³⁷Cs was released into the atmosphere where it is absorbed readily into solution and is returned to the surface of the earth as a component of radioactive fallout. Once ¹³⁷Cs enters the ground water, it is deposited on soil surfaces and removed from the landscape primarily by particle transport. As a result, the input function of these isotopes can be estimated as a function of time. Caesium-137 has a half-life of 30.17 years. It decomposes to barium-137m (a short-lived product of decay) then to a form of nonradioactive barium.

Precautions

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 (even 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.

See also

• Cs-137
• Goiânia accident - a major radioactive contamination incident involving a small rod of caesium chloride.
• Caesium compounds

References

Wikipedia information about Caesium This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Caesium". It may not have been reviewed by professional editors (see full disclaimer). Help contribute to Americola and edit this article.