Caesium

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xenon ← caesium → barium
Rb
↑
Cs
↓
Fr
Appearance
silvery gold
Csmetal.jpg.jpg
General properties
Name, symbol, number caesium, Cs, 55
Element category alkali metal
Group, period, block 1, 6, s
Standard atomic weight 132.9054519(2) g·mol−1
Electron configuration [Xe] 6s1
Electrons per shell 2, 8, 18, 18, 8, 1 (Image)
Physical properties
Phase solid
Density (near r.t.) 1.93 g·cm−3
Liquid density at m.p. 1.843 g·cm−3
Melting point 301.59 K, 28.44 °C, 83.19 °F
Boiling point 944 K, 671 °C, 1240 °F
Critical point 1938 K, 9.4 MPa
Heat of fusion 2.09 kJ·mol−1
Heat of vaporization 63.9 kJ·mol−1
Specific heat capacity (25 °C) 32.210 J·mol−1·K−1
Vapor pressure
P/Pa 1 10 100 1 k 10 k 100 k
at T/K 418 469 534 623 750 940
Atomic properties
Oxidation states 1
(strongly basic oxide)
Electronegativity 0.79 (Pauling scale)
Ionization energies 1st: 375.7 kJ·mol−1
2nd: 2234.3 kJ·mol−1
3rd: 3400 kJ·mol−1
Atomic radius 265 pm
Covalent radius 244±11 pm
Van der Waals radius 343 pm
Miscellanea
Crystal structure body-centered cubic
Magnetic ordering paramagnetic[1]
Electrical resistivity (20 °C) 205 Ω·m
Thermal conductivity (300 K) 35.9 W·m−1·K−1
Thermal expansion (25 °C) 97 µm·m−1·K−1
Young's modulus 1.7 GPa
Bulk modulus 1.6 GPa
Mohs hardness 0.2
Brinell hardness 0.14 MPa
CAS registry number 7440-46-2
Most stable isotopes
Main article: Isotopes of caesium
iso NA half-life DM DE (MeV) DP
133Cs 100% 133Cs is stable with 78 neutrons
134Cs syn 2.0648 y ε 1.229 134Xe
β− 2.059 134Ba
135Cs trace 2.3×106 y β− 0.269 135Ba
137Cs trace 30.07 y β− 1.176 137Ba
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Caesium or cesium (pronounced /ˈsiːziəm/, SEE-zee-əm) 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 only five metals that are liquid at or near room temperature.[2] Caesium has physical and chemical properties similar to those of rubidium and potassium. The metal is extremely reactive and pyrophoric, triggering explosions even at −116 °C upon contact with water. It is the least electronegative element that has stable isotopes, of which it has only one: 133Cs. This is mined mostly from pollucite, while the radioisotopes, especially 137Cs, are extracted from waste produced by nuclear reactors.

The two German chemists Robert Bunsen and Gustav Kirchhoff discovered it in 1860 by the newly developed method of flame spectroscopy. The first small scale applications for caesium have been as "getter" in vacuum tubes and in photoelectric cells. In 1967, a frequency of 133Cs was used to define the second by the International System of Units. Since then it has been widely used in atomic clocks. Since the 1990s, the largest application of the element has been as caesium formate for drilling fluids. It has a range of applications in the production of electricity, in electronics, and in chemistry. The radioactive isotope 137Cs, with a half-life of about 30 years, is used in medical applications, industrial gauges, and hydrology. While the element has a mild chemical toxicity, the radioisotopes present a high health risk in case of radiation leaks.

Contents

[edit] Characteristics

[edit] Properties

Caesium is a soft, ductile, silvery-white metal which develops a silvery-gold hue in the presence of even trace amounts of oxygen.[3] It has a melting point of 28.4 °C, making it one of the few metals that are liquid near room temperature. Mercury is the only metal with a melting point lower than caesium (The radioactive element francium may also have lower melting point). In its solid form, caesium is a very soft and ductile metal.[4] Caesium compounds burn with a blue color.

Caesium forms alloys with the other alkali metals as well as with gold, and amalgams with mercury. At temperatures below 650 °C, it alloys with cobalt, iron, molybdenum, nickel, platinum, tantalum or tungsten.[3] On the other hand, it is known to form photosensitive intermetallic compounds with antimony, gallium, indium and thorium.[3]

High purity caesium-133 preserved under argon

Isolated caesium is extremely reactive and very pyrophoric. In addition to igniting spontaneously in air, it reacts explosively with water, even more so than the other members of the first group of the periodic table.[5] The reaction with water is explosive even at temperatures as low as −116 °C. Because of its high reactivity, caesium metal is classified as a hazardous material and must be stored and transported in isolation from possible reactants. It is stored and shipped in dry mineral oil or in other dry saturated hydrocarbons, or in an inert atmosphere (such as argon or nitrogen) or vacuum, in sealed borosilicate glass ampoules which are shipped wrapped in foil and packed in an inert cushioning material, such as vermiculite, each in a metal can. In quantities above 100 grams, caesium is shipped in hermetically sealed stainless steel containers.[3]

The chemistry of caesium is very similar to that of other alkaline metals, its chemistry being particularly closely associated to that of rubidium, the element above caesium in the periodic table.[3] Some small differences arise from the fact that Cs is heavier and more electropositive than other (non-radioactive) alkali metals. Caesium is the most electropositive stable chemical element, and of all the known elements, only francium may be more electropositive (as francium is highly radioactive, it cannot be isolated in observable quantities yet). Relativistic effects can lower the reactivity and raise the electronegativity of francium, as suggested by its value of the first ionization energy.

[edit] Compounds

See also: Category:Caesium compounds
Ball-and-stick model of the cubic coordination of Cs and Cl in CsCl

The vast majority of caesium compounds contain the element as the cation Cs+. These compounds are colorless, unless the color is generated by the anion. A few alkalides containing a Cs− anion have been studied.[6]

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 but are not classic hydroxide bases and are destroyed by water.

Cs11O3 cluster

Caesium chloride is an important source of caesium ions in a variety of applications. Noteworthy, it crystallizes in the simple cubic crystal system, which is also called the "caesium chloride" structure. This is composed of a primitive cubic lattice with a two atom basis, where both atoms have eightfold 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 center of the cubes. This structure is shared with CsBr and CsI and many intermetallic compounds. In contrast, most other alkaline halides have the sodium chloride structure. When both ions are similar in size (Cs+ ionic radius 174 pm for this coordination number, Cl− 181 pm) the CsCl structure is formed, while when they are different (Na+ ionic radius 102 pm, Cl− 181 pm), the sodium chloride structure is adopted.[7]

As with the other heavier elements of the alkali metals group, caesium results in numerous compounds with oxygen. When caesium burns in air, the superoxide CsO2 is the main product.[8] The "normal" caesium oxide (Cs2O) forms yellow-orange hexagonal crystals,[9] and is the only oxide of the anti CdCl2 type.[10] Aside from the superoxide and the ozonide CsO3,[11][12] several brightly colored suboxides have also been studied.[13] These include Cs7O, Cs4O, Cs11O3, the dark-green Cs3O, CsO, Cs3O2,[14][15] as well as Cs7O2.[16][17] The latter may be heated under high vacuum to generate Cs2O.[10]

[edit] Isotopes

Main article: Isotopes of caesium

Caesium has at least 39 known isotopes ranging in atomic mass from 112 to 151. Only one of those, 133Cs, is stable. Radioactive 135Cs has a long half-life of about 2.3 million years; 137Cs and 134Cs have half-lives of 30 and 2 years, respectively, while the isotopes with atomic masses of 129, 131, 132 and 136, have half-times between a day and two weeks. Most of the other isotopes have half-lives from a few seconds to fractions of a second. 137Cs decomposes to a short-lived 137mBa and then to non-radioactive barium. There are at least 21 metastable nuclear isomers; other than 134mCs (with a half-life of just under 3 hours), they all have half-lives of a few minutes or less.[18][19]

Decay scheme of caesium-137

135Cs is one of seven long-lived fission products of uranium which form in nuclear reactors. In most reactors, its fission product yield is reduced because its predecessor 135Xe is an extremely potent neutron poison and often transmutes to stable 136Xe before it can decay to 135Cs. 137Cs is one of the two principal medium-lived fission products, along with 90Sr, which are responsible for most of the radioactivity of spent nuclear fuel after several years of cooling, up to several hundred years after use. It constitutes most of the radioactivity still left from the Chernobyl disaster. 137Cs beta decays to 137mBa (a short-lived nuclear isomer) then to non-radioactive 137Ba, and is also a strong emitter of gamma radiation. 137Cs has a very low rate of neutron capture and cannot be feasibly disposed of in this way, but must be allowed to decay. Almost all caesium produced from nuclear fission comes from beta decay of originally more neutron-rich fission products, passing through isotopes of iodine then isotopes of xenon. Because these elements are volatile and can diffuse through nuclear fuel or air, caesium is often created far from the original site of fission.

With the commencement of nuclear weapons testing around 1945, 137Cs was released into the atmosphere where it is not 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 cannot be estimated as a function of time.[20]

[edit] Occurrence

Pollucite, a caesium mineral
See also: Category:Caesium minerals

Caesium is a relatively rare element as it is estimated to average approximately 3 parts per million in the Earth’s crust.[21] This makes it the 45th most abundant of all elements and the 36th of the metals. Nevertheless, it is more abundant than such elements as antimony, cadmium, tin and tungsten, and two orders of magnitude more abundant than mercury or silver, but is 30 times less abundant than rubidium, with which it is so closely associated chemically.[3]

Because of its large ionic radius, caesium is one of the incompatible elements. During magma crystallization, it 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. Caesium does not substitute for potassium as readily as does rubidium; thus the alkali evaporite minerals sylvite (KCl) and carnallite (KMgCl3·6H2O) may contain only 0.002% caesium. Caesium is found in significant quantities in only in a few minerals. Percent amounts of caesium may be found in some beryl (Be3Al2(SiO3)6) and up to 15 wt% Cs2O in the closely related mineral pezzottaite (Cs(Be2Li)Al2Si6O18), up to 8.4 wt% Cs2O in the rare berylloborate mineral londonite ((Cs,K)Al4Be4(B,Be)12O28) and less in the more widespread K-dominant analogue rhodizite; a few percent in avogadrite ((K,Cs)BF4) and a few other mineral species. The only economically important source mineral for caesium is pollucite Cs[AlSi2O6] which is found in a few places around the world in zoned pegmatites, and is 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 creates high grade ore for mining.[22]

One of the world's most significant and rich sources of this metal is the Tanco mine at Bernic Lake in Manitoba. The deposits there are estimated to contain 350,000 metric tons[22] of pollucite ore with an average caesium content of 24 wt%.[20][23] Although the stoichiometric content of caesium in pollucite is 42.6%, pure pollucite samples from the deposit at Bernic Lake, Canada contain only about 34% caesium. Commercial pollucite contains over 19% caesium.[3][24] The Bikita pegmatite deposit in Zimbabwe is mined for its petalite but it also contains significant amount of pollucite.[20] Notable amounts of pollucite are also mined in the Karibib Desert, Namibia, but more than two-thirds of the world’s reserve base is at Bernic Lake, Canada.[23] At the present rate of world mine production, that is between 5,000 and 10,000 kg/yr, reserves would last thousands of years.[3]

[edit] Production

The mining of pollucite ore, as with other zoned pegmatites, is a selective process and is conducted on a small scale in comparison with most metal mining operations. The ore is crushed, hand-sorted, but not usually concentrated, and then ground to prepare it for conversion to caesium metal or compounds.[3] Caesium is then extracted from pollucite mainly by three methods: acid digestion, alkaline decomposition, and direct reduction.[25]

Acid digestion is the principal commercial method used and usually employs hydrochloric (HCl), sulfuric (H2SO4), hydrobromic (HBr), or hydrofluoric (HF) acids (the latter two are not used commercially in the United States owing to processing difficulties). Hydrochloric acid digestion of pollucite is performed at elevated temperatures and yields an impure caesium chloride (CsCl) solution that is converted to double chloride salts, such as caesium antimony chloride (Cs4SbCl7), caesium iodine chloride (Cs2ICl, or caesium hexachlorocerate (Cs2(CeCl6)), which are purified and then decomposed by hydrolysis to yield purified CsCl. Digestion of pollucite in hot sulfuric acid (35% to 45%) yields a solution from which caesium alum (CsAl(SO4)2·12H2O) is precipitated. The alum is roasted with 4% carbon and then leached to yield a Cs2SO4 solution, which may then be converted to CsCl.[3]

Alkaline decomposition consists of roasting of pollucite with either a CaCO3-CaCl2 mixture, or a Na2CO3-NaCl mixture. The calcine obtained is leached with water or dilute ammonia (NH4OH) to extract a dilute CsCl solution, followed by conversion of the chloride to caesium alum or carbonate (Cs2CO3. Direct reduction involves heating the ore mineral with calcium, potassium, or sodium metal in a vacuum or an inert atmosphere, thus yielding an impure caesium metal. Nevertheless, because of low yield, impurity of the product, and engineering difficulties, this method is not used commercially.[3]

Most of the mined caesium is directly converted into 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, Canada, with a capacity of 12,000 barrels per year of caesium formate solution.[26] The primary smaller-scale commercial compounds of caesium are caesium chloride and its nitrate.[27]

Alternatively, caesium metal may be obtained from the purified compounds derived from the ore. Caesium chloride, and the other caesium halides as well, can be reduced at 700 to 800 °C with calcium or barium, followed by distillation of the caesium metal. In the same way, the aluminate, carbonate, or hydroxide may be reduced by magnesium.[3] The metal can also be isolated by electrolysis of fused caesium cyanide (CsCN). Exceptionally pure and gas-free caesium can be made by the thermal decomposition at 390 °C of caesium azide CsN3, which is produced from aqueous caesium sulfate and barium azide.[25] In vacuum applications, caesium dichromate can be reacted with zirconium forming pure caesium without other gaseous products.[27]

Cs2Cr2O7 + 2 Zr → 2 Cs + 2 ZrO2+ Cr2O3

The price of 99.8% pure caesium (metal basis) in 2009 was about US$10 per gram, but its compounds are significantly cheaper.[23]

[edit] History

Black-and-white image of two middle-aged men, either one leaning with one elbow on a wooden column in the middle. Both wear long jackets, and the shorter man on the left has a beard.
Gustav Kirchhoff (left) and Robert Bunsen (right)

Caesium (Latin caesius meaning "bluish gray")[28][29] was spectroscopically discovered by Robert Bunsen and Gustav Kirchhoff in 1860 in mineral water from Dürkheim, Germany, as described below.[30]

The residue after evaporation of 44,000 liters of mineral water yielded 240 kg of concentrated salt solution. The alkaline earth metals were precipitated either as sulfates or oxalates, leaving only the alkali metal in the solution. After conversion to the nitrates 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. These salts show a slight difference in solubility in hot water, and therefore the less-soluble caesium and rubidium hexachloroplatinate (Rb2PtCl6) could be obtained by fractional crystallization. After reduction of the hexachloroplatinate with hydrogen, caesium and rubidium could be separated by the difference in solubility of the carbonates in alcohol. The process yielded 9.2 g of rubidium chloride and 7.3 g of caesium chloride from the 44,000 liters of mineral water.[31]

Identification of caesium was based upon the bright blue lines in its emission spectrum, and it was the first element discovered by spectrum analysis.[31] The German chemist Carl Setterberg first produced caesium metal in 1882 by electrolysis of caesium chloride.[32] Setterberg received his PhD from Kekule and Bunsen for this work.[30]

NIST-F1 caesium fountain atomic clock, serving as the US time and frequency standard, with an uncertainty of 5.10×10−16 (as of 2005).

Historically, the most important use for caesium has been in research and development, primarily in chemical and electrical fields. It found no significant application until it was added into radio vacuum tubes in the 1920s as a getter, a scavenger of the trace amounts of oxygen remaining in the tube after manufacture, and as a coating on the heated cathode to increase the amount of electric current that could flow through the tube. Caesium became recognized as a functional, high-performance industrial metal in electronics in the 1950s.[33] Applications of non-radioactive caesium included photoelectric cells, photomultiplier tubes, optical components (Cs salts) of infrared spectrophotometers, catalysts for several organic reactions, crystals for scintillation counters, and in magnetohydrodynamic power generators.[3]

A second was defined as: the duration of 9,192,631,770 cycles of microwave light absorbed or emitted by the hyperfine transition of cesium-133 atoms in their ground state undisturbed by external fields
—13th General Conference on Weights and Measures, 1967

Since 1967, the International System of Measurements has based its unit of time, the second, on the properties of caesium. The International System of Units (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.[34] Since 1993, IUPAC accepts the alternative spelling cesium, but recommends the spelling caesium.[35]

[edit] Applications

[edit] Petroleum exploration

The largest end-use of nonradioactive caesium today is in caesium formate based drilling fluids for the oil industry.[20] Aqueous solutions of caesium formate (HCOO-Cs)+, which is made by reacting caesium hydroxide with formic acid) were developed in the mid-1990s for use as oil well drilling and completion fluids and are especially suitable for use when downhole temperatures and pressures are high. The function of caesium formate as 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. As a completion fluid (which refers to the emplacement of control hardware after drilling but prior to production), maintaining the pressure is the most important.[3]

The high density of the caesium formate brine (up to 2.3 g/cm³, or 19.2 pounds per gallon), coupled with the relatively benign nature of most caesium compounds, reduces the requirement for toxic high-density suspended solids in the drilling fluid, which is a significant technological, engineering and environmental advantage. Note that contrary to components of many other heavy liquids, caesium has minimal radioactivity because it is almost entirely composed of a stable isotope and is relatively environment friendly.[36] The caesium formate brine can be blended with potassium and sodium formates to decrease the density of the fluids down to that of water (1.0 g/cm³. Furthermore, it is biodegradable and reclaimable, and may be recycled, which is important in view of its high cost (about $4,000 per barrel in 2001)[37]. The alkali formates are safe to handle and do not damage the producing formation or downhole metals as its corrosive alternative high-density brines (such as zinc bromide (ZNBr2 solutions) sometimes do, and they require less cleanup and disposal costs.[20]

[edit] Atomic clocks

Atomic clock ensemble at the U.S. Naval Observatory may be accessed by telephone (202-762-1401) or via internet NTP servers[38]

Caesium is also used in atomic clocks, which use the resonant vibration frequency of 133Cs atoms as a reference point. Caesium clocks, which have been improved repeatedly over the past half century, form the basis for world’s timekeeping system. Precise caesium clocks today measure frequency with an accuracy of from 2 to 3 parts in 1014, which would correspond to a time measurement accuracy of 2 nanoseconds per day, or one second in 1,400,000 years. The latest versions in the United States and France are accurate to 1.7 parts in 1015, or 1 second in 17 million years,[20] which has been regarded as "the most accurate realization of a unit that mankind has yet achieved."[34]

Due to their extreme precision, they are used at the United States Naval Observatory Time Center in Washington, D.C., and in the aircraft, satellites, and ground systems that track the space shuttle.[39] Caesium clocks are also used in networks that control the timing of cell phone transmissions, and caesium devices help control and regulate information flow on the Internet.[40]

[edit] Electric power and electronics

Magnetohydrodyamic (MHD) power-generating systems were researched but failed the gain acceptance for widespread use, and funding from the U.S. Department of Energy was stopped in the early 1990s.[41] Cesium metal has also been considered as the working fluid in high-temperature Rankine cycleturboelectric generators.[42] Caesium has been used in caesium vapor thermionic generators, which are low-power devices that convert heat energy to electrical energy. In the two-electrode vacuum tube converter, it neutralizes the space charge that builds up near the cathode and in so doing enhances current flow.[43]

Caesium is also important for its photoemissive properties by which in which light energy is converted to electron flow. It is used in photoelectric cells because caesium-based cathodes such as intermetallic compound K2CsSb, have low threshold voltage for emission of electrons.[44] The range of photoemissive devices using caesium include optical character recognition devices, photomultiplier tubes, and television camera image tubes.[45][20][46] Nevertheless, germanium, rubidium, selenium, silicon, tellurium, and several other elements and compounds can substitute for caesium as photosensitive materials.[20]

Caesium iodide (CsI) and bromide (CsBr) crystals are used in scintillation counters which are widely used in mineral exploration and particle physics research. They are well suited for the detection of gamma and x-ray radiation.[20] Caesium vapor is used in many common magnetometers.[47] Caesium is also used as an internal standard in spectrophotometry.[48] Like other alkali metals, caesium has a great affinity for oxygen and is used as a "getter" in vacuum tubes.[49] Other uses of the metal include high-energy lasers, vapor glow lamps, and vapor rectifiers.[20]

[edit] Chemical and medical

A sample of caesium floride

Chemical applications are also another important use of caesium.[50] Liquid caesium can be used as a catalyst in the hydrogenation of certain organic compounds.[51] Doping with caesium compounds is used to enhance the effectiveness of several metal-ion catalysts used in the production of chemicals, such as acrylic acid, anthraquinone, ethylene oxide, methanol, phthalic anhydride, styrene, methyl methacrylate monomers, and various olefins. It is also used in the catalytic conversion of sulfur dioxide into sulfur trioxide in the production of sulfuric acid. Caesium metal absorbs gases and other impurities in ferrous and nonferrous metallurgy and in the purification of carbon dioxide. Molten CsOH has been used in the desulfurizing of heavy crude oil.[20]

Caesium fluoride is widely used in organic chemistry as a base,[52] or as a source of anhydrous fluoride ion.[53] Caesium salts sometimes are used to replace potassium or sodium salts in many organic syntheses, such as cyclization, esterification, and polymerization.[20] Because of their high density, caesium chloride (CsCl), sulfate (Cs2SO4), and trifluoroacetate (Cs(O2CCF3)) solutions are commonly used in molecular biology for density gradient ultracentrifugation, primarily for the isolation of viralparticles, subcellular organelles and fractions, and nucleic acids from biological samples.[54] Caesium salts have been evaluated as antishock reagents to be used following the administration of 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.[20]

[edit] Nuclear applications

137Cs is an very common radioisotope used as a gamma-emitter in industrial applications. Its advantages include a half live of roughly 30 years, its availability from the nuclear fuel cycle, and having 137Ba as stable end product. The high water solubility is a disadvantage making caesium-137 incompatible with irradiation of food and medical supplies.[55] It has been used in agriculture, cancer treatment, and sterilization of food, sewage sludge, and surgical equipment.[20][56] Radioactive isotopes of caesium in radiation devices were used in the medical field to treat certain types of cancer, but emergence of better alternatives and the use of water-soluble caesium chloride in the sources, which would create wide range contamination, gradually put these Cs sources out of use.[57][58] 137Cs has been employed in a variety of industrial measurement gauges, including moisture, density, leveling, and thickness gauges.[59][59] 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.[60]

137Cs has also been used in hydrologic studies, analogous to the use of tritium. It is produced from detonation of nuclear weapons and emissions from nuclear power plants. With the commencement of nuclear testing around 1945, continuing through the mid-1980s, 137Cs was released into the atmosphere where it is absorbed readily into solution. Known year-to-year variation within that period allows correlation with soil and sediment layers. 134Cs, and to a lesser extent 134Cs and 135Cs, have also been used in hydrology as a measure of cesium output by the nuclear power industry. These isotope are used because, while they are less prevalent than either 133Cs or 137Cs, they can be produced solely by anthropogenic sources.[61]

[edit] Other uses

Schematics of an electrostatic ion thruster which were initially developed for use with caesium or mercury

Caesium and mercury were used as a propellant in early ion engines for spacecraft propulsion on very long interplanetary or extraplanetary missions. It used a method of ionization to strip the outer electron from the propellant by simple contact with tungsten. Concerns about the corrosive action of caesium on spacecraft components, have pushed development in the direction of use of inert gas propellants, such as xenon, which is easier to handle in ground-based tests and has less potential to interfere with the spacecraft. Eventually, xenon was used in the experimental spacecraft Deep Space 1 launched in 1998.[62][63] Nevertheless, field emission electric propulsion thrusters which use a simple system of accelerating liquid metal ions such as of caesium to create thrust have been built.[64]

Caesium nitrate is used as an oxidizer and pyrotechnic colorant to burn silicon in infraredflares[65] such as the LUU-19 flare,[66] because it emits much of its light in the near infrared spectrum.[67] Caesium has been used to reduce the radar signature of exhaust plumes in the SR-71 Blackbird military aircraft.[68] Caesium, along with rubidium, has been added as carbonates to glass because it reduces electrical conductivity and improves stability and durability. Such special glasses are used in fiber optics and night vision devices. Caesium fluoride or caesium aluminium fluoride are used in fluxes formulated for the brazing of aluminium alloys that contain magnesium.[20]

[edit] Precautions

The portion of the total radiation dose (in air) contributed by each isotope versus time after the Chernobyl disaster.[69]

Caesium is one of the most reactive elements and is highly explosive when it comes in contact with 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 even be triggered by cold water or ice at temperatures down to −116 °C.[3] Caesium metal is highly pyrophoric, and ignites spontaneously in air to form caesium hydroxide and various oxides. Caesium hydroxide is an extremely strong base, and can rapidly corrode glass.

Caesium compounds are rarely encountered by most persons; most caesium compounds are mildly toxic because of chemical similarity of caesium to potassium. Exposure to large amounts of Cs compounds can cause hyperirritability and spasms, but as such amounts would not ordinarily be encountered in natural sources, Cs is not a major chemical environmental pollutant.[70] The median lethal dose (LD50) value for caesium chloride in mice is 2.3 g/kg which is comparable to the LD50 values of potassium chloride and sodium chloride.[71]

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). As with other alkali metals, radiocaesium washes out of the body relatively quickly with the sweat and urine. However, radiocaesium follows potassium and tends to accumulate in plant tissues, including fruits and vegetables.[72][73] Accumulation of 137Cs in lakes has been a high concern after the Chernobyl disaster.[74] Even in small amounts, 137Cs may cause infertility, cancer and even death.[75] The International Atomic Energy Agency and other sources have indicated that radioactive materials, such as 137Cs, may be used in radiological dispersion devices, or “dirty bombs”.[76]

[edit] See also

Search Wikisource Wikisource has the text of the 1911 Encyclopædia Britannica article Caesium.

[edit] References

  1. ^ Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics 81st edition, CRC press.
  2. ^ Along with rubidium (39°C [102°F]), francium (27°C [81°F]), mercury (−39°C [−38°F]), and gallium (30°C [86°F]). Bromine is also liquid at room temperature (melting at −7.2 °C, 19 °F) but it is not a metal, but a halogen.
  3. ^ a b c d e f g h i j k l m n o Brooks, William C.. "Mineral Commodity Profile: Cesium" (PDF). United States Geological Survey. http://pubs.usgs.gov/of/2004/1432/2004-1432.pdf. Retrieved 2009-12-27. 
  4. ^ Heiserman, David L. (1992). Exploring Chemical Elements and their Compounds. McGraw-Hill. pp. 201–203. ISBN 0-8306-3015-5. 
  5. ^ See link for a video.
  6. ^ Dye, J. L. (1979). "Compounds of Alkali Metal Anions". Angewandte Chemie International Edition 18 (8): 587–598. doi:10.1002/anie.197905871. 
  7. ^ Wells, A.F. (1984). Structural Inorganic Chemistry (5 ed.). Oxford Science Publications. ISBN 0-19-855370-6. 
  8. ^ Cotton, F. Albert; Wilkinson, G. (1962). Advanced Inorganic Chemistry. Jon Wiley & sons, Inc.. p. 318. 
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v • d • e
Alkali metals
   

Lithium
Li
Atomic Number: 3
Atomic Weight: 6.941
Melting Point: 453.69 K
Boiling Point: 1615 K
Specific mass: 0.534 g/cm3
Electronegativity: 0.98

Sodium
Na
Atomic Number: 11
Atomic Weight: 22.990
Melting Point: 370.87 K
Boiling Point: 1156 K
Specific mass: 0.97 g/cm3
Electronegativity: 0.96

Potassium
K
Atomic Number: 19
Atomic Weight: 39.098
Melting Point: 336.58 K
Boiling Point: 1032 K
Specific mass: 0.86 g/cm3
Electronegativity: 0.82

Rubidium
Rb
Atomic Number: 37
Atomic Weight: 85.468
Melting Point: 312.46 K
Boiling Point: 961 K
Specific mass: 1.53 g/cm3
Electronegativity: 0.82

Caesium
Cs
Atomic Number: 55
Atomic Weight: 132.905
Melting Point: 301.59 K
Boiling Point: 944 K
Specific mass: 1.93 g/cm3
Electronegativity: 0.79

Francium
Fr
Atomic Number: 87
Atomic Weight: (223)
Melting Point: 295(?) K
Boiling Point: 950(?) K
Specific mass: ? g/cm3
Electronegativity: 0.7

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