The Full Wiki

Caesium: Wikis


Note: Many of our articles have direct quotes from sources you can cite, within the Wikipedia article! This article doesn't yet, but we're working on it! See more info or our list of citable articles.

Did you know ...

More interesting facts on Caesium

Include this on your site/blog:


From Wikipedia, the free encyclopedia



silvery gold
Some silvery-gold metal, with a liquid-like texture and lustre, sealed in a glass ampoule
General properties
Name, symbol, number caesium, Cs, 55
Pronunciation /ˈsiːziəm/, SEE-zee-əm
Element category alkali metal
Group, period, block 16, 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
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 1.174 137Ba

Caesium or cesium 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.[note 1] Caesium has physical and chemical properties similar to those of rubidium and potassium. The metal is extremely reactive and pyrophoric, reacting with water even at −116 °C. It is the least electronegative element that has stable isotopes, of which it has only one, caesium-133. This is mined mostly from pollucite, while the radioisotopes, especially caesium-137, 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 caesium-133 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 caesium-137, 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 and has been named a hazardous material .




Caesium is a very soft, very ductile, silvery-white metal, which develops a silvery-gold hue in the presence of even trace amounts of oxygen.[2][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 known melting point lower than caesium.[note 2][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. On the other hand, it is known to form intermetallic compounds with antimony, gallium, indium and thorium, which are known to be photosensitive.[2] The alloy of 41% caesium, 47% potassium, and 12% sodium has the lowest melting point of any known metal alloy, at −78 °C.[4] A couple of amalgams have been studied, including the black-metallic, purple shining CsHg2 and the golden-metallic CsHg.[5]


 Y shaped yellowish crystal in glass ampoule, looking like the branch of a pine tree
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 cold), even more so than the other members of the first group of the periodic table.[note 3] The reaction with solid water occurs even at temperatures as low as −116 °C.[4] Because of its high reactivity, the metal is classified as a hazardous material. 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. In quantities of more than about 100 grams, caesium is shipped in hermetically sealed stainless steel containers. When glass ampoules are used, they are shipped wrapped in foil and packed in an inert cushioning material, such as vermiculite, each in a metal can.[2]

The chemistry of caesium is very similar to that of other alkaline metals, and is particularly closely associated to that of rubidium, the element above caesium in the periodic table.[2] Some small differences arise from the fact that it has a higher atomic mass and is more electropositive than other (non-radioactive) alkali metals.[6] Caesium is the most electropositive stable chemical element.[note 4][4]


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.
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+, which binds ionically to a wide variety of anions. These compounds are usually colorless and many are hygroscopic. Compounds like acetate, carbonate, halides, oxide, nitrate, and sulfate are water-soluble, while double halides with antimony, bismuth, cadmium, copper, iron, and lead are insoluble.[2] There are very few examples of caesium cation (Cs+) forming covalent bonds with ligands, like with crown ethers such as 18-crown-6, where it may bind either one or two crown ethers.[7] A few alkalides containing a Cs anion have been studied.[8]

Caesium hydroxide (CsOH) is hygroscopic and a very strong base,[9] and will rapidly etch the surface of semiconductors such as silicon.[10] It has been regarded as the "strongest base", but in reality, many compounds that are not classic hydroxide bases such as n-butyllithium and sodium amide are stronger, and are destroyed by water.

 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 oxygen, and at the center of each octahedron is a small red sphere representing caesium.
Cs11O3 cluster

Caesium chloride is an important source of caesium ions in a variety of applications and it crystallizes in the simple cubic crystal system, which is also called the "caesium chloride structure".[6] This is composed of a primitive cubic lattice with a two-atom basis, each with an 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 prefer the sodium chloride structure.[6] 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.[11]

As with the other heavy elements of the alkali metals group, caesium forms numerous binary compounds with oxygen. When caesium burns in air, the superoxide CsO2 is the main product.[12] The "normal" caesium oxide (Cs2O) forms yellow-orange hexagonal crystals,[13] and is the only oxide of the anti-CdCl2 type.[14] It vaporizes at 250 °C, and decomposes to caesium metal and the peroxide Cs2O2 at temperatures above 400 °C.[15] Aside from the superoxide and the ozonide CsO3,[16][17] several brightly colored suboxides have also been studied.[18] These include Cs7O, Cs4O, Cs11O3, the dark-green Cs3O, CsO, Cs3O2,[19][20] as well as Cs7O2.[21][22] The latter may be heated under high vacuum to generate Cs2O.[14] Binary compounds with sulfur, selenium, and tellurium are also known.[2]


Caesium has at least 39 known isotopes ranging in atomic mass from 112 to 151. Only one of these, 133Cs, is stable. NMR studies can be done with this isotope at a resonating frequency of 11.70 MHz.[7] 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. 137Cs decomposes to a short-lived 137mBa and then to non-radioactive barium. The isotopes with atomic masses of 129, 131, 132 and 136, have half-times between a day and two weeks, while most of the other isotopes have half-lives from a few seconds to fractions of a second. 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.[23][24]

A graph showing the energetics behind the caesium-137 (nuclear spin: I=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=3/2+). Alternatively, caesium-137 may decay directly into barium-137 by a 0.4% probably beta emission.
Decay scheme of caesium-137

135Cs is one of long-lived fission products of uranium which form in nuclear reactors.[25] 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.[26][27] 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.[28] It is currently the largest source of radioactivity generated due to the Chernobyl disaster.[29] 137Cs beta decays to 137mBa (a short-lived nuclear isomer) then to non-radioactive 137Ba, and is also a strong emitter of gamma radiation.[30] 137Cs has a very low rate of neutron capture and cannot be feasibly disposed of in this way, but must be allowed to decay.[31] 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.[32]

With the commencement of nuclear weapons testing around 1945, 137Cs was released into the atmosphere where it is not readily absorbed 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.[2]


A white mineral, from which white and pale pink crystals protrude
Pollucite, a caesium mineral

Caesium is a relatively rare element as it is estimated to average approximately 3 parts per million in the Earth’s crust.[33] 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 30 times less abundant than rubidium—with which it is so closely chemically associated.[2]

Because of its large ionic radius, caesium is one of the incompatible elements.[34] 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. As a result, it is found in only a few minerals in significant quantities. Percent amounts of caesium may be found in beryl (Be3Al2(SiO3)6) and avogadrite ((K,Cs)BF4), up to 15 wt% Cs2O in the closely related mineral pezzottaite (Cs(Be2Li)Al2Si6O18), up to 8.4 wt% Cs2O in the rare mineral londonite ((Cs,K)Al4Be4(B,Be)12O28), and less in the more widespread rhodizite.[2] 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 create high-grade ore for mining.[35]

One of the world's most significant and rich sources of the metal is the Tanco mine at Bernic Lake in Manitoba. The deposits there are estimated to contain 350,000 metric tons of pollucite ore,[35] with an average caesium content of 24 wt%.[36] 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.[37] The Bikita pegmatite deposit in Zimbabwe is mined for its petalite but it also contains significant amount of pollucite. 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.[36] At the present rate of world mine production, that is between 5,000 and 10,000 kg/yr, reserves will last thousands of years.[2]


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.[2] Caesium is then extracted from pollucite mainly by three methods: acid digestion, alkaline decomposition, and direct reduction.[38]

The silicate pollucite is soluble in strong acids and either hydrochloric (HCl), sulfuric (H2SO4), hydrobromic (HBr), or hydrofluoric (HF) acids is used for digestion. In the reaction with hydrochloric acid, a mixture of soluble chlorides is produced. As caesium forms several insoluble double chloride salts it is possible to precipitate it as caesium antimony chloride (Cs4SbCl7), caesium iodine chloride (Cs2ICl), or caesium hexachlorocerate (Cs2(CeCl6)). After separation the pure precipitated double salt is decomposed and the CsCl is obtained by evaporation of the water. The method using sulfuric acid yield the insoluble double salt directly without the need of another compound, it is a caesium alum (CsAl(SO4)2·12H2O). The aluminium sulfate in the alum is converted to the insoluble aluminium oxide by roasting the alum with carbon. The resulting product is than leached with water to yield a Cs2SO4 solution.[2]

The roasting of pollucite with calcium carbonate and calcium chloride yields insoluble calcium silicates and soluble caesium chloride. Leaching with water or dilute ammonia (NH4OH) yields a dilute chloride (CsCl) solution. This solution can be evaporated to produce caesium chloride or transformed into caesium alum or caesium carbonate. Albeit not commercially feasible direct reduction of the ore with potassium, sodium or calcium in vacuum would produce caesium metal.[2]

Most of the mined caesium is directly converted into caesium formate (HCOOCs+) 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.[39] The primary smaller-scale commercial compounds of caesium are caesium chloride and its nitrate.[40]

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.[2] 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.[38] In vacuum applications, caesium dichromate can be reacted with zirconium forming pure caesium without other gaseous products.[40]

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.[36]


 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.
Gustav Kirchhoff (left) and Robert Bunsen (center) discovered caesium spectroscopically

In 1860, Robert Bunsen and Gustav Kirchhoff discovered caesium in mineral water from Dürkheim, Germany. The name derived from the Latin word caesius[41][42] meaning "bluish gray", based on the bright blue lines in its emission spectrum.[43][44] Caesium was the first element to be discovered by spectroscopically, only one year after the invention of the spectroscope by Bunsen and Kirchhoff.[4]

Caesium was only present as a minor component in the mineral water. To obtain a pure sample of caesium 44,000 liters of mineral water had to be evaporated. The residue yielded 240 kilograms 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 ((Cs,Rb)2PtCl6) 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 grams of rubidium chloride and 7.3 grams of caesium chloride from the 44,000 liters of mineral water.[43]

The German chemist Carl Setterberg first produced caesium metal in 1882 by electrolysis of caesium chloride.[45] Setterberg received his PhD from Kekule and Bunsen for this work.[44]

 A laboratory table with some optical devices on it.
FOCS-1, a continuous cold caesium fountain atomic clock in Switzerland, started operating in 2004 at an uncertainty of one second in 30 million years

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. It was then used in radio vacuum tubes in two functions: as a getter to remove excess oxygen after manufacture and as a coating on the heated cathode to increase its electrical conductivity. Caesium did not become recognized as a high-performance industrial metal until the 1950s.[46] 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.[2]

A second was defined 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

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 caesium-133 atom.[47]

Since 1993, IUPAC accepts the alternative spelling cesium, but recommends the spelling caesium.[48]


Petroleum exploration

The largest end-use of nonradioactive caesium today is in caesium formate based drilling fluids for the 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 fluids. 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, and while as completion fluid (which refers to the emplacement of control hardware after drilling but prior to production) is to maintain the pressure.[2]

The high density of the caesium formate brine (up to 2.3 g/cm3, or 19.2 pounds per gallon),[49] 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.[49] 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).[50] Alkali formates are safe to handle and do not damage the producing formation or downhole metals as their corrosive alternative, high-density brines (such as zinc bromide ZnBr2 solutions), sometimes do, and they require less cleanup and disposal costs.[2]

Atomic clocks

 A room with a black box in the foreground and six control cabinet with space for five to six racks each. Most, but not all, of the cabinets are filled with white boxes.
Atomic clock ensemble at the U.S. Naval Observatory, which may be accessed by telephone (202-762-1401) or via Internet NTP servers[51]

Caesium is also used in atomic clocks, which use the resonant vibration frequency of caesium-133 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 1.4 million 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,[2] which has been regarded as "the most accurate realization of a unit that mankind has yet achieved."[47]

Because of their extreme precision, atomic clocks 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.[52] 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.[53]

Electric power and electronics

Magnetohydrodynamic (MHD) power-generating systems were researched but failed to gain acceptance for widespread use, and funding from the U.S. Department of Energy was stopped in the early 1990s.[54] Caesium metal has also been considered as the working fluid in high-temperature Rankine cycleturboelectric generators.[55] 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.[56]

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.[57] The range of photoemissive devices using caesium include optical character recognition devices, photomultiplier tubes, and video camera tubes.[58][59] Nevertheless, germanium, rubidium, selenium, silicon, tellurium, and several other elements can substitute caesium in photosensitive materials.[2]

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. Caesium vapor is used in many common magnetometers.[60] Caesium is also used as an internal standard in spectrophotometry.[61] Like other alkali metals, caesium has a great affinity for oxygen and is used as a "getter" in vacuum tubes.[62] Other uses of the metal include high-energy lasers, vapor glow lamps, and vapor rectifiers.[2]

Chemical and medical

Some fine white powder on a laboratory watch glass
A sample of caesium fluoride

Chemical applications are also another important use of caesium.[63] Liquid caesium can be used as a catalyst in the hydrogenation of certain organic compounds.[64] 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 is also used in ferrous and nonferrous metallurgy and in the purification of carbon dioxide as it absorbs gases and other impurities, while molten hydroxide (CsOH) has been used in the desulfurizing of heavy crude oil.[2]

Caesium fluoride is widely used in organic chemistry as a base,[65] or as a source of anhydrous fluoride ion.[66] Caesium salts sometimes are used to replace potassium or sodium salts in many organic syntheses, such as cyclization, esterification, and polymerization. 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 viral particles, sub-cellular organelles and fractions, and nucleic acids from biological samples.[67] 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.[2]


Caesium-137 is a 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.[68] It has been used in agriculture, cancer treatment, and sterilization of food, sewage sludge, and surgical equipment.[2][69] Radioactive isotopes of caesium in radiation devices were used in the medical field to treat certain types of cancer,[70] but emergence of better alternatives and the use of water-soluble caesium chloride in the sources, which would create wide range contamination, gradually put some of these caesium sources out of use.[71][72] Caesium-137 has been employed in a variety of industrial measurement gauges, including moisture, density, leveling, and thickness gauges.[73] 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.[74]

Isotope 137 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 caesium output by the nuclear power industry. These isotopes are used because, while they are less prevalent than either 133Cs or 137Cs, they can be produced solely by anthropogenic sources.[75]

Other uses

Electrons shooting out of an electron gun hit neutral fuel nuclei which leads to their ionization; in a chamber surrounded by magnets, the positive ions are directed towards a negative grid which accelerates them; once out of this chamber, the positive ions are neutralized from another electron gun leaving the chamber behind with a significant momentum thus propelling the previous chamber in the opposite direction.
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.[2] Eventually, xenon was used in the experimental spacecraft Deep Space 1 launched in 1998.[76][77] 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.[78]

Caesium nitrate is used as an oxidizer and pyrotechnic colorant to burn silicon in infrared flares[79] such as the LUU-19 flare,[80] because it emits much of its light in the near infrared spectrum.[81] Caesium has been used to reduce the radar signature of exhaust plumes in the SR-71 Blackbird military aircraft.[82] Caesium, along with rubidium, has been added as carbonates to glass because it reduces electrical conductivity and improves stability and durability, thus 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.[2]


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 colors, 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 versus time after the Chernobyl disaster[83]

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.[2] Caesium metal is highly pyrophoric, and ignites spontaneously in air to form caesium hydroxide and various oxides. Caesium hydroxide is a very 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.[84] 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.[85]

The isotopes 134 and 137 (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.[86][87][88] Accumulation of caesium-137 in lakes has been a high concern after the Chernobyl disaster.[89][90] Even in small amounts, it may cause infertility, cancer and even death.[91] The International Atomic Energy Agency and other sources have indicated that radioactive materials, such as caesium-137, may be used in radiological dispersion devices, or “dirty bombs”.[92]

See also


  1. ^ Along with rubidium (39 °C [102 °F]), francium (27 °C [81 °F]), mercury (−39 °C [−38 °F]), and gallium (30 °C [86 °F]). Bromine and mercury are also liquid at room temperature (melting at −7.2 °C, 19 °F) but it is not a metal, but a halogen.
  2. ^ The radioactive element francium may also have a lower melting point, but its radioactivity prevents enough of it from being isolated for direct testing.
    "Francium". Retrieved 2010-02-23. 
  3. ^ See Cesium and water, by Philip Evans, for a video.
  4. ^ Francium may be more electropositive, but this has not been experimentally measured due to its high radioactivity. Measurements of the first ionization energy of francium suggest that relativistic effects may lower its reactivity and raise its electronegativity above that expected from periodic trends.
    Andreev, S. V.; Letokhov, V. S.; Mishin, V. I. (1987). "Laser resonance photoionization spectroscopy of Rydberg levels in Fr". Physical Review Letters 59: 1274–76. doi:10.1103/PhysRevLett.59.1274. 


  1. ^ Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics 81st edition, CRC press.
  2. ^ a b c d e f g h i j k l m n o p q r s t u v w x y z Butterman, William C.; Brooks, William E.; Reese, Jr., Robert G. (2004). "Mineral Commodity Profile: Cesium" (PDF). United States Geological Survey. Retrieved 2009-12-27. 
  3. ^ Heiserman, David L. (1992). Exploring Chemical Elements and their Compounds. McGraw-Hill. pp. 201–203. ISBN 0-8306-3015-5. 
  4. ^ a b c d e Kaner, Richard (2003). "C&EN: It's Elemental: The Periodic Table - Cesium". American Chemical Society. Retrieved 2010-02-25. 
  5. ^ Deiseroth, H. J. (1997). "Alkali metal amalgams, a group of unusual alloys". Progress in Solid State Chemistry 25: 73–123. doi:10.1016/S0079-6786(97)81004-7. 
  6. ^ a b c Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils (1985). "Vergleichende Übersicht über die Gruppe der Alkalimetalle" (in German). Lehrbuch der Anorganischen Chemie (91–100 ed.). Walter de Gruyter. pp. 953–955. ISBN 3110075113. 
  7. ^ a b . doi:10.1016/0277-5387(96)00018-6. 
  8. ^ Dye, J. L. (1979). "Compounds of Alkali Metal Anions". Angewandte Chemie International Edition 18 (8): 587–598. doi:10.1002/anie.197905871. 
  9. ^ "Cesium Hydroxide (ICSC)". International Programme on Chemical Safety. April 2006. Retrieved 2010-02-24. 
  10. ^ Köhler, Michael J. (1999). Etching in microsystem technology. Wiley-VCH. p. 90. ISBN 3-527-29561-5. 
  11. ^ Wells, A.F. (1984). Structural Inorganic Chemistry (5 ed.). Oxford Science Publications. ISBN 0-19-855370-6. 
  12. ^ Cotton, F. Albert; Wilkinson, G. (1962). Advanced Inorganic Chemistry. Jon Wiley & sons, Inc.. p. 318. 
  13. ^ Lide, David R., ed. (2006), CRC Handbook of Chemistry and Physics (87th ed.), Boca Raton, FL: CRC Press, pp. 451, 514, ISBN 0-8493-0487-3 
  14. ^ a b Tsai, Khi-Ruey; Harris, P. M.; Lassettre, E. N. (1956). "The Crystal Structure of Cesium Monoxide". Journal of Physical Chemistry 60: 338–344. doi:10.1021/j150537a022. 
  15. ^ "Information Bridge: DOE Scientific and Technical Information - Sponsored by OSTI". 2009-11-23. Retrieved 2010-02-15. 
  16. ^ Vol'nov, I. I.; Matveev, V. V. (1963). "Synthesis of cesium ozonide through cesium superoxide". Bulletin of the Academy of Sciences, USSR Division of Chemical Science 12: 1040–1043. doi:10.1007/BF00845494. 
  17. ^ Tokareva, S. A. (1971). "Alkali and Alkaline Earth Metal Ozonides". Russian Chemical Reviews 40: 165–174. doi:10.1070/RC1971v040n02ABEH001903. 
  18. ^ Simon, A. (1997). "Group 1 and 2 Suboxides and Subnitrides — Metals with Atomic Size Holes and Tunnels". Coordination Chemistry Reviews 163: 253–270. doi:10.1016/S0010-8545(97)00013-1. 
  19. ^ Okamoto, H. (2009). "Cs-O (Cesium-Oxygen)". Journal of Phase Equilibria and Diffusion. doi:10.1007/s11669-009-9636-5. 
  20. ^ Tsai, Khi-Ruey; Harris, P. M.; Lassettre, E. N. (1956). Journal of Physical Chemistry 60: 345–347. doi:10.1021/j150537a023. 
  21. ^ Band, A.; Albu-Yaron, A.; Livneh, T.; Cohen, H.; Feldman, Y.; Shimon, L.; Popovitz-Biro, R.; Lyahovitskaya, V. et al. (2004). "Characterization of Oxides of Cesium". The Journal of Physical Chemistry B 108: 12360–12367. doi:10.1021/jp036432o. 
  22. ^ Brauer, G. (1947). "Untersuchungen ber das System Csium-Sauerstoff". Zeitschrift fr anorganische Chemie 255: 101. doi:10.1002/zaac.19472550110. 
  23. ^ Brown, F.; Hall, G.R.; Walter, A.J. (1955). "The half-life of Cs137". Journal of Inorganic and Nuclear Chemistry 1: 241–247. doi:10.1016/0022-1902(55)80027-9. 
  24. ^ Sonzogni, Alejandro. "Interactive Chart of Nuclides". National Nuclear Data Center: Brookhaven National Laboratory. Retrieved 2008-06-06. 
  25. ^ Ohki, Shigeo; Takaki, Naoyuki (14–16 October 2002). "Transmutation of Cesium-135 with Fast Reactors". Seventh Information Exchange Meeting on Actinide and Fission Product Partitioning and Transmutation. Jeju, Korea. 
  26. ^ "Objectives" (PDF). Retrieved 2010-02-15. 
  27. ^ Taylor, V. F.; Evans, R. D.; Cornett, R. J. (2008). "Preliminary evaluation of 135Cs/137Cs as a forensic tool for identifying source of radioactive contamination". Journal of Environmental Radioactivity 99 (1). doi:10.1016/j.jenvrad.2007.07.006. 
  28. ^ "IEER Report: Transmutation - Nuclear Alchemy Gamble". 2000-05-24. Retrieved 2010-02-15. 
  29. ^ "Chernobyl's Legacy: Health, Environmental and Socia-Economic Impacts and Recommendations to the Governments of Belarus, Russian Federation and Ukraine" (PDF). Retrieved 2010-02-18. 
  30. ^ "Cesium | Radiation Protection | US EPA". 2006-06-28. Retrieved 2010-02-15. 
  31. ^ Kase, Takeshi; Konashi, Kenji; Takahashi, Hiroshi; Hirao, Yasuo (1993). "Transmutation of Cesium-137 Using Proton Accelerator". Journal of Nuclear Science and Technology 30 (9): 911–918. doi:10.3327/jnst.30.911. 
  32. ^ Knief, Ronald Allen (1992). "Fission Fragments". Nuclear engineering: theory and technology of commercial nuclear power. Taylor & Francis. ISBN 9781560320883. 
  33. ^ Turekian, K.K.; Wedepohl, K. H. (1961). "Distribution of the elements in some major units of the Earth’s crust". Geological Society of America Bulletin 72 (2): 175–192. doi:10.1130/0016-7606(1961)72[175:DOTEIS2.0.CO;2]. 
  34. ^ "Artemis Project: Cesium as a Raw Material: Occurrence and Uses". Retrieved 2010-02-15. 
  35. ^ a b Černý, Petr; Simpson, F. M. (1978). "The Tanco Pegmatite at Bernic Lake, Manitoba: X. Pollucite" (PDF). Canadian Mineralogist 16: 325–333. 
  36. ^ a b c Polyak, Désirée E. "Cesium" (PDF). United States Geological Survey. Retrieved 2009-10-17. 
  37. ^ Norton, J. J. (1973). "Lithium, cesium, and rubidium—The rare alkali metals". in Brobst, D. A., and Pratt, W. P.. United States mineral resources. Paper 820. U.S. Geological Survey Professional. pp. 365–378. 
  38. ^ a b Burt, R. O. (1993). "Caesium and cesium compounds". Kirk-Othmer encyclopedia of chemical technology. 5 (4th ed.). New York: John Wiley & Sons, Inc.. pp. 749–764. ISBN 9780471484943. 
  39. ^ Benton, William; Turner, Jim (2000). "Cesium formate fluid succeeds in North Sea HPHT field trials" (PDF). Drilling Contractor (May/June): 38–41. 
  40. ^ a b transl. and rev. by Mary Eagleson. (1994). Concise encyclopedia chemistry. Berlin: de Gruyter. p. 198. ISBN 9783110114515. 
  41. ^ Bunsen quotes 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.
  42. ^ Oxford English Dictionary, 2nd Edition
  43. ^ a b Kirchhoff, G.; Bunsen, R. (1861). "Chemische Analyse durch Spectralbeobachtungen". Annalen der Physik und Chemie 189 (7): 337–381. doi:10.1002/andp.18611890702. 
  44. ^ a b Weeks, Mary Elvira (1932). "The discovery of the elements. XIII. Some spectroscopic discoveries". Journal of Chemical Education 9 (8): 1413–1434. doi:10.1021/ed009p1413. 
  45. ^ Setterberg, Carl (1882). "Ueber die Darstellung von Rubidium- und Cäsiumverbindungen und über die Gewinnung der Metalle selbst". Justus Liebig's Annalen der Chemie 211: 100–116. doi:10.1002/jlac.18822110105. 
  46. ^ Strod, A.J. (1957). "Cesium—A new industrial metal". American Ceramic Bulletin 36 (6): 212–213. 
  47. ^ a b "Cesium Atoms at Work". Retrieved 2009-12-20. 
  48. ^ International Union of Pure and Applied Chemistry. 2000-06-07. Retrieved 2010-01-25. 
  49. ^ a b Downs, J. D.; Blaszczynski, M.; Turner, J.; Harris, M. (February 2006). "Drilling and Completing Difficult HP/HT Wells With the Aid of Cesium Formate Brines-A Performance Review". IADC/SPE Drilling Conference. Miami, Florida, USASociety of Petroleum Engineers. doi:10.2118/99068-MS. 
  50. ^ Flatern, Rick (2001). "Keeping cool in the HPHT environment". Offshore Engineer (February): 33–37. 
  51. ^ "Usno Network Time Servers". Retrieved 2010-01-25. 
  52. ^ Breakiron, L. A. (2003). "Cesium atomic clocks". Retrieved 2003-02-24. 
  53. ^ Reel, Monte. "Where timing truly is everything". Washington Post: p. B1. Retrieved 2020-01-26. 
  54. ^ National Research Council (U.S.) (2001). Energy research at DOE—Was it worth it?. National Academy Press. pp. 190–194. ISBN 9780309074483. Retrieved 2001-11-16. 
  55. ^ Roskill Information Services (1984). Economics of Caesium and Rubidium (Reports on Metals & Minerals). London, United Kingdom: Roskill Information Services. p. 51. ISBN 9780862142506. 
  56. ^ Rasor, Ned S.; Warner, Charles (September 1964). "Correlation of Emission Processes for Adsorbed Alkali Films on Metal Surfaces". Journal of Applied Physics 35 (9): 2589–2600. doi:10.1063/1.1713806. 
  57. ^ "Cesium Supplier & Technical Information". American Elements. Retrieved 2010-01-25. 
  58. ^ Smedley, John; Rao, Triveni; Wang, Erdong (2009). "K 2CsSb Cathode Development". American Institute of Physics Conference Proceedings 1149: 1062–1066. doi:10.1063/1.3215593. 
  59. ^ Görlich, P. (1936). "Über zusammengesetzte, durchsichtige Photokathoden". Zeitschrift für Physik 101: 335–342. doi:10.1007/BF01342330. 
  60. ^ Groeger, S.; Pazgalev, A. S.; Weis, A. (2005). "Comparison of discharge lamp and laser pumped cesium magnetometers". Applied Physics B 80: 645–654. doi:10.1007/s00340-005-1773-x. 
  61. ^ "Internal Standards". Laboratory instrumentation. New York: Wiley. 1995. p. 108. ISBN 9780471285724. 
  62. ^ McGee, James D. (1969). Photo-electronic image devices: proceedings of the fourth symposium held at Imperial College, London, September 16–20, 1968. 1. Academic Press. p. 391. Retrieved 2010-01-25. 
  63. ^ Burt, R. O. (1993). "Cesium and cesium compounds". Kirk-Othmer encyclopedia of chemical technology. 5 (4th ed.). New York: John Wiley & Sons. p. 759. ISBN 978-0-471-15158-6. 
  64. ^ Pulham, R. J.; Turner, G. M. (1990). "The catalytic hydrogenation of ethene on the surface of liquid caesium". Journal of Catalysis 125 (1): 89–94. doi:10.1016/0021-9517(90)90080-4. 
  65. ^ Greenwood, N.N.; Earnshaw, A. (1984). Chemistry of the Elements. Oxford, UK: Pergamon Press. 
  66. ^ Evans, F. W.; Litt, M. H.; Weidler-Kubanek, A. M.; Avonda, F. P. (1968). "Reactions Catalyzed by Potassium Fluoride. 111. The Knoevenagel Reaction". Journal of Organic Chemistry 33: 1837–1839. doi:10.1021/jo01269a028. 
  67. ^ Desai, Mohamed A., ed (2000). "Gradient Materials". Downstream processing methods. Totowa, N.J.: Humana Press. pp. 61–62. ISBN 9780896035645. 
  68. ^ Okumura, Takeshi. "The material flow of radioactive cesium-137 in the U.S. 2000" (PDF). United States Environmental Protection Agency. Retrieved 2009-12-20. 
  69. ^ Jensen, N. L. (1985). Cesium, in Mineral facts and problems. Bulletin 675. U.S. Bureau of Mines. pp. 133–138. 
  70. ^ "IsoRay's Cesium-131 Medical Isotope Used In Milestone Procedure Treating Eye Cancers At Tufts-New England Medical Center". Retrieved 2010-02-15. 
  71. ^ Bentel, Gunilla Carleson (1996). "Caesium-137 Machines". Radiation therapy planning. McGraw-Hill Professional. p. 22. ISBN 9780070051157. 
  72. ^ National Research Council (U.S.). Committee on Radiation Source Use and Replacement (2008). Radiation source use and replacement: abbreviated version. National Academies Press. ISBN 9780309110143. 
  73. ^ Loxton, R., ed (1995). "Level and density measurement using non-contact nuclear gauges". Instrumentation : a reader ; for the instrumentation course at the Open University. London: Chapman & Hall. pp. 82–85. ISBN 9780412534003. 
  74. ^ Timur, A.; Toksoz, M. N. (1985). "Downhole Geophysical Logging". Annual Review of Earth and Planetary Sciences 13: 315. doi:10.1146/annurev.ea.13.050185.001531. 
  75. ^ Carol Kendall. USGS – Isotope Tracers – Resources. Retrieved 2010-01-25. 
  76. ^ Marcucci, M. G.; Polk, J. E. (2000). "NSTAR Xenon Ion Thruster on Deep Space 1: Ground and flight tests (invited)". Review of Scientific Instruments 71: 1389–1400. doi:10.1063/1.1150468. 
  77. ^ Sovey, James S.; Rawlin, Vincent K.; Patterson, Michael J. "A Synopsis of Ion Propulsion Development Projects in the United States: SERT I to Deep Space I" (PDF). NASA. Retrieved 2009-12-12. 
  78. ^ "In-FEEP Thruster Ion Beam Neutralization with Thermionic and Field Emission Cathodes" (PDF). 27th International Electric Propulsion Conference. Pasadena, California. October 2001. Retrieved 2010-01-25. 
  79. ^ "United States Patent 6230628: Infrared illumination compositions and articles containing the same". Retrieved 2010-01-25. 
  80. ^ "LUU-19 Flare". Federation of American Scientists. 2000-04-23. Retrieved 2009-12-12. 
  81. ^ Charrier, E.; Charsley, E.L.; Laye, P.G.; Markham, H.M.; Berger, B.; Griffiths, T.T. (2006). "Determination of the temperature and enthalpy of the solid–solid phase transition of caesium nitrate by differential scanning calorimetry". Thermochimica Acta 445: 36–39. doi:10.1016/j.tca.2006.04.002. 
  82. ^ Crickmore, Paul F. (2000). Lockheed SR-71: the secret missions exposed. Osprey. p. 47. ISBN 9781841760988. 
  83. ^ Data from the OECD report and The radiochemical Manual (2nd ed.) B.J. Wilson (1966).
  84. ^ Pinsky, Carl; Bose, Ranjan; Taylor, J. R.; McKee, Jasper; Lapointe, Claude; Birchall, James (1981). "Cesium in mammals: Acute toxicity, organ changes and tissue accumulation". Journal of Environmental Science and Health, Part A 16: 549– 567. doi:10.1080/10934528109375003. 
  85. ^ Johnson, Garland T.; Lewis, Trent R.; Wagner, D. Wagner (1975). "Acute toxicity of cesium and rubidium compounds". Toxicology and Applied Pharmacology 32 (2): 239–245. doi:10.1016/0041-008X(75)90216-1. PMID 1154391. 
  86. ^ Nishita, H.; Dixon, D.; Larson, K. H. (1962). "Accumulation of Cs and K and growth of bean plants in nutrient solution and soils". Plant and Soil 17: 221–242. doi:10.1007/BF01376226. 
  87. ^ Avery, S. (1996). "Fate of caesium in the environment: Distribution between the abiotic and biotic components of aquatic and terrestrial ecosystems". Journal of Environmental Radioactivity 30: 139–171. doi:10.1016/0265-931X(96)89276-9. 
  88. ^ Salbu, Brit; Østby, Georg; Garmo, Torstein H.; Hove, Knut (1992). "Availability of caesium isotopes in vegetation estimated from incubation and extraction experiments". Analyst 117: 487–491. doi:10.1039/AN9921700487. 
  89. ^ Smith, Jim T.; Beresford, Nicholas A.. Chernobyl: Catastrophe and Consequences. Berlin: Springer. ISBN 3540238662. 
  90. ^ Eremeev, V. N.; Chudinovskikh, T. V.; Batrakov, G. F.; Ivanova, T. M. (1991). "Radioactive isotopes of caesium in the waters and near-water atmospheric layer of the Black Sea". Physical Oceanography 2 (1): 57–64. doi:10.1007/BF02197418. 
  91. ^ "Asia-Pacific | Chinese 'find' radioactive ball". BBC News. 2009-03-27. Retrieved 2010-01-25. 
  92. ^ Charbonneau, Louis. "IAEA director warns of “dirty bomb” risk". Washington Post (Reuters): p. A15. 

External links

1911 encyclopedia

Up to date as of January 14, 2010

From LoveToKnow 1911

CAESIUM (symbol Cs, atomic weight 132.9), one of the alkali metals. Its name is derived from 'the Lat. caestus, sky-blue, from two bright blue lines of its spectrum. It is of historical importance, since it was the first metal to be discovered by the aid of the spectroscope (R. Bunsen, Berlin Acad. Ber., 1860), although caesium salts had undoubtedly been examined before, but had been mistaken for potassium salts (see C. F. Plattner, Pog. Ann., 1846, p. 443, on the analysis of pollux and the subsequent work of F. Pisani, Comptes Rendus, 1864, 58, p. 714). Caesium is found in the mineral springs of Frankenhausen, Montecatini, di Val di Nievole, Tuscany, and Wheal Clifford near Redruth, Cornwall (W. A. Miller, Chem. News, 1864, 10, p. 181), and, associated with rubidium, at Diirkheim; it is also found in lepidolite, leucite, petalite, triphylline and in the carnallite from Stassfurt. The separation of caesium from the minerals which contain it is an exceedingly difficult and laborious process. According to R. Bunsen, the best source of rubidium and caesium salts is the residue left after extraction of lithium salts from lepidolite. This residue consists of sodium, potassium and lithium chlorides, with small quantities of caesium and rubidium chlorides. The caesium and rubidium are separated from this by repeated fractional crystallization of their double platinum chlorides, which are much less soluble in water than those of the other alkali metals (R. Bunsen, Ann., 1862, 122, p. 347; 1863, 125, p. 367). The platino-chlorides are reduced by hydrogen, and the caesium and rubidium chlorides extracted by water. See also A. Schrdtter (Jour. prak. Chem., 1864, 93, p. 2075) and W. Heintz (Journ. prak. Chem., 1862, 87, p. 310). W. Feit and K. Kubierschky (Chem. Zeit., 1892, 16, p. 335) separate rubidium and caesium from the other alkali metals by converting them into double chlorides with stannic chloride; whilst J. Redtenbacher (Jour. prak. Chem., 1865, 94, p. 44 2) separates them from potassium by conversion into alums, which C. Setterberg (Ann., 1882, 211, p. loo) has shown are very slightly soluble in a solution of potash alum. In order to separate caesium from rubidium, use is made of the different solubilities of their various salts. The bitartrates RbHC 4 H 4 0 6 and CsHC 4 H 4 O 6 have been employed, as have also the alums (see above). The double chloride of caesium and antimony 3CsC1 2SbC1 3 (R. Godeffroy, Ber., 1874, 7, p. 375; Ann., 1876, 181, p. 176) has been used, the corresponding compound not being formed by rubidium. The metal has been obtained by electrolysis of a mixture of caesium and barium cyanides (C. Setterberg, Ann., 1882, 211, p. loo) and by heating the hydroxide with magnesium or aluminium (N. Beketoff, Chem. Centralblatt, 1889, 2, p. 245). L. Hackspill (Comptes Rendus, 5905, 141, p. 101) finds that metallic caesium can be obtained more readily by heating the chloride with metallic calcium. A special V-shaped tube is used in the operation, and the reaction commences between 400° C. and 500° C.

It is a silvery white metal which burns on heating in air. It melts at 26° to 27° C. and has a specific gravity of 1.88 (15°C.). The atomic weight of caesium has been determined by the analysis of its chloride and bromide. Richards and Archibald (Zeit. anorg. Chem., 1903, 34, p. 353) obtained 132.879 (O =16).

Caesium hydroxide, Cs(OH) 2, obtained by the decomposition of the sulphate with baryta water,is a greyish-white deliquescent solid,which melts at a red heat and absorbs carbon dioxide rapidly. It readily dissolves in water, with evolution of much heat. Caesium chloride,. CsCl, is obtained by the direct action of chlorine on caesium, or by solution of the hydroxide in hydrochloric acid. It forms small cubes which melt at a red heat and volatilize readily. It deliquesces in moist air. Many double chlorides are known, and may be prepared by mixing solutions of the two components in the requisite proportions. The bromide, CsBr, and iodide, CsI, resemble the corresponding potassium salts. Many trihaloid salts of caesium are also known, such as CsBr 3, CsC1Br 2, Cs13, CsBrI 2, CsBr 2 I, &c. (H. L. Wells and S. L. Penfield, Zeit. fiir anorg. Chem., 1892, 1, p. 85). Caesium sulphate, Cs 2 SO 4, may be prepared by dissolving the hydroxide or carbonate in sulphuric acid. It crystallizes in short hard prisms, which are readily soluble in water but insoluble in alcohol. It combines with many metallic sulphates (silver, zinc, cobalt, nickel, &c.) to form double sulphates of the type Cs2S04 RS04.6H20. It also, forms a caesium-alum Cs 2 S04-Al 2 (S04) 3 24H 2 O. Caesium nitrate, CsNO 3, is obtained by dissolving the carbonate in nitric acid, and crystallizes in glittering prisms, which melt readily, and on heating evolve oxygen and leave a residue of caesium nitrite. The carbonate, Cs2C03, silicofluoride, Cs2SiF6, borate, Cs20.3B 2 O 3, and the sulphides. Cs 2 S 4H 2 0, Cs 2 S 2 H 2 O, Cs 2 S 3 H 2 0, Cs 2 S 4 and ,Cs 2 S 5 H 2 0, are also known.

. Caesium compounds can be readily recognized by the two bright. blue lines (of wave length 4555 and 4593) in their flame spectrum, but these are not present in the spark spectrum. The other lines, include three in the green, two in the yellow, and two in the orange..

<< Caesarea Philippi

Caespitose >>


Up to date as of January 15, 2010
(Redirected to caesium article)

Definition from Wiktionary, a free dictionary



Chemical element
Cs Previous: xenon (Xe)
Next: barium (Ba)

Alternative spellings





caesium (uncountable)

Wikipedia has an article on:


  1. A metallic chemical element (symbol Cs) with an atomic number of 55.

Derived terms

  • caesium auride, cesium auride
  • caesium azide, cesium azide
  • caesium cell, cesium cell
  • caesium chloride, cesium chloride
  • caesium clock, cesium clock
  • caesium fluoride, cesium fluoride
  • caesium hydroxide, cesium hydroxide
  • caesium iodide, cesium iodide
  • caesium nitrate, cesium nitrate

Related terms


External links

For etymology and more information refer to (a lot of the translations were taken from that site with permission from the author).

See also


Simple English

The gold-colored metal in the tubes is Caesium.

Caesium (or cesium) is the chemical element with the atomic number 55 on the periodic table. Its symbol is Cs.

Caesium is an alkali metal. Its melting point is low (28 °C). It is extremely reactive. Because of its high reactivity, it is a dangerous chemical. It may set itself on fire (ignite) in air. It explodes on contact with water. It reacts more violently than the other alkali metals with water. Consequently, caesium is stored in mineral oil.[1]

Caesium is a rare element. Since there is little caesium on the Earth, it is rather expensive. The human body does not need caesium. In large quantities, it is mildly poisonous because it resembles potassium, which the body does need.


Caesium was first described in 1861, by Gustav Robert Kirchhoff and Robert Wilhelm Bunsen. They were testing mineral water, from Bad Dürkheim. After they separated calcium, strontium, magnesium and lithium, they saw two lines in the "blue" range of the spectometry. Because of these lines, they concluded that in addition to the elements already found, there must be another unknown substance in the mineral water. They named this substance caesium, after the color blue.[2]

Isotopes and compounds

Caesium has at least 39 known isotopes ranging in atomic mass from 112 to 151. Only one of these, 133Cs, is stable. Therefore, the naturally-occurring isotope of caesium is 133Cs, which is not radioactive. 133Cs is used in atomic clocks, its vibration frequency used to define the length of the second. Another isotope, 137Cs does not occur naturally but is produced as a byproduct of nuclear fission. It is intensely radioactive and used as an industrial gamma ray source.

Caesium forms compounds with many other chemical elements. Caesium formate is used in oil drilling because of its high density.


  1. William C. Butterman, William E. Brooks, and Robert G. Reese Jr.2009. Butterman, William C.; Brooks, William E.; Reese, Jr., Robert G. (2004). "Mineral Commodity Profile: Cesium" (PDF). United States Geological Survey. Retrieved 2009-12-27. 
  2. G. Kirchhoff, R. Bunsen: Chemische Analyse durch Spectralbeobachtungen. In: Annalen der Physik und Chemie. 1861, 189, 7, S. 337–381 (doi:10.1002/andp.18611890702).

Got something to say? Make a comment.
Your name
Your email address