Caesium

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Not to be confused with cerium.

Caesium, 55Cs
Some pale gold metal, with a liquid-like texture and lustre, sealed in a glass ampoule
Caesium
Pronunciation/ˈsiːziəm/ ​(SEE-zee-əm)
Alternative namecesium (US)
Appearancepale gold
Standard atomic weight Ar°(Cs)
Caesium in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
Rb
↑
Cs
↓
Fr
xenon ← caesium → barium
kJ/mol
Heat of vaporization63.9 kJ/mol
Molar heat capacity32.210 J/(mol·K)
Vapour pressure
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K) 418 469 534 623 750 940
Atomic properties
Discovery
Robert Bunsen and Gustav Kirchhoff (1860)
First isolationCarl Setterberg (1882)
Isotopes of caesium
Main isotopes[7] Decay
abun­dance half-life (t1/2) mode pro­duct
131Cs synth 9.7 d ε
131Xe
133Cs 100%
stable
134Cs synth 2.0648 y ε
134Xe
β−
134Ba
135Cs trace 1.33×106 y β−
135Ba
137Cs synth 30.17 y[8] β−
137Ba
 Category: Caesium
| references

Caesium (

picometers
.

The German chemist Robert Bunsen and physicist Gustav Kirchhoff discovered caesium in 1860 by the newly developed method of flame spectroscopy. The first small-scale applications for caesium were as a "getter" in vacuum tubes and in photoelectric cells. Caesium is widely used in highly accurate atomic clocks. In 1967, the International System of Units began using a specific hyperfine transition of neutral caesium-133 atoms to define the basic unit of time, the second.

Since the 1990s, the largest

radioisotopes
present a significant health and environmental hazard.

Characteristics

Physical properties

Y-shaped yellowish crystal in glass ampoule, looking like the branch of a pine tree
High-purity caesium-133 stored in argon.

Of all elements that are solid at room temperature, caesium is the softest: it has a hardness of 0.2 Mohs. It is a very

lowest of all metals other than mercury.[19] Its compounds burn with a blue[20][21] or violet[21]
colour.

Caesium crystals (golden) compared to rubidium crystals (silvery)

Caesium forms

photosensitive.[14] It mixes with all the other alkali metals (except lithium); the alloy with a molar distribution of 41% caesium, 47% potassium, and 12% sodium has the lowest melting point of any known metal alloy, at −78 °C (−108 °F).[18][22] A few amalgams have been studied: CsHg
2 is black with a purple metallic lustre, while CsHg is golden-coloured, also with a metallic lustre.[23]

The golden colour of caesium comes from the decreasing frequency of light required to excite electrons of the alkali metals as the group is descended. For lithium through rubidium this frequency is in the ultraviolet, but for caesium it enters the blue–violet end of the spectrum; in other words, the plasmonic frequency of the alkali metals becomes lower from lithium to caesium. Thus caesium transmits and partially absorbs violet light preferentially while other colours (having lower frequency) are reflected; hence it appears yellowish.[24]

Allotropes

Caesium exists in the form of different allotropes, one of them a dimer called dicaesium.[25]

Chemical properties

Addition of a small amount of caesium to cold water is explosive.

Caesium metal is highly reactive and

hazardous material. It is stored and shipped in dry, saturated hydrocarbons such as mineral oil. It can be handled only under inert gas, such as argon. However, a caesium-water explosion is often less powerful than a sodium-water explosion with a similar amount of sodium. This is because caesium explodes instantly upon contact with water, leaving little time for hydrogen to accumulate.[26] Caesium can be stored in vacuum-sealed borosilicate glass ampoules. In quantities of more than about 100 grams (3.5 oz), caesium is shipped in hermetically sealed, stainless steel containers.[14]

The chemistry of caesium is similar to that of other alkali metals, in particular

alkali metals
.

Compounds

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

Most caesium compounds contain the element as the

p-block element and capable of forming higher fluorides with higher oxidation states (i.e., CsFn with n > 1) under high pressure.[32] This prediction needs to be validated by further experiments.[33]

Salts of Cs+ are usually colourless unless the anion itself is coloured. Many of the simple salts are

soluble.[14]

polar aprotic solvents, are far more basic on the basis of the Brønsted–Lowry acid–base theory.[27]

A stoichiometric mixture of caesium and gold will react to form yellow caesium auride (Cs+Au−) upon heating. The auride anion here behaves as a pseudohalogen. The compound reacts violently with water, yielding caesium hydroxide, metallic gold, and hydrogen gas; in liquid ammonia it can be reacted with a caesium-specific ion exchange resin to produce tetramethylammonium auride. The analogous platinum compound, red caesium platinide (Cs2Pt), contains the platinide ion that behaves as a pseudochalcogen.[36]

Complexes

Like all metal cations, Cs+ forms complexes with

softness (tendency to form covalent bonds) are properties exploited in separating Cs+ from other cations in the remediation of nuclear wastes, where 137Cs+ must be separated from large amounts of nonradioactive K+.[37]

Halides

Monatomic caesium halide wires grown inside double-wall carbon nanotubes (TEM image).[38]

cubic closest packed array as do Na+ and Cl− in sodium chloride.[27] Notably, caesium and fluorine have the lowest and highest electronegativities
, respectively, among all the known elements.

pm and Cl−
 181 pm.[40]

Oxides

The stick and ball diagram shows three regular octahedra, which are connected to the next one by one surface and the last one shares one surface with the first. All three have one edge in common. All eleven vertices are purple spheres representing caesium, and at the center of each octahedron is a small red sphere representing oxygen.
Cs
11O
3 cluster

More so than the other alkali metals, caesium forms numerous binary compounds with

hexagonal crystals,[42] and is the only oxide of the anti-CdCl
2 type.[43] It vaporizes at 250 °C (482 °F), and decomposes to caesium metal and the peroxide Cs
2O
2 at temperatures above 400 °C (752 °F). In addition to the superoxide and the ozonide CsO
3,[44][45] several brightly coloured suboxides have also been studied.[46] These include Cs
7O, Cs
4O, Cs
11O
3, Cs
3O (dark-green[47]), CsO, Cs
3O
2,[48] as well as Cs
7O
2.[49][50] The latter may be heated in a vacuum to generate Cs
2O.[43] Binary compounds with sulfur, selenium, and tellurium also exist.[14]

Isotopes

Main article: Isotopes of caesium

Caesium has 41 known

nuclear spin (7/2+), nuclear magnetic resonance studies can use this isotope at a resonating frequency of 11.7 MHz.[53]

A graph showing the energetics of 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% probability beta emission.
Decay of caesium-137

The radioactive

137mBa by beta decay, and then to nonradioactive barium, while 134Cs transforms into 134Ba directly. The isotopes with mass numbers of 129, 131, 132 and 136, have half-lives between a day and two weeks, while most of the other isotopes have half-lives from a few seconds to fractions of a second. At least 21 metastable nuclear isomers exist. Other than 134mCs (with a half-life of just under 3 hours), all are very unstable and decay with half-lives of a few minutes or less.[54][55]

The isotope 135Cs is one of the

136Xe before it can decay to 135Cs.[57][58]

The

medium-lived products of nuclear fission, and the prime sources of radioactivity from spent nuclear fuel after several years of cooling, lasting several hundred years.[60] Those two isotopes are the largest source of residual radioactivity in the area of the Chernobyl disaster.[61] Because of the low capture rate, disposing of 137Cs through neutron capture is not feasible and the only current solution is to allow it to decay over time.[62]

Almost all caesium produced from nuclear fission comes from the beta decay of originally more neutron-rich fission products, passing through various isotopes of iodine and xenon.[63] Because iodine and xenon are volatile and can diffuse through nuclear fuel or air, radioactive caesium is often created far from the original site of fission.[64] With nuclear weapons testing in the 1950s through the 1980s, 137Cs was released into the atmosphere and returned to the surface of the earth as a component of radioactive fallout. It is a ready marker of the movement of soil and sediment from those times.[14]

Occurrence

A white mineral, from which white and pale pink crystals protrude
Pollucite, a caesium mineral
See also: Caesium minerals

Caesium is a relatively rare element, estimated to average 3 

parts per million in the Earth's crust.[65] Nevertheless, it is more abundant than such elements as antimony, cadmium, tin, and tungsten, and two orders of magnitude more abundant than mercury and silver; it is 3.3% as abundant as rubidium, with which it is closely associated, chemically.[14]

Due to its large

rhodizite.[14] The only economically important ore for caesium is pollucite Cs(AlSi
2O
6), which is found in a few places around the world in zoned pegmatites, associated with the more commercially important lithium minerals, lepidolite and petalite. Within the pegmatites, the large grain size and the strong separation of the minerals results in high-grade ore for mining.[67]

The world's most significant and richest known source of caesium is the Tanco Mine at Bernic Lake in Manitoba, Canada, estimated to contain 350,000 metric tons of pollucite ore, representing more than two-thirds of the world's reserve base.[67][68] Although the stoichiometric content of caesium in pollucite is 42.6%, pure pollucite samples from this deposit contain only about 34% caesium, while the average content is 24 wt%.[68] Commercial pollucite contains more than 19% caesium.[69] The Bikita pegmatite deposit in Zimbabwe is mined for its petalite, but it also contains a significant amount of pollucite. Another notable source of pollucite is in the Karibib Desert, Namibia.[68] At the present rate of world mine production of 5 to 10 metric tons per year, reserves will last for thousands of years.[14]

Production

Mining and refining pollucite ore is a selective process and is conducted on a smaller scale than for most other metals. The ore is crushed, hand-sorted, but not usually concentrated, and then ground. Caesium is then extracted from pollucite primarily by three methods: acid digestion, alkaline decomposition, and direct reduction.[14][70]

In the acid digestion, the silicate pollucite rock is dissolved with strong acids, such as hydrochloric (HCl), sulfuric (H
2SO
4), hydrobromic (HBr), or hydrofluoric (HF) acids. With hydrochloric acid, a mixture of soluble chlorides is produced, and the insoluble chloride double salts of caesium are precipitated as caesium antimony chloride (Cs
4SbCl
7), caesium iodine chloride (Cs
2ICl), or caesium hexachlorocerate (Cs
2(CeCl
6)). After separation, the pure precipitated double salt is decomposed, and pure CsCl is precipitated by evaporating the water.

The sulfuric acid method yields the insoluble double salt directly as caesium alum (CsAl(SO
4)
2·12H
2O). The aluminium sulfate component is converted to insoluble aluminium oxide by roasting the alum with carbon, and the resulting product is leached with water to yield a Cs
2SO
4 solution.[14]

Roasting pollucite with calcium carbonate and calcium chloride yields insoluble calcium silicates and soluble caesium chloride. Leaching with water or dilute ammonia (NH
4OH) yields a dilute chloride (CsCl) solution. This solution can be evaporated to produce caesium chloride or transformed into caesium alum or caesium carbonate. Though not commercially feasible, the ore can be directly reduced with potassium, sodium, or calcium in vacuum to produce caesium metal directly.[14]

Most of the mined caesium (as salts) is directly converted into

Tanco mine near Bernic Lake in Manitoba, with a capacity of 12,000 barrels (1,900 m3) per year of caesium formate solution.[71] The primary smaller-scale commercial compounds of caesium are caesium chloride and nitrate.[72]

Alternatively, caesium metal may be obtained from the purified compounds derived from the ore. Caesium chloride and the other caesium halides can be reduced at 700 to 800 °C (1,292 to 1,472 °F) with calcium or barium, and caesium metal distilled from the result. In the same way, the aluminate, carbonate, or hydroxide may be reduced by magnesium.[14]

The metal can also be isolated by

dichromate can be reacted with zirconium to produce pure caesium metal without other gaseous products.[72]

Cs
2Cr
2O
7 + 2 Zr → 2 Cs + 2 ZrO
2+ Cr
2O
3

The price of 99.8% pure caesium (metal basis) in 2009 was about $10 per gram ($280/oz), but the compounds are significantly cheaper.[68]

History

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 (centre) discovered caesium with their newly invented spectroscope.

In 1860, Robert Bunsen and Gustav Kirchhoff discovered caesium in the mineral water from Dürkheim, Germany. Because of the bright blue lines in the emission spectrum, they derived the name from the Latin word caesius, meaning 'bluish grey'.[note 6][73][74][75] Caesium was the first element to be discovered with a spectroscope, which had been invented by Bunsen and Kirchhoff only a year previously.[18]

To obtain a pure sample of caesium, 44,000 litres (9,700 imp gal; 12,000 US gal) of mineral water had to be evaporated to yield 240 kilograms (530 lb) of concentrated salt solution. The alkaline earth metals were precipitated either as sulfates or oxalates, leaving 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, but these salts show a slight difference in solubility in hot water, and the less-soluble caesium and rubidium hexachloroplatinate ((Cs,Rb)2PtCl6) were obtained by fractional crystallization. After reduction of the hexachloroplatinate with hydrogen, caesium and rubidium were separated by the difference in solubility of their carbonates in alcohol. The process yielded 9.2 grams (0.32 oz) of rubidium chloride and 7.3 grams (0.26 oz) of caesium chloride from the initial 44,000 litres of mineral water.[74]

From the caesium chloride, the two scientists estimated the

Kekulé and Bunsen.[75] In 1882, he produced caesium metal by electrolysing caesium cyanide, avoiding the problems with the chloride.[77]

Historically, the most important use for caesium has been in research and development, primarily in chemical and electrical fields. Very few applications existed for caesium until the 1920s, when it came into use in radio

magnetohydrodynamic power generators.[14] Caesium is also used as a source of positive ions in secondary ion mass spectrometry
(SIMS).

Since 1967, the

caesium-133.[79] The 13th General Conference on Weights and Measures
of 1967 defined a second as: "the duration of 9,192,631,770 cycles of microwave light absorbed or emitted by the hyperfine transition of caesium-133 atoms in their ground state undisturbed by external fields".

Applications

Petroleum exploration

The largest present-day use of nonradioactive caesium is in

extractive oil industry.[14] 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 a drilling fluid is to lubricate drill bits, to bring rock cuttings to the surface, and to maintain pressure on the formation during drilling of the well. Completion fluids assist the emplacement of control hardware after drilling but prior to production by maintaining the pressure.[14]

The high density of the caesium formate brine (up to 2.3 g/cm3, or 19.2 pounds per gallon),

barrel in 2001).[81] Alkali formates are safe to handle and do not damage the producing formation or downhole metals as corrosive alternative, high-density brines (such as zinc bromide ZnBr
2 solutions) sometimes do; they also require less cleanup and reduce disposal costs.[14]

Atomic clocks

A room with a black box in the foreground and six control cabinets with space for five to six racks each. Most, but not all, of the cabinets are filled with white boxes.
Atomic clock ensemble at the U.S. Naval Observatory

Caesium-based

National Physical Laboratory in the UK.[82] Caesium clocks have improved over the past half-century and are regarded as "the most accurate realization of a unit that mankind has yet achieved."[79] These clocks measure frequency with an error of 2 to 3 parts in 1014, which corresponds to an accuracy of 2 nanoseconds per day, or one second in 1.4 million years. The latest versions are more accurate than 1 part in 1015, about 1 second in 20 million years.[14] The caesium standard is the primary standard for standards-compliant time and frequency measurements.[83] Caesium clocks regulate the timing of cell phone networks and the Internet.[84]

Definition of the second

The second, symbol s, is the SI unit of time. The

Hz, which is equal to s−1."[85]

Electric power and electronics

Caesium vapour thermionic generators are low-power devices that convert heat energy to electrical energy. In the two-electrode vacuum tube converter, caesium neutralizes the space charge near the cathode and enhances the current flow.[86]

Caesium is also important for its photoemissive properties, converting light to electron flow. It is used in photoelectric cells because caesium-based cathodes, such as the intermetallic compound K
2CsSb, have a low threshold voltage for emission of electrons.[87] The range of photoemissive devices using caesium include optical character recognition devices, photomultiplier tubes, and video camera tubes.[88][89] Nevertheless, germanium, rubidium, selenium, silicon, tellurium, and several other elements can be substituted for caesium in photosensitive materials.[14]

Caesium iodide (CsI), bromide (CsBr) and fluoride (CsF) crystals are employed for scintillators in scintillation counters widely used in mineral exploration and particle physics research to detect gamma and X-ray radiation. Being a heavy element, caesium provides good stopping power with better detection. Caesium compounds may provide a faster response (CsF) and be less hygroscopic (CsI).

Caesium vapour is used in many common magnetometers.[90]

The element is used as an internal standard in spectrophotometry.[91] Like other alkali metals, caesium has a great affinity for oxygen and is used as a "getter" in vacuum tubes.[92] Other uses of the metal include high-energy lasers, vapour glow lamps, and vapour rectifiers.[14]

Centrifugation fluids

The high density of the caesium ion makes solutions of caesium chloride, caesium sulfate, and caesium trifluoroacetate (Cs(O
2CCF
3)) useful in molecular biology for density gradient ultracentrifugation.[93] This technology is used primarily in the isolation of viral particles, subcellular organelles and fractions, and nucleic acids from biological samples.[94]

Chemical and medical use

Some fine white powder on a laboratory watch glass
Caesium chloride powder

Relatively few chemical applications use caesium.[95] Doping with caesium compounds enhances the effectiveness of several metal-ion catalysts for chemical synthesis, such as acrylic acid, anthraquinone, ethylene oxide, methanol, phthalic anhydride, styrene, methyl methacrylate monomers, and various olefins. It is also used in the catalytic conversion of sulfur dioxide into sulfur trioxide in the production of sulfuric acid.[14]

esterification, and polymerization. Caesium has also been used in thermoluminescent radiation dosimetry (TLD): When exposed to radiation, it acquires crystal defects that, when heated, revert with emission of light proportionate to the received dose. Thus, measuring the light pulse with a photomultiplier tube
can allow the accumulated radiation dose to be quantified.

Nuclear and isotope applications

Caesium-137 is a radioisotope commonly used as a gamma-emitter in industrial applications. Its advantages include a half-life of roughly 30 years, its availability from the nuclear fuel cycle, and having 137Ba as a stable end product. The high water solubility is a disadvantage which makes it incompatible with large pool irradiators for food and medical supplies.[97] It has been used in agriculture, cancer treatment, and the sterilization of food, sewage sludge, and surgical equipment.[14][98] Radioactive isotopes of caesium in radiation devices were used in the medical field to treat certain types of cancer,[99] but emergence of better alternatives and the use of water-soluble caesium chloride in the sources, which could create wide-ranging contamination, gradually put some of these caesium sources out of use.[100][101] Caesium-137 has been employed in a variety of industrial measurement gauges, including moisture, density, levelling, and thickness gauges.[102] 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.[103]

Caesium-137 has been used in hydrologic studies analogous to those with tritium. As a daughter product of fission bomb testing from the 1950s through the mid-1980s, caesium-137 was released into the atmosphere, where it was absorbed readily into solution. Known year-to-year variation within that period allows correlation with soil and sediment layers. Caesium-134, and to a lesser extent caesium-135, have also been used in hydrology to measure the caesium output by the nuclear power industry. While they are less prevalent than either caesium-133 or caesium-137, these bellwether isotopes are produced solely from anthropogenic sources.[104]

Other uses

Electrons beamed from an electron gun hit and ionize neutral fuel atoms; in a chamber surrounded by magnets, the positive ions are directed toward a negative grid that accelerates them. The force of the engine is created by expelling the ions from the rear at high velocity. On exiting, the positive ions are neutralized from another electron gun, ensuring that neither the ship nor the exhaust is electrically charged and are not attracted.
Schematics of an electrostatic ion thruster developed for use with caesium or mercury fuel

Caesium and mercury were used as a propellant in early ion engines designed for spacecraft propulsion on very long interplanetary or extraplanetary missions. The fuel was ionized by contact with a charged tungsten electrode. But corrosion by caesium on spacecraft components has pushed development in the direction of inert gas propellants, such as xenon, which are easier to handle in ground-based tests and do less potential damage to the spacecraft.[14] Xenon was used in the experimental spacecraft Deep Space 1 launched in 1998.[105][106] Nevertheless, field-emission electric propulsion thrusters that accelerate liquid metal ions such as caesium have been built.[107]

CIA reconnaissance aircraft.[111] Caesium and rubidium have been added as a carbonate to glass because they reduce electrical conductivity and improve stability and durability of fibre optics and night vision devices. Caesium fluoride or caesium aluminium fluoride are used in fluxes formulated for brazing aluminium alloys that contain magnesium.[14]

Magnetohydrodynamic (MHD) power-generating systems were researched, but failed to gain widespread acceptance.[112] Caesium metal has also been considered as the working fluid in high-temperature Rankine cycle turboelectric generators.[113]

Caesium salts have been evaluated as antishock reagents 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.[14]

Caesium-133 can be laser cooled and used to probe fundamental and technological problems in quantum physics. It has a particularly convenient Feshbach spectrum to enable studies of ultracold atoms requiring tunable interactions.[114]

Health and safety hazards

Caesium
Hazards
GHS labelling:[115]
GHS02: Flammable GHS05: Corrosive
Danger
H260, H314
P223, P231+P232, P280, P305+P351+P338, P370+P378, P422
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 4: Will rapidly or completely vaporize at normal atmospheric pressure and temperature, or is readily dispersed in air and will burn readily. Flash point below 23 °C (73 °F). E.g. propaneInstability 3: Capable of detonation or explosive decomposition but requires a strong initiating source, must be heated under confinement before initiation, reacts explosively with water, or will detonate if severely shocked. E.g. hydrogen peroxideSpecial hazard W: Reacts with water in an unusual or dangerous manner. E.g. sodium, sulfuric acid
3
4
3
W
Graph of percentage of the radioactive output by each nuclide that form after a nuclear fallout vs. logarithm of time after the incident. In curves of various colours, the predominant source of radiation are depicted in order: Te-132/I-132 for the first five or so days; I-131 for the next five; Ba-140/La-140 briefly; Zr-95/Nb-95 from day 10 until about day 200; and finally Cs-137. Other nuclides producing radioactivity, but not peaking as a major component are Ru, peaking at about 50 days, and Cs-134 at around 600 days.
The portion of the total radiation dose (in air) contributed by each isotope plotted against time after the Chernobyl disaster. Caesium-137 became the primary source of radiation about 200 days after the accident.[116]

Nonradioactive caesium compounds are only mildly toxic, and nonradioactive caesium is not a significant environmental hazard. Because biochemical processes can confuse and substitute caesium with

arrhythmia, and acute cardiac arrest, but such amounts would not ordinarily be encountered in natural sources.[117][118]

The median lethal dose (LD50) for caesium chloride in mice is 2.3 g per kilogram, which is comparable to the LD50 values of potassium chloride and sodium chloride.[119] The principal use of nonradioactive caesium is as caesium formate in petroleum drilling fluids because it is much less toxic than alternatives, though it is more costly.[80]

Caesium metal is one of the most reactive elements and is highly

explosive in the presence of water. The hydrogen gas produced by the reaction is heated by the thermal energy released at the same time, causing ignition and a violent explosion. This can occur with other alkali metals, but caesium is so potent that this explosive reaction can be triggered even by cold water.[14]

It is highly pyrophoric: the autoignition temperature of caesium is −116 °C (−177 °F), and it ignites explosively in air to form caesium hydroxide and various oxides. Caesium hydroxide is a very strong base, and will rapidly corrode glass.[19]

The

sporocarps.[124] Accumulation of caesium-137 in lakes has been a great concern after the Chernobyl disaster.[125][126] Experiments with dogs showed that a single dose of 3.8 millicuries (140 MBq, 4.1 μg of caesium-137) per kilogram is lethal within three weeks;[127] smaller amounts may cause infertility and cancer.[128] The International Atomic Energy Agency and other sources have warned that radioactive materials, such as caesium-137, could be used in radiological dispersion devices, or "dirty bombs".[129]

See also

Notes

  1. ligature æ as cæsius; hence, an alternative but now old-fashioned orthography is cæsium. More spelling explanation at ae/oe vs e
    .
  2. ^ Along with rubidium (39 °C [102 °F]), francium (estimated at 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 [19.0 °F]), but it is a halogen and not a metal. Preliminary work with copernicium and flerovium suggests that they are gaseous metals at room temperature.
  3. ^ The radioactive element francium may also have a lower melting point, but its radioactivity prevents enough of it from being isolated for direct testing.[17] Copernicium and flerovium may also have lower melting points.
  4. ^ It differs from this value in caesides, which contain the Cs− anion and thus have caesium in the −1 oxidation state.[28] Additionally, 2013 calculations by Mao-sheng Miao indicate that under conditions of extreme pressure (greater than 30 GPa), the inner 5p electrons could form chemical bonds, where caesium would behave as the seventh 5p element. This discovery indicates that higher caesium fluorides with caesium in oxidation states from +2 to +6 could exist under such conditions.[29]
  5. periodic trends.[31]
  6. ^ 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.

References

  1. ^ "Standard Atomic Weights: Caesium". CIAAW. 2013.
  2. ISSN 1365-3075
    .
  3. ^
    ISBN 978-1-62708-155-9
    .
  4. ISBN 1-4398-5511-0
    .
  5. doi:10.1002/anie.197905871
    .
  6. ISBN 0-8493-0487-3
    . Retrieved 26 September 2010.
  7. doi:10.1088/1674-1137/abddae
    .
  8. NIST
    . Retrieved 13 March 2011.
  9. ^ "IUPAC Periodic Table of Elements". International Union of Pure and Applied Chemistry.
  10. ISBN 0-85404-438-8. pp. 248–49. Electronic version.
    .
  11. ISBN 978-0-8412-3999-9
    .
  12. S2CID 96729044. Archived
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