Caesium

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 ductile, pale metal, which darkens in the presence of trace amounts of oxygen.[12][13][14] When in the presence of mineral oil (where it is best kept during transport), it loses its metallic lustre and takes on a duller, grey appearance. It has a melting point of 28.5 °C (83.3 °F), making it one of the few elemental metals that are liquid near room temperature. Mercury is the only stable elemental metal with a known melting point lower than caesium.[note 3][16] In addition, the metal has a rather low boiling point, 641 °C (1,186 °F), the lowest of all metals other than mercury.[17] Its compounds burn with a blue[18][19] or violet[19] colour.

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

Caesium forms alloys with the other alkali metals, gold, and mercury (amalgams). At temperatures below 650 °C (1,202 °F), it does not alloy with cobalt, iron, molybdenum, nickel, platinum, tantalum, or tungsten. It forms well-defined intermetallic compounds with antimony, gallium, indium, and thorium, which are photosensitive.[12] 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).[16][20] 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.[21]

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.[22]

Caesium metal is highly reactive and pyrophoric. It ignites spontaneously in air, and reacts explosively with water even at low temperatures, more so than the other alkali metals.[12] It reacts with ice at temperatures as low as −116 °C (−177 °F).[16] Because of this high reactivity, caesium metal is classified as a 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.[23] 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.[12]

The chemistry of caesium is similar to that of other alkali metals, in particular rubidium, the element above caesium in the periodic table.[24] As expected for an alkali metal, the only common oxidation state is +1.[note 4] Some slight differences arise from the fact that it has a higher atomic mass and is more electropositive than other (nonradioactive) alkali metals.[27] Caesium is the most electropositive chemical element.[note 5][16] The caesium ion is also larger and less "hard" than those of the lighter alkali metals.

Ball-and-stick model of the cubic coordination of Cs and Cl in CsCl

Most caesium compounds contain the element as the cation Cs⁺
, which binds ionically to a wide variety of anions. One noteworthy exception is the caeside anion (Cs⁻
),[25] and others are the several suboxides (see section on oxides below). More recently, caesium is predicted to behave as a p-block element and capable of forming higher fluorides with higher oxidation states (i.e., CsF_n with n > 1) under high pressure.[29] This prediction needs to be validated by further experiments.[30]

Salts of Cs⁺ are usually colourless unless the anion itself is coloured. Many of the simple salts are hygroscopic, but less so than the corresponding salts of lighter alkali metals. The phosphate,[31] acetate, carbonate, halides, oxide, nitrate, and sulfate salts are water-soluble. Its double salts are often less soluble, and the low solubility of caesium aluminium sulfate is exploited in refining Cs from ores. The double salts with antimony (such as CsSbCl
₄), bismuth, cadmium, copper, iron, and lead are also poorly soluble.[12]

Caesium hydroxide (CsOH) is hygroscopic and strongly basic.[24] It rapidly etches the surface of semiconductors such as silicon.[32] CsOH has been previously regarded by chemists as the "strongest base", reflecting the relatively weak attraction between the large Cs⁺ ion and OH⁻;[18] it is indeed the strongest Arrhenius base; however, a number of compounds such as n-butyllithium, sodium amide, sodium hydride, caesium hydride, etc., which cannot be dissolved in water as reacting violently with it but rather only used in some anhydrous polar aprotic solvents, are far more basic on the basis of the Brønsted–Lowry acid–base theory.[24]

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.[33]

Like all metal cations, Cs⁺ forms complexes with Lewis bases in solution. Because of its large size, Cs⁺ usually adopts coordination numbers greater than 6, the number typical for the smaller alkali metal cations. This difference is apparent in the 8-coordination of CsCl. This high coordination number and 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⁺.[34]

Monatomic caesium halide wires grown inside double-wall carbon nanotubes (TEM image).[35]

Caesium fluoride (CsF) is a hygroscopic white solid that is widely used in organofluorine chemistry as a source of fluoride anions.[36] Caesium fluoride has the halite structure, which means that the Cs⁺ and F⁻ pack in a cubic closest packed array as do Na⁺ and Cl⁻ in sodium chloride.[24] Notably, caesium and fluorine have the lowest and highest electronegativities, respectively, among all the known elements.

Caesium chloride (CsCl) crystallizes in the simple cubic crystal system. Also called the "caesium chloride structure",[27] this structural motif 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 centre of the cubes. This structure is shared with CsBr and CsI, and many other compounds that do not contain Cs. In contrast, most other alkaline halides have the sodium chloride (NaCl) structure.[27] The CsCl structure is preferred because Cs⁺ has an ionic radius of 174 pm and Cl⁻
181 pm.[37]

Cs
₁₁O
₃ cluster

More so than the other alkali metals, caesium forms numerous binary compounds with oxygen. When caesium burns in air, the superoxide CsO
₂ is the main product.[38] The "normal" caesium oxide (Cs
₂O) forms yellow-orange hexagonal crystals,[39] and is the only oxide of the anti-CdCl
₂ type.[40] 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,[41][42] several brightly coloured suboxides have also been studied.[43] These include Cs
₇O, Cs
₄O, Cs
₁₁O
₃, Cs
₃O (dark-green[44]), CsO, Cs
₃O
₂,[45] as well as Cs
₇O
₂.[46][47] The latter may be heated in a vacuum to generate Cs
₂O.[40] Binary compounds with sulfur, selenium, and tellurium also exist.[12]

Caesium has 40 known isotopes, ranging in mass number (i.e. number of nucleons in the nucleus) from 112 to 151. Several of these are synthesized from lighter elements by the slow neutron capture process (S-process) inside old stars[48] and by the R-process in supernova explosions.[49] The only stable caesium isotope is ¹³³Cs, with 78 neutrons. Although it has a large nuclear spin (7/2+), nuclear magnetic resonance studies can use this isotope at a resonating frequency of 11.7 MHz.[50]

Decay of caesium-137

The radioactive ¹³⁵Cs has a very long half-life of about 2.3 million years, the longest of all radioactive isotopes of caesium. ¹³⁷Cs and ¹³⁴Cs have half-lives of 30 and two years, respectively. ¹³⁷Cs decomposes to a short-lived ^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.[51][52]

The isotope ¹³⁵Cs is one of the long-lived fission products of uranium produced in nuclear reactors.[53] However, this fission product yield is reduced in most reactors because the predecessor, ¹³⁵Xe, is a potent neutron poison and frequently transmutes to stable ¹³⁶Xe before it can decay to ¹³⁵Cs.[54][55]

The beta decay from ¹³⁷Cs to ^137mBa results in gamma radiation as the ^137mBa relaxes to ground state ¹³⁷Ba, with the emitted photons having an energy of 0.6617 MeV.[56] ¹³⁷Cs and ⁹⁰Sr are the principal 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.[57] Those two isotopes are the largest source of residual radioactivity in the area of the Chernobyl disaster.[58] 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.[59]

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.[60] 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.[61] 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.[12]

Pollucite, a caesium mineral

Caesium is a relatively rare element, estimated to average 3 parts per million in the Earth's crust.[62] It is the 45th most abundant element and the 36th among 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 and silver; it is 3.3% as abundant as rubidium, with which it is closely associated, chemically.[12]

Due to its large ionic radius, caesium is one of the "incompatible elements".[63] During magma crystallization, caesium 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. Because caesium does not substitute for potassium as readily as rubidium does, the alkali evaporite minerals sylvite (KCl) and carnallite (KMgCl
₃·6H
₂O) may contain only 0.002% caesium. Consequently, caesium is found in few minerals. Percentage amounts of caesium may be found in beryl (Be
₃Al
₂(SiO
₃)
₆) and avogadrite ((K,Cs)BF
₄), up to 15 wt% Cs₂O in the closely related mineral pezzottaite (Cs(Be
₂Li)Al
₂Si
₆O
₁₈), up to 8.4 wt% Cs₂O in the rare mineral londonite ((Cs,K)Al
₄Be
₄(B,Be)
₁₂O
₂₈), and less in the more widespread rhodizite.[12] 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.[64]

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.[64][65] 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%.[65] Commercial pollucite contains more than 19% caesium.[66] 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.[65] At the present rate of world mine production of 5 to 10 metric tons per year, reserves will last for thousands of years.[12]

The largest present-day use of nonradioactive caesium is in caesium formate drilling fluids for the extractive oil industry.[12] 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.[12]

The high density of the caesium formate brine (up to 2.3 g/cm³, or 19.2 pounds per gallon),[77] 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.[77] Caesium formate brine can be blended with potassium and sodium formates to decrease the density of the fluids to that of water (1.0 g/cm³, or 8.3 pounds per gallon). Furthermore, it is biodegradable and may be recycled, which is important in view of its high cost (about $4,000 per barrel in 2001).[78] 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.[12]

Atomic clock ensemble at the U.S. Naval Observatory

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

Caesium-based atomic clocks use the electromagnetic transitions in the hyperfine structure of caesium-133 atoms as a reference point. The first accurate caesium clock was built by Louis Essen in 1955 at the National Physical Laboratory in the UK.[79] 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."[76] 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.[12] The caesium standard is the primary standard for standards-compliant time and frequency measurements.[80] Caesium clocks regulate the timing of cell phone networks and the Internet.[81]

The second, symbol s, is the SI unit of time. The BIPM restated its definition at its 26th conference in 2018: "[The second] is defined by taking the fixed numerical value of the caesium frequency Δν_Cs, the unperturbed ground-state hyperfine transition frequency of the caesium-133 atom, to be 9192631770 when expressed in the unit Hz, which is equal to s⁻¹."[82]

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.[83]

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.[84] The range of photoemissive devices using caesium include optical character recognition devices, photomultiplier tubes, and video camera tubes.[85][86] Nevertheless, germanium, rubidium, selenium, silicon, tellurium, and several other elements can be substituted for caesium in photosensitive materials.[12]

Caesium iodide (CsI), bromide (CsBr) and caesium 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.[87]

The element is used as an internal standard in spectrophotometry.[88] Like other alkali metals, caesium has a great affinity for oxygen and is used as a "getter" in vacuum tubes.[89] Other uses of the metal include high-energy lasers, vapour glow lamps, and vapour rectifiers.[12]

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.[90] This technology is used primarily in the isolation of viral particles, subcellular organelles and fractions, and nucleic acids from biological samples.[91]

Caesium chloride powder

Relatively few chemical applications use caesium.[92] 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.[12]

Caesium fluoride enjoys a niche use in organic chemistry as a base[24] and as an anhydrous source of fluoride ion.[93] Caesium salts sometimes replace potassium or sodium salts in organic synthesis, such as cyclization, 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.

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.[94] It has been used in agriculture, cancer treatment, and the sterilization of food, sewage sludge, and surgical equipment.[12][95] Radioactive isotopes of caesium in radiation devices were used in the medical field to treat certain types of cancer,[96] 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.[97][98] Caesium-137 has been employed in a variety of industrial measurement gauges, including moisture, density, levelling, and thickness gauges.[99] 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.[100]

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.[101]

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.[12] Xenon was used in the experimental spacecraft Deep Space 1 launched in 1998.[102][103] Nevertheless, field-emission electric propulsion thrusters that accelerate liquid metal ions such as caesium have been built.[104]

Caesium nitrate is used as an oxidizer and pyrotechnic colorant to burn silicon in infrared flares,[105] such as the LUU-19 flare,[106] because it emits much of its light in the near infrared spectrum.[107] Caesium compounds may have been used as fuel additives to reduce the radar signature of exhaust plumes in the Lockheed A-12 CIA reconnaissance aircraft.[108] 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.[12]

Magnetohydrodynamic (MHD) power-generating systems were researched, but failed to gain widespread acceptance.[109] Caesium metal has also been considered as the working fluid in high-temperature Rankine cycle turboelectric generators.[110]

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.[12]

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.[111]