Since the 1990s, the largestapplication of the element has been ascaesium formate fordrilling fluids, but it has a range of applications in the production of electricity, in electronics, and in chemistry. The radioactive isotope caesium-137 has ahalf-life of about 30 years and is used in medical applications, industrial gauges, and hydrology.
Caesium is the spelling recommended by theInternational Union of Pure and Applied Chemistry (IUPAC).[10] TheAmerican Chemical Society (ACS) has used the spellingcesium since 1921,[11][12] followingWebster's New International Dictionary. The element was named after the Latin wordcaesius, meaning "bluish grey".[13] In medieval and early modern writingscaesius was spelled with theligatureæ ascæsius; hence, an alternative but now old-fashioned orthography iscæsium. More spelling explanation atae/oe vs e.
Of all elements that are solid at room temperature, caesium is the softest: it has a hardness ofMohs 0.2. It is a veryductile, pale metal, which darkens in the presence of trace amounts ofoxygen.[14][15][16] When in the presence ofmineral oil (where it is best kept during transport), it loses its metalliclustre and takes on a duller, grey appearance. It has amelting point of 28.5 °C (83.3 °F), making it one of the few elemental metals that are liquid nearroom temperature. The others arerubidium (39 °C [102 °F]),francium (estimated at 27 °C [81 °F]),mercury (−39 °C [−38 °F]), andgallium (30 °C [86 °F]); bromine is also liquid at room temperature (melting at −7.2 °C [19.0 °F]), but it is ahalogen and not a metal.Mercury is the only stable elemental metal with a known melting point lower than caesium.[17] In addition, caesium has a rather lowboiling point, 641 °C (1186 °F), thelowest of all stable metals other than mercury.[18]Copernicium andflerovium have been predicted to have lower boiling points than mercury and caesium, but they are extremely radioactive and it is not certain that they are metals.[19][20]
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, theplasmonic 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.[23] Its compounds burn with a blue[24][25] or violet[25] colour.
Addition of a small amount of caesium to cold water is explosive.
Caesium metal is highly reactive andpyrophoric. It ignites spontaneously in air, and reacts explosively with water even at low temperatures, more so than the otheralkali metals.[14] It reacts with ice at temperatures as low as −116 °C (−177 °F).[17] Because of this high reactivity, caesium metal is classified as ahazardous material. It is stored and shipped in dry, saturated hydrocarbons such asmineral oil. It can be handled only underinert gas, such asargon. However, a caesium-water explosion is often less powerful than asodium-water explosion with a similar amount of sodium. This is because caesium explodes instantly upon contact with water, leaving little time forhydrogen to accumulate.[27] Caesium can be stored in vacuum-sealedborosilicate glassampoules. 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 particularrubidium, the element above caesium in the periodic table.[28] As expected for an alkali metal, the only common oxidation state is +1. It differs from this value in caesides, which contain theCs− anion and thus have caesium in the −1 oxidation state.[5] Under conditions of extreme pressure (greater than 30 GPa), theoretical studies indicate that the inner 5p electrons could form chemical bonds, where caesium would behave as the seventh 5p element, suggesting that higher caesium fluorides with caesium in oxidation states from +2 to +6 could exist under such conditions.[29][30] Some slight differences arise from the fact that it has a higheratomic mass and is moreelectropositive than other (nonradioactive) alkali metals.[31] Caesium is the most electropositive chemical element.[17] The caesium ion is also larger andless "hard" than those of the lighteralkali metals.
Ball-and-stick model of the cubic coordination of Cs and Cl in CsCl
Most caesium compounds contain the element as thecationCs+ , whichbinds ionically to a wide variety ofanions. One noteworthy exception is thecaeside anion (Cs− ),[5] and others are the several suboxides (see section§ Oxides below). More recently, caesium is predicted to behave as ap-block element and capable of forming higher fluorides with higheroxidation states (i.e.,CsF n withn > 1) under high pressure.[32] This prediction needs to be validated by further experiments.[33]
Salts ofCs+ are usually colourless unless the anion itself is coloured. Many of the simple salts arehygroscopic, but less so than the corresponding salts of lighter alkali metals. Thephosphate,[34]acetate,carbonate,halides,oxide,nitrate, andsulfate salts are water-soluble. Itsdouble 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 asCsSbCl 4),bismuth,cadmium,copper,iron, andlead are also poorlysoluble.[14]
Astoichiometric mixture of caesium and gold will react to form yellowcaesium auride (Cs+ Au− ) upon heating. The auride anion here behaves as apseudohalogen. The compound reacts violently with water, yieldingcaesium hydroxide, metallic gold, and hydrogen gas; it dissolves in liquid ammonia and then can be reacted with a caesium-specific ion exchange resin to producetetramethylammonium auride. The analogousplatinum compound, red caesium platinide (Cs2Pt), contains the platinide ion that behaves as a pseudochalcogen.[36]
Like all metal cations,Cs+ forms complexes withLewis bases in solution. Because of its large size,Cs+ usually adoptscoordination 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 andsoftness (tendency to form covalent bonds) are properties exploited in separatingCs+ from other cations in the remediation of nuclear wastes, where137 Cs+ must be separated from large amounts of nonradioactiveK+ .[37]
Caesium chloride (CsCl) crystallizes in the simplecubic crystal system. Also called the "caesium chloride structure",[31] this structural motif is composed of aprimitive cubic lattice with a two-atom basis, each with an eightfoldcoordination; 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 withCsBr andCsI, and many other compounds that do not contain Cs. In contrast, most other alkaline halides have thesodium chloride (NaCl) structure.[31] The CsCl structure is preferred becauseCs+ has anionic radius of 174 pm andCl− 181 pm.[40]
More so than the other alkali metals, caesium forms numerous binary compounds withoxygen. When caesium burns in air, thesuperoxideCsO 2 is the main product.[41] The "normal"caesium oxide (Cs 2O) forms yellow-orangehexagonal 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 theperoxideCs 2O 2 at temperatures above 400 °C (752 °F). In addition to the superoxide and theozonideCsO 3,[44][45] several brightly colouredsuboxides have also been studied.[46] These includeCs 7O,Cs 4O,Cs 11O 3,Cs 3O (dark-green[47]), CsO,Cs 3O 2,[48] as well asCs 7O 2.[49][50] The latter may be heated in a vacuum to generateCs 2O.[43] Binary compounds withsulfur,selenium, andtellurium also exist.[14]
Caesium has 41 knownisotopes, ranging inmass number from 112 to 152. The onlystable caesium isotope is133 Cs, with 78neutrons. The radioactive135Cs has a very long half-life of about 1.33 million years, the longest of all radioactive isotopes of caesium.137Cs and134Cs have half-lives of 30.04 and 2.065 years, respectively. 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 metastablenuclear isomers exist. Other than134mCs (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]
Thebeta decay from137Cs to137mBa results ingamma radiation as the137mBa relaxes to ground state137Ba, with the emitted photons having an energy of 0.6617 MeV.[55]137 Cs and90Sr are the principalmedium-lived products ofnuclear fission, and the prime sources ofradioactivity fromspent nuclear fuel from several years to several hundred years after removal.[56] Those two isotopes are the largest source of residual radioactivity in the area of theChernobyl disaster.[57] Because of the low capture rate, disposing of137 Cs throughneutron capture is not feasible and the only current solution is to allow it to decay over time.[58]
Almost all caesium produced from nuclear fission comes from thebeta decay of originally more neutron-rich fission products, passing through variousisotopes of iodine andxenon.[59] 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.[60] Withnuclear weapons testing in the 1950s through the 1980s,137 Cs was released into theatmosphere and returned to the surface of the earth as a component ofradioactive fallout, becoming a marker of the movement of soil and sediment from those times.[14]
Caesium is a relatively rare element, estimated to average 3 parts per million in theEarth's crust.[62] It is the 45th most abundant element and 36th among the metals.[63] Caesium is 30 times less abundant thanrubidium, with which it is closely associated, chemically.[14]
Due to its largeionic radius, caesium is one of the "incompatible elements".[64] Duringmagma crystallization, caesium is concentrated in the liquid phase and crystallizes last. Therefore, the largest deposits of caesium are zonepegmatite ore bodies formed by this enrichment process. Because caesium does not substitute forpotassium as readily as rubidium does, the alkali evaporite mineralssylvite (KCl) andcarnallite (KMgCl 3·6H 2O) may contain only 0.002% caesium. Consequently, caesium is found in few minerals. Percentage amounts of caesium may be found inberyl (Be 3Al 2(SiO 3) 6) andavogadrite ((K,Cs)BF 4), up to 15 wt% Cs2O in the closely related mineralpezzottaite (Cs(Be 2Li)Al 2Si 6O 18), up to 8.4 wt% Cs2O in the rare minerallondonite ((Cs,K)Al 4Be 4(B,Be) 12O 28), and less in the more widespreadrhodizite.[14] The only economically important ore for caesium ispolluciteCs(AlSi 2O 6), which is found in a few places around the world in zoned pegmatites, associated with the more commercially importantlithium minerals,lepidolite andpetalite. Within the pegmatites, the large grain size and the strong separation of the minerals results in high-grade ore for mining.[65]
The world's most significant and richest known source of caesium is theTanco Mine atBernic Lake inManitoba, Canada, estimated to contain 350,000 metric tons of pollucite ore, representing more than two-thirds of the world's reserve base.[65][66] 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%.[66] Commercial pollucite contains more than 19% caesium.[67] TheBikita pegmatite deposit inZimbabwe is mined for its petalite, but it also contains a significant amount of pollucite. Another notable source of pollucite is in theKaribib Desert,Namibia.[66] At the present rate of world mine production of 5 to 10 metric tons per year, reserves will last for thousands of years.[14]
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][68]
In the acid digestion, thesilicate pollucite rock is dissolved with strong acids, such ashydrochloric (HCl),sulfuric (H 2SO 4),hydrobromic (HBr), orhydrofluoric (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 caesiumalum (CsAl(SO 4) 2·12H 2O). Thealuminium sulfate component is converted to insolublealuminium oxide by roasting the alum withcarbon, and the resulting product isleached with water to yield aCs 2SO 4 solution.[14]
Roasting pollucite withcalcium carbonate andcalcium chloride yields insoluble calcium silicates and soluble caesium chloride. Leaching with water or diluteammonia (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 intocaesium formate (HCOO− Cs+ ) for applications such asoil drilling. To supply the developing market,Cabot Corporation built a production plant in 1997 at theTanco mine nearBernic Lake inManitoba, with a capacity of 12,000 barrels (1,900 m3) per year of caesium formate solution.[69] The primary smaller-scale commercial compounds of caesium arecaesium chloride andnitrate.[70]
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 orbarium, and caesium metal distilled from the result. In the same way, the aluminate, carbonate, or hydroxide may be reduced bymagnesium.[14]
The metal can also be isolated byelectrolysis of fused caesiumcyanide (CsCN). Exceptionally pure and gas-free caesium can be produced by 390 °C (734 °F) thermal decomposition of caesiumazideCsN 3, which can be produced from aqueouscaesium sulfate andbarium azide.[68] In vacuum applications, caesiumdichromate can be reacted withzirconium to produce pure caesium metal without other gaseous products.[70]
Cs 2Cr 2O 7 + 2Zr → 2Cs + 2ZrO 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.[66]
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. Thealkaline earth metals were precipitated either as sulfates oroxalates, leaving the alkali metal in the solution. After conversion to thenitrates and extraction withethanol, a sodium-free mixture was obtained. From this mixture, the lithium was precipitated byammonium carbonate. Potassium, rubidium, and caesium form insoluble salts withchloroplatinic 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 byfractional crystallization. After reduction of the hexachloroplatinate withhydrogen, caesium and rubidium were separated by the difference in solubility of their carbonates in alcohol. The process yielded 9.2 grams (0.32 oz) ofrubidium chloride and 7.3 grams (0.26 oz) of caesium chloride from the initial 44,000 litres of mineral water.[72]
From the caesium chloride, the two scientists estimated theatomic weight of the new element at 123.35 (compared to the currently accepted one of 132.9).[72] They tried to generate elemental caesium by electrolysis of molten caesium chloride, but instead of a metal, they obtained a blue homogeneous substance which "neither under the naked eye nor under the microscope showed the slightest trace of metallic substance"; as a result, they assigned it as asubchloride (Cs 2Cl). In reality, the product was probably acolloidal mixture of the metal and caesium chloride.[74] The electrolysis of the aqueous solution of chloride with a mercury cathode produced a caesium amalgam which readily decomposed under the aqueous conditions.[72] The pure metal was eventually isolated by the Swedish chemistCarl Setterberg while working on his doctorate withKekulé and Bunsen.[73] In 1882, he produced caesium metal by electrolysingcaesium cyanide, avoiding the problems with the chloride.[75]
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 radiovacuum tubes, where it had two functions; as agetter, it removed excess oxygen after manufacture, and as a coating on the heatedcathode, it increased theelectrical conductivity. Caesium was not recognized as a high-performance industrial metal until the 1950s.[76] Applications for nonradioactive caesium includedphotoelectric cells,photomultiplier tubes, optical components ofinfrared spectrophotometers, catalysts for several organic reactions, crystals forscintillation counters, and inmagnetohydrodynamic power generators.[14] Caesium is also used as a source of positive ions insecondary ion mass spectrometry (SIMS).
The largest present-day use of nonradioactive caesium is incaesium formatedrilling fluids for theextractive oil industry.[14] Aqueous solutions of caesium formate (HCOO− Cs+ )—made by reacting caesium hydroxide withformic acid—were developed in the mid-1990s for use as oil well drilling andcompletion 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),[78] 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.[78] 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 perbarrel in 2001).[79] Alkali formates are safe to handle and do not damage the producing formation or downhole metals as corrosive alternative, high-density brines (such aszinc bromideZnBr 2 solutions) sometimes do; they also require less cleanup and reduce disposal costs.[14]
Atomic clock ensemble at the U.S. Naval Observatory
Caesium-basedatomic clocks use theelectromagnetic transitions in thehyperfine structure of caesium-133 atoms as a reference point. The first accurate caesium clock was built byLouis Essen in 1955 at theNational Physical Laboratory in the UK.[80] 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."[77] 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] Thecaesium standard is the primary standard for standards-compliant time and frequency measurements.[81] Caesium clocks regulate the timing of cell phone networks and the Internet.[82]
The second, symbol s, is the SI unit of time. TheBIPM restated its definition at its 26th conference in 2018: "[The second] is defined by taking the fixed numerical value of the caesium frequencyΔνCs, the unperturbed ground-state hyperfine transition frequency of the caesium-133 atom, to be9192631770 when expressed in the unitHz, which is equal to s−1."[83]
Caesium vapourthermionic generators are low-power devices that convert heat energy to electrical energy. In the two-electrodevacuum tube converter, caesium neutralizes the space charge near the cathode and enhances the current flow.[84]
Caesium is also important for itsphotoemissive properties, converting light to electron flow. It is used inphotoelectric cells because caesium-based cathodes, such as the intermetallic compoundK 2CsSb, have a low threshold voltage for emission ofelectrons.[85] The range of photoemissive devices using caesium includeoptical character recognition devices,photomultiplier tubes, andvideo camera tubes.[86][87] Nevertheless,germanium, rubidium, selenium, silicon, tellurium, and several other elements can be substituted for caesium in photosensitive materials.[14]
Caesium iodide (CsI),bromide (CsBr) andfluoride (CsF) crystals are employed forscintillators inscintillation counters widely used in mineral exploration and particle physics research to detectgamma andX-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).
The high density of the caesium ion makes solutions of caesium chloride, caesium sulfate, and caesiumtrifluoroacetate (Cs(O 2CCF 3)) useful in molecular biology for density gradientultracentrifugation.[91] This technology is used primarily in the isolation ofviral particles, subcellularorganelles and fractions, andnucleic acids from biological samples.[92]
Caesium-137 is aradioisotope commonly used as agamma-emitter in industrial applications. Its advantages include a half-life of roughly 30 years, its availability from thenuclear fuel cycle, and having137Ba 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.[95] It has been used in agriculture, cancer treatment, and thesterilization of food, sewage sludge, and surgical equipment.[14][96] Radioactiveisotopes of caesium inradiation devices were used in the medical field to treat certain types of cancer,[97] 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.[98][99] Caesium-137 has been employed in a variety of industrial measurement gauges, including moisture, density, levelling, and thickness gauges.[100] It has also been used inwell logging devices for measuring theelectron density of the rock formations, which is analogous to the bulk density of the formations.[101]
Caesium-137 has been used inhydrologic studies analogous to those withtritium. 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.[102]
Schematics of an electrostatic ion thruster developed for use with caesium or mercury fuel
Caesium and mercury were used as a propellant in earlyion engines designed forspacecraft propulsion on very long interplanetary or extraplanetary missions. The fuel was ionized by contact with a chargedtungsten electrode. But corrosion by caesium on spacecraft components has pushed development in the direction of inert gas propellants, such asxenon, which are easier to handle in ground-based tests and do less potential damage to the spacecraft.[14] Xenon was used in the experimental spacecraftDeep Space 1 launched in 1998.[103][104] Nevertheless,field-emission electric propulsion thrusters that accelerate liquid metal ions such as caesium have been built.[105]
Magnetohydrodynamic (MHD) power-generating systems were researched, but failed to gain widespread acceptance.[110] Caesium metal has also been considered as the working fluid in high-temperatureRankine cycle turboelectric generators.[111]
Caesium salts have been evaluated as antishock reagents following the administration ofarsenical 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 treatepilepsy.[14]
The portion of the total radiation dose (in air) contributed by each isotope plotted against time after theChernobyl disaster. Caesium-137 became the primary source of radiation about 200 days after the accident.[114]
Nonradioactive caesium compounds are only mildly toxic, and nonradioactive caesium is not a significant environmental hazard. Because biochemical processes can confuse and substitute caesium withpotassium, excess caesium can lead tohypokalemia,arrhythmia, and acutecardiac arrest, but such amounts would not ordinarily be encountered in natural sources.[115][116]
Elemental caesium is one of the most reactive elements and is highlyexplosive 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 highlypyrophoric: theautoignition temperature of caesium is −116 °C (−177 °F), and it ignites explosively in air to formcaesium hydroxide and various oxides. Caesium hydroxide is a very strongbase, and will rapidly corrode glass.[18]
Theisotopes134 and 137 are present in thebiosphere in small amounts from human activities, differing by location. Radiocaesium does not accumulate in the body as readily as other fission products (such as radioiodine and radiostrontium). About 10% of absorbed radiocaesium washes out of the body relatively quickly in sweat and urine. The remaining 90% has a biological half-life between 50 and 150 days.[118] Radiocaesium follows potassium and tends to accumulate in plant tissues, including fruits and vegetables.[119][120][121] Plants vary widely in the absorption of caesium, sometimes displaying great resistance to it. It is also well-documented that mushrooms from contaminated forests accumulate radiocaesium (caesium-137) in the fungalsporocarps.[122] Accumulation of caesium-137 in lakes has been a great concern after theChernobyl disaster.[123][124] Experiments with dogs showed that a single dose of 3.8millicuries (140 MBq, 4.1 μg of caesium-137) per kilogram is lethal within three weeks;[125] smaller amounts may cause infertility and cancer.[126] TheInternational 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".[127]
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