Chemistry:Caesium chloride

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Caesium chloride
Caesium chloride.jpg
CsCl polyhedra.png
Caesium-chloride-3D-ionic.png
Names
IUPAC name
Caesium chloride
Other names
Cesium chloride
Identifiers
3D model (JSmol)
ChemSpider
EC Number
  • 231-600-2
UNII
Properties
CsCl
Molar mass 168.36 g/mol
Appearance white solid
hygroscopic
Density 3.988 g/cm3[1]
Melting point 646 °C (1,195 °F; 919 K)[1]
Boiling point 1,297 °C (2,367 °F; 1,570 K)[1]
1910 g/L (25 °C)[1]
Solubility soluble in ethanol[1]
Band gap 8.35 eV (80 K)[2]
-56.7·10−6 cm3/mol[3]
1.712 (0.3 μm)
1.640 (0.59 μm)
1.631 (0.75 μm)
1.626 (1 μm)
1.616 (5 μm)
1.563 (20 μm)[4]
Structure
CsCl, cP2
Pm3m, No. 221[5]
a = 0.4119 nm
0.0699 nm3
1
Cubic (Cs+)
Cubic (Cl−)
Hazards
GHS pictograms GHS07: HarmfulGHS08: Health hazard
GHS Signal word Warning
H302, H341, H361, H373
P201, P202, P260, P264, P270, P281, P301+312, P308+313, P314, P330, P405, P501
Lethal dose or concentration (LD, LC):
2600 mg/kg (oral, rat)[6]
Related compounds
Other anions
 Caesium fluoride
Caesium bromide
Caesium iodide
Caesium astatide
Other cations
 Lithium chloride
Sodium chloride
Potassium chloride
Rubidium chloride
Francium chloride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references
Tracking categories (test):

Caesium chloride or cesium chloride is the inorganic compound with the formula CsCl. This colorless salt is an important source of caesium ions in a variety of niche applications. Its crystal structure forms a major structural type where each caesium ion is coordinated by 8 chloride ions. Caesium chloride dissolves in water. CsCl changes to NaCl structure on heating. Caesium chloride occurs naturally as impurities in carnallite (up to 0.002%), sylvite and kainite. Less than 20 tonnes of CsCl is produced annually worldwide, mostly from a caesium-bearing mineral pollucite.[7]

Caesium chloride is widely used medicine structure in isopycnic centrifugation for separating various types of DNA. It is a reagent in analytical chemistry, where it is used to identify ions by the color and morphology of the precipitate. When enriched in radioisotopes, such as 137CsCl or 131CsCl, caesium chloride is used in nuclear medicine applications such as treatment of cancer and diagnosis of myocardial infarction. Another form of cancer treatment was studied using conventional non-radioactive CsCl. Whereas conventional caesium chloride has a rather low toxicity to humans and animals, the radioactive form easily contaminates the environment due to the high solubility of CsCl in water. Spread of 137CsCl powder from a 93-gram container in 1987 in Goiânia, Brazil, resulted in one of the worst-ever radiation spill accidents killing four and directly affecting 249 people.

Crystal structure

Main page: Cubic crystal system

The caesium chloride structure adopts a primitive cubic lattice with a two-atom basis, where both atoms have eightfold coordination. The chloride atoms lie upon the lattice points at the corners of the cube, while the caesium atoms lie in the holes in the center of the cubes; an alternative and exactly equivalent 'setting' has the caesium ions at the corners and the chloride ion in the center. This structure is shared with CsBr and CsI and many binary metallic alloys. In contrast, the other alkaline halides have the sodium chloride (rocksalt) structure.[8] When both ions are similar in size (Cs+ ionic radius 174 pm for this coordination number, Cl− 181 pm) the CsCl structure is adopted, when they are different (Na+ ionic radius 102 pm, Cl− 181 pm) the sodium chloride structure is adopted. Upon heating to above 445 °C, the normal caesium chloride structure (α-CsCl) converts to the β-CsCl form with the rocksalt structure (space group Fm3m).[5] The rocksalt structure is also observed at ambient conditions in nanometer-thin CsCl films grown on mica, LiF, KBr and NaCl substrates.[9]

Physical properties

Caesium chloride is colorless in the form of large crystals and white when powdered. It readily dissolves in water with the maximum solubility increasing from 1865 g/L at 20 °C to 2705 g/L at 100 °C.[10] The crystals are very hygroscopic and gradually disintegrate at ambient conditions.[11] Caesium chloride does not form hydrates.[12]

Solubility of CsCl in water[13]
Т (°C) 0 10 20 25 30 40 50 60 70 80 90 100
S (wt%) 61.83 63.48 64.96 65.64 66.29 67.50 68.60 69.61 70.54 71.40 72.21 72.96

In contrast to sodium chloride and potassium chloride, caesium chloride readily dissolves in concentrated hydrochloric acid.[14][15] Caesium chloride has also a relatively high solubility in formic acid (1077 g/L at 18 °C) and hydrazine; medium solubility in methanol (31.7 g/L at 25 °C) and low solubility in ethanol (7.6 g/L at 25 °C),[12][15][16] sulfur dioxide (2.95 g/L at 25 °C), ammonia (3.8 g/L at 0 °C), acetone (0.004% at 18 °C), acetonitrile (0.083 g/L at 18 °C),[15] ethylacetate and other complex ethers, butanone, acetophenone, pyridine and chlorobenzene.[17]

Despite its wide band gap of about 8.35 eV at 80 K,[2] caesium chloride weakly conducts electricity, and the conductivity is not electronic but ionic. The conductivity has a value of the order 10−7 S/cm at 300 °C. It occurs through nearest-neighbor jumps of lattice vacancies, and the mobility is much higher for the Cl− than Cs+ vacancies. The conductivity increases with temperature up to about 450 °C, with an activation energy changing from 0.6 to 1.3 eV at about 260 °C. It then sharply drops by two orders of magnitude because of the phase transition from the α-CsCl to β-CsCl phase. The conductivity is also suppressed by application of pressure (about 10 times decrease at 0.4 GPa) which reduces the mobility of lattice vacancies.[18]

Reactions

Caesium chloride completely dissociates upon dissolution in water, and the Cs+ cations are solvated in dilute solution. CsCl converts to caesium sulfate upon being heated in concentrated sulfuric acid or heated with caesium hydrogen sulfate at 550–700 °C:[21]

2 CsCl + H2SO4 → Cs2SO4 + 2 HCl
CsCl + CsHSO4 → Cs2SO4 + HCl

Caesium chloride forms a variety of double salts with other chlorides. Examples include 2CsCl·BaCl2,[22] 2CsCl·CuCl2, CsCl·2CuCl and CsCl·LiCl,[23] and with interhalogen compounds:[24]

[math]\ce{ CsCl + ICl3 -> Cs[ICl4] }[/math]

Occurrence and production

Monatomic caesium halide wires grown inside double-wall carbon nanotubes.[25]

Caesium chloride occurs naturally as an impurity in the halide minerals carnallite (KMgCl3·6H2O with up to 0.002% CsCl),[26] sylvite (KCl) and kainite (MgSO4·KCl·3H2O),[27] and in mineral waters. For example, the water of Bad Dürkheim spa, which was used in isolation of caesium, contained about 0.17 mg/L of CsCl.[28] None of these minerals are commercially important.

On industrial scale, CsCl is produced from the mineral pollucite, which is powdered and treated with hydrochloric acid at elevated temperature. The extract is treated with antimony chloride, iodine monochloride, or cerium(IV) chloride to give the poorly soluble double salt, e.g.:[29]

CsCl + SbCl3 → CsSbCl4

Treatment of the double salt with hydrogen sulfide gives CsCl:[29]

2 CsSbCl4 + 3 H2S → 2 CsCl + Sb2S3 + 8 HCl

High-purity CsCl is also produced from recrystallized <ce>Cs[ICl2]</ce> (and <ce>Cs[ICl4]</ce>) by thermal decomposition:[30]

Only about 20 tonnes of caesium compounds, with a major contribution from CsCl, were being produced annually around the 1970s[31] and 2000s worldwide.[32] Caesium chloride enriched with caesium-137 for radiation therapy applications is produced at a single facility Mayak in the Ural Region of Russia[33] and is sold internationally through a UK dealer. The salt is synthesized at 200 °C because of its hygroscopic nature and sealed in a thimble-shaped steel container which is then enclosed into another steel casing. The sealing is required to protect the salt from moisture.[34]

Laboratory methods

In the laboratory, CsCl can be obtained by treating caesium hydroxide, carbonate, caesium bicarbonate, or caesium sulfide with hydrochloric acid:

CsOH + HCl → CsCl + H2O
Cs2CO3 + 2 HCl → 2 CsCl + 2 H2O + CO2

Uses

Precursor to Cs metal

Caesium chloride is the main precursor to caesium metal by high-temperature reduction:[31]

2 CsCl (l) + Mg (l) → MgCl2 (s) + 2 Cs (g)

A similar reaction – heating CsCl with calcium in vacuum in presence of phosphorus – was first reported in 1905 by the French chemist M. L. Hackspill[35] and is still used industrially.[31]

Caesium hydroxide is obtained by electrolysis of aqueous caesium chloride solution:[36]

2 CsCl + 2 H2O → 2 CsOH + Cl2 + H2

Solute for ultracentrifugation

Caesium chloride is widely used in centrifugation in a technique known as isopycnic centrifugation. Centripetal and diffusive forces establish a density gradient that allow separation of mixtures on the basis of their molecular density. This technique allows separation of DNA of different densities (e.g. DNA fragments with differing A-T or G-C content).[31] This application requires a solution with high density and yet relatively low viscosity, and CsCl suits it because of its high solubility in water, high density owing to the large mass of Cs, as well as low viscosity and high stability of CsCl solutions.[29]

Organic chemistry

Caesium chloride is rarely used in organic chemistry. It can act as a phase transfer catalyst reagent in selected reactions. One of these reactions is the synthesis of glutamic acid derivatives

[math]\displaystyle{ \overbrace{\ce{CH2=CHCOOCH3}}^\text{Methyl acrylate} + \ce{ArCH=N-CH(CH3)-COOC(CH3)3 -\gt [\ce{TBAB,\ CsCl,\ K2CO3}][\ce{CPME,\ 0^\circ C}] {ArCH=N-C(C2H4COOCH3)(CH3)-COOC(CH3)3}} }[/math]

where TBAB is tetrabutylammonium bromide (interphase catalyst) and CPME is a cyclopentyl methyl ether (solvent).[37]

Another reaction is substitution of tetranitromethane[38]

[math]\displaystyle{ \overbrace{\ce{C(NO2)4}}^\text{tetranitromethane} + \ce{CsCl -\gt [\ce{DMF}] {C(NO2)3Cl} + CsNO2} }[/math]

where DMF is dimethylformamide (solvent).

Analytical chemistry

Caesium chloride is a reagent in traditional analytical chemistry used for detecting inorganic ions via the color and morphology of the precipitates. Quantitative concentration measurement of some of these ions, e.g. Mg2+, with inductively coupled plasma mass spectrometry, is used to evaluate the hardness of water.[39]

It is also used for detection of the following ions:

Ion Accompanying reagents Detection Detection limit (μg/mL)
Al3+ K2SO4 Colorless crystals form in neutral media after evaporation 0.01
Ga3+ KHSO4 Colorless crystals form upon heating 0.5
Cr3+ KHSO4 Pale-violet crystals precipitate in slightly acidic media 0.06

Medicine

The American Cancer Society states that "available scientific evidence does not support claims that non-radioactive cesium chloride supplements have any effect on tumors."[40] The Food and Drug Administration has warned about safety risks, including significant heart toxicity and death, associated with the use of cesium chloride in naturopathic medicine.[41][42]

Nuclear medicine and radiography

Caesium chloride composed of radioisotopes such as 137CsCl and 131CsCl,[43] is used in nuclear medicine, including treatment of cancer (brachytherapy) and diagnosis of myocardial infarction.[44][45] In the production of radioactive sources, it is normal to choose a chemical form of the radioisotope which would not be readily dispersed in the environment in the event of an accident. For instance, radiothermal generators (RTGs) often use strontium titanate, which is insoluble in water. For teletherapy sources, however, the radioactive density (Ci in a given volume) needs to be very high, which is not possible with known insoluble caesium compounds. A thimble-shaped container of radioactive caesium chloride provides the active source.

Miscellaneous applications

Caesium chloride is used in the preparation of electrically conducting glasses[43][46] and screens of cathode ray tubes.[31] In conjunction with rare gases CsCl is used in excimer lamps[47][48] and excimer lasers. Other uses include activation of electrodes in welding;[49] manufacture of mineral water, beer[50] and drilling muds;[51] and high-temperature solders.[52] High-quality CsCl single crystals have a wide transparency range from UV to the infrared and therefore had been used for cuvettes, prisms and windows in optical spectrometers;[31] this use was discontinued with the development of less hygroscopic materials.

CsCl is a potent inhibitor of HCN channels, which carry the h-current in excitable cells such as neurons.[53] Therefore, it can be useful in electrophysiology experiments in neuroscience.

Toxicity

Caesium chloride has a low toxicity to humans and animals.[54] Its median lethal dose (LD50) in mice is 2300 mg per kilogram of body weight for oral administration and 910 mg/kg for intravenous injection.[55] The mild toxicity of CsCl is related to its ability to lower the concentration of potassium in the body and partly substitute it in biochemical processes.[56] When taken in large quantities, however, can cause a significant imbalance in potassium and lead to hypokalemia, arrythmia, and acute cardiac arrest.[57] However, caesium chloride powder can irritate the mucous membranes and cause asthma.[51]

Because of its high solubility in water, caesium chloride is highly mobile and can even diffuse through concrete. This is a drawback for its radioactive form which urges a search for less chemically mobile radioisotope materials. Commercial sources of radioactive caesium chloride are well sealed in a double steel enclosure.[34] However, in the Goiânia accident in Brazil , such a source containing about 93 grams of 137CsCl, was stolen from an abandoned hospital and forced open by two scavengers. The blue glow emitted in the dark by the radioactive caesium chloride attracted the thieves and their relatives who were unaware of the associated dangers and spread the powder. This resulted in one of the worst radiation spill accidents in which 4 people died within a month from the exposure, 20 showed signs of radiation sickness, 249 people were contaminated with radioactive caesium chloride, and about a thousand received a dose exceeding a yearly amount of background radiation. More than 110,000 people overwhelmed the local hospitals, and several city blocks had to be demolished in the cleanup operations. In the first days of the contamination, stomach disorders and nausea due to radiation sickness were experienced by several people, but only after several days one person associated the symptoms with the powder and brought a sample to the authorities.[58][59]

See also

  • List of ineffective cancer treatments

References

  1. ↑ 1.0 1.1 1.2 1.3 1.4 Haynes, p. 4.57
  2. ↑ 2.0 2.1 Lushchik, A; Feldbach, E; Frorip, A; Ibragimov, K; Kuusmann, I; Lushchik, C (1994). "Relaxation of excitons in wide-gap CsCl crystals". Journal of Physics: Condensed Matter 6 (12): 2357–2366. doi:10.1088/0953-8984/6/12/009. Bibcode: 1994JPCM....6.2357L. 
  3. ↑ Haynes, p. 4.132
  4. ↑ Haynes, p. 10.240
  5. ↑ 5.0 5.1 Watanabe, M.; Tokonami, M.; Morimoto, N. (1977). "The transition mechanism between the CsCl-type and NaCl-type structures in CsCl". Acta Crystallographica Section A 33 (2): 294. doi:10.1107/S0567739477000722. Bibcode: 1977AcCrA..33..294W. 
  6. ↑ Cesium chloride. nlm.nih.gov
  7. ↑ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8. 
  8. ↑ Wells A.F. (1984) Structural Inorganic Chemistry 5th edition Oxford Science Publications ISBN:0-19-855370-6
  9. ↑ Schulz, L. G. (1951). "Polymorphism of cesium and thallium halides". Acta Crystallographica 4 (6): 487–489. doi:10.1107/S0365110X51001641. Bibcode: 1951AcCry...4..487S. 
  10. ↑ Lidin, p. 620
  11. ↑ "ЭСБЕ/Цезий". Brockhaus and Efron Encyclopedic Dictionary. 1890–1907. http://ru.wikisource.org/wiki/%D0%AD%D0%A1%D0%91%D0%95/%D0%A6%D0%B5%D0%B7%D0%B8%D0%B9. 
  12. ↑ 12.0 12.1 Knunyants, I. L, ed (1998). "Цезия галогениды". Химическая энциклопедия (Chemical encyclopedia). 5. Moscow: Soviet Encyclopedia. p. 657. ISBN 978-5-85270-310-1. 
  13. ↑ Haynes, p. 5.191
  14. ↑ Turova, N. Ya. (1997). Неорганическая химия в таблицах. Moscow. p. 85. 
  15. ↑ 15.0 15.1 15.2 Plyushev, V.E.; Stepin, B. D (1975). Аналитическая химия рубидия и цезия. Moscow: Nauka. pp. 22–26. 
  16. ↑ Plyushev, p. 97
  17. ↑ Plyushev, V.E. (1976). Bolshakov, K. A.. ed. Химия и технология редких и рассеянных элементов. 1 (2 ed.). Moscow: Vysshaya Shkola. pp. 101–103. 
  18. ↑ Ehrenreich, Henry (1984). Solid state physics: advances in research and applications. Academic Press. pp. 29–31. ISBN 978-0-12-607738-4. https://books.google.com/books?id=ljr450_uYlsC&pg=PA30. 
  19. ↑ Haynes, p. 5.126
  20. ↑ Lidin, p. 645
  21. ↑ Lidin, R. A; Molochko V.; Andreeva, L. L. A. (2000). Химические свойства неорганических веществ (3 ed.). Moscow: Khimiya. p. 49. ISBN 978-5-7245-1163-6. 
  22. ↑ Knunyants, I. L, ed (1988). "Бария хлорид". Химическая энциклопедия. 1. Moscow: Soviet Encyclopedia. p. 463. 
  23. ↑ National Research Council (U.S.). Office of Critical Tables, ed (1962). Consolidated Index of Selected Property Values: Physical Chemistry and Thermodynamics (Publication 976 ed.). Washington, D.C.: National Academy of Science. p. 271. 
  24. ↑ Knunyants, I. L, ed (1992). "Полигалогениды". Химическая энциклопедия. 3. Moscow: Soviet encyclopedia. pp. 1237–1238. ISBN 978-5-85270-039-1. 
  25. ↑ Senga, Ryosuke; Komsa, Hannu-Pekka; Liu, Zheng; Hirose-Takai, Kaori; Krasheninnikov, Arkady V.; Suenaga, Kazu (2014). "Atomic structure and dynamic behaviour of truly one-dimensional ionic chains inside carbon nanotubes". Nature Materials 13 (11): 1050–4. doi:10.1038/nmat4069. PMID 25218060. Bibcode: 2014NatMa..13.1050S. 
  26. ↑ Knunyants, I. L, ed (1998). "Цезий". Химическая энциклопедия (Chemical encyclopedia). 5. Moscow: Soviet Encyclopedia. pp. 654–656. ISBN 978-5-85270-310-1. 
  27. ↑ Plyushev, pp. 210–211
  28. ↑ Plyushev, p. 206
  29. ↑ 29.0 29.1 29.2 "Cesium and Cesium Compounds". Kirk-Othmer Encyclopedia of Chemical Technology. 5 (4th ed.). New York: John Wiley & Sons. 1994. pp. 375–376. 
  30. ↑ Plsyushev, pp. 357–358
  31. ↑ 31.0 31.1 31.2 31.3 31.4 31.5 Bick, Manfred and Prinz, Horst (2002) "Cesium and Cesium Compounds" in Ullmann's Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim. Vol. A6, pp. 153–156. doi:10.1002/14356007.a06_153
  32. ↑ Halka M.; Nordstrom B. (2010). Alkali and Alkaline Earth Metals. Infobase Publishing. p. 52. ISBN 978-0-8160-7369-6. https://books.google.com/books?id=oZHKEG5V4_0C&pg=PA52. 
  33. ↑ Enrique Lima "Cesium: Radionuclide" in Encyclopedia of Inorganic Chemistry, 2006, Wiley-VCH, Weinheim. doi:10.1002/0470862106.ia712
  34. ↑ 34.0 34.1 National Research Council (U.S.). Committee on Radiation Source Use and Replacement; Nuclear and Radiation Studies Board (January 2008). Radiation source use and replacement: abbreviated version. National Academies Press. pp. 28–. ISBN 978-0-309-11014-3. https://books.google.com/books?id=3cT2REdXJ98C&pg=PA28. 
  35. ↑ Hackspill, M. L. (1905). "Sur une nouvelle prepapratíon du rubidium et du cæsium" (in fr). Comptes Rendus Hebdomadaires des Séances de l'Académie des Sciences 141: 106. http://gallica.bnf.fr/ark:/12148/bpt6k3095m/f106. 
  36. ↑ Plyushev, p. 90
  37. ↑ Kano T.; Kumano T.; Maruoka K. (2009). "Rate Enhancement of Phase Transfer Catalyzed Conjugate Additions by CsCl". Organic Letters 11 (9): 2023–2025. doi:10.1021/ol900476e. PMID 19348469. 
  38. ↑ Katritzky A. R.; Meth-Cohn O.; Rees Ch. W. (1995). Gilchrist, T. L.. ed. Synthesis: Carbon with Three or Four Attached Heteroatoms. Comprehensive Organic Functional Group Transformations. 6 (First ed.). New York: Elsevier. p. 283. ISBN 978-0-08-040604-6. https://archive.org/details/comprehensiveorg0000unse/page/283. 
  39. ↑ ГОСТ 52407-2005. Вода питьевая. Методы определения жесткости. Moscow: Стандартинформ. 2006. 
  40. ↑ "Cesium Chloride". Complementary and Alternative Medicine: Herbs, Vitamins, and Minerals. American Cancer Society. 30 November 2008. http://www.cancer.org/Treatment/TreatmentsandSideEffects/ComplementaryandAlternativeMedicine/HerbsVitaminsandMinerals/cesium-chloride. 
  41. ↑ "FDA alerts health care professionals of significant safety risks associated with cesium chloride". Food and Drug Administration. July 23, 2018. https://www.fda.gov/Drugs/GuidanceComplianceRegulatoryInformation/PharmacyCompounding/ucm614211.htm. 
  42. ↑ "FDA blacklists cesium chloride, ineffective and dangerous naturopathic cancer treatment". August 2, 2018. https://sciencebasedmedicine.org/fda-blacklists-cesium-chloride-ineffective-and-dangerous-naturopathic-cancer-treatment/. 
  43. ↑ 43.0 43.1 Cesium. Mineral Commodity Summaries January 2010. U.S. Geological Survey
  44. ↑ Carrea, JR; Gleason, G; Shaw, J; Krontz, B (1964). "The direct diagnosis of myocardial infarction by photoscanning after administration of cesium-131". American Heart Journal 68 (5): 627–36. doi:10.1016/0002-8703(64)90271-6. PMID 14222401. http://deepblue.lib.umich.edu/bitstream/2027.42/32170/1/0000225.pdf. 
  45. ↑ McGeehan, John T. (1968). "Cesium 131 Photoscan: Aid in the Diagnosis of Myocardial Infarction". JAMA: The Journal of the American Medical Association 204 (7): 585–589. doi:10.1001/jama.1968.03140200025006. PMID 5694480. 
  46. ↑ Tver'yanovich, Y. S. (1998). "Optical absorption and composition of the nearest environment of neodymium in glasses based on the gallium-germanium-chalcogen system". Glass Phys. Chem. 24: 446. 
  47. ↑ Klenovskii, M.S.; Kel'man, V.A.; Zhmenyak, Yu.V.; Shpenik, Yu.O. (2010). "Electric-discharge UV radiation source based on a Xe-CsCl vapor-gas mixture". Technical Physics 55 (5): 709–714. doi:10.1134/S1063784210050178. Bibcode: 2010JTePh..55..709K. 
  48. ↑ Klenovskii, M.S.; Kel'man, V.A.; Zhmenyak, Yu.V.; Shpenik, Yu.O. (2013). "Luminescence of XeCl* and XeBr* exciplex molecules initiated by a longitudinal pulsed discharge in a three-component mixture of Xe with CsCl and CsBr vapors". Optics and Spectroscopy 114 (2): 197–204. doi:10.1134/S0030400X13010141. Bibcode: 2013OptSp.114..197K. 
  49. ↑ "Тугоплавкие и химически активные металлы". Migatronic. http://www.migatronic.ru/content/view/52/46/. 
  50. ↑ Morris, Ch. G., ed (1992). "Cesium chloride". Academic Press Dictionary of Science and Technology. San Diego: Academic Press. p. 395. ISBN 978-0-12-200400-1. https://archive.org/details/academicpressdic00morr. 
  51. ↑ 51.0 51.1 "Cesium Chloride MSDS". Cesium Fine Chemicals. Cabot Corporation. http://www.cabot-corp.com/wcm/msds/en-us/CE/CESCHLORDRY-NA-EN.pdf. 
  52. ↑ Industrial Minerals & Rocks: Commodities, Markets, and Uses (7th ed.). Littleton: Society for Mining, Metallurgy, and Exploration. 2006. p. 1430. ISBN 978-0-87335-233-8. 
  53. ↑ Biel, Martin; Christian Wahl-Schott; Stylianos Michalakis; Xiangang Zong (2009). "Hyperpolarization-Activated Cation Channels: From Genes to Function". Physiological Reviews 89 (3): 847–85. doi:10.1152/physrev.00029.2008. PMID 19584315. 
  54. ↑ "Chemical Safety Data: Caesium chloride". Hands-on Science (H-Sci) Project: Chemical Safety Database. Physical and Theoretical Chemistry Laboratory, Oxford University. http://cartwright.chem.ox.ac.uk/hsci/chemicals/caesium_chloride.html. 
  55. ↑ "Safety data for caesium chloride". Chemical and Other Safety Information. The Physical and Theoretical Chemistry Laboratory Oxford University. http://msds.chem.ox.ac.uk/CA/caesium_chloride.html. 
  56. ↑ Lazarev N.V. and Gadaskina, I.D., ed (1977) (in ru). 3 (7 ed.). St. Petersburg: Khimiya. pp. 328–329. 
  57. ↑ Melnikov, P; Zanoni, LZ (June 2010). "Clinical effects of cesium intake". Biological Trace Element Research 135 (1–3): 1–9. doi:10.1007/s12011-009-8486-7. PMID 19655100. 
  58. ↑ The Radiological Accident in Goiânia. Vienna: IAEA. 1988. ISBN 978-92-0-129088-5. http://www-pub.iaea.org/MTCD/publications/PubDetAR.asp?pubId=3684. . See pp. 1–6 for summary and p. 22 for the source description
  59. ↑ "The Worst Nuclear Disasters". Time (magazine) . 2009.

Bibliography



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