Caesium or cesium is the chemical element with the symbol Cs and atomic number 55. There are characteristics, applications (other applications and nuclear applications), history, occurrence, isotopes and precautions about Caesium in this summary.

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Subject: Caesium

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1 Characteristics

2 Applications

2.1 Other applications

2.1.1 As the free element

2.1.2 As the salt

2.2 Nuclear applications

3 History

4 Occurrence

5 Isotopes

6 Precautions




Caesium or cesium (pronounced /?si?zi?m/, SEE-zee-?m) is the chemical element with the symbol Cs and atomic number 55. It is a soft, silvery-gold alkali metal with a melting point of 28 °C (83 °F), which makes it one of only five metals that are liquid at or near room temperature.[2] Caesium is most notably used in atomic clocks.

Caesium is the international spelling standardized by the IUPAC, but in the United States it is more commonly spelled as cesium.[3]

Figure. 0 - Silvery gold


The emission spectrum of caesium has two bright lines in the blue area of the spectrum along with several other lines in the red, yellow, and green areas. This metal is silvery gold in color and is both soft and ductile. Caesium has the lowest ionization energy of the chemical elements. Caesium is the least abundant of the five non-radioactive alkali metals. (Francium is the least common alkali metal but it has no stable isotopes. [4]).

Caesium, gallium, francium, rubidium, and mercury are the only pure metals liquid at or near room temperature. (Some sodium-potassium alloys are also liquid at room temperature.) Caesium reacts explosively in cold water and also reacts with ice at temperatures above -116 °C (-177 °F, 157 K).

Caesium hydroxide (CsOH) is a very strong base and will rapidly etch the surface of glass. CsOH is often stated to be the «strongest base», but in fact many compounds such as n-butyllithium and sodium amide are stronger but are not classic hydroxide bases and are destroyed by water.

Figure 1 - An ampoule containing high purity caesium metal


Probably the most widespread use of caesium today is in caesium formate-based drilling fluids for the oil industry. The high density of the caesium formate brine (up to 2.3 sg), coupled with the relatively benign nature of natural caesium (which has minimal radioactivity because it is almost entirely composed of a stable istotope), reduces the requirement for toxic high-density suspended solids in the drilling fluid, which is a significant technological, engineering and environmental advantage.[5][6]

Caesium is also used in atomic clocks, which are accurate to seconds over many thousands of years. Since 1967, the International System of Measurements has based its unit of time, the second, on the properties of caesium. The International System of Units (SI) defines the second as 9,192,631,770 cycles of the radiation, which corresponds to the transition between two hyperfine energy levels of the ground state of the 133Cs atom.

2.1 Other applications

2.1.1 As the free element

Like other elements of group 1, caesium has a great affinity for oxygen and is used as a «getter» in vacuum tubes.

The metal is also used in photoelectric cells due to its ready emission of electrons.

Caesium was used as a propellant in early ion engines. It used a method of ionization to strip the outer electron from the propellant by simple contact with tungsten. Caesium use as a propellant was discontinued when Hughes Research Laboratory conducted a study finding xenon gas as a suitable replacement.

2.1.2 As the salt

Caesium is used as a catalyst in the hydrogenation of certain organic compounds.

Caesium fluoride is widely used in organic chemistry as a base and as a source of anhydrous fluoride ion.

Caesium vapor is used in many common magnetometers.[7]

Because of their high density, caesium chloride solutions are commonly used in molecular biology for density gradient ultracentrifugation, primarily for the isolation of viral particles, subcellular organelles and fractions, and nucleic acids from biological samples.[8]

Caesium nitrate is used as an oxidizer and pyrotechnic colorant to burn silicon in infrared flares[9] such as the LUU-19 flare,[10] because it emits much of its light in the near infrared spectrum.[11]

Caesium is also used as an internal standard in spectrophotometry.[12]

Caesium has been used to reduce the radar signature of exhaust plumes in military aircraft.[citation needed]

2.2 Nuclear applications

134Cs has been used in hydrology as a measure of caesium output by the nuclear power industry. This isotope is used because, while it is less prevalent than either 133Cs or 137Cs, 134Cs can be produced solely by nuclear reactions. 135Cs has also been used in this function.

Caesium-137 is an extremely common radioisotope used as a gamma-emitter in industrial applications such as: moisture/density gauges; leveling gauges; thickness gauges; well logging devices which are used to measure the electron density, which is analogous to the bulk density, of the rock formations.

Radioactive isotopes of caesium are used in the medical field to treat certain types of cancer.


Caesium (Latin caesius meaning «blueish grey») [13][14] was spectroscopically discovered by Robert Bunsen and Gustav Kirchhoff in 1860 in mineral water from Dьrkheim, Germany. The residues of 44,000 liters of mineral water yielded several grams of caesium salt for further analysis. Its identification was based upon the bright blue lines in its spectrum and it was the first element discovered by spectrum analysis.[15] The first caesium metal was produced in 1882 by electrolysis of caesium chloride by Carl Setterberg.[16] Setterberg received his PhD from Kekule and Bunsen for this work. Historically, the most important use for caesium has been in research and development, primarily in chemical and electrical applications.


An alkali metal, caesium occurs in lepidolite, pollucite (hydrated silicate of aluminium and caesium) and within other sources. One of the world's most significant and rich sources of this metal is the Tanco mine at Bernic Lake in Manitoba.[17] The deposits there are estimated to contain 350,000 metric tons[17] of pollucite ore at an average of composition of 24% caesium by weight.[18][19]

It can be isolated by electrolysis of fused caesium cyanide and in a number of other ways. Exceptionally pure and gas-free caesium can be made by the thermal decomposition of caesium azide. The primary compounds of caesium are caesium chloride and its nitrate. The price of caesium metal in 1997 was about US$30 per gram, but its compounds are much cheaper.

Figure 2 - Pollucite, a caesium mineral


Caesium has at least 39 known isotopes, which is more than any other element except francium. The atomic masses of these isotopes range from 112 to 151. Even though this element has a large number of isotopes, it has only one naturally occurring stable isotope, 133Cs. Most of the other isotopes have half-lives from a few days to fractions of a second. The radiogenic isotope 137Cs has been used in hydrologic studies, analogous to the use of 3H. 137Cs is produced from the detonation of nuclear weapons and is produced in nuclear power plants, and was released to the atmosphere most notably from the 1986 Chernobyl disaster. This isotope (137Cs) is one of the numerous products of fission, directly issued from the fission of uranium.

Beginning in 1945 with the commencement of nuclear weapons testing, 137Cs was released into the atmosphere where it is not absorbed readily into solution and is returned to the surface of the earth as a component of radioactive fallout. Once 137Cs enters the ground water, it is deposited on soil surfaces and removed from the landscape primarily by particle transport. As a result, the input function of these isotopes cannot be estimated as a function of time. Caesium-137 has a half-life of 30.17 years. It decomposes to barium-137m (a short-lived product of decay) then to a form of nonradioactive barium.


All alkali metals are highly reactive. Caesium, being one of the heavier alkali metals, is also one of the most reactive and is highly explosive when it comes in contact with water. The hydrogen gas produced by the reaction is heated by the thermal energy released at the same time, causing ignition and a violent explosion (the same as all alkali metals)--but caesium is so reactive that this explosive reaction can even be triggered by cold water or ice at temperatures down to ?116°C. Caesium is highly pyrophoric and ignites spontaneously in air to form caesium hydroxide and various oxides. Caesium hydroxide is an extremely strong base, and can rapidly corrode glass.

Caesium compounds are rarely encountered by most persons. All caesium compounds should be regarded as mildly toxic because of its chemical similarity to potassium. Large amounts cause hyperirritability and spasms, but such amounts would not ordinarily be encountered in natural sources, so Cs is not a major chemical environmental pollutant.[20]

The median lethal dose (LD50) value for caesium chloride in mice was determined to be 2300 mg/kg which is comparable to the LD50 values of potassium chloride and sodium chloride.[21]

The isotopes 134Cs and 137Cs (present in the biosphere in small amounts as a result of radiation leaks) represent a radioactivity burden which varies depending on location. Radiocaesium does not accumulate in the body as effectively as many other fission products (such as radioiodine and radiostrontium), which are actively accumulated by the body. As with other alkali metals, radiocaesium washes out of the body relatively quickly in the sweat and urine. However, radiocaesium follows potassium and tends to accumulate in plant tissues, including fruits and vegetables.[citation needed]


As an individual representative of the periodic table of chemical elements Dmitry Ivanovich Mendeleyev, the element has unique chemical and physical properties.

Element is of great economic importance and plays a major role in world culture.


1. Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics 81th edition, CRC press.

2. Along with rubidium (39 °C [102 °F]), francium (27 °C [81 °F]), mercury (?39 °C [?38 °F]), and gallium (30 °C [86 °F]).

3. Bromine is also liquid at room temperature (-7.2?°C, 19?°F) but it is not a metal, but a halogen.

4. IUPAC Periodic Table of the Elements

5. Adloff, Jean-Pierre; George B. Kauffman (09/23 2005). «Francium (Atomic Number 87), the Last Discovered Natural Element». The Chemical Educator 10 (5). doi:10.1333/s00897050956a. Retrieved 2006-05-16.

6. Drilling and Completing Difficult HP/HT Wells With the Aid of Cesium Formate Brines-A Performance Review

7. Overview: Cesium Formate Fluids

8. Groeger, S. (2005). «Comparison of discharge lamp and laser pumped cesium magnetometers». Applied Physics B 80: 645. doi:10.1007/s00340-005-1773-x.

9. edited by Mohamed A. Desai. (2000). «Gradient Materials». Downstream processing methods. Totowa, N.J.: Humana Press. pp. 61-62. ISBN 9780896035645.

10. United States Patent 6230628: Infrared illumination compositions and articles containing the same

11. LUU-19 Flare

12. Charrier, E. (2006). «Determination of the temperature and enthalpy of the solid-solid phase transition of caesium nitrate by differential scanning calorimetry». Thermochimica Acta 445: 36. doi:10.1016/j.tca.2006.04.002.

13. «Internal Standarts». Laboratory instrumentation.. New York: Wiley. 1995. p. 108. ISBN 9780471285724.

14. 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.

15. Oxford English Dictionary, 2nd Edition

16. G. Kirchhoff, R. Bunsen (1861). «Chemische Analyse durch Spectralbeobachtungen». Annalen der Physik und Chemie 189 (7): 337-381. doi:10.1002/andp.18611890702.

17. Setterberg, Carl (1882). «10.1002/jlac.18822110105». Justus Liebig's Annalen der Chemie 211: 100. doi:10.1002/jlac.18822110105.

18. a b Иernэ, Petr; Simpson, F. M. (1978). «The Tanco Pegmatite at Bernic Lake, Manitoba: X. Pollucite». Canadian Mineralogist 16: 325-333.

19. Polyak, Dйsirйe E.. «Cesium». United States Geological Survey. Retrieved 2009-10-17.

20. Butterman, William C.; Brooks, William E.; Reese, Jr., Robert G.. «Cesium». United States Geological Survey. Retrieved 2009-10-17.

21. Pinsky, Carl (1981). «Cesium in mammals: Acute toxicity, organ changes and tissue accumulation». Journal of Environmental Science and Health, Part A 16: 549. doi:10.1080/10934528109375003.

22. Johnson, G (1975). «Acute toxicity of cesium and rubidium compounds». Toxicology and Applied Pharmacology 32 (2): 239. doi:10.1016/0041-008X(75)90216-1. PMID 1154391.

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