Difference between revisions of "Cesium" - New World Encyclopedia

Vapor pressure
55	xenon ← cesium → barium
	periodic table
General
Name, Symbol, Number	cesium, Cs, 55
Chemical series	alkali metals
Group, Period, Block	1, 6, s
Appearance	silvery gold ;
Atomic mass	132.9054519(2) g/mol
Electron configuration	[Xe] 6s1
Electrons per shell	2, 8, 18, 18, 8, 1
Physical properties
Phase	solid
Density (near r.t.)	1.93 g/cm³
Liquid density at m.p.	1.843 g/cm³
Melting point	301.59 K; (28.44 °C, 83.19 °F)
Boiling point	944 K; (671 °C, 1240 °F)
Critical point	1938 K, 9.4 MPa
Heat of fusion	2.09 kJ/mol
Heat of vaporization	63.9 kJ/mol
Heat capacity	(25 °C) 32.210 J/(mol·K)
Atomic properties
Crystal structure	cubic body centered
Oxidation states	1; (strongly basic oxide)
Electronegativity	0.79 (Pauling scale)
Ionization energies	1st: 375.7 kJ/mol
	2nd: 2234.3 kJ/mol
	3rd: 3400 kJ/mol
Atomic radius	260 pm
Atomic radius (calc.)	298 pm
Covalent radius	225 pm
Miscellaneous
Magnetic ordering	no data
Electrical resistivity	(20 °C) 205 nΩ·m
Thermal conductivity	(300 K) 35.9 W/(m·K)
Thermal expansion	(25 °C) 97 µm/(m·K)
Speed of sound (thin rod)	(r.t.) 1.7 m/s
Bulk modulus	1.6 GPa
Mohs hardness	0.2
Brinell hardness	0.14 MPa
CAS registry number	7440-46-2
Notable isotopes

Latest revision as of 00:09, 4 December 2023

Cesium (also spelled caesium, chemical symbol Cs, atomic number 55) is a member of the group of chemical elements known as alkali metals. Soft and silvery gold in color, it is one of a handful of metals that liquefy near room temperature. It is the most alkaline of the stable elements, reacting explosively with water to produce the alkali cesium hydroxide, which can corrode glass. Its atoms readily lose electrons to produce ions, and this property makes cesium useful for photoelectric cells. Cesium is a catalyst for certain reactions in organic chemistry.

This element has many isotopes, of which the single stable isotope (cesium-133) has become the primary standard for measuring seconds of time and is therefore the basis of atomic clocks. Some of its radioactive isotopes are used to treat certain types of cancer, and cesium-134 helps measure cesium output by the nuclear power industry. Cesium-137 is commonly used in industry for such applications as moisture density gauges, leveling gauges, and thickness gauges. Cesium compounds are mildly toxic.

Discovery

Cesium was discovered by Robert Bunsen and Gustav Kirchhoff in 1860, when they analyzed the spectrum of mineral water obtained from Dürkheim, Germany. They held a drop of the water in a flame and observed the spectral lines of several elements. Among these lines they observed two blue lines that did not come from any of the known elements. They concluded that the blue lines were produced by a hitherto unknown element, which they named cesium (from the Latin word caesius, meaning "sky blue" or "heavenly blue"). Thus cesium became the first element discovered by spectral analysis. Bunsen isolated cesium salts from the spring water, and the metal itself was isolated in 1881 by Carl Setterberg, who worked in Bunsen's laboratory.

Occurrence

Pollucite, a cesium mineral.

Cesium occurs in several minerals, particularly lepidolite and pollucite (a hydrated silicate of aluminum and cesium). One of the world's most significant and rich sources of this metal is located at Bernic Lake in Manitoba. The deposits there are estimated to contain 300,000 metric tons of pollucite, at an average of 20 percent cesium.

Metallic cesium can be isolated by the electrolysis of fused (molten) cesium cyanide, as well as in several other ways. Exceptionally pure, gas-free cesium can be made by decomposing cesium azide with heat.

Notable characteristics

As an alkali metal, cesium is part of group one of the periodic table, between rubidium and francium. Also, it lies in period six, just before barium. This silvery gold metal is soft and ductile (it can be readily drawn into wires).

As noted above, the spectrum of cesium has two bright lines in the blue part of the electromagnetic spectrum. In addition, its spectrum has several other lines in the red, yellow, and green regions.

Cesium is the most electropositive and most alkaline of the stable chemical elements. Aside from francium, it has the lowest ionization potential of all the elements, which means that it readily loses its outermost electron to become an ion. (Ionization potential is the energy needed to remove an electron from the atom's outermost shell.)

Cesium is the least abundant of the five nonradioactive alkali metals. Technically, francium is the least common alkali metal, but because it is highly radioactive and its total worldwide quantity is estimated to be only 340 to 550 grams, its abundance can be considered zero in practical terms.^[1]

Along with gallium, francium, and mercury, cesium is among the few metals that are liquid at or near room temperature. Its melting point is 28.44°C.

Cesium reacts explosively when it comes in contact with water (even cold water), and also reacts with ice at temperatures above −116°C. The reaction with water produces cesium hydroxide (CsOH), an extremely strong chemical base that will rapidly etch the surface of glass. In addition cesium reacts violently with chlorine gas to produce cesium chloride (CsCl).

Isotopes

Cesium 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. Yet it has only one naturally occurring stable isotope: ¹³³Cs. Most of the other isotopes (except for those noted in the table) have half-lives from a few days to fractions of a second.

Since 1967, the International System of Units (SI units) has defined the second based on oscillations between two energy levels of the cesium-133 atom. Thus the cesium oscillator has become the primary standard for time and frequency measurements.

The isotope ¹³⁷Cs, with a half-life of 30.17 years, has been used in hydrologic studies, analogous to the use of tritium (³H). It decomposes to barium-137 (a short-lived product of decay), then to a form of nonradioactive barium.

Cesium-137 is produced during the detonation of nuclear weapons and in nuclear power plants. It was a notable product during the 1986 Chernobyl meltdown. Beginning in 1945, with the commencement of nuclear testing, ¹³⁷Cs has been released into the atmosphere, where it is absorbed by moisture and returned to the Earth's surface as a component of radioactive fallout. Once ¹³⁷Cs enters groundwater, it is deposited on soil surfaces and removed from the landscape primarily by particle transport.

Compounds of cesium

Cesium carbonate: Cesium carbonate is a white crystalline solid, with the chemical formula Cs₂CO₃. It is more soluble in organic solvents than most carbonates such as potassium carbonate. It is therefore used as a base in organic chemistry.

Cesium chloride: Cesium chloride (CsCl) is an ionic compound. It can be prepared by the reaction of cesium hydroxide or cesium carbonate with hydrochloric acid. The resulting cesium chloride salt is purified by recrystallization. The crystalline structure of this salt is composed of interlocking simple cubic lattices of anions and cations.

Cesium hydroxide: Cesium hydroxide (CsOH) is formed when cesium reacts with water. It is a powerful base, much like other alkali metal hydroxides such as sodium hydroxide and potassium hydroxide. It can corrode glass.

Cesium hydroxide is extremely hygroscopic, which means that it readily absorbs moisture from the environment. Laboratory cesium hydroxide is typically a hydrate.

This compound is not usually used in experiments, because the extraction of cesium is very expensive and it behaves very much like rubidium hydroxide and potassium hydroxide.

Applications

Cesium is most notably used in atomic clocks, which are accurate to seconds in many thousands of years.
Cesium-134 has been used as a measure of cesium output by the nuclear power industry. This isotope is used because, while it is less prevalent than either ¹³³Cs or ¹³⁷Cs, ¹³⁴Cs is produced solely by nuclear reactions. ¹³⁵Cs has also been used for this purpose.
Like other group one elements, cesium has great affinity for oxygen and is used as a "getter" in vacuum tubes.
This metal is also used in photoelectric cells because it readily emits electrons.
Cesium is used as a catalyst in the hydrogenation of certain organic compounds.
Cesium hydroxide is used to etch silicon, exposing octahedral planes. This technique can create pyramids and regularly shaped etch pits for uses such as MEMS (microelectromechanical systems).
Cesium chloride is used in the production of electrically conducting glasses.
Radioisotopes of cesium (in the form of cesium chloride) are used in nuclear medicine, including the treatment of cancer.
Cesium fluoride is widely used in organic chemistry as a base and as a source of anhydrous fluoride ion.
Cesium vapor is used in many common magnetometers.
Because of their high density, cesium chloride solutions are commonly used in molecular biology experiments for the isolation of nucleic acids by a technique known as "density gradient ultracentrifugation."
Recently, this metal has been used in ion propulsion systems.
Cesium-137 is an extremely common isotope in industrial applications such as:
- moisture density gauges;
- leveling gauges;
- thickness gauges;
- well-logging devices (used to measure the thickness of rock strata).

Precautions

All alkali metals are extremely reactive. As one of the heavier alkali metals, cesium is also one of the most reactive. As noted above, it is highly explosive when it comes in contact with water or ice. Cesium hydroxide is an extremely strong base and can attack glass.

All cesium compounds should be regarded as mildly toxic, because of its chemical similarity to potassium. Contact with large amounts can cause hyperirritability and spasms. On the other hand, such amounts would not ordinarily be encountered in nature, so cesium is not a major chemical environmental pollutant. Rats fed with cesium in place of potassium in their diet were found to die, so this element cannot replace potassium in function.

The isotopes ¹³⁴Cs and ¹³⁷Cs (present in the biosphere in small amounts as a result of radiation leaks) represent a radioactivity burden that varies depending on location. Radioactive cesium does not accumulate in the body as effectively as many other fission products, such as radioactive iodine or strontium.

Notes

↑ Adloff, Jean-Pierre and George B. Kauffman. (September 23, 2005). Francium (Atomic Number 87), the Last Discovered Natural Element. The Chemical Educator 10 (5). Retrieved December 9, 2007.

References
ISBN links support NWE through referral fees

Los Alamos National Laboratory - Cesium Retrieved December 9, 2007.

External links

All links retrieved December 3, 2023.

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[1] Adloff, Jean-Pierre and George B. Kauffman. (September 23, 2005). Francium (Atomic Number 87), the Last Discovered Natural Element. The Chemical Educator 10 (5). Retrieved December 9, 2007.

[1]

@@ Line 1: / Line 1: @@
-{{Elementbox_header | number=55 | symbol=Cs | name=caesium | left=[[xenon]] | right=[[barium]] | above=[[rubidium|Rb]] | below=[[francium|Fr]] | color1=#ff6666 | color2=black }}
+{{Copyedited}}{{Paid}}{{Images OK}}{{Submitted}}{{Approved}}
+{{Elementbox_header | number=55 | symbol=Cs | name=cesium | left=[[xenon]] | right=[[barium]] | above=[[rubidium|Rb]] | below=[[francium|Fr]] | color1=#ff6666 | color2=black }}
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 {{Elementbox_brinellhardness_mpa | 0.14 }}
 {{Elementbox_cas_number | 7440-46-2 }}
-{{Elementbox_isotopes_begin | isotopesof=caesium | color1=#ff6666 | color2=black }}
+{{Elementbox_isotopes_begin | isotopesof=cesium | color1=#ff6666 | color2=black }}
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-'''Caesium''' (also spelled '''cesium'''; pronounced {{IPA|/ˈsiːziəm/}}) is a [[chemical element]] in the [[periodic table]] that has the symbol '''Cs''' and [[atomic number]] 55. It is a soft silvery-gold [[alkali metal]] which is one of metals that are liquid at or near [[room temperature]]: [[Rubidium]] (Mp: 38,89°C), [[Francium]] (Mp: 27°C), [[Mercury]] (Mp: -38,84°C), [[Gallium]] (Mp: 29,78°C) and [[Potassium]] (Mp: 63,65°C). This element is most notably used in [[atomic clock]]s.
+'''Cesium''' (also spelled '''caesium''', chemical symbol '''Cs''', [[atomic number]] 55) is a member of the group of [[chemical element]]s known as alkali metals. Soft and silvery gold in color, it is one of a handful of [[metal]]s that liquefy near room [[temperature]]. It is the most [[alkaline]] of the stable elements, reacting explosively with [[water]] to produce the alkali cesium hydroxide, which can corrode glass. Its [[atom]]s readily lose [[electron]]s to produce [[ion]]s, and this property makes cesium useful for [[photoelectric cell]]s. Cesium is a [[catalyst]] for certain reactions in [[organic chemistry]].
+{{toc}}
+This element has many [[isotopes]], of which the single stable isotope (cesium-133) has become the primary standard for measuring seconds of time and is therefore the basis of [[atomic clock]]s. Some of its radioactive isotopes are used to treat certain types of [[cancer]], and cesium-134 helps measure cesium output by the [[nuclear power]] industry. Cesium-137 is commonly used in industry for such applications as moisture density gauges, leveling gauges, and thickness gauges. Cesium compounds are mildly toxic.
+==Discovery==
+Cesium was discovered by [[Robert Bunsen]] and [[Gustav Kirchhoff]] in 1860, when they analyzed the spectrum of [[mineral water]] obtained from Dürkheim, [[Germany]]. They held a drop of the water in a flame and observed the spectral lines of several elements. Among these lines they observed two blue lines that did not come from any of the known elements. They concluded that the blue lines were produced by a hitherto unknown element, which they named cesium (from the [[Latin]] word ''caesius'', meaning "sky blue" or "heavenly blue"). Thus cesium became the first element discovered by spectral analysis. Bunsen isolated cesium [[salt]]s from the spring water, and the metal itself was isolated in 1881 by Carl Setterberg, who worked in Bunsen's laboratory.
+==Occurrence==
+[[Image:Pollucite(CesiumMineral)USGOV.jpg|thumb|left|Pollucite, a cesium mineral.]]
+Cesium occurs in several minerals, particularly [[lepidolite]] and [[pollucite]] (a hydrated [[silicate]] of [[aluminum]] and cesium). One of the world's most significant and rich sources of this metal is located at [[Bernic Lake]] in [[Manitoba]]. The deposits there are estimated to contain 300,000 [[tonne|metric tons]] of pollucite, at an average of 20 percent cesium.
-The variant spelling ''cesium'' is sometimes used, especially in [[North American English]], but ''caesium'' is the spelling used by the [[IUPAC]], although since [[1993]] it has recognized ''cesium'' as a variant as well.
+Metallic cesium can be isolated by the [[electrolysis]] of fused (molten) cesium [[cyanide]], as well as in several other ways. Exceptionally pure, gas-free cesium can be made by decomposing cesium [[azide]] with heat.
 ==Notable characteristics==
-The [[electromagnetic spectrum]] of caesium has two bright lines in the [[blue]] part of the spectrum along with several other lines in the [[red]], [[yellow]], and [[green]]. This metal is silvery gold in [[color]] and is both soft and ductile. Caesium is also the most [[electropositive]] and most [[alkaline]] of the stable [[chemical element]]s and also has the lowest [[ionization potential]] of all the elements, except for [[francium]]. Caesium is the least abundant of the five non-radioactive alkali metals.  (Technically, [[francium]] is the least common alkali metal, but since it is highly radioactive with an estimated 340 to 550 grams<ref>{{cite journal| url=http://chemeducator.org/sbibs/s0010005/spapers/1050387gk.htm| journal=The Chemical Educator| volume=10| issue=5| month=09/23| year=2005| {{doi|10.1333/s00897050956a}}| title=Francium (Atomic Number 87), the Last Discovered Natural Element| first=Jean-Pierre| last=Adloff| coauthors=George B. Kauffman| accessdate=2006-05-16}}</ref> in the entire earth at one time, its abundance can be considered zero in practical terms.)
-Along with [[gallium]], francium and [[mercury (element)|mercury]], caesium is among the only metals that are liquid at or near room [[temperature]]. Caesium reacts explosively in cold [[water_(molecule)|water]] and also reacts with ice at temperatures above &minus;116°C. Caesium [[hydroxide]] (CsOH) is a very strong [[base (chemistry)|base]] and will rapidly etch the surface of [[glass]].
+As an alkali metal, cesium is part of group one of the periodic table, between [[rubidium]] and [[francium]]. Also, it lies in period six, just before [[barium]]. This silvery gold metal is soft and ductile (it can be readily drawn into wires).
+As noted above, the spectrum of cesium has two bright lines in the blue part of the [[electromagnetic spectrum]]. In addition, its spectrum has several other lines in the red, yellow, and green regions.
+Cesium is the most [[Electronegativity|electropositive]] and most [[alkaline]] of the stable [[chemical element]]s. Aside from francium, it has the lowest [[ionization potential]] of all the elements, which means that it readily loses its outermost [[electron]] to become an [[ion]]. (Ionization potential is the energy needed to remove an electron from the atom's outermost shell.)
+Cesium is the least abundant of the five nonradioactive alkali metals. Technically, [[francium]] is the least common alkali metal, but because it is highly radioactive and its total worldwide quantity is estimated to be only 340 to 550 grams, its abundance can be considered zero in practical terms.<ref>{{cite journal| journal=The Chemical Educator| volume=10| issue=5| month=September 23,| year=2005| {{doi|10.1333/s00897050956a}}| title=Francium (Atomic Number 87), the Last Discovered Natural Element| first=Jean-Pierre| last=Adloff| coauthors=George B. Kauffman.}} Retrieved December 9, 2007.</ref>
+Along with [[gallium]], francium, and [[mercury (element)|mercury]], cesium is among the few metals that are liquid at or near room [[temperature]]. Its melting point is 28.44&deg;C.
+Cesium reacts explosively when it comes in contact with [[water]] (even cold water), and also reacts with ice at temperatures above &minus;116°C. The reaction with water produces cesium [[hydroxide]] (CsOH), an extremely strong chemical [[base (chemistry)|base]] that will rapidly etch the surface of [[glass]]. In addition cesium reacts violently with [[chlorine]] gas to produce cesium chloride (CsCl).
+===Isotopes===
+Cesium has at least 39 known [[isotope]]s, which is more than any other element except [[francium]]. The [[atomic mass]]es of these isotopes range from 112 to 151. Yet it has only one naturally occurring stable isotope: <sup>133</sup>Cs. Most of the other isotopes (except for those noted in the table) have [[half-life|half-lives]] from a few days to fractions of a second.
+Since 1967, the International System of Units ([[SI]] units) has defined the second based on oscillations between two energy levels of the cesium-133 atom. Thus the cesium oscillator has become the primary standard for time and frequency measurements.
+The isotope <sup>137</sup>Cs, with a half-life of 30.17 years, has been used in hydrologic studies, analogous to the use of tritium (<sup>3</sup>H). It decomposes to barium-137 (a short-lived product of decay), then to a form of nonradioactive barium.
+Cesium-137 is produced during the detonation of [[nuclear weapon]]s and in [[nuclear power plant]]s. It was a notable product during the 1986 [[Chernobyl]] meltdown. Beginning in 1945, with the commencement of [[nuclear testing]], <sup>137</sup>Cs has been released into the [[Earth's atmosphere|atmosphere]], where it is absorbed by moisture and returned to the [[Earth]]'s surface as a component of [[radioactive fallout]]. Once <sup>137</sup>Cs enters groundwater, it is deposited on soil surfaces and removed from the landscape primarily by [[particle transport]].
-CsOH is often stated to be a "strongest base". That is not correct. Although Caesium might be a strongest inorganic base (after FrOH), it is certainly not *the* strongest base. There are a number of organic compounds that act as a base. For instance, a hexane - cyclic molecule consisted of six C and twelve H atoms is incredibly strong base (pKa value: 52).
+== Compounds of cesium ==
-There is an account that caesium, reacting with [[fluorine]], takes up more fluorine than it [[stoichometry|stoichometrically]] should. It is possible that, after the salt Cs<sup>+</sup>F<sup>&minus;</sup> has formed, the Cs<sup>+</sup> ion, which has the same electronic structure as elemental [[xenon]], can like xenon be oxidised further by [[fluorine]] and form traces of a higher fluoride such as CsF<sub>3</sub>, analogous to XeF<sub>2</sub>.
+''Cesium carbonate'': Cesium carbonate is a white crystalline [[solid]], with the chemical formula Cs<sub>2</sub>CO<sub>3</sub>. It is more soluble in organic solvents than most carbonates such as [[potassium carbonate]]. It is therefore used as a [[base (chemistry)|base]] in [[organic chemistry]].
-==Applications==
+''Cesium chloride'': Cesium chloride (CsCl) is an ionic compound. It can be prepared by the reaction of [[cesium hydroxide]] or [[cesium carbonate]] with [[hydrochloric acid]]. The resulting cesium chloride salt is purified by [[recrystallization]]. The crystalline structure of this salt is composed of interlocking simple cubic lattices of anions and cations.
-Caesium is most notably used in [[atomic clock]]s, which are accurate to seconds in many thousands of years. Since 1967, the [[SI|International System of Measurements]] bases its unit of [[time]], the [[second]], on the properties of caesium. SI defines the [[second]] as 9,192,631,770 cycles of the [[Radioactive decay|radiation]] which corresponds to the transition between two [[energy level]]s of the [[ground state]] of the <sup>133</sup>Cs [[atom]].
-*<sup>134</sup>Cs 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 <sup>133</sup>Cs or [[Cs-137|<sup>137</sup>Cs]], <sup>134</sup>Cs can be produced solely by nuclear reactions. <sup>135</sup>Cs has also been used in this function.
+''Cesium hydroxide'': Cesium hydroxide (CsOH) is formed when cesium reacts with water. It is a powerful [[base (chemistry)|base]], much like other alkali metal hydroxides such as [[sodium hydroxide]] and [[potassium hydroxide]]. It can corrode glass.
-*Like other group 1 elements, caesium has a great affinity for [[oxygen]] and is used as a "[[getter]]" in [[vacuum tube]]s.
-*This metal is also used in [[photoelectric cell]]s due to its ready emission of [[electron]]s.
-*In addition, caesium is used as a [[catalyst]] in the [[hydrogenation]] of certain [[organic compound]]s.
-*[[Radioactive decay|Radioactive]] isotopes of caesium are used in the medical field to treat certain types of [[cancer]].
-*[[Caesium fluoride]] is widely used in [[organic chemistry]] as a [[Base (chemistry)|base]] and as a source of [[anhydrous]] [[fluoride]] ion.
-*Caesium vapor is used in many common [[magnetometers]].
-*Because of their high density, [[Caesium chloride]] solutions are commonly used in molecular biology for density gradient ultracentrifugation, primarily for the isolation of nucleic acids from biological samples.
-*More recently this metal has been used in [[ion propulsion]] systems.
-*Cesium-137 is an extremely common isotope in industrial applications such as:
-**moisture density gauges
-**leveling gauges
-**thickness gauges
-**well-logging devices (used to measure the thickness of rock-strata)
-==History==
+Cesium hydroxide is extremely [[hygroscopic]], which means that it readily absorbs moisture from the environment. Laboratory cesium hydroxide is typically a [[hydrate]].
-Caesium ([[Latin]] ''caesius'' meaning "sky blue" or "light blue") was [[spectroscopy|spectroscopically]] discovered by [[Robert Bunsen]] and [[Gustav Kirchhoff]] in [[1860]] in [[mineral water]] from Dürkheim, Germany. Its identification was based upon the bright blue lines in its spectrum and it was the first element discovered by spectrum analysis. The first caesium metal was produced in [[1881]]. Historically, the most important use for caesium has been in research and development, primarily in chemical and electrical applications.
-==Occurrence==
+This compound is not usually used in experiments, because the extraction of cesium is very expensive and it behaves very much like [[rubidium hydroxide]] and [[potassium hydroxide]].
-[[Image:Pollucite(CesiumMineral)USGOV.jpg|thumb|left|Pollucite, a caesium mineral]]
-An alkali metal, caesium occurs in [[lepidolite]], [[pollucite]] ([[hydrate]]d [[silicate]] of [[aluminium]] and caesium) and within other sources. One of the world's most significant and rich sources of this metal is located at [[Bernic Lake]] in [[Manitoba]]. The deposits there are estimated to contain 300,000 [[tonne|metric tons]] of pollucite at an average of 20% caesium.
-It can be isolated by [[electrolysis]] of fused [[cyanide]] and in a number of other ways.
+==Applications==
-Exceptionally pure and gas-free caesium can be made by the thermal decomposition of caesium [[azide]].
-The primary compounds of caesium are its [[chloride]] and its [[nitrate]]. The price of caesium in 1997 was about $US 30 per gram.
-; See also [[:category:Caesium minerals|Caesium minerals]].
-==Isotopes==
+* Cesium is most notably used in [[atomic clock]]s, which are accurate to seconds in many thousands of years.
-Caesium has at least 39 known isotopes which is more than any other element, except [[francium]]. The [[atomic mass]]es of these isotopes range from 112 to 151.
+*Cesium-134 has been used as a measure of cesium output by the [[nuclear power]] industry. This isotope is used because, while it is less prevalent than either <sup>133</sup>Cs or <sup>137</sup>Cs, <sup>134</sup>Cs is produced solely by nuclear reactions. <sup>135</sup>Cs has also been used for this purpose.
-Even though this element has the largest number of isotopes, it has only one naturally occurring stable [[isotope]], <sup>133</sup>Cs, the other isotopes (except for the isotopes noted on this page) have [[half-life|half-lives]] from a few days to fractions of a second. The [[radiogenic]] isotope [[Caesium-137|<sup>137</sup>Cs]] has been used in hydrologic studies, analogous to the use of [[tritium|<sup>3</sup>H]]. <sup>137</sup>Cs is produced from the detonation of [[nuclear weapon]]s and is produced in [[nuclear power plant]]s, and notably from the 1986 [[Chernobyl]] meltdown.  Beginning in 1945 with the commencement of [[nuclear testing]], <sup>137</sup>Cs was released into the [[Earth's atmosphere|atmosphere]] where it is absorbed readily into solution and is returned to the surface of the earth as a component of [[radioactive fallout]]. Once <sup>137</sup>Cs 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 can be estimated as a function of time.
+*Like other group one elements, cesium has great affinity for [[oxygen]] and is used as a "[[getter]]" in [[vacuum tube]]s.
-Cesium-137 has a half-life of 30.17 years.  It decomposes to
+*This metal is also used in [[photoelectric cell]]s because it readily emits [[electron]]s.
-barium-137m (a short-lived product of decay) then to a form of
+*Cesium is used as a [[catalyst]] in the [[hydrogenation]] of certain [[organic compound]]s.
-nonradioactive barium.
+*Cesium hydroxide is used to etch [[silicon]], exposing [[octahedron|octahedral]] planes. This technique can create pyramids and regularly shaped etch pits for uses such as [[MEMS]] (microelectromechanical systems).
+*Cesium chloride is used in the production of electrically conducting [[glass]]es.
+*[[Radioisotope]]s of cesium (in the form of cesium chloride) are used in [[nuclear medicine]], including the treatment of [[cancer]].
+*[[Cesium fluoride]] is widely used in [[organic chemistry]] as a [[Base (chemistry)|base]] and as a source of [[anhydrous]] [[fluoride]] ion.
+*Cesium vapor is used in many common [[magnetometers]].
+*Because of their high density, [[cesium chloride]] solutions are commonly used in molecular biology experiments for the isolation of nucleic acids by a technique known as "density gradient ultracentrifugation."
+*Recently, this metal has been used in [[ion propulsion]] systems.
+*Cesium-137 is an extremely common isotope in industrial applications such as:
+**moisture density gauges;
+**leveling gauges;
+**thickness gauges;
+**well-logging devices (used to measure the thickness of rock strata).
 ==Precautions==
-All alkaline metals are highly reactive. Caesium, being one of the heavier [[alkaline]] [[metal]]s, is also one of the most reactive and is highly [[explosive]] when it comes in contact with [[water]] (even cold water, or ice). [[Caesium hydroxide]] is an extremely strong [[Base (chemistry)|base]], and can attack [[glass]].
-Caesium compounds are encountered rarely by most people.  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. Rats fed caesium in place of potassium in their diet die, so this element cannot replace potassium in function.
+All alkali metals are extremely reactive. As one of the heavier alkali [[metal]]s, cesium is also one of the most reactive. As noted above, it is highly [[explosive]] when it comes in contact with [[water]] or ice. [[Cesium hydroxide]] is an extremely strong base and can attack [[glass]].
+All cesium compounds should be regarded as mildly toxic, because of its chemical similarity to [[potassium]]. Contact with large amounts can cause [[hyperirritability]] and [[spasms]]. On the other hand, such amounts would not ordinarily be encountered in nature, so cesium is not a major chemical environmental pollutant. Rats fed with cesium in place of potassium in their diet were found to die, so this element cannot replace potassium in function.
-The [[isotope]]s <sup>134</sup>Cs and <sup>137</sup>Cs (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.
+The [[isotope]]s <sup>134</sup>Cs and <sup>137</sup>Cs (present in the [[biosphere]] in small amounts as a result of radiation leaks) represent a radioactivity burden that varies depending on location. Radioactive cesium does not accumulate in the body as effectively as many other fission products, such as radioactive iodine or strontium.
 ==See also==
-*[[Cs-137]]
-*[[Goiânia accident]] - a major radioactive contamination incident involving a small rod of caesium chloride.
+*[[Periodic table]]
-*[[:Category:Caesium compounds|Caesium compounds]]
+*[[Chemical element]]
+==Notes==
+<references/>
 ==References==
-<div class="references-small"><references/>
+*[http://periodic.lanl.gov/elements/55.html Los Alamos National Laboratory - Cesium] Retrieved December 9, 2007.
-*[http://periodic.lanl.gov/elements/55.html Los Alamos National Laboratory - Cesium]</div>
 ==External links==
-{{Commons|Caesium}}
+All links retrieved December 3, 2023.
-{{wiktionary|caesium}}
-*[http://www.wm-rsc.co.uk/students/interactive/gcse/alkali/AlkaliMetals2.wmv Video of alkaline metals reacting with water]
-*[http://www.webelements.com/webelements/elements/text/Cs/index.html WebElements.com &ndash; Caesium]
-*[http://www.cs.rochester.edu/users/faculty/nelson/cesium/cesium_faq.html FAQ from alt.cesium newsgroup]
-[[Category:Chemical elements]]
+*[http://www.webelements.com/webelements/elements/text/Cs/index.html WebElements.com &ndash; Caesium]
-[[Category:Alkali metals]]
+*[http://www.cs.rochester.edu/users/faculty/nelson/cesium/cesium_faq.html FAQ from alt.cesium newsgroup]
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+[[Category:Chemistry]]
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