From en.wikipedia.org:
[cerium] [Featured article] [date=January 2018] [date=November 2022]
[engvar=en-OED] CAESIUM (IUPAC spelling;[1] also spelled CESIUM in American
English) is a chemical element; it has symbol CS and atomic number 55. It is
a soft, silvery-golden alkali metal with a melting point of [28.5], which
makes it one of only five elemental metals that are liquid at or near room
temperature. Caesium has physical and chemical properties similar to those of
rubidium and potassium. It is pyrophoric and reacts with water even at
[−116]. It is the least electronegative stable element,<!--This is sourced as
most electropostive element below, do not change without a reliable modern
secondary source.--> with a value of 0.79 on the Pauling scale. It has only
one stable isotope, caesium-133. Caesium is mined mostly from pollucite.
Caesium-137, a fission product, is extracted from waste produced by nuclear
reactors. It has the largest atomic radius of all elements whose radii have
been measured or calculated, at about 260 picometres.
The German chemist Robert Bunsen and physicist Gustav Kirchhoff discovered
caesium in 1860 by the newly developed method of flame spectroscopy. The
first small-scale applications for caesium were as a "getter" in vacuum tubes
and in photoelectric cells. Caesium is widely used in highly accurate atomic
clocks. In 1967, the International System of Units began using a specific
hyperfine transition of neutral caesium-133 atoms to define the basic unit of
time, the second.
Since the 1990s, the largest application of the element has been as caesium
formate for drilling fluids, but it has a range of applications in the
production of electricity, in electronics, and in chemistry. The radioactive
isotope caesium-137 has a half-life of about 30 years and is used in medical
applications, industrial gauges, and hydrology. Nonradioactive caesium
compounds are only mildly toxic, but the pure metal's tendency to react
explosively with water means that it is considered a hazardous material, and
the radioisotopes present a significant health and environmental hazard.
** Spelling
_Caesium_ is the spelling recommended by the International Union of Pure and
Applied Chemistry (IUPAC).[2] The American Chemical Society (ACS) has used the
spelling _cesium_ since 1921,[3][4] following _Webster's New International
Dictionary_. The element was named after the Latin word _caesius_, meaning
"bluish grey".[5] In medieval and early modern writings _caesius_ was spelled
with the ligature _æ_ as _cæsius_; hence, an alternative but now old-fashioned
orthography is _cæsium_. More spelling explanation at ae/oe vs e.
** Characteristics
*** Physical properties
left Of all elements that are solid at room temperature, caesium is the
softest: it has a hardness of Mohs 0.2. It is a very ductile, pale metal,
which darkens in the presence of trace amounts of oxygen.[6][7][8] When in the
presence of mineral oil (where it is best kept during transport), it loses its
metallic lustre and takes on a duller, grey appearance. It has a melting
point of [28.5], making it one of the few elemental metals that are liquid
near room temperature. The others are rubidium ([39]), francium (estimated at
[27]), mercury ([−39]), and gallium ([30]); bromine is also liquid at room
temperature (melting at [−7.2]), but it is a halogen and not a metal. Mercury
is the only stable elemental metal with a known melting point lower than
caesium.[9] In addition, the metal has a rather low boiling point, [641], the
lowest of all stable metals other than mercury.<ref name="RSC"/> Copernicium
and flerovium 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.<ref name=CRNL>[last1=Mewes ]<ref name=liquid>[last1=Mewes ]
left Caesium forms alloys with the other alkali metals, gold, and mercury
(amalgams). At temperatures below [650], it does not alloy with cobalt, iron,
molybdenum, nickel, platinum, tantalum, or tungsten. It forms well-defined
intermetallic compounds with antimony, gallium, indium, and thorium, which are
photosensitive.<ref name="USGS"/> It mixes with all the other alkali metals
(except lithium); the alloy with a molar distribution of 41% caesium, 47%
potassium, and 12% sodium has the lowest melting point of any known metal
alloy, at [−78].<ref name="KanerACS"/>[10] A few amalgams have been studied:
[CsHg] is black with a purple metallic lustre, while CsHg is golden-coloured,
also with a metallic lustre.[11]
The golden colour of caesium comes from the decreasing frequency of light
required to excite electrons of the alkali metals as the group is descended.
For lithium through rubidium this frequency is in the ultraviolet, but for
caesium it enters the blue–violet end of the spectrum; in other words, the
plasmonic frequency of the alkali metals becomes lower from lithium to
caesium. Thus caesium transmits and partially absorbs violet light
preferentially while other colours (having lower frequency) are reflected;
hence it appears yellowish.[12] Its compounds burn with a blue[13]<ref
name="flametests"/> or violet[14] colour.
*** Allotropes
Caesium exists in the form of different allotropes; one of them is a dimer,
called dicaesium.[15]<!-- More exist... See
https://www.knowledgedoor.com/2/elements_handbook/allotropes.html#cesium -->
*** Chemical properties
left Caesium metal is highly reactive and pyrophoric. It ignites
spontaneously in air, and reacts explosively with water even at low
temperatures, more so than the other alkali metals.<!--YES INCLUDING
FRANCIUM--><ref name="USGS"/> It reacts with ice at temperatures as low as
[−116].<ref name="KanerACS"/> Because of this high reactivity, caesium metal
is classified as a hazardous material. It is stored and shipped in dry,
saturated hydrocarbons such as mineral oil. It can be handled only under
inert gas, such as argon. However, a caesium-water explosion is often less
powerful than a sodium-water explosion with a similar amount of sodium. This
is because caesium explodes instantly upon contact with water, leaving little
time for hydrogen to accumulate.[16] Caesium can be stored in vacuum-sealed
borosilicate glass ampoules. In quantities of more than about [100], caesium
is shipped in hermetically sealed, stainless steel containers.<ref
name="USGS"/>
The chemistry of caesium is similar to that of other alkali metals, in
particular rubidium, the element above caesium in the periodic table.<ref
name="greenwood"/> As expected for an alkali metal, the only common oxidation
state is +1. It differs from this value in caesides, which contain the [Cs]
anion and thus have caesium in the −1 oxidation state.<ref name="caeside2"/>
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.[17][18] Some slight differences arise from the fact that it
has a higher atomic mass and is more electropositive than other
(nonradioactive) alkali metals.[19] Caesium is the most electropositive
chemical element.<ref name="KanerACS"/> The caesium ion is also larger and
less "hard" than those of the lighter alkali metals.
*** Compounds
Most caesium compounds contain the element as the cation [Cs], which binds
ionically to a wide variety of anions. One noteworthy exception is the
caeside anion ([Cs]),<ref name="caeside2"/> and others are the several
suboxides (see section [#Oxides] below). More recently, caesium is predicted
to behave as a p-block element and capable of forming higher fluorides with
higher oxidation states (i.e., [Cs] with _n_ > 1) under high pressure.[20]
This prediction needs to be validated by further experiments.[21]
Salts of [Cs] are usually colourless unless the anion itself is coloured.
Many of the simple salts are hygroscopic, but less so than the corresponding
salts of lighter alkali metals. The phosphate,[22] acetate, carbonate,
halides, oxide, nitrate, and sulfate salts are water-soluble. Its double
salts are often less soluble, and the low solubility of caesium aluminium
sulfate is exploited in refining Cs from ores. The double salts with antimony
(such as [CsSbCl]), bismuth, cadmium, copper, iron, and lead are also poorly
soluble.<ref name="USGS"/>
Caesium hydroxide (CsOH) is hygroscopic and strongly basic.<ref
name="greenwood"/> It rapidly etches the surface of semiconductors such as
silicon.[23] CsOH has been previously regarded by chemists as the "strongest
base", reflecting the relatively weak attraction between the large [Cs] ion
and [O];<ref name="CRC74"/> it is indeed the strongest Arrhenius base;
however, a number of compounds such as _n_ -butyllithium, sodium amide, sodium
hydride, caesium hydride, etc., which cannot be dissolved in water as reacting
violently with it but rather only used in some anhydrous polar aprotic
solvents, are far more basic on the basis of the Brønsted–Lowry acid–base
theory.<ref name="greenwood"/>
A stoichiometric mixture of caesium and gold will react to form yellow caesium
auride ([Cs]) upon heating. The auride anion here behaves as a pseudohalogen.
The compound reacts violently with water, yielding caesium hydroxide, metallic
gold, and hydrogen gas; it dissolves in liquid ammonia and then can be reacted
with a caesium-specific ion exchange resin to produce tetramethylammonium
auride. The analogous platinum compound, red caesium platinide ([Cs2Pt]),
contains the platinide ion that behaves as a [pseudo chalcogen].[24]
**** Complexes
Like all metal cations, [Cs] forms complexes with Lewis bases in solution.
Because of its large size, [Cs] usually adopts coordination numbers greater
than 6, the number typical for the smaller alkali metal cations. This
difference is apparent in the 8-coordination of CsCl. This high coordination
number and softness (tendency to form covalent bonds) are properties exploited
in separating [Cs] from other cations in the remediation of nuclear wastes,
where [137] must be separated from large amounts of nonradioactive [K].[25]
**** Halides
[[File:CsX@DWNT.jpg|thumb|upright|Monatomic caesium halide wires grown inside
double-wall carbon nanotubes (TEM image).[26]]] Caesium fluoride (CsF) is a
hygroscopic white solid that is widely used in organofluorine chemistry as a
source of fluoride anions.[27] Caesium fluoride has the halite structure,
which means that the [Cs] and [F] pack in a cubic closest packed array as do
[Na] and [Cl] in sodium chloride.<ref name="greenwood"/> Notably, caesium and
fluorine have the lowest and highest electronegativities, respectively, among
all the known elements.
Caesium chloride (CsCl) crystallizes in the simple cubic crystal system. Also
called the "caesium chloride structure",<ref name="HollemanAF"/> this
structural motif is composed of a primitive cubic lattice with a two-atom
basis, each with an eightfold coordination; the chloride atoms lie upon the
lattice points at the edges of the cube, while the caesium atoms lie in the
holes in the centre of the cubes. This structure is shared with CsBr and CsI,
and many other compounds that do not contain Cs. In contrast, most other
alkaline halides have the sodium chloride (NaCl) structure.<ref
name="HollemanAF"/> The CsCl structure is preferred because [Cs] has an ionic
radius of 174 pm and [Cl] 181 pm.[28]
**** Oxides
More so than the other alkali metals, caesium forms numerous binary compounds
with oxygen. When caesium burns in air, the superoxide [CsO] is the main
product.[29] The "normal" caesium oxide ([Cs]) forms yellow-orange hexagonal
crystals,[30] and is the only oxide of the anti-[CdCl] type.[31] It vaporizes
at [250], and decomposes to caesium metal and the peroxide [Cs] at
temperatures above [400]. In addition to the superoxide and the ozonide
[CsO],[32][33] several brightly coloured suboxides have also been studied.[34]
These include [Cs], [Cs], [Cs], [Cs] (dark-green[35]), CsO, [Cs],[36] as well
as [Cs].[37][38] The latter may be heated in a vacuum to generate [Cs].<ref
name="ReferenceA"/> Binary compounds with sulfur, selenium, and tellurium also
exist.<ref name="USGS"/>
*** Isotopes
[Isotopes of caesium] Caesium has 41 known isotopes, ranging in mass number
from 112 to 152. The only stable caesium isotope is [133], with 78 neutrons.
The radioactive <sup> 135 </sup> Cs has a very long half-life of about 1.33
million years, the longest of all radioactive isotopes of caesium. <sup> 137
</sup> Cs and <sup> 134 </sup> Cs 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
metastable nuclear isomers exist. Other than <sup>134m</sup>Cs (with a
half-life of just under 3 hours), all are very unstable and decay with
half-lives of a few minutes or less.[39]
The isotope <sup>135</sup>Cs is one of the long-lived fission products of
uranium produced in nuclear reactors.[40] However, this fission product yield
is reduced in most reactors because the predecessor, <sup> 135 </sup> Xe, is a
potent neutron poison and frequently transmutes to stable <sup> 136 </sup> Xe
before it can decay to <sup>135</sup>Cs.[41][42]
The beta decay from <sup>137</sup>Cs to <sup>137m</sup>Ba results in gamma
radiation as the <sup>137m</sup>Ba relaxes to ground state <sup>137</sup>Ba,
with the emitted photons having an energy of 0.6617 MeV.[43] [137] and <sup>
90 </sup> Sr are the principal medium-lived products of nuclear fission, and
the prime sources of radioactivity from spent nuclear fuel from several years
to several hundred years after removal.[44] Those two isotopes are the largest
source of residual radioactivity in the area of the Chernobyl disaster.[45]
Because of the low capture rate, disposing of [137] through neutron capture is
not feasible and the only current solution is to allow it to decay over
time.[46]
Almost all caesium produced from nuclear fission comes from the beta decay of
originally more neutron-rich fission products, passing through various
isotopes of iodine and xenon.[47] 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.[48] With nuclear weapons testing in the
1950s through the 1980s, [137] was released into the atmosphere and returned
to the surface of the earth as a component of radioactive fallout, becoming a
marker of the movement of soil and sediment from those times.<ref
name="USGS"/>
Although it has a large nuclear spin ([7]+), nuclear magnetic resonance can
use the stable [133] isotope.[49]
*** Occurrence
[:Category:Caesium minerals] Caesium is a relatively rare element, estimated
to average 3 parts per million in the Earth's crust.[50] It is the 45th most
abundant element and 36th among the metals.[51] Caesium is 30 times less
abundant than rubidium, with which it is closely associated, chemically.<ref
name="USGS"/>
Due to its large ionic radius, caesium is one of the "incompatible
elements".[52] During magma crystallization, caesium is concentrated in the
liquid phase and crystallizes last. Therefore, the largest deposits of
caesium are zone pegmatite ore bodies formed by this enrichment process.
Because caesium does not substitute for potassium as readily as rubidium does,
the alkali evaporite minerals sylvite (KCl) and carnallite ([KMgCl]) may
contain only 0.002% caesium. Consequently, caesium is found in few minerals.
Percentage amounts of caesium may be found in beryl ([Be]) and avogadrite
([(K,Cs)BF]), up to 15 wt% Cs<sub>2</sub>O in the closely related mineral
pezzottaite ([Cs]), up to 8.4 wt% Cs<sub>2</sub>O in the rare mineral
londonite ([(Cs,K)Al]), and less in the more widespread rhodizite.<ref
name="USGS"/> The only economically important ore for caesium is pollucite
[Cs(AlSi], which is found in a few places around the world in zoned
pegmatites, associated with the more commercially important lithium minerals,
lepidolite and petalite. Within the pegmatites, the large grain size and the
strong separation of the minerals results in high-grade ore for mining.[53]
The world's most significant and richest known source of caesium is the Tanco
Mine at Bernic Lake in Manitoba, Canada, estimated to contain 350,000 metric
tons of pollucite ore, representing more than two-thirds of the world's
reserve base.<ref name="Cerny"/><ref name="USGS-Cs2"/> 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%.[54] Commercial pollucite contains more than 19%
caesium.[55] The Bikita pegmatite deposit in Zimbabwe is mined for its
petalite, but it also contains a significant amount of pollucite. Another
notable source of pollucite is in the Karibib Desert, Namibia.<ref
name="USGS-Cs2"/> At the present rate of world mine production of 5 to 10
metric tons per year, reserves will last for thousands of years.<ref
name="USGS"/>
** Production
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.<ref name="USGS"/>[56]
In the acid digestion, the silicate pollucite rock is dissolved with strong
acids, such as hydrochloric (HCl), sulfuric ([H]), hydrobromic (HBr), or
hydrofluoric (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]), caesium iodine chloride
([Cs]), or caesium hexachlorocerate ([Cs]). 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 caesium
alum ([CsAl(SO]). The aluminium sulfate component is converted to insoluble
aluminium oxide by roasting the alum with carbon, and the resulting product is
leached with water to yield a [Cs] solution.<ref name="USGS"/>
Roasting pollucite with calcium carbonate and calcium chloride yields
insoluble calcium silicates and soluble caesium chloride. Leaching with water
or dilute ammonia ([NH]) 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.<ref name="USGS"/>
Most of the mined caesium (as salts) is directly converted into caesium
formate ([H]) for applications such as oil drilling. To supply the developing
market, Cabot Corporation built a production plant in 1997 at the Tanco mine
near Bernic Lake in Manitoba, with a capacity of [12000] per year of caesium
formate solution.[57] The primary smaller-scale commercial compounds of
caesium are caesium chloride and nitrate.<ref name="CEC"/>
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] with calcium or barium, and caesium metal distilled from the
result. In the same way, the aluminate, carbonate, or hydroxide may be
reduced by magnesium.<ref name="USGS"/>
The metal can also be isolated by electrolysis of fused caesium cyanide
(CsCN). Exceptionally pure and gas-free caesium can be produced by [390]
thermal decomposition of caesium azide [CsN], which can be produced from
aqueous caesium sulfate and barium azide.<ref name="Burt"/> In vacuum
applications, caesium dichromate can be reacted with zirconium to produce pure
caesium metal without other gaseous products.[58]
[Cs] + 2 [Zr] → 2 [Cs] + 2 [ZrO] + [Cr]
The price of 99.8% pure caesium (metal basis) in 2009 was about [10], but the
compounds are significantly cheaper.<ref name="USGS-Cs2"/>
** History
In 1860, Robert Bunsen and Gustav Kirchhoff discovered caesium in the mineral
water from Dürkheim, Germany. Because of the bright blue lines in the
emission spectrum, they derived the name from the Latin word [la], meaning
[bluish grey].<ref group=note>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._[59]<ref
name="BuKi1861"/>[60] Caesium was the first element to be discovered with a
spectroscope, which had been invented by Bunsen and Kirchhoff only a year
previously.<ref name="KanerACS"/>
To obtain a pure sample of caesium, [44,000] of mineral water had to be
evaporated to yield [240] of concentrated salt solution. The alkaline earth
metals were precipitated either as sulfates or oxalates, leaving the alkali
metal in the solution. After conversion to the nitrates and extraction with
ethanol, a sodium-free mixture was obtained. From this mixture, the lithium
was precipitated by ammonium carbonate. Potassium, rubidium, and caesium form
insoluble salts with chloroplatinic 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 by fractional
crystallization. After reduction of the hexachloroplatinate with hydrogen,
caesium and rubidium were separated by the difference in solubility of their
carbonates in alcohol. The process yielded [9.2] of rubidium chloride and
[7.3] of caesium chloride from the initial 44,000 litres of mineral water.[61]
From the caesium chloride, the two scientists estimated the atomic weight of
the new element at 123.35 (compared to the currently accepted one of
132.9).<ref name="BuKi1861"/> 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 a subchloride ([Cs]). In reality, the product was
probably a colloidal mixture of the metal and caesium chloride.[62] The
electrolysis of the aqueous solution of chloride with a mercury cathode
produced a caesium amalgam which readily decomposed under the aqueous
conditions.<ref name="BuKi1861"/> The pure metal was eventually isolated by
the Swedish chemist Carl Setterberg while working on his doctorate with Kekulé
and Bunsen.<ref name="Weeks"/> In 1882, he produced caesium metal by
electrolysing caesium cyanide, avoiding the problems with the chloride.[63]
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
radio vacuum tubes, where it had two functions; as a getter, it removed excess
oxygen after manufacture, and as a coating on the heated cathode, it increased
the electrical conductivity. Caesium was not recognized as a high-performance
industrial metal until the 1950s.[64] Applications for nonradioactive caesium
included photoelectric cells, photomultiplier tubes, optical components of
infrared spectrophotometers, catalysts for several organic reactions, crystals
for scintillation counters, and in magnetohydrodynamic power generators.<ref
name="USGS"/> Caesium is also used as a source of positive ions in secondary
ion mass spectrometry (SIMS).
Since 1967, the International System of Measurements has based the primary
unit of time, the second, on the properties of caesium. The International
System of Units (SI) defines the second as the duration of [9,192,631,770]
cycles at the microwave frequency of the spectral line corresponding to the
transition between two hyperfine energy levels of the ground state of
caesium-133.[65] The 13th General Conference on Weights and Measures of 1967
defined a second as: "the duration of [9,192,631,770] cycles of microwave
light absorbed or emitted by the hyperfine transition of caesium-133 atoms in
their ground state undisturbed by external fields".
** Applications
*** Petroleum exploration
<!-- linked from "Formate" article --> The largest present-day use of
nonradioactive caesium is in caesium formate drilling fluids for the
extractive oil industry.<ref name="USGS"/> Aqueous solutions of caesium
formate ([H])—made by reacting caesium hydroxide with formic acid—were
developed in the mid-1990s for use as oil well drilling and completion fluids.
The function of a drilling fluid is to lubricate drill bits, to bring rock
cuttings to the surface, and to maintain pressure on the formation during
drilling of the well. Completion fluids assist the emplacement of control
hardware after drilling but prior to production by maintaining the
pressure.<ref name="USGS"/>
The high density of the caesium formate brine (up to 2.3 g/cm<sup>3</sup>, or
19.2 pounds per gallon),[66] 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.<ref name="Down"/> 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/cm<sup>3</sup>, or 8.3
pounds per gallon). Furthermore, it is biodegradable and may be recycled,
which is important in view of its high cost (about $4,000 per barrel in
2001).[67] Alkali formates are safe to handle and do not damage the producing
formation or downhole metals as corrosive alternative, high-density brines
(such as zinc bromide [ZnBr] solutions) sometimes do; they also require less
cleanup and reduce disposal costs.<ref name="USGS"/>
*** Atomic clocks
Caesium-based atomic clocks use the electromagnetic transitions in the
hyperfine structure of caesium-133 atoms as a reference point. The first
accurate caesium clock was built by Louis Essen in 1955 at the National
Physical Laboratory in the UK.[68] 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."<ref name="USNO"/> These clocks measure frequency
with an error of 2 to 3 parts in 10<sup>14</sup>, 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 10<sup>15</sup>, about 1
second in 20 million years.<ref name="USGS"/> The caesium standard is the
primary standard for standards-compliant time and frequency measurements.[69]
Caesium clocks regulate the timing of cell phone networks and the
Internet.[70]
**** Definition of the second
The second, symbol s, is the SI unit of time. The BIPM restated its
definition at its 26th conference in 2018: "[The second] is defined by taking
the fixed numerical value of the caesium frequency [Δ_ν_ <sub> Cs </sub>], the
unperturbed ground-state hyperfine transition frequency of the caesium-133
atom, to be [9192631770] when expressed in the unit Hz, which is equal to
s<sup>−1</sup>."[71]
*** Electric power and electronics
Caesium vapour thermionic generators are low-power devices that convert heat
energy to electrical energy. In the two-electrode vacuum tube converter,
caesium neutralizes the space charge near the cathode and enhances the current
flow.[72]
Caesium is also important for its photoemissive properties, converting light
to electron flow. It is used in photoelectric cells because caesium-based
cathodes, such as the intermetallic compound [K], have a low threshold voltage
for emission of electrons.[73] The range of photoemissive devices using
caesium include optical character recognition devices, photomultiplier tubes,
and video camera tubes.[74][75] Nevertheless, germanium, rubidium, selenium,
silicon, tellurium, and several other elements can be substituted for caesium
in photosensitive materials.<ref name="USGS"/>
Caesium iodide (CsI), bromide (CsBr) and fluoride (CsF) crystals are employed
for scintillators in scintillation counters widely used in mineral exploration
and particle physics research to detect gamma and X-ray radiation. Being a
heavy element, caesium provides good stopping power with better detection.
Caesium compounds may provide a faster response (CsF) and be less hygroscopic
(CsI).
Caesium vapour is used in many common magnetometers.[76]
The element is used as an internal standard in spectrophotometry.[77] Like
other alkali metals, caesium has a great affinity for oxygen and is used as a
"getter" in vacuum tubes.[78] Other uses of the metal include high-energy
lasers, vapour glow lamps, and vapour rectifiers.<ref name="USGS"/>
*** Centrifugation fluids
The high density of the caesium ion makes solutions of caesium chloride,
caesium sulfate, and caesium trifluoroacetate ([Cs(O]) useful in molecular
biology for density gradient ultracentrifugation.[79] This technology is used
primarily in the isolation of viral particles, subcellular organelles and
fractions, and nucleic acids from biological samples.[80]
*** Chemical and medical use
Relatively few chemical applications use caesium.[81] Doping with caesium
compounds enhances the effectiveness of several metal-ion catalysts for
chemical synthesis, such as acrylic acid, anthraquinone, ethylene oxide,
methanol, phthalic anhydride, styrene, methyl methacrylate monomers, and
various olefins. It is also used in the catalytic conversion of sulfur
dioxide into sulfur trioxide in the production of sulfuric acid.<ref
name="USGS"/> <!--No way: Caesium metal is also used in ferrous and nonferrous
metallurgy and in the purification of carbon dioxide as it absorbs gases and
other impurities, while molten hydroxide (CsOH) has been used in the
desulfurizing of heavy crude oil.<ref name="USGS"/> -->
Caesium fluoride enjoys a niche use in organic chemistry as a base[82] and as
an anhydrous source of fluoride ion.[83] Caesium salts sometimes replace
potassium or sodium salts in organic synthesis, such as cyclization,
esterification, and polymerization. Caesium has also been used in
thermoluminescent radiation dosimetry <small>(TLD)</small>: When exposed to
radiation, it acquires crystal defects that, when heated, revert with emission
of light proportionate to the received dose. Thus, measuring the light pulse
with a photomultiplier tube can allow the accumulated radiation dose to be
quantified.
*** Nuclear and isotope applications
Caesium-137 is a radioisotope commonly used as a gamma-emitter in industrial
applications. Its advantages include a half-life of roughly 30 years, its
availability from the nuclear fuel cycle, and having <sup> 137 </sup> Ba 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.[84]
It has been used in agriculture, cancer treatment, and the sterilization of
food, sewage sludge, and surgical equipment.<ref name="USGS"/>[85] Radioactive
isotopes of caesium in radiation devices were used in the medical field to
treat certain types of cancer,[86] 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.[87][88] Caesium-137 has been employed in a variety of industrial
measurement gauges, including moisture, density, levelling, and thickness
gauges.[89] It has also been used in well logging devices for measuring the
electron density of the rock formations, which is analogous to the bulk
density of the formations.[90]
Caesium-137 has been used in hydrologic studies analogous to those with
tritium. As a daughter product of fission bomb testing from the 1950s through
the mid-1980s, caesium-137 was released into the atmosphere, where it was
absorbed readily into solution. Known year-to-year variation within that
period allows correlation with soil and sediment layers. Caesium-134, and to
a lesser extent caesium-135, have also been used in hydrology to measure the
caesium output by the nuclear power industry. While they are less prevalent
than either caesium-133 or caesium-137, these bellwether isotopes are produced
solely from anthropogenic
sources.[91]<!--https://books.google.com/books?id=pWDQnxd-r1UC&pg=PT360
&pg=PT12 -->
*** Other uses
Caesium and mercury were used as a propellant in early ion engines designed
for spacecraft propulsion on very long interplanetary or extraplanetary
missions. The fuel was ionized by contact with a charged tungsten electrode.
But corrosion by caesium on spacecraft components has pushed development in
the direction of inert gas propellants, such as xenon, which are easier to
handle in ground-based tests and do less potential damage to the
spacecraft.<ref name="USGS"/> Xenon was used in the experimental spacecraft
Deep Space 1 launched in 1998.[92][93] Nevertheless, field-emission electric
propulsion thrusters that accelerate liquid metal ions such as caesium have
been built.[94]
Caesium nitrate is used as an oxidizer and pyrotechnic colorant to burn
silicon in infrared flares,[95] such as the LUU-19 flare,[96] because it emits
much of its light in the near infrared spectrum.[97] Caesium compounds may
have been used as fuel additives to reduce the radar signature of exhaust
plumes in the Lockheed A-12 CIA reconnaissance aircraft.[98] Caesium and
rubidium have been added as a carbonate to glass because they reduce
electrical conductivity and improve stability and durability of fibre optics
and night vision devices. Caesium fluoride or caesium aluminium fluoride are
used in fluxes formulated for brazing aluminium alloys that contain
magnesium.<ref name="USGS"/>
Magnetohydrodynamic (MHD) power-generating systems were researched, but failed
to gain widespread acceptance.[99] Caesium metal has also been considered as
the working fluid in high-temperature Rankine cycle turboelectric
generators.[100]
Caesium salts have been evaluated as antishock reagents following the
administration of arsenical drugs. Because of their effect on heart rhythms,
however, they are less likely to be used than potassium or rubidium salts.
They have also been used to treat epilepsy.<ref name="USGS"/>
Caesium-133 can be laser cooled and used to probe fundamental and
technological problems in quantum physics. It has a particularly convenient
Feshbach spectrum to enable studies of ultracold atoms requiring tunable
interactions.[101]
** Health and safety hazards
{{chembox |container_only = yes |Section7 = {{chembox Hazards | ExternalSDS =
| GHSPictograms = [GHS flame] [GHS corrosion] | GHSSignalWord = Danger |
HPhrases = [H260] | PPhrases = [P223] | GHS_ref = [102] | NFPA-H = 3 | NFPA-F
= 4 | NFPA-R = 3 | NFPA-S = w | NFPA_ref =
}}
}}
[[File:AirDoseChernobylVector.svg|thumb|right|upright=1.4|alt=Graph of percentage of the radioactive output by each nuclide that form after a nuclear fallout vs. logarithm of time after the incident. In curves of various colours, the predominant source of radiation are depicted in order: Te-132/I-132 for the first five or so days; I-131 for the next five; Ba-140/La-140 briefly; Zr-95/Nb-95 from day 10 until about day 200; and finally Cs-137. Other nuclides producing radioactivity, but not peaking as a major component are Ru, peaking at about 50 days, and Cs-134 at around 600 days.|The portion of the total radiation dose (in air) contributed by each isotope plotted against time after the Chernobyl disaster. Caesium-137 became the primary source of radiation about 200 days after the accident.[103]]]
Nonradioactive caesium compounds are only mildly toxic, and nonradioactive
caesium is not a significant environmental hazard. Because biochemical
processes can confuse and substitute caesium with potassium, excess caesium
can lead to hypokalemia, arrhythmia, and acute cardiac arrest, but such
amounts would not ordinarily be encountered in natural sources.[104][105]
The median lethal dose (LD<sub>50</sub>) for caesium chloride in mice is 2.3
g/kg, which is comparable to the LD<sub>50</sub> values of potassium chloride
and sodium chloride.[106] The principal use of nonradioactive caesium is as
caesium formate in petroleum drilling fluids because it is much less toxic
than alternatives, though it is more costly.<ref name="Down"/>
Elemental caesium is one of the most reactive elements and is highly explosive
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.<ref
name="USGS"/>
It is highly pyrophoric: the autoignition temperature of caesium is [−116],
and it ignites explosively in air to form caesium hydroxide and various
oxides. Caesium hydroxide is a very strong base, and will rapidly corrode
glass.[107]
The isotopes 134 and 137 are present in the biosphere 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.[108] Radiocaesium follows potassium and
tends to accumulate in plant tissues, including fruits and
vegetables.[109][110][111] 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 fungal sporocarps.[112] Accumulation of caesium-137 in lakes has been a
great concern after the Chernobyl disaster.[113][114] Experiments with dogs
showed that a single dose of 3.8 millicuries (140 MBq, 4.1 μg of caesium-137)
per kilogram is lethal within three weeks;[115] smaller amounts may cause
infertility and cancer.[116] The International 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".[117]<!--
10.1016/S0098-8472(01)00124-1-->
** See also
- [Caesium-137#Incidents and accidents]
- Acerinox accident , a caesium-137 contamination accident in 1998
- Goiânia accident , a major radioactive contamination incident in 1987 involving caesium-137
- Kramatorsk radiological accident , a <sup> 137 </sup> Cs lost-source incident between 1980 and 1989
** Notes
[group="note" ]
** References
[30em]
** External links
[Caesium.ogg]
- Caesium or Cesium (see https://www.periodicvideos.com/videos/055.htm) at _The Periodic Table of Videos_ (University of Nottingham)
- View the reaction of Caesium (most reactive metal in the periodic table) with Fluorine (most reactive non-metal) (see https://web.archive.org/web/20171104215850/http://richannel.org/the-modern-alchemist-reacting-fluorine-with-caesium) courtesy of The Royal Institution.
- [title=Molecular CsF <sub> 5 </sub> and CsF <sub> 2 </sub> <sup> + </sup> ]
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