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= Cesium =
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Introduction
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Caesium (IUPAC spelling; 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 C, 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 C. It is the least electronegative stable element,
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 caesium is considered a hazardous material, and the radioisotopes
present a significant health and environmental hazard.
Spelling
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'Caesium' is the spelling recommended by the International Union of
Pure and Applied Chemistry (IUPAC). The American Chemical Society
(ACS) has used the spelling 'cesium' since 1921, following 'Webster's
New International Dictionary'. The element was named after the Latin
word 'caesius', meaning "bluish grey". 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.
Physical properties
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Of all elements that are solid at room temperature, caesium is the
softest: it has a hardness of 0.2 Mohs. It is a very ductile, pale
metal, which darkens in the presence of trace amounts of oxygen. 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 C, making it one of the few
elemental metals that are liquid near room temperature. The others are
rubidium (39 C), francium (estimated at 27 C), mercury (−39 C), and
gallium (30 C); bromine is also liquid at room temperature (melting at
−7.2 C), but it is a halogen and not a metal. Mercury is the only
stable elemental metal with a known melting point lower than caesium.
In addition, the metal has a rather low boiling point, 641 C, the
lowest of all stable metals other than mercury. 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 if they are metals.
Caesium forms alloys with the other alkali metals, gold, and mercury
(amalgams). At temperatures below 650 °C, 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. 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 C. A few
amalgams have been studied: is black with a purple metallic lustre,
while CsHg is golden-coloured, also with a metallic lustre.
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.
Its compounds burn with a blue or violet colour.
Allotropes
============
Caesium exists in the form of different allotropes, one of them a
dimer called dicaesium.
Chemical properties
=====================
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. It reacts with ice
at temperatures as low as −116 C. 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. Caesium can be stored in vacuum-sealed borosilicate glass
ampoules. In quantities of more than about 100 g, caesium is shipped
in hermetically sealed, stainless steel containers.
The chemistry of caesium is similar to that of other alkali metals, in
particular rubidium, the element above caesium in the periodic table.
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. 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. Some slight
differences arise from the fact that it has a higher atomic mass and
is more electropositive than other (nonradioactive) alkali metals.
Caesium is the most electropositive chemical element. The caesium ion
is also larger and less "hard" than those of the lighter alkali
metals.
Compounds
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Most caesium compounds contain the element as the cation , which binds
ionically to a wide variety of anions. One noteworthy exception is the
caeside anion (), and others are the several suboxides (see section on
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., CsFn with n > 1) under high pressure. This
prediction needs to be validated by further experiments.
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,
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 ), bismuth,
cadmium, copper, iron, and lead are also poorly soluble.
Caesium hydroxide (CsOH) is hygroscopic and strongly basic. It rapidly
etches the surface of semiconductors such as silicon. CsOH has been
previously regarded by chemists as the "strongest base", reflecting
the relatively weak attraction between the large Cs+ ion and OH−; 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.
A stoichiometric mixture of caesium and gold will react to form yellow
caesium auride (Cs+Au−) upon heating. The auride anion here behaves as
a pseudohalogen. The compound reacts violently with water, yielding
caesium hydroxide, metallic gold, and hydrogen gas; in liquid ammonia
it can be reacted with a caesium-specific ion exchange resin to
produce tetramethylammonium auride. The analogous platinum compound,
red caesium platinide (), contains the platinide ion that behaves as a
.
Complexes
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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 137Cs+ must be
separated from large amounts of nonradioactive K+.
Halides
=========
Caesium fluoride (CsF) is a hygroscopic white solid that is widely
used in organofluorine chemistry as a source of fluoride anions.
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. 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", 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. The CsCl structure is preferred because Cs+ has an ionic
radius of 174 pm and 181 pm.
Oxides
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More so than the other alkali metals, caesium forms numerous binary
compounds with oxygen. When caesium burns in air, the superoxide is
the main product. The "normal" caesium oxide () forms yellow-orange
hexagonal crystals, and is the only oxide of the anti-cadmium chloride
type. It vaporizes at 250 °C, and decomposes to caesium metal and the
peroxide caesium peroxide at temperatures above 400 °C. In addition to
the superoxide and the ozonide caesium ozonide, several brightly
coloured suboxides have also been studied. These include , , ,
(dark-green), CsO, , as well as . The latter may be heated in a vacuum
to generate . Binary compounds with sulfur, selenium, and tellurium
also exist.
Isotopes
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Caesium has 41 known isotopes, ranging in mass number (i.e. number of
nucleons in the nucleus) from 112 to 152. Several of these are
synthesized from lighter elements by the slow neutron capture process
(S-process) inside old stars and by the R-process in supernova
explosions. The only stable caesium isotope is 133Cs, with 78
neutrons. Although it has a large nuclear spin (+), nuclear magnetic
resonance studies can use this isotope.
The radioactive 135Cs has a very long half-life of about 2.3 million
years, the longest of all radioactive isotopes of caesium. 137Cs and
134Cs have half-lives of 30 and two years, respectively. 137Cs
decomposes to a short-lived 137mBa by beta decay, and then to
nonradioactive barium, while 134Cs transforms into 134Ba directly. 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 134mCs (with a half-life
of just under 3 hours), all are very unstable and decay with
half-lives of a few minutes or less.
The isotope 135Cs is one of the long-lived fission products of uranium
produced in nuclear reactors. However, this fission product yield is
reduced in most reactors because the predecessor, 135Xe, is a potent
neutron poison and frequently transmutes to stable 136Xe before it can
decay to 135Cs.
The beta decay from 137Cs to 137mBa results in gamma radiation as the
137mBa relaxes to ground state 137Ba, with the emitted photons having
an energy of 0.6617 MeV. 137Cs and 90Sr are the principal medium-lived
products of nuclear fission, and the prime sources of radioactivity
from spent nuclear fuel after several years of cooling, lasting
several hundred years. Those two isotopes are the largest source of
residual radioactivity in the area of the Chernobyl disaster. Because
of the low capture rate, disposing of 137Cs through neutron capture is
not feasible and the only current solution is to allow it to decay
over time.
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. 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. With
nuclear weapons testing in the 1950s through the 1980s, 137Cs was
released into the atmosphere and returned to the surface of the earth
as a component of radioactive fallout. It is a ready marker of the
movement of soil and sediment from those times.
Occurrence
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Caesium is a relatively rare element, estimated to average 3 parts per
million in the Earth's crust. It is the 45th most abundant element and
36th among the metals. Caesium is 30 times less abundant than
rubidium, with which it is closely associated, chemically.
Due to its large ionic radius, caesium is one of the "incompatible
elements". 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 () may contain only 0.002% caesium. Consequently,
caesium is found in few minerals. Percentage amounts of caesium may be
found in beryl () and avogadrite (), up to 15 wt% Cs2O in the closely
related mineral pezzottaite (), up to 8.4 wt% Cs2O in the rare mineral
londonite (), and less in the more widespread rhodizite. The only
economically important ore for caesium is pollucite , 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.
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. 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%. Commercial pollucite contains more than 19% caesium. 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. At the present
rate of world mine production of 5 to 10 metric tons per year,
reserves will last for thousands of years.
Production
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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.
In the acid digestion, the silicate pollucite rock is dissolved with
strong acids, such as hydrochloric (HCl), sulfuric (), 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 (),
caesium iodine chloride (), or caesium hexachlorocerate (). 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 (). 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 solution.
Roasting pollucite with calcium carbonate and calcium chloride yields
insoluble calcium silicates and soluble caesium chloride. Leaching
with water or dilute ammonia () 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.
Most of the mined caesium (as salts) is directly converted into
caesium formate (HCOO−Cs+) 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 oilbbl per year of caesium formate solution. The
primary smaller-scale commercial compounds of caesium are caesium
chloride and nitrate.
Alternatively, caesium metal may be obtained from the purified
compounds derived from the ore. Caesium chloride and the other caesium
halides can be reduced at 700 to 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.
The metal can also be isolated by electrolysis of fused caesium
cyanide (CsCN). Exceptionally pure and gas-free caesium can be
produced by 390 °C thermal decomposition of caesium azide , which can
be produced from aqueous caesium sulfate and barium azide. In vacuum
applications, caesium dichromate can be reacted with zirconium to
produce pure caesium metal without other gaseous products.
: + 2 → 2 + 2 +
The price of 99.8% pure caesium (metal basis) in 2009 was about 10
$/g, but the compounds are significantly cheaper.
History
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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 ,
meaning . Caesium was the first element to be discovered with a
spectroscope, which had been invented by Bunsen and Kirchhoff only a
year previously.
To obtain a pure sample of caesium, 44,000 litre of mineral water had
to be evaporated to yield 240 kg 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 () 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 g of rubidium chloride
and 7.3 g of caesium chloride from the initial 44,000 litres of
mineral water.
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). 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 (). In
reality, the product was probably a colloidal mixture of the metal and
caesium chloride. The electrolysis of the aqueous solution of chloride
with a mercury cathode produced a caesium amalgam which readily
decomposed under the aqueous conditions. The pure metal was eventually
isolated by the Swedish chemist Carl Setterberg while working on his
doctorate with Kekulé and Bunsen. In 1882, he produced caesium metal
by electrolysing caesium cyanide, avoiding the problems with the
chloride.
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.
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. 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. 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".
Petroleum exploration
=======================
The largest present-day use of nonradioactive caesium is in caesium
formate drilling fluids for the extractive oil industry. Aqueous
solutions of caesium formate (HCOO−Cs+)--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.
The high density of the caesium formate brine (up to 2.3 g/cm3, or
19.2 pounds per gallon), 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. Caesium formate brine can be blended with
potassium and sodium formates to decrease the density of the fluids to
that of water (1.0 g/cm3, or 8.3 pounds per gallon). Furthermore, it
is biodegradable and may be recycled, which is important in view of
its high cost (about $4,000 per barrel in 2001). 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 solutions) sometimes do; they also require less cleanup and
reduce disposal costs.
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. 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." These clocks
measure frequency with an error of 2 to 3 parts in 1014, which
corresponds to an accuracy of 2 nanoseconds per day, or one second in
1.4 million years. The latest versions are more accurate than 1 part
in 1015, about 1 second in 20 million years. The caesium standard is
the primary standard for standards-compliant time and frequency
measurements. Caesium clocks regulate the timing of cell phone
networks and the Internet.
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 , the
unperturbed ground-state hyperfine transition frequency of the
caesium-133 atom, to be when expressed in the unit Hz, which is equal
to s−1."
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.
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 , have a
low threshold voltage for emission of electrons. The range of
photoemissive devices using caesium include optical character
recognition devices, photomultiplier tubes, and video camera tubes.
Nevertheless, germanium, rubidium, selenium, silicon, tellurium, and
several other elements can be substituted for caesium in
photosensitive materials.
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.
The element is used as an internal standard in spectrophotometry. Like
other alkali metals, caesium has a great affinity for oxygen and is
used as a "getter" in vacuum tubes. Other uses of the metal include
high-energy lasers, vapour glow lamps, and vapour rectifiers.
Centrifugation fluids
=======================
The high density of the caesium ion makes solutions of caesium
chloride, caesium sulfate, and caesium trifluoroacetate () useful in
molecular biology for density gradient ultracentrifugation. This
technology is used primarily in the isolation of viral particles,
subcellular organelles and fractions, and nucleic acids from
biological samples.
Chemical and medical use
==========================
Relatively few chemical applications use caesium. 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.
Caesium fluoride enjoys a niche use in organic chemistry as a base and
as an anhydrous source of fluoride ion.
Friestad, Gregory K.; Branchaud, Bruce P.; Navarrini, Walter and
Sansotera, Maurizio (2007) "Cesium Fluoride" in 'Encyclopedia of
Reagents for Organic Synthesis', John Wiley & Sons. 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 (TLD): 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
137Ba 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. It has been used in agriculture, cancer
treatment, and the sterilization of food, sewage sludge, and surgical
equipment. Radioactive isotopes of caesium in radiation devices were
used in the medical field to treat certain types of cancer, 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.
Caesium-137 has been employed in a variety of industrial measurement
gauges, including moisture, density, levelling, and thickness gauges.
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.
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.
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. Xenon was used
in the experimental spacecraft Deep Space 1 launched in 1998.
Nevertheless, field-emission electric propulsion thrusters that
accelerate liquid metal ions such as caesium have been built.
Caesium nitrate is used as an oxidizer and pyrotechnic colorant to
burn silicon in infrared flares, such as the LUU-19 flare, because it
emits much of its light in the near infrared spectrum. 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. 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.
Magnetohydrodynamic (MHD) power-generating systems were researched,
but failed to gain widespread acceptance. Caesium metal has also been
considered as the working fluid in high-temperature Rankine cycle
turboelectric generators.
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.
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.
Health and safety hazards
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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.
The median lethal dose (LD50) for caesium chloride in mice is 2.3 g
per kilogram, which is comparable to the LD50 values of potassium
chloride and sodium chloride. 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.
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.
It is highly pyrophoric: the autoignition temperature of caesium is
−116 C, 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.
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.
Radiocaesium follows potassium and tends to accumulate in plant
tissues, including fruits and vegetables. 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.
Accumulation of caesium-137 in lakes has been a great concern after
the Chernobyl disaster. 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; smaller amounts may cause infertility
and cancer. 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".
See also
======================================================================
*
* 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 137Cs lost-source incident
between 1980 and 1989
External links
======================================================================
* [
https://www.periodicvideos.com/videos/055.htm Caesium or Cesium] at
'The Periodic Table of Videos' (University of Nottingham)
*
[
https://web.archive.org/web/20171104215850/http://richannel.org/the-modern-alchemist-reacting-fluorine-with-caesium
View the reaction of Caesium (most reactive metal in the periodic
table) with Fluorine (most reactive non-metal)] courtesy of The Royal
Institution.
*
License
=========
All content on Gopherpedia comes from Wikipedia, and is licensed under CC-BY-SA
License URL:
http://creativecommons.org/licenses/by-sa/3.0/
Original Article:
http://en.wikipedia.org/wiki/Cesium
