Caesium (British spelling and IUPAC spelling) or cesium (American
spelling)[note 1] is a chemical element with symbol Cs and atomic
number 55. It is a soft, silvery-gold alkali metal with a melting
point of 28.5 °C (83.3 °F), which makes it one of only
five elemental metals that are liquid at or near room
Caesium has physical and chemical properties
similar to those of rubidium and potassium. The most reactive of all
metals, it is pyrophoric and reacts with water even at −116 °C
(−177 °F). It is the least electronegative element, with a
value of 0.79 on the Pauling scale. It has only one stable isotope,
Caesium is mined mostly from pollucite, while the
radioisotopes, especially caesium-137, a fission product, are
extracted from waste produced by nuclear reactors.
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. In 1967, acting
on Einstein's proof that the speed of light is the most constant
dimension in the universe, the
International System of Units
International System of Units used two
specific wave counts from an emission spectrum of caesium-133 to
co-define the second and the metre. Since then, caesium has been
widely used in highly accurate atomic clocks.
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 ecological hazard in the environment.
1.1 Physical properties
1.2 Chemical properties
4.1 Petroleum exploration
4.2 Atomic clocks
4.3 Electric power and electronics
4.4 Centrifugation fluids
4.5 Chemical and medical use
4.6 Nuclear and isotope applications
4.7 Other uses
5 Health and safety hazards
6 See also
9 External links
High-purity caesium-133 stored in argon.
Caesium is the softest element (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 (83.3 °F), making it one of the few elemental
metals that are liquid near room temperature. Mercury is the only
elemental metal with a known melting point lower than caesium.[note
3] In addition, the metal has a rather low boiling point,
641 °C (1,186 °F), the lowest of all metals other than
mercury. Its compounds burn with a blue or violet
Caesium crystals (golden) compared to rubidium crystals (silvery)
Caesium forms alloys with the other alkali metals, gold, and mercury
(amalgams). At temperatures below 650 °C (1,202 °F), 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 (−108 °F). A few
amalgams have been studied: CsHg
2 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. Thus caesium absorbs violet light preferentially and hence
Addition of a small amount of caesium to cold water is explosive.
Caesium metal is highly reactive and very pyrophoric. It ignites
spontaneously in air, and reacts explosively with water even at low
temperatures, more so than the other alkali metals (first group of the
periodic table). It reacts with solid water at temperatures as low
as −116 °C (−177 °F). 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 grams (3.5 oz), 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.[note 4] Some small 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.[note 5] The caesium ion is also
larger and less "hard" than those of the lighter alkali metals.
Ball-and-stick model of the cubic coordination of Cs and Cl in CsCl
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−), and others are the several suboxides
(see section on oxides below).
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. Double salts are often less soluble, and the low
solubility of caesium aluminium sulfate is exploited in refining Cs
from ores. The double salt with antimony (such as CsSbCl
4), bismuth, cadmium, copper, iron, and lead are also poorly
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, but a number
of compounds that do not dissolve in water, such as n-butyllithium and
sodium amide, are more basic.
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 (Cs2Pt), contains the platinide ion that behaves
as a pseudochalcogen.
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+.
Monatomic caesium halide wires grown inside double-wall carbon
nanotubes (TEM image).
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
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 Cl−
More so than the other alkali metals, caesium forms numerous binary
compounds with oxygen. When caesium burns in air, the superoxide CsO
2 is the main product. The "normal" caesium oxide (Cs
2O) forms yellow-orange hexagonal crystals, and is the only oxide
of the anti-CdCl
2 type. It vaporizes at 250 °C (482 °F), and
decomposes to caesium metal and the peroxide Cs
2 at temperatures above 400 °C (752 °F). In addition
to the superoxide and the ozonide CsO
3, several brightly coloured suboxides have also been
studied. These include Cs
3O (dark-green), CsO, Cs
2, as well as Cs
2. The latter may be heated in a vacuum to generate Cs
2O. Binary compounds with sulfur, selenium, and tellurium also
Main article: Isotopes of caesium
Caesium has 39 known isotopes, ranging in mass number (i.e. number of
nucleons in the nucleus) from 112 to 151. 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 (7/2+), nuclear
magnetic resonance studies can use this isotope at a resonating
frequency of 11.7 MHz.
Decay of caesium-137
The radioactive 135Cs has a very long half-life of about
2.3 million years, 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 is a strong emission of gamma
radiation. 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 allowed to decay over
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
Pollucite, a caesium mineral
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 the 36th among the metals. Nevertheless, it is more
abundant than such elements as antimony, cadmium, tin, and tungsten,
and two orders of magnitude more abundant than mercury and silver; it
is 3.3% as abundant as rubidium, with which it is closely associated,
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, the alkali evaporite minerals sylvite (KCl)
and carnallite (KMgCl
2O) may contain only 0.002% caesium. Consequently, Cs is found in few
minerals. Percentage amounts of caesium may be found in beryl (Be
6) and avogadrite ((K,Cs)BF
4), up to 15 wt% Cs2O in the closely related mineral pezzottaite
(Cs(Be2Li)Al2Si6O18), up to 8.4 wt% Cs2O in the rare mineral
28), and less in the more widespread rhodizite. The only
economically important ore for caesium is pollucite Cs(AlSi
6), 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.
One of the world's most significant and richest sources of caesium is
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.
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 (H
4), 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
7), caesium iodine chloride (Cs
2ICl), or caesium hexachlorocerate (Cs
6)). 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
2O). 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
Roasting pollucite with calcium carbonate and calcium chloride yields
insoluble calcium silicates and soluble caesium chloride. Leaching
with water or dilute ammonia (NH
4OH) 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
can 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 12,000 barrels (1,900 m3) 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 800 °C (1,292 to
1,472 °F) 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 (734 °F) thermal decomposition of
caesium azide CsN
3, 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
7 + 2 Zr → 2 Cs + 2 ZrO
The price of 99.8% pure caesium (metal basis) in 2009 was about US$10
per gram ($280 per ounce), but the compounds are significantly
Gustav Kirchhoff (left) and
Robert Bunsen (centre) discovered caesium
with their newly invented spectroscope.
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 caesius, meaning sky-blue.[note 6]
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 litres
(9,700 imp gal; 12,000 US gal) of mineral water
had to be evaporated to yield 240 kilograms (530 lb) 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 grams (0.32 oz) of rubidium chloride and 7.3
grams (0.26 oz) 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
2Cl). 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 anode produced a caesium amalgam
which readily decomposed under the aqueous conditions. The pure
metal was eventually isolated by the German 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.
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
Caesium also was, and still is, 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
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".
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
The high density of the caesium formate brine (up to
2.3 g·cm−3, 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·cm−3, 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 ZnBr
2 solutions) sometimes do; they also require less cleanup and reduce
Atomic clock ensemble at the U.S. Naval Observatory
FOCS-1, a continuous cold caesium fountain atomic clock in
Switzerland, started operating in 2004 at an uncertainty of one second
in 30 million years
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 corresponding 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
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 K
2CsSb, 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 caesium 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.
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
The high density of the caesium ion makes solutions of caesium
chloride, caesium sulfate, and caesium trifluoroacetate (Cs(O
3)) 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
Caesium chloride powder
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.
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
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.
Schematics of an electrostatic ion thruster developed for use with
caesium or mercury fuel
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.
used in the experimental spacecraft
Deep Space 1
Deep Space 1 launched in
1998. Nevertheless, field-emission electric propulsion
thrusters that accelerate liquid metal ions such as caesium have been
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 is used to reduce the radar signature of exhaust plumes in the
SR-71 Blackbird military aircraft.
Caesium and rubidium have been
added as a carbonate to glass because they reduces 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
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
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
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
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, as caesium formate in petroleum drilling
fluids because it is much less toxic than alternatives, though it is
The fire diamond hazard sign for caesium metal
Caesium metal 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
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".
Goiânia accident, a major radioactive contamination incident
involving a rod of caesium-137 chloride
Acerinox accident, a caesium-137 contamination accident
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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
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.
^ Along with rubidium (39 °C [102 °F]), francium
(estimated at 27 °C [81 °F]), mercury (−39 °C
[−38 °F]), and gallium (30 °C [86 °F]); bromine is
also liquid at room temperature (melting at −7.2 °C,
19 °F), but it is a halogen, not a metal. Preliminary work with
copernicium and flerovium suggests that they are gaseous metals at
^ The radioactive element francium may also have a lower melting
point, but its radioactivity prevents enough of it from being isolated
for direct testing.
^ It differs from this value in caesides, which contain the Cs−
anion and thus have caesium in the −1 oxidation state.
Additionally, 2013 calculations by Mao-sheng Miao indicate that under
conditions of extreme pressure (greater than 30 GPa), the inner
5p electrons could form chemical bonds, where caesium would behave as
the seventh 5p element. This discovery indicates that higher caesium
fluorides with caesium in oxidation states from +2 to +6 could exist
under such conditions.
^ Francium's electropositivity has not been experimentally measured
due to its high radioactivity. Measurements of the first ionization
energy of francium suggest that its relativistic effects may lower its
reactivity and raise its electronegativity above that expected from
^ 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.
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