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Caesium

From Archania
Caesium
Atomic number 55
Symbol Cs
Group 1 (alkali metals)
Boiling point 671 °C
Electron configuration [Xe] 6s1
Melting point 28.5 °C
Period 6
Main isotopes 133Cs, 137Cs
Phase STP Solid
Block s
Oxidation states +1
Wikidata Q1108

Caesium (symbol Cs, atomic number 55) is a soft, silvery-golden alkali metal in group 1 of the periodic table. Near room temperature it is one of only a few elemental metals that melt just above ambient conditions (melting point ~28.5 °C). It is highly reactive – more so than any stable metal – tarnishing in air and erupting in flames on contact with water. These extreme chemical properties have made caesium both useful (for example, in atomic clocks and photoelectric devices) and hazardous (handling the pure metal is challenging). Its only stable isotope is Cs-133, whose precise hyperfine energy transition defines the second in the International System of Units. A notable radioactive cousin, Cs-137 (half-life ~30.2 years), is a common byproduct of nuclear fission and has important roles in industry, medicine, and environmental monitoring. Caesium compounds have found industrial applications (today the largest use is in high-density caesium formate drilling fluids) and occasional laboratory uses. Overall, caesium is a scarce, heavy alkali metal distinguished by its large atomic radius, strong reducing power, and key role in modern technology.

Overview and Key Facts

Caesium is an alkali metal (group 1, s-block) in period 6 of the periodic table. Its chemical symbol is Cs and its atomic number is 55. It is the heaviest stable alkali and is classified as a metallic element. At standard conditions (solid at 20 °C), it appears as a soft, silvery-golden metal that can be cut with a knife. It is so soft that it may deform; it is also very light for its size (density ≈ 1.9 g/cm³, about half that of iron). The only stable oxidation state of caesium is +1 (Cs⁺), which it attains by losing its single valence electron. Caesium metal is pyrophoric, meaning it ignites spontaneously in air, and it reacts explosively with water to form caesium hydroxide (CsOH) and hydrogen gas. Because of these hazards, metallic caesium must be handled under inert gas or oil.

Some key facts at a glance:

  • Category: Alkali metal (Group 1, Period 6, s-block).
  • Atomic number: 55.
  • Standard atomic weight: 132.905 (average of natural isotopic mixture, essentially 133 for pure Cs-133).
  • Electron configuration: [Xe] 6s¹ (with one valence electron in the 6s orbital).
  • Common oxidation state: +1 (caesium form Cs⁺).
  • Phase at STP: Solid (silvery-golden metal). Melting point ~28.5 °C (very near room temperature) and boiling point ~671 °C.
  • Cassio physical state: Extremely soft metal (Mohs hardness ≈ 0.2) that floats on water.
  • Density: ~\(1.93\) g·cm⁻³ at 20 °C (less dense than water).
  • Color: Silvery or pale gold when freshly cut (turns grayish quickly in air).

Caesium is one of the least abundant elements in the Earth’s crust (~3 parts per million by weight, around the 45th most common element). It is typically found in certain rare mineral ores (pollucite, lepidolite, etc.) rather than broadly distributed like sodium or potassium.

Atomic Structure and Electron Configuration

As element 55, caesium atoms contain 55 protons in the nucleus and 55 electrons in neutral form. The ground-state electron configuration is [Xe] 6s¹, meaning it has the filled shell structure of xenon plus one electron in the 6s orbital. Because of this single valence electron, caesium behaves similarly to other alkali metals (lithium, sodium, potassium, etc.) and tends to form Cs⁺ ions easily by losing that electron. The valence shell (the sixth energy level) has configuration 6s¹; the filled inner shells match xenon’s configuration, leading to chemical behavior dominated by that outer electron.

Several periodic trends characterize caesium’s atomic properties:

  • Atomic radius: Caesium has an exceptionally large atomic radius. The van der Waals radius is about 265 picometers (pm), and its metallic atomic radius is ~243 pm. This is the largest atomic radius measured among all elements. In group 1 metals, atomic radius increases down the group (Cs is larger than K or Rb), so the outer electron is very far from the nucleus.
  • Ionization energy: The first ionization energy of caesium is extremely low (≈ 376 kJ/mol, or about 3.89 eV per atom). This low energy reflects how weakly the outer 6s electron is held. Caesium has the smallest first ionization energy of any stable element (only the radioactive francium might be lower, but francium is too short-lived to study easily). This makes Cs extremely easy to ionize (lose an electron).
  • Electronegativity: on the Pauling scale caesium’s electronegativity is only about 0.79, the lowest of all stable elements. This very low electronegativity means Cs has little tendency to attract electrons in bonds; it readily gives up its valence electron instead.

These trends (large size, weakly bound valence electron, low electronegativity) are why caesium is such a potent reducing agent and reactive metal. Structurally, in the solid metal caesium atoms adopt a body-centered cubic (bcc) lattice at standard conditions, reflecting the loosely packed arrangement of large atoms each contributing one conduction electron.

Isotopes and Nuclear Properties

Naturally occurring caesium is almost entirely the isotope Cs-133 (with 78 neutrons). This isotope is stable and has 100% natural abundance. Caesium-133 has nuclear spin 7/2, which gives rise to hyperfine structure in its ground-state energy levels. In fact, the hyperfine transition of ground-state Cs-133 atoms at microwave frequency (9 192 631 770 Hz) is used to define the SI second. In nuclear magnetic resonance (NMR), ^133Cs (spin 7/2) is an NMR-active nucleus (though less commonly used than ^19F or ^1H).

No other stable isotopes of Cs exist, so all other isotopes are radioactive (radioisotopes). Key radioisotopes include:

  • Cs-137 (80 neutrons): Half-life ~30.17 years. It is a major fission product of uranium and plutonium. Decays by beta emission to metastable barium-137m, then emits a 0.662 MeV gamma-ray to reach stable ^137Ba. Because of its half-life and gamma emissions, Cs-137 is important in medical applications (radiotherapy), industrial gauges, and for tracing and dating (see below).
  • Cs-134 (79 neutrons): Half-life ~2.07 years. Produced in nuclear reactors (neutron capture on ^133Cs). Decays by beta emission with a complex gamma spectrum.
  • Cs-135 (80 neutrons): Half-life ~2.3 million years. Also a fission product; decays by beta to ^135Ba. Its long half-life makes it a concern in nuclear waste management (long-lived stockpiling).
  • Cs-136 and others exist with shorter half-lives (minutes to days), but the above are most significant.

Cs-137’s uses as a gamma source exploit its penetrating radiation. It also serves as a radiotracer in environmental studies (e.g. tracking sediment deposition since the 1940s) and even wine authentication (irradiation from nuclear testing gives a known baseline of ^137Cs). Biologically, Cs chemistry resembles potassium and rubidium, so radioactive Cs tends to spread uniformly through soft tissues if ingested. Fortunately, it is excreted relatively quickly (biological half-life tens of days). Anti-radiation treatments like Prussian blue can trap Cs⁺ in the gut and accelerate its removal if contaminated.

One important note: in nuclear reactors or weapons, caesium can be produced but not separated easily. Used nuclear fuel contains mixed Cs isotopes along with stable Cs-133. Specific isotope separation of ^137Cs or ^134Cs is difficult and rarely economically done. Therefore, sources of cesium radiation (e.g. Cs-137) often come from processing reactor waste or decommissioned sources.

Allotropes and Typical Compounds

Caesium does not have allotropes in the same sense as carbon or sulfur (distinct molecular forms). The elemental metal has one main crystal form at ambient pressure (body-centered cubic) and may adopt other metal structures under extreme pressure, but that is generally of academic interest. At high temperature in the gas phase, caesium can exist briefly as a diatomic molecule Cs₂ before dissociating; this diatomic “dicaesium” species is sometimes noted in theoretical discussions, but it is not a stable form at ordinary conditions. In practice, elemental caesium is just the metal or vapor.

The chemistry of caesium is dominated by its +1 oxidation state. Typical bonding is ionic, with Cs⁺ pairing with anions. Major compounds include:

  • Oxides and related: Caesium forms several oxygen compounds. Pure caesium oxide (Cs₂O) is known (yellowish solid) but Cs easily oxidizes further to superoxide CsO₂ (the usual product of burning Cs in air) and even ozonide CsO₃ under special conditions. Cs₂O₂ (peroxide) and various “suboxides” (CsₙO, Csₙ₊₂O₂, etc.) have been characterized. These are all highly reactive (especially peroxides and superoxides with water).
  • Hydroxide: The caesium hydroxide CsOH is a colorless crystal or glass that is extremely hygroscopic (water-absorbing) and a very strong base. It dissolves in water to give the Cs⁺ and OH⁻ ions; its large cation means its hydroxide is one of the strongest alkali bases (it even etches glass). CsOH is so caustic that it was once called a “superbase” of the alkali group.
  • Halides: Caesium forms stable ionic halides: CsF, CsCl, CsBr, CsI. These salts are colorless or white solids. Notably, CsCl crystallizes in the simple cubic CsCl structure (each Cs⁺ in a cube of eight Cl⁻ neighbors), rather than the sodium chloride structure. Many of these halides have high solubility in water. Caesium chloride has a very high density and is used in laboratory applications (density gradient centrifugation of DNA and viruses). CsI crystals are widely used in scintillation detectors (Cesium iodide scintillator for X-rays and gamma rays).
  • Other salts: Caesium carbonate (Cs₂CO₃) and nitrate (CsNO₃) are common salts; Cs₂CO₃ is sometimes used in chemical synthesis. Caesium sulfate, phosphate, and others can be made. Because Cs⁺ is a large, monovalent cation, many double salts and complexes are known — for example, Cs-alum (CsAl(SO₄)₂), or cryptates/crown ether complexes (Cs⁺ fits well into certain large crown ether rings).
  • Hydride: CsH (cesium hydride) is an ionic solid (Cs⁺H⁻) that forms when Cs metal reacts with hydrogen under pressure. It is stable but decomposes easily back to H₂.
  • Nitrides: The nitride Cs₃N is reported when Cs metal reacts violently with nitrogen. (Alkali metals can form M₃N nitrides, although they oxidize quickly in air.) The azide CsN₃ (cesium azide) is a well-known explosive, analogous to potassium azide.

In summary, caesium’s chemistry is straightforward: it produces ionic compounds with characteristic group-1 behavior, often with water-soluble, high-melting ionic solids. The large Cs⁺ size leads to relatively low lattice energies, so many Cs salts are soluble and easily hydrated. Because of its reactivity, many caesium compounds (especially oxides and hydroxides) must be handled carefully to avoid moisture and air.

Physical Properties

Caesium metal exhibits properties typical of a heavy alkali metal, plus some extremes. At room temperature (around 20–25 °C), caesium is metallic and silvery-golden in color. It is extremely soft: it can be cut with a knife and will even sort of flow or deform under slight pressure. It is ductile and malleable like other alkali metals. Its density is around 1.93 g/cm³ at ambient conditions (some sources give 1.93–1.94 at 20 °C), which is surprisingly low for a metal (it is less dense than water).

Because of the low melting point (28.5 °C), pure caesium metal melts at just slightly above room temperature, becoming a silvery liquid when warmed. (Thus, it is sometimes listed as a liquid metal at just above 285 K.) It boils at about 671 °C under atmospheric pressure. This very low melting point (lowest of all elemental metals except mercury) and low boiling point give caesium one of the lowest latent heats of fusion and vaporization among metals (on the order of 2.1 kJ/mol for fusion, 65–70 kJ/mol for vaporization). These values reflect the weak metallic bonding of large Cs atoms.

In the solid state, caesium adopts a body-centered cubic (bcc) crystal structure at room temperature and atmospheric pressure (the same structure as sodium and potassium metal). At sufficiently low temperature under 14 K (very cold cryogenic), it undergoes a slight distortion (tetragonal) before reverting to bcc on warming – but for most practical purposes it is simply bcc (space group Im3m). Under high pressure (tens of gigapascals), caesium even becomes orthorhombic and eventually close-packed, but these are exotic phases.

Thermally, caesium metal is a good conductor of heat (thermal conductivity ~36 W/(m·K) at room temperature, somewhat lower than lighter alkalis but still respectable) and an excellent electrical conductor (electrical resistivity ~0.205 μΩ·m at 20 °C). It is paramagnetic in the solid state (slightly attracted to magnetic fields, unlike heavier groups of metals which may be ferromagnetic or diamagnetic).

Optically and spectroscopically, caesium has notable features: it has a very low work function (~2.135 eV), among the lowest of any element, which makes Cs surfaces extremely efficient at emitting electrons when illuminated (this was crucial in early photoelectric effect experiments and in photocathodes). Atomic vapor of caesium has bright emission lines in the red and near-infrared; the strongest lines are the “D-lines” at 852.1 nm and 894.3 nm (the transitions 6²P₃/₂→6²S₁/₂ and 6²P₁/₂→6²S₁/₂, respectively). These transitions are commonly used in laser cooling and spectroscopy of Cs atoms. The 6²S₁/₂ ground state itself is split into two hyperfine levels (F=3 and F=4), whose microwave transition defines the SI second (at 9,192,631,770 Hz). In solid compounds, Cs-containing crystals can be used as scintillators (e.g. CsI emits light when struck by high-energy radiation).

Overall, caesium’s physical constants are: atomic weight ~132.9055, atomic or ionic radius large (~265 pm van der Waals), melting point 28.5 °C, boiling point 671 °C, density ~1.90–1.93 g/cm³, and high thermal/electrical conductivity, with a single 6s electron contributing to metallic bonding. Its near-room-temperature melting makes it one of only five elemental metals (with rubidium, francium, gallium, and mercury) that are liquid close to ordinary temperatures.

Caesium is the most reactive of the stable alkali metals, and indeed one of the most reactive elements known. This follows down-group periodic trends: the larger size and lower ionization energy make Cs the strongest reductant in its group. In practice, Cs metal reacts vigorously and often explosively with almost every substrate:

  • Water: Cs reacts violently even at very low temperatures. Dropping caesium into water produces CsOH (a caustic solution) and hydrogen gas, in an exothermic reaction so vigorous that the hydrogen usually ignites. Even solid ice (-116°C) can be attacked by Cs. The reaction is: \[2 Cs + 2 H₂O → 2 CsOH + H₂↑.\]
  • Air/Oxygen: In moist air, Cs metal rapidly oxidizes. Unlike light alkalis that produce full oxides (Na₂O) or peroxides (K₂O₂), burning caesium primarily yields the superoxide CsO₂ (due to easy one-electron oxidation of O₂). Pure Cs₂O (oxide) can form under dry conditions but often decomposes on heating. Even ozone (O₃) can react with Cs to form odd Cs–O species. Caesium metal is spontaneously flammable in air and must be stored under oil or inert atmosphere.
  • Halogens: Caesium reacts with all halogens (F₂, Cl₂, Br₂, I₂) to form the corresponding halide (CsF, CsCl, CsBr, CsI). The reactions are typically vigorous (cesium guzzles chlorine to form white CsCl). Because Cs has a very negative standard reduction potential (E° ≈ –2.92 V for Cs⁺/Cs), it reduces most halogens readily.
  • Nonmetals: Caesium metal can reduce many nonmetals and oxoacids. For example, it reacts with nitrogen to form caesium nitride (Cs₃N) and with sulfur, selenium, and tellurium to form sulphides (Cs₂S), selenides, etc. With hydrogen under certain conditions it forms the ionic hydride CsH. It even reduces some metal oxides under strong conditions, e.g. reacting with the PbO₂ (lead dioxide) in glass to produce metallic lead (formerly used in glass manufacturing).

In aqueous solution, the Cs⁺ ion is very stable. Caesium salts are generally water-soluble; in water Cs behaves like a very strong (but typical) alkali: CsOH is strongly basic (pH ~10 for low concentrations) and fully dissociates to Cs⁺ and OH⁻, but because Cs⁺ is so large, CsOH is actually a stronger base than KOH or NaOH. Caesium salts do not hydrolyze (Cs⁺ does not form acidity in water) and do not form precipitates with OH⁻ (unlike some smaller Group 1 cations that might interact with structure-inducing agents). Caesium carbonate (Cs₂CO₃) and acetate (CsOAc) are stronger bases in solution than their lithium or sodium analogues.

An interesting aspect of group trends: Cs is so large that chemists often compare CsOH to “one of the strongest bases.” It is indeed extremely caustic and deliquescent (absorbing moisture from air). Caesium oxide (Cs₂O and CsO₂) can react with water even more exothermically. This means that everyday contact with CsOH is dangerous, as it can burn skin. Similarly, alkali metals form solvated electrons in liquid ammonia; liquid Cs in ammonia yields a deep blue solution containing [Cs⁺ + e⁻] that is even more reactive than Na/NH₃.

Chemists also note how Cs compares to Rb and K. In reactivity series, Cs tops the stable list. It displaces all the others from their compounds (for instance, halides or nitrates of K, Rb, Na). Alkali-metal reduction potentials show Cs as the hardest reductant (e.g. can reduce water and most other metals). In terms of acid–base chemistry, CsOH is stronger than KOH in aqueous strength (though all are very strong bases). Caesium’s behavior consistently extends group-1 trends downward: largest atomic/ionic radius, lowest ionization energy, highest reactivity. Francium (Fr) would theoretically be even more reactive, but it is almost never encountered outside of physics labs because of its extreme radioactivity and scarcity.

In terms of complexes, Cs⁺, as a large nonpolarizing cation, does not form covalent complexes easily. However, it fits snugly in certain cavity ligands. For example, 18-crown-6 (an 18-atom cyclic polyether) binds Cs⁺ quite well because its size fits the Cs ionic radius, forming stable crown ether complexes. Cryptands and calixarenes have been designed to host Cs⁺ selectively. These complexation tendencies are similar to other alkali metals, scaled by ionic radius.

Finally, in a broad reactivity context, caesium metal is often kept in comparison with potassium or rubidium: a drop of caesium will ignite where a drop of rubidium might only glow, and a drop of potassium might react less violently. Memorably, in demonstration videos Cs vapor can turn nitrogen gas to nitride easily (N₂ becomes black Ceſium nitride). All of this fits with simple redox chemistry: Cs → Cs⁺ + e⁻ is very favorable.

Occurrence and Production

Caesium is a relatively rare element in nature. Its cosmic and terrestrial abundances are low: about 0.8 parts per billion (ppb) by weight in the universe and ~8 ppb in the Sun, but around 3–4 ppm in the Earth’s continental crust. It does not concentrate on its own and is not found in abundance like sodium or potassium. Most naturally occurring Cs is found locked in a few mineral types.

The primary ore of caesium is pollucite, a hydrated aluminosilicate with formula roughly 2Cs₂O·Al₂O₃·4SiO₂ (with water). Pollucite can contain up to ~19–34% Cs by weight. It is usually found in zoned pegmatite geological veins rich in lithium minerals (lepidolite, petalite) and in albite. The world’s largest known deposit of pollucite is the Tanco Mine at Bernic Lake in Manitoba, Canada, which owns most of the global economically extractable reserve (on the order of hundreds of thousands of tons of ore). Other notable pollucite sites include Bikita in Zimbabwe, and a deposit in Namibia. (In these pegmatites, Cs occurs alongside trace rubidium and lithium. Pump- or dump-cyanide processes are unsuitable; instead, the ore is mined and concentrated for caesium directly.)

Occasionally, caesium is also obtained from small amounts in leucite or nepheline (potassium and sodium aluminosilicates), or rare borate minerals like rhodizite. However, these sources contain only minuscule Cs content (often <<1%), so they are not primary producers. Sea water and brines contain only ~0.5 ppb Cs, so extraction from seawater is not practical.

Globally, only a few tonnes of caesium are mined or produced per year – perhaps 5–10 metric tons annually. In terms of supply, China, Canada, and Zimbabwe are major producers (Canada’s Tanco mine is paramount; Zimbabwe’s Bikita pegmatite yields some Cs as a byproduct of petalite mining). Historically small production also came from Germany and, at times, from an alumine-brine (KCl/CsCl) refinery in Europe or Japan. Because Cs is used in relatively small amounts and is expensive, its mining is done on a specialized scale.

Extraction and Refining: To extract caesium from pollucite ore, the mineral is usually crushed and subjected to either acid or alkaline treatments:

  • Acid digestion: The crushed ore is attacked with strong acids (e.g. hydrochloric, sulfuric, or hydrofluoric acid). This dissolves most components, leaving some insoluble residues. Typically, a hydrochloric digestion yields a solution of CsCl along with other metal chlorides. Frequently a double salt such as CsSbCl₆ or Cs₂Na(HfCl₆) precipitates and can be purified. Alternatively, sulfuric acid leads to formation of “caesium alum” (Cs₂SO₄·Al₂(SO₄)₃·24H₂O), which then can be roasted and leached to produce Cs₂SO₄ and ultimately CsCl or Cs₂CO₃. After such steps, CsCl or Cs₂CO₃ is purified by evaporation/crystallization.
  • Alkaline processing: One method is to roaster the ore with calcium carbonate (CaCO₃) and calcium chloride (CaCl₂). This step forms Ca-silicates (insoluble) and leaves CsCl in solution when water is added. The CsCl solution is evaporated to crystals. (Similar to how potassium is extracted from ore, this lump CaCO₃/CaCl₂ roast is essentially a metathesis reaction.) Then CsCl can be boiled with sodium or potassium hydroxide to transform it to CsOH, if the hydroxide is the desired product.

After obtaining CsCl (cesium chloride), the metallic Cs is rarely produced at large scale because the metal must be handled carefully. Industrially, Cs metal can be obtained by reducing molten CsCl with sodium metal (less common than the analogous K reduction). Alternatively, electrolysis of CsCl fused with CsOH has been used on a small scale. Some sources also report direct vacuum reduction of ore with potassium or sodium metal to get Cs metal, but this is not common commercially.

Isotopic Sources: Caesium-137 is produced in nuclear reactors and weapons testing. It can be separated from spent fuel by solvent extraction or ion-exchange processes, but typically it remains mixed with other fission products. In practice, Cs-137 sources for medical and industrial use come from reprocessing of fuel or from dried-down solutions made from waste. Bulk Cs-133 (stable) is rarely “produced” industrially — it is just what is mined. In other words, aside from small synthetic routes, we obtain Cs from its minerals. The amounts used by laboratories (for atomic clocks or research) are tiny compared to those needed for drilling fluids or fission byproducts.

Because pollucite ore reserves are limited, some efforts look at extracting Cs from geologic brines or nuclear waste, but traditional mining remains primary. Given the low production rate and high demand in high-tech applications, caesium is a relatively expensive element. However, Cisco OPEC-level discussions note that current reserves would last for thousands of years at present usage rates.

Applications and Technology

Caesium’s striking properties have led to several important applications, especially in specialized fields. The breadth of uses is wide, reflecting its unique chemistry:

  • Atomic Clocks: The premier use of caesium is in atomic timekeeping. Since 1967, the official definition of the second has been based on the resonance frequency of the Cs-133 hyperfine transition (the exact 9 192 631 770 Hz frequency of the ground-state spin flipping). Caesium beam clocks and fountain clocks achieve accuracies of ~10⁻¹⁶ or better. All GPS satellites carry caesium or rubidium clocks derived from Cs standards. Essentially, every precise time standard on Earth ultimately relies on caesium’s atomic properties. This application stems from Cs’s convenient microwave transition in the ground state and its insensitivity to external fields when shielded.
  • Photoelectric Devices: Owing to its very low work function, caesium metal or cesium compounds (like Cs–Sb in photomultiplier tubes, known as a “Cs–Sb photocathode”) are ideal for converting photons to electrons. Early photoelectric cells and modern photomultiplier detectors often use caesium or caesium-antimony surfaces. Instruments in nuclear physics and astronomy still use Cs-based cathodes for detecting UV and visible light. Caesium-enriched glass (e.g. in vacuum tubes) was used in older television and radar tubes to improve electron emission.
  • Drilling and Oilfield Fluids: Since the late 1990s, the largest bulk industrial use of caesium has been in caesium formate brines (approximately Cs(HCOO)·H₂O in solution). These high-density fluids (density >2.0 g/cm³) are used as weighting agents in high-pressure oil well drilling and workover operations. A solution of caesium formate remains liquid at low temperatures and is less toxic than some other heavy brines, improving wellbore stability. In effect, Cs-formate fluid replaces heavier (like zinc bromide) brines for controlling downhole pressure. The production of Cs-formate fluid consumes many tons of Cs annually.
  • Scintillators and Detectors: Cesium iodide (CsI) crystals are widely used as scintillators: they emit visible light when struck by ionizing radiation (gamma rays or X-rays). CsI is used in medical imaging (e.g. CT scanners), security scanners, and high-energy physics detectors because of its high density and good light yield. Caesium fluoride is used in some precision optical components and as a source of fluoride ions in inorganic syntheses.
  • Radiation Sources: Radioactive Cs-137 sources are used in industrial gauges (moisture-density gauges, level gauges), in medical radiotherapy (cancer treatment machines, chiefly older models), and in calibration of radiation detectors. As a gamma source (0.662 MeV line), Cs-137 can penetrate materials and is convenient to handle (as long as the encapsulation is intact). It has largely been replaced by Cobalt-60 or X-ray tubes in some applications, but still finds use in fixed monitors and irradiators.
  • Analytical Techniques: Caesium chloride (CsCl) is used in density-gradient centrifugation (e.g. to purify DNA, viruses, or subcellular organelles). By spinning a tube of CsCl solution, biomolecules can be separated by buoyant density. This classic technique (dating back to the Meselson–Stahl experiment) relies on the high density of CsCl. Similarly, CsNO₃ and other salts can be used to adjust solution densities in separation processes.
  • Catalysis and Chemistry: Some caesium salts are used as catalysts or reagents in organic synthesis. For example, Cs₂CO₃ is a strong base in certain coupling reactions. CsF is used as a fluorinating agent (it is a source of F⁻ for SN2 reactions in synthetic chemistry). Occasionally, Cs compounds are used in molecular beam epitaxy to dope materials.
  • Space and Research: Caesium has been used in ion engines and space thrusters. Early electric propulsion tests (e.g. Kaufman-type ion thrusters) sometimes used Cs as the propellant because of its low ionization energy (making it easier to ionize). Today, xenon is more common, but some experimental thrusters still consider Cs or Rb. In fundamental research, ultracold cesium atoms have been used in Bose–Einstein condensate experiments and atom interferometry (since Cs atoms are heavy and have convenient optical transitions).
  • Other Uses: CsBr crystals can be used for X-ray diffraction in protein crystallography (as cryoprotectants/dyes). In some special industrial seals or vacuum tubes, Cs getters are used (tiny amounts of Cs or Rb to scavenge residual gases). And CsCl has been used as a cesium source for doping ionization chambers in neutrino detectors.

A notable historical use was as a “getter” in vacuum tubes. A small amount of caesium (or barium) film was evaporated onto hot surfaces to remove residual gases, improving vacuum. This application is largely obsolete now (solid state devices replaced tubes).

In summary, caesium’s main technological roles exploit either its chemical reactivity (as in drilling fluids or gettering) or its atomic physics (as in clocks and detectors). Many other uses are niche or specialized, because the metal’s danger and cost limit widespread handling.

Biology, Environment, and Safety

Biology: Cesium has no known essential biological role. It behaves chemically much like potassium (K⁺), so organisms do not distinguish it and it can substitute (poorly) for K⁺ in some cell processes. However, it is not metabolically useful, and at high levels it can disrupt cellular functions by perturbing potassium channels. Non-radioactive cesium salts (e.g. CsCl) have relatively low acute toxicity compared to heavy metals, but large doses can have physiological effects: for example, CsCl was once (irresponsibly) sold as an alternative medicine, but ingestion of several grams can cause nausea, arrhythmias, or even heart failure by interfering with the heart’s potassium-dependent electrochemistry. The threshold for poisons is high – rat LD₅₀ for CsCl is around 2 g/kg, so it’s not extremely toxic chemically – but it is certainly toxic at high concentrations. Caesium can accumulate in tissues – experiments show accumulation particularly in muscle, but with a biological half-life on the order of 30–70 days (meaning it is gradually excreted via kidneys). Unlike strontium or plutonium, Cs does not strongly bind bone or organs.

Environmentally, stable cesium compounds are moderately mobile. Caesium-133 in soil does not pose a hazard beyond standard chemical concerns. However, radioactive Cs-137 and Cs-134 represent the main environmental worry. These isotopes are water-soluble and were dispersed worldwide by atmospheric nuclear testing (mid-20th century) and accidents like Chernobyl (1986) and Fukushima (2011). Once deposited, ^137Cs can enter food chains (e.g. mushrooms, wild boars) because plants absorb Cs⁺ from soil as if it were K⁺. Contamination of agricultural land with Cs-137 requires precautions (monitoring soils, limits on food). Thankfully, ^137Cs decays, reducing its radiological hazard over decades. Measures like removing topsoil or adding potassium fertilizers can mitigate uptake by plants.

Health and Safety: For the non-radioactive element and its salts, the biggest hazards are chemical. Metallic caesium is extremely dangerous: it ignites explosively in air and water, and burns skin on contact. Any handling of Cs metal requires glovebox or mineral oil immersion, strict exclusion of moisture and air, and proper fire-suppression materials (dry sand class D extinguishers). The metal reacts with glass and most ceramics, so glass equipment is poor choice. Caesium hydroxide is also very dangerous: it is strongly caustic (pH ~ 13–14) and will cause severe chemical burns. Skin contact or inhalation of CsOH dust must be treated as with other strong bases (wash thoroughly, neutralize if needed).

Exposure limits: Non-radioactive Cs compounds are regulated like other strong alkalis, so workplace exposure limits relate to causticity rather than element-specific concerns. However, any cesium-137 or cesium-134 source has radiation safety rules – for example, 1 µSv/hour dose near a container or 1 mSv/year occupational limit, etc., established by nuclear safety authorities. A cesium-137 ingestion is treated by chelation therapy (Prussian blue pills accelerate elimination to perhaps 30 days effective half-life).

Environmental fate: Elemental Cs does not occur freely in nature, so there is no concern of spills except from lab accidents. Soluble Cs salts (like CsCl from a broken source) could, if released, contaminate groundwater as radiological hazard if radioactive, or cause local soil alkalinization if massive (unlikely). Regulatory limits in water for radioactive Cs are stringent (e.g. for Cs-137, on the order of Bq/L being digits for safe drinking water). Cleanup of Cs-137 contamination often involves removal of contaminated soil or using plants that can uptake Cs (phytoremediation).

In summary, non-radioactive cesium (metal and compounds) is moderately hazardous chemically (caustic, reactive) but not especially bioaccumulative. Its radioactive isotopes, mainly Cs-137, are highly regulated due to long half-life gamma emission. The greatest risks with cesium are handled through standard precautions (avoid skin contact, control sources). Unlike heavy metals (e.g. lead, mercury), Cs doesn’t bioaccumulate in bones, but it can be lumped with group 2 mimic like radium: the key worry is whole-body gamma exposure for the radioactive forms.

History and Etymology

Caesium’s story begins in 1860, when German chemists Robert Bunsen and Gustav Kirchhoff discovered it using flame spectroscopy. They vaporized mineral samples and observed two bright spectral lines of a distinctive sky-blue color (in Latin, caesius means “blue-grey”). This gave the element its name: caesium (often spelled “cesium” in American English; the IUPAC recommended spelling is caesium). The name reflects the blue lines in its emission spectrum. Bunsen and Kirchhoff isolated the new element from mineral water samples by reduction of caesium salts using potassium.

In the early 20th century, experiments with caesium further illuminated atomic physics. For example, the photoelectric effect was first observed by Heinrich Hertz in 1887 on metallic surfaces including platinum, but caesium (and later rubidium) coatings became important in phototubes around 1900; these materials exhibited strong photoemission, aiding studies by Einstein and others. Antoine Henri Becquerel and the Curies did not use caesium, but later caesium-137 was identified as one of the “hot” fission products in nuclear research after WWII.

The half-century after its discovery saw only modest use for caesium beyond vacuum technology. Small amounts of caesium metal were used as getters in evacuated tubes (in light bulbs, vacuum tubes, etc.) to maintain vacuum. Caesium-antimony cathodes were used in some early photodetectors (photoelectric cells for television cameras in the mid-20th century often used Cs₂Te or Cs–Sb photocathodes).

The first caesium atomic clock was built in the early 1950s (Kenneth (Weldon) Powles at the National Physical Laboratory in the UK or the NIST in the US), and by 1967 the international community adopted a Cs-based frequency as the definition of the second. This was an enormous milestone: Cs went from a obscure metal to the heartbeat of global timekeeping. Since then, caesium fountain clocks (using laser cooling and microwave interrogation) have progressively improved timing standards.

Nuclear technology also brought caesium to prominence. Cs-137 contamination from nuclear weapons tests and accidents became a public concern by the 1960s (e.g. Strontium-90 and Cs-137 in fallout). The Chernobyl accident in 1986 released an estimated 85 petabecquerels of ^137Cs into the environment, making “cesium-137” a household name in affected countries (Ukraine, Belarus, parts of Scandinavia). After these events, environmental monitoring of Cs isotopes became standard practice.

The modern era saw a pivot: in the 1990s, petrochemical industries began exploiting caesium’s high density solubility. By the 2000s, caesium formate fluid emerged as a breakthrough drilling fluid (first patents around 1997). This gave caesium a sudden value beyond laboratory curiosities. At the same time, research into cold-atom physics and quantum information started using cesium’s convenient level structure for experiments in the 1990s and 2000s.

Therefore, the milestones of caesium history include its spectroscopic discovery in 1860, its early adoption in electronics, its fundamental role in atomic clocks (mid-20th century onward), and its newer roles in oil-drilling technology and advanced physics. The name caesium, meaning “sky blue”, endures as a nod to the method of discovery. Cast of characters: Bunsen and Kirchhoff (discoverers), later scientists like Norman Ramsey and teams at NIST (clock developers), and engineers who harnessed Cs in industry.

Data Table

Property Value
Symbol Cs
Element category Alkali metal (Group 1, s-block)
Atomic number (Z) 55
Standard atomic weight 132.905 (6)
Electron configuration [Xe] 6s¹
Period, group Period 6, Group 1
Phase at STP Solid (silvery-golden metal)
Density (20 °C) 1.93 g·cm⁻³
Melting point 28.44 °C
Boiling point 671 °C
Crystal structure (RT) Body-centered cubic
Oxidation state(s) +1 (chief)
Electronegativity (Pauling) 0.79
First ionization energy 375.7 kJ/mol (≈3.89 eV)
Atomic radius ~265 pm (van der Waals); 244 pm (metallic)
Stable isotope ¹³³Cs (100% natural abundance)
Nuclear spin (¹³³Cs) 7/2
Thermal conductivity 36 W·m⁻¹·K⁻¹ (RT)
Heat of fusion 2.09 kJ/mol
Heat of vaporization 63.9 kJ/mol (at 300 K)

Notes: Values are given for the most common conditions; atomic weight is of naturally occurring Cs (essentially all ¹³³Cs). Radii are empirical; the “largest atomic radius” claim refers to theoretical and van der Waals radii. Mnemonic: Cs is the soft, golden metal that defined the sec.

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