7440-46-2

silvery, light ductile metal

In photoelectric cells, as a "getter" in vacuum tubes; in photoemitter devices, scintillation counters. Adsorbent in CO2 purifn; scavenger of gases and impurities in metallurgy. For doping catalysts. For construction and operation of one type of atomic clock based on the vibrational frequency (9,192.76 megacycles/sec) of 133Cs. 137Cs in process control instruments, sewage and sludge sterilization.

A soft metallic solid. Melts at 85°F. Causes burns to skin and eyes.

CESIUM METAL reacts violently with oxidizing agents, even weaker ones. Reacts with boron trifluoride with incandescence when heated [Merck 11th ed. 1989]. Reacts explosively with maleic anhydride [Chem Safety Data Sheet SD-88 1962; Chem. Haz. Info. Series C-71 1960]. Burns in chlorine with a luminous flame [Mellor 2 Supp. 1:380 1956]. Reacts violently with most acids. Reacts violently with fluorine, chlorine, bromine and iodine. Reacts with incandescence with sulfur and phosphorus. Burns vigorously in air.

Highly flammable. CESIUM(7440-46-2) is spontaneously flammable in air at room temperature, if the surface is clean [Merck 11th ed. 1989]. Reacts with water to generate enough heat to ignite the hydrogen produced during the reaction, and to generate caustic CESIUM(7440-46-2) hydroxide [Mellor 2 419 1946-47].

Dangerous fire and explosion risk, ignites spontaneously in moist air, may explode in contact with sulfur or phosphorus, reacts violently with oxi- dizing materials, causes burns in contact with skin.

Inhalation or contact with vapors, substance or decomposition products may cause severe injury or death. May produce corrosive solutions on contact with water. Fire will produce irritating, corrosive and/or toxic gases. Runoff from fire control may cause pollution.

Produce flammable gases on contact with water. May ignite on contact with water or moist air. Some react vigorously or explosively on contact with water. May be ignited by heat, sparks or flames. May re-ignite after fire is extinguished. Some are transported in highly flammable liquids. Runoff may create fire or explosion hazard.

Cesium was discovered in 1860 by Robert Bunsen and Gustav Kirchoff. It is used in the most accurate atomic clocks. Cesium melts at 28.41°C (just below body temperature) and occurs in Earth’s crust at 2.6 ppm. Cesium is the rarest of the naturally occurring alkali metals as the isotope ¹³³Cs. Its compounds are correspondingly rare. Granites contain about 1 ppm cesium and sedimentary rocks contain approximately 4 ppm cesium. The most common commercial source of cesium is pollucite, which contains between 5 and 32% cesium oxide. Radioactive forms of cesium (¹³⁴Cs and ¹³⁷Cs) can also be found in the environment. They are produced during nuclear fission, and are used in cancer treatment.

Like the other alkali metals, cesium is a soft-solid silvery metal, but much softer than theothers. It is the least electronegative and most reactive of the Earth metals. Cesium has anoxidation state of +1, and because its atoms are larger than Li, Na, and K atoms, it readilygives up its single outer valence electron. The single electron in the P shell is weakly attachedto its nucleus and thus available to combine with many other elements. It is much too reactiveto be found in its metallic state on Earth.Cs has a melting point of 29°C, which is lower than the body temperature of humans(37°C), and thus a chunk of cesium will melt in a person’s hand with disastrous results. Sinceit reacts with moisture on skin as well as with the air to release hydrogen, it will burn vigorously through the palm of one’s hand.Cesium’s boiling point is 669.3°C and its density is 1.837 g/cm³. Mercury is the only metalwith a lower melting point than cesium. It is extremely dangerous when exposed to air, water,and organic compounds or to sulfur, phosphorus, and any other electronegative elements. Itmust be stored in a glass container containing an inert atmosphere or in kerosene.Cesium reacts with water in ways similar to potassium and rubidium metals. In additionto hydrogen, it forms what is known as superoxides, which are identified with the generalformula CsO₂. When these superoxides react with carbon dioxide, they release oxygen gas,which makes this reaction useful for self-contained breathing devices used by firemen andothers exposed to toxic environments.

Cs-133 is the only stable isotope of cesium, and it makes up all of the naturallyoccurring cesium found in the Earth’s crust. In addition to Cs-133 there are about 36radioactive isotopes of Cs, most of which are artificially formed in nuclear reactors. Allare produced in small numbers of atoms with relatively short half-lives. The range of Csisotopes is from Cs-113 (amu = 112.94451) to Cs-148 (amu = 147.94900). Most ofthese radioisotopes produce beta radiation as they rapidly decay, with the exception ofCs-135, which has a half-life of 3×10⁶yr, which makes it a useful research tool. Cs-137,with a half-life of 33 years, produces both beta and gamma radiation.

In 1860 Gustav Kirchhoff and Robert Bunsen named the element “Cesium,” using the Latin word caesius, which means bluish-gray.

The stable form of Cs-133 is the 48th most abundant element on Earth, but because it isso reactive, it is always in compound form. The Earth’s crust contains only about 7 ppm ofCs-133. Like the other alkali metals, it is found in mixtures of complex minerals. Its mainsource is the mineral pollucite (CsAlSi₂O₆). It is also found in lepidolite, a potassium ore.Pollucite is found in Maine, South Dakota, Manitoba, and Elba and primarily in Rhodesia,South Africa.One problem in refining cesium is that it is usually found along with rubidium; therefore,the two elements must be separated after they are extracted from their sources. The mainprocess to produce cesium is to finely grind its ores and then heat the mix to about 600°Calong with liquid sodium, which produces an alloy of Na, Cs, and Ru, which are separatedby fractional distillation. Cesium can also be produced by the thermochemical reduction of amixture of cesium chloride (CsCl) and calcium (Cs).

Cesium is located between rubidium and francium in group 1 of the periodic table. It isthe heaviest of the stable alkali metals and has the lowest melting point. It is also the mostreactive of the alkali metals.Cesium will decompose water, producing hydrogen, which will burn as it is liberated fromH₂O. Cesium is extremely dangerous to handle and will burn spontaneously or explode whenexposed to air, water, and many organic compounds.

Although Cesium metals have been prepared by fused salt electrolysis, the highly reactive nature of the metals complicates the collection step and favors the use of other preparative methods where the metals can be removed in vapor form from the reaction mixture. The oxides, hydroxides, carbonates, halides, sulphates, chromates and nitrates of Cesium have been reduced to the metals by strong reducing metals such as sodium, calcium, magnesium, barium, iron, zirconium, aluminum or silicon at moderately high temperatures. The preferred method, however, involves the reduction of the anhydrous metal chlorides with calcium metal under vacuum. Anhydrous cesium chloride is mixed with a large excess of calcium chips and heated under vacuum at 700-800°C. As the chloride is reduced, metal vapors issue from the reaction mixture and are led under the vacuum to a cooler portion of the vessel where they condense and drop into a collection vessel.

A chemical element, cesium (symbol Cs) is theheaviest of the alkali metals in group I. It is asoft, light, very low melting temperature metal.It is the most reactive of the alkali metals andindeed is the most electropositive and the mostreactive of all the elements.
Cesium oxidizes easily in the air, ignites atordinary temperatures, and decomposes waterwith explosive violence. It can be contained invacuum, inert gas, or anhydrous liquid hydrocarbonsprotected from O2 and air. The specificgravity is 1.903, melting point 28.5°C, and boilingpoint 670°C. It is used in low-voltage tubesto scavenge the last traces of air. It is usuallymarketed in the form of its compounds such ascesium nitrate, CsNO3, cesium fluoride, CsF, orcesium carbonate, Cs2CO3. In the form ofcesium chloride, CsCl, it is used on the filamentsof radio tubes to increase sensitivity. Itinteracts with the thorium of the filament toproduce positive tons. In photoelectric cellsCsCl is used for a photosensitive deposit on thecathode, since cesium releases its outer electronunder the action of ordinary light, and its colorsensitivity is higher than that of other alkali metals. The high-voltage rectifying tube forchanging AC to DC has cesium metal coatedon the nickel cathode, and has cesium vapor forcurrent carrying. The cesium metal gives off acopious flow of electrons and is continuouslyrenewed from the vapor. Cesium vapor is alsoused in the infrared signaling lamp as it producesinfrared waves without visible light. Cesiumsalts have been used medicinally as antishockagents after administration of arsenic drugs.Cesium metal is generaly made by thermochemicalprocesses. The carbonate can bereduced by metallic magnesium, or the chloridecan be reduced by CaC. Metallic cesium volatilizesfrom the reaction mixture and is collectedby cooling the vapor.

Stable cesium was shown to affect various central nervous system functions, mainly involving displacing potassium, with which it competes for transport through the potassium channel, and it can also activate sodium pump and subsequent transport into the cell across membranes. Thus, this resulted in potassium deficiency.
Radioactive isotopes of cesium, such as ¹³⁴Cs and ¹³⁷Cs, are a greater health concern than stable cesium. These radioactive isotopes of cesium are formed during nuclear fission. Both ¹³⁴Cs and ¹³⁷Cs emit beta and gamma radiations. Beta radiation travels short distances and can penetrate the skin and superficial body tissues, whereas gamma radiation can travel great distances and penetrate the entire body. Both beta and gamma radiations may induce tissue damage and disruption of cellular function.

Naturally occurring cesium can enter the environment mostly from the erosion and weathering of rocks and minerals. The production and use of cesium compounds may also result in their release to the environment through various waste streams. However, there are relatively few commercial uses for cesium compounds, such as cesium radioactive isotopes (¹³⁴Cs and ¹³⁷Cs), and they have been released into the environment by human activities such as the atmospheric testing of nuclear weapons (1945–80) and leakages at nuclear power plants. Cesium compounds can travel long distances in the air before being brought back to the earth by rainfall and gravitational settling. If released to water, cesium compounds are deposited on land and water via wet and dry deposition. These deposited-cesium particles may be resuspended into the atmosphere from soil and dust. If released to soil, cesium compounds have low mobility and do not migrate below 40 cm in depth. The majority of cesium ions are retained in the upper 20 cm of the soil surface. Clay and zeolite minerals strongly bind cesium cations irreversibly. Soils rich in organic matter also adsorb cesium ions. However, cesium compounds are readily exchangeable and highly available for plant uptake in these soils. If released into water, cesium compounds are very water soluble and exist primarily as cesium cations. Because most cesium compounds are ionic, they will not volatilize from water surfaces. Most cesium compounds released to water adsorb to suspended solids in the water column and ultimately they are deposited in sediments. Cesium compounds bioconcentrate and have been shown to bioaccumulate in both terrestrial and aquatic food chains. The half-life of ¹³⁴Cs is ～2 years and that of ¹³⁷Cs is～30 years.