Description

cobalt (Co)chemical element, ferromagnetic metal of Group 9 (VIIIb) of the periodic table, used especially for heat-resistant and magnetic alloys.

cobalt
cobalt

The metal was isolated (c. 1735) by Swedish chemist Georg Brandt, though cobalt compounds had been used for centuries to impart a blue colour to glazes and ceramics. Cobalt has been detected in Egyptian statuettes and Persian necklace beads of the 3rd millennium BCE, in glass found in the Pompeii ruins, and in China as early as the Tang dynasty (618–907 CE) and later in the blue porcelain of the Ming dynasty (1368–1644). The name kobold was first applied (16th century) to ores thought to contain copper but eventually found to be poisonous arsenic-bearing cobalt ores. Brandt finally determined (1742) that the blue colour of those ores was due to the presence of cobalt.BRITANNICA QUIZ118 Names and Symbols of the Periodic Table QuizAn elementary quiz about fundamental matters.

atomic number27
atomic weight58.933194
melting point1,495 °C (2,723 °F)
boiling point2,870 °C (5,198 °F)
density8.9 gram/cm3 at 20 °C (68 °F)
oxidation states+2, +3
electron configuration[Ar]3d74s2

Occurrence, properties, and uses

erythrite; skutterudite
erythrite; skutterudite

Cobalt, though widely dispersed, makes up only 0.001 percent of Earth’s crust. It is found in small quantities in terrestrial and meteoritic native nickel-iron, in the Sun and stellar atmospheres, and in combination with other elements in natural waters, in ferromanganese crusts deep in the oceans, in soils, in plants and animals, and in minerals such as cobaltitelinnaeiteskutteruditesmaltite, heterogenite, and erythrite. In animals, cobalt is a trace element essential in the nutrition of ruminants (cattle, sheep) and in the maturation of human red blood cells in the form of vitamin B12, the only vitamin known to contain such a heavy element.

With few exceptions, cobalt ore is not usually mined for the cobalt content. Rather, it is often recovered as a by-product from the mining of ores of ironnickelcoppersilvermanganesezinc, and arsenic, which contain traces of cobalt. Complex processing is required to concentrate and extract cobalt from these ores. By the second decade of the 21st century, the Democratic Republic of the Congo (DRC), China, Canada, and Russia were the world’s leading producers of mined cobalt. The largest producer of refined cobalt, however, was China, which imported vast additional amounts of cobalt mineral resources from the DRC. (For additional information on the mining, refining, and recovery of cobalt, see cobalt processing.)

Polished cobalt is silver-white with a faint bluish tinge. Two allotropes are known: the hexagonal close-packed structure, stable below 417 °C (783 °F), and the face-centred cubic, stable at high temperatures. It is ferromagnetic up to 1,121 °C (2,050 °F, the highest known Curie point of any metal or alloy) and may find application where magnetic properties are needed at elevated temperatures.

Cobalt is one of the three metals that are ferromagnetic at room temperature. It dissolves slowly in dilute mineral acids, does not combine directly with either hydrogen or nitrogen, but will combine, on heating, with carbonphosphorus, or sulfur. Cobalt is also attacked by oxygen and by water vapour at elevated temperatures, with the result that cobaltous oxide, CoO (with the metal in the +2 state), is produced.

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Natural cobalt is all stable isotope cobalt-59, from which the longest-lived artificial radioactive isotope cobalt-60 (5.3-year half-life) is produced by neutron irradiation in a nuclear reactor. Gamma radiation from cobalt-60 has been used in place of X-rays or alpha rays from radium in the inspection of industrial materials to reveal internal structure, flaws, or foreign objects. It has also been used in cancer therapy, in sterilization studies, and in biology and industry as a radioactive tracer.

Most of the cobalt produced is used for special alloys. A relatively large percentage of the world’s production goes into magnetic alloys such as the Alnicos for permanent magnets. Sizable quantities are utilized for alloys that retain their properties at high temperatures and superalloys that are used near their melting points (where steels would become too soft). Cobalt is also employed for hard-facing alloys, tool steels, low-expansion alloys (for glass-to-metal seals), and constant-modulus (elastic) alloys (for precision hairsprings). Cobalt is the most satisfactory matrix for cemented carbides.

Finely divided cobalt ignites spontaneously. Larger pieces are relatively inert in air, but above 300 °C (570 °F) extensive oxidation occurs.

Compounds

In its compounds cobalt nearly always exhibits a +2 or +3 oxidation state, although states of +4, +1, 0, and −1 are known. The compounds in which cobalt exhibits the +2 oxidation state (Co2+, the ion being stable in water) are called cobaltous, while those in which cobalt exhibits the +3 oxidation state (Co3+) are called cobaltic.

Both Co2+ and Co3+ form numerous coordination compounds, or complexes. Co3+ forms more known complex ions than any other metal except platinum. The coordination number of the complexes is generally six.

Cobalt forms two well-defined binary compounds with oxygen: cobaltous oxide, CoO, and tricobalt textroxide, or cobalto-cobaltic oxide, Co3O4. The latter contains cobalt in both +2 and +3 oxidation states and constitutes up to 40 percent of the commercial cobalt oxide used in the manufacture of ceramics, glass, and enamel and in the preparation of catalysts and cobalt metal powder.

One of the more important salts of cobalt is the sulfate CoSO4, which is employed in electroplating, in preparing drying agents, and for pasture top-dressing in agriculture. Other cobaltous salts have significant applications in the production of catalysts, driers, cobalt metal powders, and other salts. Cobaltous chloride (CoCl2∙6H2O in commercial form), a pink solid that changes to blue as it dehydrates, is utilized in catalyst preparation and as an indicator of humidity. Cobaltous phosphate, Co3(PO4)2∙8H2O, is used in painting porcelain and colouring glass.

This article was most recently revised and updated by Erik Gregersen.

xenon

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Alternate titles: Xe

By The Editors of Encyclopaedia Britannica Last Updated: Sep 28, 2022 Edit History

xenon

See all mediaKey People: Sir William RamsayRelated Topics: chemical elementnoble gasairxenon-129

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xenon (Xe)chemical element, a heavy and extremely rare gas of Group 18 (noble gases) of the periodic table. It was the first noble gas found to form true chemical compounds. More than 4.5 times heavier than air, xenon is colourless, odourless, and tasteless. Solid xenon belongs to the face-centred cubic crystal system, which implies that its molecules, which consist of single atoms, behave as spheres packed together as closely as possible. The name xenon is derived from the Greek word xenos, “strange” or “foreign.”

atomic number54
atomic weight131.29
melting point−111.9 °C (−169.4 °F)
boiling point−108.0 °C (−162.4 °F)
density (1 atm, 0 °C [32 °F])5.887 g/litre (0.078 ounce/gallon)
oxidation states0, +2, +4, +6, +8
electron config.(Kr)4d105s25p6

Properties of the element

Xenon occurs in slight traces in gases within Earth and is present to an extent of about 0.0000086 percent, or about 1 part in 10 million by volume of dry air. Like several other noble gases, xenon is present in meteorites. Xenon is manufactured on a small scale by the fractional distillation of liquid air. It is the least volatile (boiling point, −108.0 °C [−162.4 °F]) of the noble gases obtainable from the air. The British chemists Sir William Ramsay and Morris W. Travers isolated the element in 1898 by repeated fractional distillation of the noble gas krypton, which they had discovered six weeks previously.BRITANNICA QUIZ118 Names and Symbols of the Periodic Table QuizAn elementary quiz about fundamental matters.

The element xenon is used in lamps that produce extremely short and intense flashes of light, such as stroboscopes and lights for high-speed photography. When a charge of electricity is passed through the gas at low pressure, it emits a flash of bluish-white light; at higher pressures, white light resembling daylight is emitted. Xenon flashlamps are used to activate ruby lasers.

Natural xenon is a mixture of nine stable isotopes in the following percentages: xenon-124 (0.096), xenon-126 (0.090), xenon-128 (1.92), xenon-129 (26.44), xenon-130 (4.08), xenon-131 (21.18), xenon-132 (26.89), xenon-134 (10.44), and xenon-136 (8.87). The mass numbers of the known isotopes of xenon range from 118 to 144. The xenon found in some stony meteorites shows a large proportion of xenon-129, believed to be a product of radioactive decay of iodine-129, whose half-life is 17,000,000 years. Measuring the xenon-129 content of meteorites casts light on the history of the solar system. More than a dozen radioactive xenon isotopes produced by fission of uranium and other nuclear reactions are known. For example, xenon-135 (9.2-hour half-life) is produced by uranium fission in nuclear reactors, where it is troublesome because it absorbs fission-producing neutrons. Xenon-129 is of particular importance because this isotope can be observed by nuclear magnetic resonance spectroscopy, which makes it useful for the structural characterization of xenon compounds. The xenon isotopes produced in the greatest amount by nuclear fission are xenon-131, -132, -134, and -136, which are stable, and xenon-133, which is radioactive, with a half-life of 5.27 days.

Compounds

Noble gases were thought to be chemically inert until 1962, when British chemist Neil Bartlett produced the first noble-gas compound, a yellow-orange solid that can best be formulated as a mixture of [XeF+][PtF6], [XeF+][Pt2F11], and PtF5. Xenon has the most extensive chemistry in Group 18 and exhibits the oxidation states +1/2, +2, +4, +6, and +8 in the compounds it forms. Since the discovery of noble-gas reactivity, xenon compounds, including halides, oxides, oxofluorides, oxo salts, and numerous covalent derivatives with a number of compounds covalently bonded to other polyatomic ligands, have been synthesized and structurally characterized. As might be predicted from the position of xenon in the periodic table, xenon compounds are poorer oxidizing agents than krypton compounds. Hence, much of currently known xenon chemistry involves its fluorides and oxofluorides in their reactions with strong Lewis acid acceptors and fluoride-ion donors to form a variety of fluoro- and oxofluorocations and anions, respectively. Examples of xenon covalently bonded to fluorineoxygennitrogen, and carbon are now known.

Three fluorides of xenon are known, XeF2 (the easiest to prepare), XeF4, and XeF6. They are stable, colourless, crystalline solids that can be sublimed under vacuum at 25 °C (77 °F). Like KrF2, XeF2 is a linear symmetric molecule. Xenon tetrafluoride (XeF4) is a square planar molecule, and XeF6 in the gas phase is a distorted octahedral molecule arising from the presence of an “extra” pair of nonbonding electrons in the xenon valence shell. Higher halides such as XeCl2, XeClF, XeBr2, and XeCl4 are thermodynamically unstable and have been detected only in small amounts. The unstable and short-lived monohalides XeF, XeCl, XeBr, and XeI have been produced in the gas phase and are of considerable importance as light-emitting species in gas lasers.

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Two oxides of xenon are known: xenon trioxide (XeO3) and xenon tetroxide (XeO4), and both are unstable, highly explosive solids that must be handled with the greatest care. The oxide fluorides XeO3F2, XeO2F4, XeOF4, XeO2F2, and XeOF2 are known and, with the exception of XeOF4, all are thermodynamically unstable.

Xenon difluoride behaves as a simple fluoride-ion donor toward many metal pentafluorides to form complex salts containing the XeF+ and Xe2F3+[F(XeF)2]+ cations by analogy with KrF2 (see krypton: compounds). Mixtures of xenon and fluorine gases react spontaneously with liquid antimony pentafluoride in the dark to form solutions of XeF+Sb2 F11, in which Xe2+ is formed as an intermediate product that is subsequently oxidized by fluorine to the XeF+ cation. The bright emerald green, paramagnetic dixenon cation, Xe2+, is the only example of xenon in a fractional oxidation state, +1/2.

Xenon tetrafluoride is a much weaker fluoride-ion donor than XeF2 and only forms stable complex salts with the strongest fluoride-ion acceptors to give compounds such as [XeF3+][SbF6] and [XeF3+][Sb2F11]. Xenon tetrafluoride has also been shown to behave as a weak fluoride-ion acceptor toward the fluoride ion to give salts of the pentagonal planar XeF5 anion. Xenon oxide difluoride is also a fluoride-ion acceptor, forming the only other anion containing xenon in the +4 oxidation state, the XeOF3 anion in Cs+XeOF3.

Xenon hexafluoride is both a strong fluoride-ion donor and a strong fluoride-ion acceptor. Examples of salts containing the XeF5+ cation are numerous, with counter anions such as PtF6 and AuF6. Examples of salts containing the fluoride bridged Xe2F11+ cation are also known. Xenon hexafluoride behaves as a fluoride-ion acceptor, reacting with alkali metal fluorides to form salts that contain the XeF7 and XeF82− anions. Several nonalkali metal salts have been shown to contain the anions XeF7 and XeF82− and include [NF4+][XeF7] and [NO+]2[XeF82−].

The oxofluorides of xenon +6, XeOF4 and XeO2F2, exhibit analogous fluoride-ion donor and acceptor properties. Salts of both the XeOF3+ and XeO2F+ cations, as well as a salt of the fluoride-bridged cation Xe2O4F3+, are known. These include [XeOF3+][SbF6] and [Xe2O4F3+][AsF6]. Several alkali metal fluoride complexes with XeOF4 are known, such as 3KF∙XeOF4 and CsF∙3XeOF4. Structural studies show that the CsF and N(CH3)4F complexes are best formulated as [Cs+][XeOF5], [N(CH3)4+][XeOF5], and [Cs+][(XeOF4)3F]. In these compounds, XeOF4 behaves as a fluoride acceptor. The only complexes between XeO2F2 and a strong fluoride-ion donor are the salts [Cs+][XeO2F3] and [NO2+][XeO2F3∙XeO2F2].

When XeF6 is hydrolyzed in a strongly alkaline solution, part of the xenon is lost as gas (reduced to the 0 oxidation state), but a large fraction precipitates as a perxenate (XeO64−) salt in which xenon is in the +8 oxidation state. The salts are kinetically very stable and lose water gradually when heated; for example, Na4XeO6∙6H2O becomes anhydrous at 100 °C (212 °F) and decomposes at 360 °C (680 °F).

Alkali metal xenates of composition MHXeO4∙1.5H2O, where M is sodiumpotassiumrubidium, or cesium and xenon is in the +6 oxidation state, have been prepared. The xenates are unstable explosive solids. Alkali metal fluoroxenates [K+][XeO3F], [Rb+][XeO3F], [Cs+][XeO3F] (which decomposes above 200 °C [392 °F]), and the chloroxenate [Cs+][XeO3Cl] (which decomposes above 150 °C [302 °F]) have been prepared by evaporating aqueous solutions of XeO3 and the corresponding alkali metal fluorides and chlorides. The alkali metal fluoroxenates are the most stable solid oxygen compounds of xenon(+6) known. However, CsXeO3Br is unstable even at room temperature.

A number of polyatomic ligands of high effective group electronegativities form compounds with xenon. The greatest variety of polyatomic ligand groups bonded to xenon occurs for xenon in its +2 oxidation state, and those groups bonded through oxygen are most plentiful. Both mono- and disubstituted derivatives having the formulations FXeL and XeL2 are known where L = OTeF5 and OSeF5, for example.

The highly electronegative OTeF5  group closely mimics the ability of F to stabilize the oxidation states of xenon, with stable OTeF5 derivatives also existing for the +4 and +6 oxidation states of xenon. Cations that contain the (OTeF5)+ group also are known.

Several ligand groups form compounds containing xenon-nitrogen bonds. Among the first xenon-nitrogen bonded compounds to be prepared were FXe[N(SO2F)2] and Xe[N(SO2F)2]2. Like XeF2 and KrF2, FXe[N(SO2F)2] is a fluoride-ion donor toward AsF5, forming [XeN(SO2F)2+][AsF6]. Like KrF+, the XeF+ cation behaves as a electron pair acceptor toward nitrogen Lewis bases, but because XeF+ is not as powerful an oxidant as KrF+, the range of ligands that can be coordinated to XeF+ is more extensive. These include HCN and (CH3)3CCN, which interact with XeF+ to form the HCNXeF+ and (CH3)3CCNXeF+ cations, respectively.

A number of compounds containing Xe-C bonds are known. These compounds are salts of cations containing xenon(+2) coordinated to carbon and include cations such as (C6F5)Xe+ and (m-CF3C6H4)Xe+. An example of xenon(+4) bonded to carbon is also known. The (C6F5)XeF2+ cation has been prepared as the BF4 salt.

This article was most recently revised and updated by Adam Augustyn.Load Next Page 

Cobalt is a chemical element with the symbol Co and atomic number 27. As with nickel, cobalt is found in the Earth’s crust only in a chemically combined form, save for small deposits found in alloys of natural meteoric iron. The free element, produced by reductive smelting, is a hard, lustrous, silver-gray metal.

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