Cobalt: Wikis


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hard lustrous gray metal
General properties
Name, symbol, number cobalt, Co, 27
Element category transition metal
Group, period, block 94, d
Standard atomic weight 58.933195(5)g·mol−1
Electron configuration [Ar] 4s2 3d7
Electrons per shell 2, 8, 15, 2 (Image)
Physical properties
Color metallic gray
Density (near r.t.) 8.90 g·cm−3
Liquid density at m.p. 7.75 g·cm−3
Melting point 1768 K, 1495 °C, 2723 °F
Boiling point 3200 K, 2927 °C, 5301 °F
Heat of fusion 16.06 kJ·mol−1
Heat of vaporization 377 kJ·mol−1
Specific heat capacity (25 °C) 24.81 J·mol−1·K−1
Vapor pressure
P/Pa 1 10 100 1 k 10 k 100 k
at T/K 1790 1960 2165 2423 2755 3198
Atomic properties
Oxidation states 5, 4 , 3, 2, 1, -1[1]
(amphoteric oxide)
Electronegativity 1.88 (Pauling scale)
Ionization energies
1st: 760.4 kJ·mol−1
2nd: 1648 kJ·mol−1
3rd: 3232 kJ·mol−1
Atomic radius 125 pm
Covalent radius 126±3 (low spin), 150±7 (high spin) pm
Crystal structure hexagonal
Magnetic ordering ferromagnetic
Electrical resistivity (20 °C) 62.4 nΩ·m
Thermal conductivity (300 K) 100 W·m−1·K−1
Thermal expansion (25 °C) 13.0 µm·m−1·K−1
Speed of sound (thin rod) (20 °C) 4720 m/s
Young's modulus 209 GPa
Shear modulus 75 GPa
Bulk modulus 180 GPa
Poisson ratio 0.31
Mohs hardness 5.0
Vickers hardness 1043 MPa
Brinell hardness 700 MPa
CAS registry number 7440-48-4
Most stable isotopes
Main article: Isotopes of cobalt
iso NA half-life DM DE (MeV) DP
56Co syn 77.27 d ε 4.566 56Fe
57Co syn 271.79 d ε 0.836 57Fe
58Co syn 70.86 d ε 2.307 58Fe
59Co 100% 59Co is stable with 32 neutrons
60Co syn 5.2714 years β,γ,γ 2.824 60Ni

Cobalt (pronounced /ˈkoʊbɒlt/ KOH-bolt)[2] is a hard, lustrous, gray metal, a chemical element with symbol Co and atomic number 27. Cobalt-based colors and pigments have been used since ancient times for jewelry and paints, and miners have long used the name kobold ore for some minerals.

Cobalt occurs in various metallic-lustered ores, for example cobaltite (CoAsS), but is mainly produced as a by-product of copper and nickel mining. The copper belt in the Democratic Republic of the Congo and Zambia yields most of the cobalt mined worldwide.

Cobalt is used in the preparation of magnetic, wear-resistant, and high-strength alloys. Smalte (cobalt silicate glass) and cobalt blue (cobalt(II) aluminate, CoAl2O4) gives a distinctive deep blue color to glass, ceramics, inks, paints, and varnishes. Cobalt-60 is a commercially important radioisotope, used as a tracer and in the production of gamma rays for industrial use.

Cobalt is an essential trace element for all multicellular organisms as the active center of coenzymes called cobalamins. These include vitamin B-12 which is essential for mammals. Cobalt is also an active nutrient for bacteria, algae, and fungi, and may be a necessary nutrient for all life.



Cobalt is a ferromagnetic metal with a specific gravity of 8.9 (20°C). Pure cobalt is not found in nature, but compounds of cobalt are common. Small amounts of it are found in most rocks, soil, plants, and animals. It has the atomic number 27. The Curie temperature is 1115 °C, and the magnetic moment is 1.6–1.7 Bohr magnetons per atom. In nature, it is frequently associated with nickel, and both are characteristic minor components of meteoric iron. Mammals require small amounts of cobalt which is the basis of vitamin B12. Cobalt-60, an artificially produced radioactive isotope of cobalt, is an important radioactive tracer and cancer-treatment agent. Cobalt has a relative permeability two thirds that of iron. Metallic cobalt occurs as two crystallographic structures: hcp and fcc. The ideal transition temperature between hcp and fcc structures is 450 °C, but in practice, the energy difference is so small that random intergrowth of the two is common.[3]


Common oxidation states of cobalt include +2 and +3, although compounds with oxidation state +1 are also known. The most common oxidation state for simple compounds is +2. Cobalt(II) salts form the red-pink [Co(H2O)6]2+ complex in aqueous solution. Adding excess chloride will change the color from pink to blue, due to the formation of [CoCl4]2−.[4]


Oxygen and chalcogen compounds

Several oxides of cobalt are known. Green cobalt(II) oxide (CoO) has rocksalt structure. It is readily oxidized with water and oxygen to brown cobalt(III) hydroxide CoO(OH). At temperatures of 400–500 °C, CoO oxidizes to the blue cobalt(II,III) oxide (Co3O4), which has spinel structure. Simple cobalt(III) oxide (Co2O3) is either extremely rare or unstable. Cobalt oxides are antiferromagnetic at low temperature: CoO (Neel temperature 291 K) and Co3O4 (Neel temperature: 40 K), which is analogous to magnetite (Fe3O4), with a mixture of +2 and +3 oxidation states. The oxide Co2O3 is probably unstable as it has not been reported yet.

The principal chalcogenides of cobalt include the black cobalt(II) sulfides, CoS2, which adopts a pyrite-like structure, and Co2S3. Pentlandite (Co9S8) is metal-rich.


The four dihalides of cobalt are known: cobalt(II) fluoride (CoF2), cobalt(II) chloride (CoCl2), cobalt(II) bromide (CoBr2), cobalt(II) iodide (CoI2). These dihalides exist as anhydrous and hydrates. Most famously, the anhydrous dichloride is blue, whereas the hydrate is red.

The reduction potential for the reaction:

Co3+ + eCo2+

is +1.92 V, far beyond the one for chlorine. As a consequence cobalt(III) fluoride is one of the few simple stable cobalt(III) compounds. Cobalt(III) fluoride, which is used in some fluorination reactions, reacts vigorously with water.[5]

cobalt(II) chloride hexahydrate

Coordination compounds

As for all metals, molecular compounds of cobalt are classified as coordination complexes, i.e molecules or ions that contain cobalt linked to several ligands. The ligands determine the oxidation state of the cobalt. For example Co+3 complexes tend to have amine ligands. Phosphine ligands tend to feature Co2+ and Co+, an example being tris(triphenylphosphine)cobalt(I) chloride ((P(C6H5)3)3CoCl). Oxide and fluoride can stabilize Co4+ derivatives, e.g. caesium hexafluorocobaltate (Cs2CoF6)) and potassium percobaltate (K3CoO4).[5]

Alfred Werner, a Nobel-prize winning pioneer in coordination chemistry, worked with compounds of empirical formula CoCl3(NH3)6. One of the isomers determined was cobalt(III) hexammine chloride. This coordination complex, a "typical" Werner-type complex, consists of a central cobalt atom coordinated by six ammine ligands orthogonal to each other, and three chloride counteranions.

Using chelating ethylenediamine ligands in place of ammonia gives tris(ethylenediamine)cobalt(III) chloride ([Co(en)3]Cl), which was one of the first coordination complexes that was resolved into optical isomers. The complex exists as both either right- or left-handed forms of a "three-bladed propeller." This complex was first isolated by Werner as yellow-gold needle-like crystals.[6]

Organometallic compounds

Cobaltocene is a stable cobalt analog to ferrocene. Cobalt carbonyl (Co2(CO)8) is a catalyst in carbonylation reactions. Vitamin B12 (see below) is a rare organometallic compound found in nature.


59Cobalt is the only stable cobalt isotope. 22 radioisotopes have been characterized with the most stable being 60Co with a half-life of 5.2714 years, 57Co with a half-life of 271.79 days, 56Co with a half-life of 77.27 days, and 58Co with a half-life of 70.86 days. All of the remaining radioactive isotopes have half-lives that are less than 18 hours, and the majority of these are less than 1 second. This element also has 4 meta states, all of which have half-lives less than 15 minutes.

The isotopes of cobalt range in atomic weight from 50 u (50Co) to 73 u (73Co). The primary decay mode for isotopes with atomic mass unit values less than that of the most abundant stable isotope, 59Co, is electron capture and the primary mode of decay for those of greater than 59 atomic mass units is beta decay. The primary decay products before 59Co are element 26 (iron) isotopes and the primary products after are element 28 (nickel) isotopes.

Cobalt radioisotopes in medicine

Cobalt-60 (Co-60 or 60Co) is a radioactive metal that is used in radiotherapy. It produces two gamma rays with energies of 1.17 MeV and 1.33 MeV. The 60Co source is about 2 cm in diameter and as a result produces a geometric penumbra, making the edge of the radiation field fuzzy. The metal has the unfortunate habit of producing a fine dust, causing problems with radiation protection. Cobalt-60 has a radioactive half-life of 5.27 years. This decrease in activity requires periodic replacement of the sources used in radiotherapy and is one reason why cobalt machines have been largely replaced by linear accelerators in modern radiation therapy. Cobalt from radiotherapy machines has been a serious hazard when not disposed of properly, and one of the worst radiation contamination accidents in North America occurred in 1984, after a discarded cobalt-60 containing radiotherapy unit was mistakenly disassembled in a junkyard in Juarez, Mexico.[7][8]

Cobalt-57 (Co-57 or 57Co) is a cobalt radioisotope most often used in medical tests, as a radiolabel for vitamin B12 uptake, and for the Schilling test.[9]

Industrial uses for radioactive isotopes

Cobalt-60 (Co-60 or 60Co) is useful as a gamma ray source because it can be produced in predictable quantity and high activity by simply exposing natural cobalt to neutrons in a reactor for a period. Its uses include sterilization of medical supplies and medical waste, radiation treatment of foods for sterilization (cold pasteurization), industrial radiography (e.g., weld integrity radiographs), density measurements (e.g., concrete density measurements), and tank fill height switches. Cobalt-57 is used as a source in Mössbauer spectroscopy and is one of several possible sources in XRF devices (Lead Paint Spectrum Analyzers).

Cobalt-60 as weapon

Nuclear weapon designs could intentionally incorporate 59Co, some of which would be activated in a nuclear explosion to produce 60Co. The 60Co, dispersed as nuclear fallout, creates what is sometimes called a dirty bomb or cobalt bomb.[10]


Cobalt compounds have been used for centuries to impart a rich blue color to glass, glazes, and ceramics. Cobalt has been detected in Egyptian sculpture and Persian jewelry from the third millennium BC, in the ruins of Pompeii (destroyed AD 79), and in China dating from the Tang dynasty (AD 618–907) and the Ming dynasty (AD 1368–1644)[11]. Cobalt glass ingots have been recovered from the Uluburun shipwreck, dating to the late 14th century BC.[12]

Swedish chemist Georg Brandt (1694–1768) is credited with isolating cobalt circa 1735.[13] He was able to show that cobalt was the source of the blue color in glass, which previously had been attributed to the bismuth found with cobalt. The word cobalt is derived from the German kobalt, from kobold meaning "goblin", a term used for the ore of cobalt by miners. The first attempts at smelting the cobalt ores to produce cobalt metal failed, yielding cobalt(II) oxide instead. Also, because the primary ores of cobalt always contain arsenic, smelting the ore oxidized into the highly toxic and volatile arsenic oxide.

During the 19th century, cobalt blue was produced at the Norwegian Blaafarveværket (70–80% of world production), led by the Prussian industrialist Benjamin Wegner.

In 1938, John Livingood and Glenn Seaborg discovered cobalt-60. This isotope was famously used at Columbia University in the 1950s to establish parity violation in beta decay.


Cobalt occurs in copper and nickel minerals and in combination with sulfur and arsenic in the sulfidic cobaltite (CoAsS), safflorite (CoAs2) and skutterudite (CoAs3) minerals.[5] The mineral cattierite is similar to pyrite and occurs together vaesite in the copper deposits in the Katanga Province.[14] Upon contact with the atmosphere weathering the sulfide minerals oxidatize to pink erythrite ('cobalt glance': Co3(AsO4)2·8H2O) and sphaerocobaltite (CoCO3).


Cobalt ore
Cobalt output in 2005
World production trend

Cobalt is not found as a native metal but is mainly obtained as a by-product of nickel and copper mining activities. The main ores of cobalt are cobaltite, erythrite, glaucodot, and skutterudite.[15][16]

In 2005, the copper deposits in the Katanga Province (former Shaba province) of the Democratic Republic of the Congo was the top producer of cobalt with almost 40% world share, reports the British Geological Survey.[17] The political situation in the Congo influences the price of cobalt significantly.[18]

Several methods exist for the separation of cobalt from copper and nickel. They depend on the concentration of cobalt and the exact composition of the used ore. One separation step involves froth flotation, in which surfactants bind to different ore components, leading to an enrichment of cobalt ores. Subsequent roasting converts the ores to the cobalt sulfate, whereas the copper and the iron are oxidized to the oxide. The leaching with water extracts the sulfate together with the arsenates. The residues are further leached with sulfuric acid yielding a solution of copper sulfate. Cobalt can also be leached from the slag of the copper smelter.[19]

The products of the above-mentioned processes are transformed into the cobalt oxide Co3O4. This oxide is reduced to the metal by the aluminothermic reaction or reduction with carbon in a blast furnace.[5]



Cobalt-based superalloys consume most of the produced cobalt. The temperature stability of these alloys makes them suitable for use in turbine blades for gas turbines and jet aircraft engines, though nickel-based single crystal alloys surpass them in this regard. Cobalt-based alloys are also corrosion and wear-resistant.[20] Special cobalt-chromium-molybdenum alloys are used for prosthetic parts such as hip and knee replacements.[21] Cobalt alloys are also used for dental prosthetics, where they are useful to avoid allergies to nickel.[22] Some high speed steels also use cobalt to increase heat and wear-resistance. The special alloys of aluminium, nickel, cobalt and iron, known as Alnico, and of samarium and cobalt (samarium-cobalt magnet) are used in permanent magnets.[23]


Lithium cobalt oxide (LiCoO2) is widely used in Lithium ion battery electrodes.[24] Nickel-cadmium (NiCd) and nickel metal hydride (NiMH) batteries also contain significant amounts of cobalt.


Several cobalt compounds are used in chemical reactions as oxidation catalysts. Cobalt acetate is used for the conversion of xylene to terephthalic acid, the precursor to the bulk polymer Polyethylene terephthalate. Typical catalysts are the cobalt carboxylates (known as cobalt soaps). They are also used in paints, varnishes, and inks as "drying agents" through the oxidation of drying oils.[24] The same carboxylates are used to improve the adhesion of the steel to rubber in steel-belted radial tires.

Cobalt-based catalysts are also important in reactions involving carbon monoxide. Steam reforming, useful in hydrogen production, uses cobalt oxide-base catalysts. Cobalt is also a catalyst in the Fischer-Tropsch process, used in the conversion of carbon monoxide into liquid fuels.[25] The hydroformylation of alkenes often rely on cobalt octacarbonyl as the catalyst,[26] although such processes have been displaced by more efficient iridium- and rhodium-based catalysts, e.g. the Cativa process.

The hydrodesulfurization of petroleum uses a catalyst derived from cobalt and molybdenum. This process helps to rid petroleum of sulfur impurities that interfere with the refining of liquid fuels.[24]

Pigments and coloring

Cobalt blue glass

Before the 19th century, the predominant use of cobalt was as pigment. Since the Middle Age, the production of smalt, a blue colored glass was known. Smalt is produced by melting a mixture of the roasted mineral smaltite, quartz and potassium carbonate, yielding a dark blue silicate glass which is grinded after the production.[27] Smalt was widely used for the coloration of glass and as pigment for paintings.[28] In 1780 Sven Rinman discovered cobalt green and in 1802 Louis Jacques Thénard discovered cobalt blue.[29] The two colors cobalt blue, a cobalt aluminate, and cobalt green, a mixture of cobalt(II) oxide and zinc oxide, were used as pigments for paintings due to their superior stability.[30][31]

Cobalt has been used to color glass since the Bronze Age. The excavation of the Uluburun shipwreck yielded an ingot of blue glass which was cast during the 14th century BC.[32] Blue glass items from Egypt are colored with copper, iron, or cobalt. The oldest cobalt-colored glass was from the time of the Eighteenth dynasty (1550–1292 BC). The location where the cobalt compounds were obtained is unknown.[33][34]

Other uses

Biological role


Cobalt is essential to all animals, including humans. It is a key constituent of cobalamin, also known as vitamin B12. A deficiency of cobalt leads to pernicious anemia, a lethal disorder. Pernicious anemia is however very rare, because trace amounts of cobalt are available in most diets. The presence of 0.13 to 0.30 mg/kg of cobalt in soils markedly improves the health of grazing animals.

The cobalamin-based proteins use corrin to hold the cobalt. Coenzyme B12 features a reactive C-Co bond, which participates in its reactions.[35] In humans, B12 exists with two types alkyl ligand, methyl and adenosyl. MeB12 promotes methyl (-CH3) group transfers. The adenosyl version of B12 catalyzes rearrangements in which a hydrogen atom is directly transferred between two adjacent atoms with concomitant exchange of the second substituent, X, which may be a carbon atom with substituents, an oxygen atom of an alcohol, or an amine. Methylmalonyl Coenzyme A mutase (MUT) converts MMl-CoA to Su-CoA, an important step in the extraction of energy from proteins and fats.

Although far less common than other metalloproteins (e.g. those of zinc and iron), cobaltoproteins are known aside from non-B12. These proteins include Methionine aminopeptidase 2 and Nitrile hydratase are two examples.[36]


Cobalt is an essential element for life in minute amounts. The LD50 values soluble cobalt salts has been estimated to be between 150 and 500 mg/kg. Thus, for a 100 kg person the LD50 would be about 20 grams.[37]

After nickel and chromium, cobalt is a major cause of contact dermatitis and is considered carcinogenic.[38]


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  23. ^ Luborsky, F. E.; Mendelsohn, L. I.; Paine, T. O. (1957). "Reproducing the Properties of Alnico Permanent Magnet Alloys with Elongated Single-Domain Cobalt-Iron Particles". Journal Applied Physics 28 (344): 344. doi:10.1063/1.1722744. 
  24. ^ a b c Hawkins, M. (2001). "Why we need cobalt". Applied Earth Science: Transactions of the Institution of Mining & Metallurgy, Section B 110 (2): 66–71. 
  25. ^ Andrei Y. Khodakov, Wei Chu, and Pascal Fongarland “Advances in the Development of Novel Cobalt Fischer-Tropsch Catalysts for Synthesis of Long-Chain Hydrocarbons and Clean Fuels” Chemical Review, 2007, volume 107, pp 1692–1744. doi:10.1021/cr050972v
  26. ^ Frdric Hebrard and Philippe Kalck “Cobalt-Catalyzed Hydroformylation of Alkenes: Generation and Recycling of the Carbonyl Species, and Catalytic Cycle” Chemical Reviews, 2009, volume 109, pp 4272–4282. doi:10.1021/cr8002533
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  30. ^ Witteveen, H. J. (1921). "Colors Developed by Cobalt Oxides". Industrial & Engineering Chemistry 13: 1061. doi:10.1021/ie50143a048. 
  31. ^ Venetskii, S. (1970). "The charge of the guns of peace". Metallurgist 14 (5): 334–336. doi:10.1007/BF00739447. 
  32. ^ Henderson, Julian (2000). "Glass". The Science and Archaeology of Materials: An Investigation of Inorganic Materials. Routledge. p. 60. ISBN 9780415199339. 
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External links

Travel guide

Up to date as of January 14, 2010

From Wikitravel


Cobalt is a small town in Northern Ontario, Canada.


Cobalt has some run-down buildings that used to be stores. You could take pictures of yourself in front of them, but the doors are locked and nobody is home.

Cobalt has some old mineshafts from back when the town started up.

There is a Classic Theater that draws people in to see live performances of struggling/starving performance artists.

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1911 encyclopedia

Up to date as of January 14, 2010

From LoveToKnow 1911

COBALT (symbol Co, atomic weight 59), one of the metallic chemical elements. The term "cobalt" is met with in the writings of Paracelsus, Agricola and Basil Valentine, being used to denote substances which, although resembling metallic ores, gave no metal on smelting. At a later date it was the name given to the mineral used for the production of a blue colour in glass. In 1735 G. Brandt prepared an impure cobalt metal, which was magnetic and very infusible. Cobalt is usually found associated with nickel, and frequently with arsenic, the chief ores being speiss-cobalt, (Co,Ni,Fe)As 2, cobaltite, wad, cobalt bloom, linnaeite, Co 3 S 4, and skutterudite, CoAs 3. Its presence has also been detected in the sun and in meteoric iron. For the technical preparation of cobalt, and its separation from nickel, see Nickel. The metal is chiefly used, as the oxide, for colouring glass and porcelain.

Metallic cobalt may be obtained by reduction of the oxide or chloride in a current of hydrogen at a red heat, or by heating the oxalate, under a layer of powdered glass. As prepared by the reduction of the oxide it is a grey powder. In the massive state it has a colour resembling polished iron, and is malleable and very tough. It has a specific gravity of 8.8, and it melts at 1530° C. (H. Copaux). Its mean specific heat between 9° and 97° C. is x 10674 (H. Kopp). It is permanent in dry air, but in the finely divided state it rapidly combines with oxygen, the compact metal requiring a strong heating to bring about this combination. It decomposes steam at a red heat, and slowly dissolves in dilute hydrochloric and sulphuric acids, but more readily in nitric acid. Cobalt burns in nitric oxide at 150° C. giving the monoxide. It may be obtained in the pure state, according to C. Winkler (Zeit. fiir anorg. Chem., 1895, 8, p. 1), by electrolysing the pure sulphate in the presence of ammonium sulphate and ammonia, using platinum electrodes, any occluded oxygen in the deposited metal being removed by heating in a current of hydrogen.

Three characteristic oxides of cobalt are known, the monoxide, CoO, the sesquioxide, C0203, and tricobalt tetroxide, C0304; besides these there are probably oxides of composition Co02, Co 8 0 9, C0607 and C0405. Cobalt monoxide, CoO, is prepared by heating the hydroxide or carbonate in a current of air, or by heating the oxide C0304 in a current of carbon dioxide. It is a brown coloured powder which is stable in air, but gives a higher oxide when heated. On heating in hydrogen, ammonia or carbon monoxide, or with carbon or sodium, it is reduced to the metallic state. It is readily soluble in warm dilute mineral acids forming cobaltous salts. Cobaltous hydroxide, Co(OH) 21 is formed when a cobaltous salt is precipitated by caustic potash in the absence of air. A blue basic salt is precipitated first, which, on boiling, rapidly changes to the rose-coloured hydroxide. It dissolves in acids forming cobaltous salts, and on exposure to air it rapidly absorbs oxygen, turning brown in colour. A. de Schulten (Comptes Rendus, 1889, 109, p. 266) has obtained it in a crystalline form; the crystals have a specific gravity of 3.597, and are easily soluble in warm ammonium chloride solution. Cobalt sesquioxide, Co 2 0 3, remains as a dark-brown powder when cobalt nitrate is gently heated. Heated at 190-300° in a current of hydrogen it gives the oxide C0304, while at higher temperatures the monoxide is formed, and ultimately cobalt is obtained. Cobaltic hydroxide, Co(OH) 31 is formed when a cobalt salt is precipitated by an alkaline hypochlorite, or on passing chlorine through water containing suspended cobaltous hydroxide or carbonate. It is a brown-black powder soluble in hydrochloric acid, chlorine being simultaneously liberated. This hydroxide is soluble in well cooled acids, forming solutions which contain cobaltic salts, one of the most stable of which is the acetate. Cobalt dioxide, Co02, has not yet been isolated in the pure state; it is probably formed when iodine and caustic soda are added to a solution of a cobaltous salt. By suspending cobaltous hydroxide in water and adding hydrogen peroxide, a strongly acid liquid is obtained (after filtering) which probably contains cobaltous acid, H2CoO 3. The barium and magnesium salts of this acid are formed when baryta and magnesia are fused with cobalt sesquioxide. Tricobalt tetroxide, C0304, is produced when the other oxides, or the nitrate, are heated in air.

By heating a mixture of cobalt oxalate and sal-ammoniac in air, it is obtained in the form of minute hard octahedra, which are not magnetic, and are only soluble in concentrated sulphuric acid.

The cobaltous salts are formed when the metal, cobaltous oxide, hydroxide or carbonate, are dissolved in acids, or, in the case of the insoluble salts, by precipitation. The insoluble salts are rose-red or violet in colour. The soluble salts are, when in the hydrated condition, also red, but in the anhydrous condition are blue. They are precipitated from their alkaline solutions as cobalt sulphide by sulphuretted hydrogen, but this precipitation is prevented by the presence of citric and tartaric acids; similarly the presence of ammonium salts hinders their precipitation by caustic alkalis. Alkaline carbonates give precipitates of basic carbonates, the formation of which is also retarded by the presence of ammonium salts. For the action of ammonia on the cobaltous salts in the presence of air see Cobaltammines (below). On the addition of potassium cyanide they give a brown precipitate of cobalt cyanide, Co(CN) 2, which dissolves in excess of potassium cyanide to a green solution.

Cobalt chloride, CoC1 2, in the anhydrous state, is formed by burning the metal in chlorine or by heating the sulphide in a current of the same gas. It is blue in colour and sublimes readily. It dissolves easily in water, forming the hydrated chloride, CoC12.6H20, which may also be prepared by dissolving the hydroxide or carbonate in hydrochloric acid. The hydrated salt forms rose-red prisms, readily soluble in water to a red solution, and in alcohol to a blue solution. Other hydrated forms of the chloride, of composition CoCl 2.2H 2 O and CoCl 2.4H 2 O have been described (P. Sabatier, Bull. Soc. Chim. 51, p. 88; Bersch, Jahresb. d. Chemie, 1867, p. 291). Double chlorides of composition CoC1 2 NH 4 C1.6H 2 O; CoC1 2 SnCl 4.6H 2 0 and CoC1 2.2CdC1 2.12H 2 O are also known. By the addition of excess of ammonia to a cobalt chloride solution in absence of air, a greenishblue precipitate is obtained which, on heating, dissolves in the solution, giving a rose-red liquid. This solution, on standing, deposits octahedra of the composition CoC1 2.6NH 3. These crystals when heated to 120° C. lose ammonia and are converted into the compound CoC1 2.2NH 3 (E. Fremy). The bromide, CoBr 2, resembles the chloride, and may be prepared by similar methods. The hydrated salt readily loses water on heating, forming at 100° C. the hydrate CoBr 2.2H 2 O, and at 130° C. passing into the anhydrous form. The iodide, Co12, is produced by heating cobalt and iodine together, and forms a greyish-green mass which dissolves readily in water forming a red solution. On evaporating this solution the hydrated salt CoI 2.6H 2 0 is obtained in hexagonal prisms. It behaves in an analogous manner to CoBr 2.6H 2 0 on heating.

Cobalt fluoride, CoF 2.2H 2 0, is formed when cobalt carbonate is evaporated with an excess of aqueous hydrofluoric acid, separating in rose-red crystalline crusts. Electrolysis of a solution in hydrofluoric acid gives cobaltic fluoride, CoF3.

Sulphides of cobalt of composition C04S3, CoS, C03S4, C02S3 and CoS 2 are known. The most common of these sulphides is cobaltous sulphide, CoS, which occurs naturally as syepoorite, and can be artificially prepared by heating cobaltous oxide with sulphur, or by fusing anhydrous cobalt sulphate with barium sulphide and common salt. By either of these methods, it is obtained in the form of bronzecoloured crystals. It may be prepared in the amorphous form by heating cobalt with sulphur dioxide, in a sealed tube, at 200° C. In the hydrated condition it is formed by the action of alkaline sulphides on cobaltous salts, or by precipitating cobalt acetate with sulphuretted hydrogen (in the absence of free acetic acid). It is a black amorphous powder soluble in concentrated sulphuric and hydrochloric acids, and when in the moist state readily oxidizes on exposure.

Cobaltous sulphate, CoSO 4.7H 2 O, is found naturally as the mineral bieberite, and is formed when cobalt, cobaltous oxide or carbonate are dissolved in dilute sulphuric acid. It forms dark red crystals isomorphous with ferrous sulphate, and readily soluble in water. By dissolving it in concentrated sulphuric acid and warming the solution, the anhydrous salt is obtained. Hydrated sulphates of composition CoS04.6H20, CoSO 4.4H 2 O and CoS04 H 2 0 are also, known. The heptahydrated salt combines with the alkaline sulphates to form double sulphates of composition CoS04 M2S04.6H20 (M = K, NH4, &c.).

The cobaltic salts corresponding to the oxide Co 2 0 3 are generally unstable compounds which exist only in solution. H. Marshall (Proc. Roy. Soc. Edin. 59, p. 760) has prepared cobaltic sulphate C02(S04)3.18H20, in the form of small needles, by the electrolysis of cobalt sulphate. In a similar way potassium and ammonium cobalt alums have been obtained. A cobaltisulphurous acid, probably H 6 [(S03)6 C02] has been obtained by E. Berglund (Berichte, 18 74, 7, p. 469), in aqueous solution, by dissolving ammonium cobaltocobaltisulphite (NH4)2C02 [(S03) 6 'C02] 14H 2 O in dilute hydrochloric or nitric acids, or by decomposition of its silver salt with hydrochloric acid. The ammonium cobalto-cobaltisulphite is prepared by saturating an air-oxidized ammoniacal solution of cobaltous chloride with sulphur dioxide. The double salts containing the metal in the cobaltic form are more stable than the corresponding single salts, and of these potassium cobaltinitrite, C02(N02) 6 '6KN02.3H20, is best known. It may be prepared by the addition of potassium nitrite to an acetic acid solution of cobalt chloride. The yellow precipitate obtained is washed with a solution of potassium acetate and finally with dilute alcohol. The reaction proceeds according to the following equation: 2CoC12+10KN02+ 4HNO 2 = C02(N02)6.6KN02+4KC1+2N0+2H20 (A. Stromeyer, Annalen, 1855, 96, p. 220). This salt may be used for the separation of cobalt and nickel, since the latter metal does not form a similar double nitrite, but it is necessary that the alkaline earth metals should be absent, for in their presence nickel forms complex nitrites containing the alkaline earth metal and the alkali metal. A sodium cobaltinitrite is also known.

Cobalt nitrate, Co(NO 3) 2.6H 2 0, is obtained in dark-red monoclinic tables by the slow evaporation of a solution of the metal, its hydroxide or carbonate, in nitric acid. It deliquesces in the air and melts readily on heating. By the addition of excess of ammonia to its aqueous solution, in the complete absence of air, a blue precipitate of a basic nitrate of the composition 6C00 N 2 0 6 5H 2 O is obtained.

By boiling a solution of cobalt carbonate in phosphoric acid, the acid phosphate CoHPO 4.3H 2 O is obtained, which when heated with water to 250° C. is converted into the neutral phosphate C03(P04)2.2H20 (H. Debray, Ann. de chimie, 1861, [3] 61, P. 438). Cobalt ammonium phosphate, CoNH4PO 4.12H 2 0, is formed when a soluble cobalt salt is digested for some time with excess of a warm solution of ammonium phosphate. It separates in the form of small rose-red crystals, which decompose on boiling with water.

Cobaltous cyanide, Co(CN)2.3H20, is obtained when the carbonate is dissolved in hydrocyanic acid or when the acetate is precipitated by potassium cyanide. It is insoluble in dilute acids, but is readily soluble in excess of potassium cyanide. The double cyanides of cobalt are analogous to those of iron. Hydrocobaltocyanic acid is not known, but its potassium salt, K4Co(CN) 6, is formed when freshly precipitated cobalt cyanide is dissolved in an ice-cold solution of potassium cyanide. The liquid is precipitated by alcohol, and the washed and dried precipitate is then dissolved in water and allowed to stand, when the salt separates in dark-coloured crystals. In alkaline solution it readily takes up oxygen and is converted into potassium cobalticyanide, K 3 Co(CN) 6, which may also be obtained by evaporating a solution of cobalt cyanide, in excess of potassium cyanide, in the presence of air, 8KCN+2Co(CN)2+H20+0= 2K 3 Co(CN) 6 +2KHO. It forms monoclinic crystals which are very soluble in water. From its aqueous solution, concentrated hydrochloric acid precipitates hydrocobalticyanic acid, H 3 Co(CN) 61 as a colourless solid which is very deliquescent, and is not attacked by concentrated hydrochloric and nitric acids. For a description of the various salts of this acid, see P. Wesselsky, Berichte, 1869, 2, p. 588.

Cobaltammines. A large number of cobalt compounds are known, of which the empirical composition represents them as salts of cobalt to which one or more molecules of ammonia have been added. These salts have been divided into the following series: Diammine Series, [Co(NH3)2]X4M. In these salts X = NO 2 and M = one atomic proportion of a monovalent metal, or the equivalent quantity of a divalent metal.

Triammine Series, [Co(NH 3)3]X3. Here X = Cl, N03, N02, 2S04, &c.

Tetrammine Series. This group may be divided into the Praseo-salts [R 2 Co(NH 3) 4 ]X, where X = Cl.

Croceo-salts [(N02)2Co(NH3)4]X, which may be considered as a subdivision of the praseo-salts.

Tetrammine purpureo-salts [RCo(NH3)4 H20]X2. Tetrammine roseo-salts [Co(NH3)4 (H20)21X3.

Fuseo-salts [Co(NH 3)4]OH X2.

Pentammine Series.

Pentammine purpureo-salts [R Co(NH 3) 5 ]X 2 where X = Cl, Br, N03, N02, 1S04, &c.

Pentammine roseo-salts [Co(NH 3) 6 H 2 O] X2.

Hexammine or Luteo Series [Co(NH3)6] X3.

The hexammine salts are formed by the oxidizing action of air on dilute ammoniacal solutions of cobaltous salts, especially in presence of a large excess of ammonium chloride. They form yellow or bronze-coloured crystals, which decompose on boiling their aqueous solution. On boiling their solution in caustic alkalis, ammonia is liberated. The pentammine purpureo-salts are formed from the luteo-salts by loss of ammonia, or from an air slowly oxidized ammoniacal cobalt salt solution, the precipitated luteosalt being filtered off and the filtrate boiled with concentrated acids. They are violet-red in colour, and on boiling or long standing with dilute acids they pass into the corresponding roseo-salts.

The pentammine nitrito salts are known as the xanthocobalt salts and have the general formula [NO 2 Co (NH 3) 3]X2. They are formed by the action of nitrous fumes on ammoniacal solutions of cobaltous salts, or purpureo-salts, or by the mutual reaction of chlorpurpureosalts and alkaline nitrites. They are soluble in water and give characteristic precipitates with platinic and auric chlorides, and with potassium ferrocyanide. The pentammine roseo-salts can be obtained from the action of concentrated acids, in the cold, on airoxidized solutions of cobaltous salts. They are of a reddish colour and usually crystallize well; on heating with concentrated acids are usually transformed into the purpureo-salts. Their alkaline solutions liberate ammonia on boiling. They give a characteristic pale red precipitate with sodium pyrophosphate, soluble in an excess of the precipitant; they also form precipitates on the addition of platinic chloride and potassium ferrocyanide. For methods of preparation of the tetrammine and triammine salts, see 0. Dammer's Handbuch der anorganischen Chemie, vol. 3 (containing a complete account of the preparation of the cobaltammine salts). The diammine salts are prepared by the action of alkaline nitrites on cobaltous salts in the presence of much ammonium chloride or nitrate; they are yellow or brown crystalline solids, not very soluble in cold water. The above series of salts show striking differences in their behaviour towards reagents; thus, aqueous solutions of the luteo chlorides are strongly ionized, as is shown by their high electric conductivity; and all their chlorine is precipitated on the addition of silver nitrate solution. The aqueous solution, however, does not show the ordinary reactions of cobalt or of ammonia, and so it is to be presumed that the salt ionizes into [Co(NH 3) 6 ] and 3C1'. The purpureo chloride has only two-thirds of its chlorine precipitated on the addition of silver nitrate, and the electric conductivity is much less than that of the luteo chloride; again in the praseosalts only one-third of the chlorine is precipitated by silver nitrate, the conductivity again falling; while in the triammine salts all ionization has disappeared. For the constitution of these salts and of the "metal ammonia" compounds generally, see A. Werner, Zeit. fiir anorg. Chemie, 1893 et seq., and Berichte, 1895, et seq.; and S. Jorgensen, Zeit. fiir anorg. Chemie, 1892 et seq.

The oxycobaltammines are a series of compounds of the general type [Co 20341 2 (NH 3) 1 o]X 4 first observed by L. Gmelin, and subsequently examined by E. Fremy, W. Gibbs and G. Vortmann (Monatshefte fur Chemie, 1885, 6, p. 404). They result from the cobaltammines by the direct taking up of oxygen and water. On heating, they decompose, forming basic tetrammine salts.

The atomic weight of cobalt has been frequently determined, the earlier results not being very concordant (see R. Schneider, Pog. Ann., 1857, 101, p. 387; C. Marignac, Arch. Phys. Nat. [2], I, p. 373; W. Gibbs, Amer. Jour. Sci. [2], 2 5, p. 483; J. B. Dumas, Ann. Chim. Phys., 18 59 [3], 55, p. 129; W. J. Russell, Jour. Chem. Soc., 1863, 16, p. 51). C. Winkler, by the analysis of the chloride, and by the action of iodine on the metal, obtained the values 59.37 and 59.07, whilst W. Hempel and H. Thiele (Zeit. f. anorg. Chem., 1896, II, p. 73), by reducing cobalto-cobaltic oxide, and by the analysis of the chloride, have obtained the values 58.56 and 58.48. G. P. Baxter and others deduced the value 58.995 (0 =16).

Cobalt salts may be readily detected by the formation of the black sulphide, in alkaline solution, and by the blue colour they produce when fused with borax. For the quantitative determination of cobalt, it is either weighed as the oxide, C0304, obtained by ignition of the precipitated monoxide, or it is reduced in a current of hydrogen and weighed as metal. For the quantitative separation of cobalt and nickel, see E. Hintz (Zeit. f. anal. Chem., 1891, 30, p. 227), and also Nickel.

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cobalt (uncountable)

  1. a chemical element (symbol Co) with an atomic number of 27.

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  • IPA: /kɔ.balt/


cobalt m (usually uncountable)

  1. (chemistry) cobalt

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[[File:|thumb|Sheet of cobalt metal]] Cobalt (chemical symbol Co) is a chemical element. It has an atomic number of 27 and an atomic mass of about 59. It is a metal.



Cobalt(II) chloride with water
Cobalt(II) chloride without water

Cobalt is a transition metal. It is shiny and conducts electricity. It is magnetic. It is a hard metal. It is moderately reactive. Iron is more reactive and copper is less reactive. It dissolves slowly in acids. This reaction makes hydrogen and a salt of cobalt. Cobalt is normally in its +2 oxidation state as an ion. Some chemical compounds contain cobalt ions in its +4 oxidation state. Cobalt(II) chloride is one of the most common cobalt compounds. Many cobalt compounds are blue or pink. One of them is black.

Occurrence and preparation

Cobalt is too reactive to occur pure in the earth. It is found in certain minerals. It is found with copper and nickel deposits. Normally the three metals are bonded to arsenic and sulfur. It is found as a byproduct (left over substance) when copper and nickel are produced. It is made by reaction with the sludge from copper and nickel processing.


Cobalt is used in some types of steel. It hardens the steel. It is also used to make very strong tough alloys. These alloys are known as superalloys. Some cobalt compounds are used in the lithium-ion battery. Cobalt compounds were used as an artificial food coloring until 1971. It was discovered that it has harmful effects. It is used to make glass blue. It is also used as a catalyst. It is also used in some medicines.

The human body needs small amounts of cobalt for certain vitamins. Cobalt compounds are used to stop cyanide from poisoning the body.


Cobalt compounds are needed in small amounts, but they are toxic in large quantities. Sometimes cobalt compounds were added to beer and poisoned people that drank it. It can make skin irritation when touched.


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