Cobalt

iron ← cobalt → nickel - ↑ Co ↓ Rh 27Co Periodic table
Appearance
hard lustrous gray metal
General properties
Name, symbol, number	cobalt, Co, 27
Pronunciation	/ˈkoʊbɒlt/ KOH-bolt[1]
Element category	transition metal
Group, period, block	9, 4, d
Standard atomic weight	58.933195g·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[2] (amphoteric oxide)
Electronegativity	1.88 (Pauling scale)
Ionization energies (more)	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
Miscellanea
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
v • d • e

Cobalt ( /ˈkoʊbɒlt/ or /ˈkoʊbɔːlt/)^[3][4] is a chemical element with symbol Co and atomic number 27. It is found naturally only in chemically combined form. The free element, produced by reductive smelting, is a hard, lustrous, silver-gray metal.

Cobalt-based blue pigments have been used since ancient times for jewelry and paints, and to impart a distinctive blue tint to glass, but the color was later thought by alchemists to be due to the known metal bismuth. Miners had long used the name Kobold ore (German for goblin ore) for some of these minerals; they were named so because they gave poisonous arsenic-containing fumes upon smelting. In 1735, such ores were found to be reducible to a new metal (the first discovered since ancient times), and this was ultimately named for the Kobold. Today, some cobalt is produced specifically from various metallic-lustered ores, for example cobaltite (CoAsS), but the main source of the element is 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 metal mined worldwide.

Cobalt is used in the preparation of magnetic, wear-resistant, and high-strength alloys. Smalt (cobalt silicate glass) and cobalt blue (cobalt(II) aluminate, CoAl₂O₄) gives a distinctive deep blue color to glass, ceramics, inks, paints, and varnishes. Cobalt occurs naturally as only one stable isotope, cobalt-59. 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 animal organisms, as the active center of coenzymes called cobalamins. These include vitamin B₁₂ which is essential for mammals. Cobalt is also an active nutrient for bacteria, algae, and fungi.

Contents 1 Characteristics 1.1 Physical 1.2 Chemical 1.3 Isotopes 2 History 2.1 Occurrence 3 Production 4 Compounds 4.1 Oxygen and chalcogen compounds 4.2 Halides 4.3 Coordination compounds 4.4 Organometallic compounds 5 Applications 5.1 Alloys 5.2 Batteries 5.3 Catalysis 5.4 Pigments and coloring 5.5 Radioisotopes in medicine 5.5.1 Cobalt-60 as weapon 5.6 Other uses 6 Biological role 7 Precautions 8 References 9 External links

[edit] Characteristics

[edit] Physical

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.^[5]

[edit] Chemical

Cobalt is weakly reducing metal that is protected from oxidation by a passivating oxide film, as is typical for most metals. It is attacked by halogens and sulfur. Heating in oxygen gives Co₃O₄ which loses oxygen at 900 °C to give the monoxide CoO.^[6]

[edit] Isotopes

⁵⁹Co is the only stable cobalt isotope and the only isotope to exist in nature. 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.^[7]

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.^[7]

[edit] History

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).^[8]

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.^[9][10] 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 in Egypt (1550–1292 BC). The location where the cobalt compounds were obtained is unknown.^[11][12]

The word cobalt is derived from the German kobalt, from kobold meaning "goblin", a superstitious term used for the ore of cobalt by miners. The first attempts at smelting these ores to produce metals such as copper or nickel failed, yielding simply powder (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, which also decreased the reputation of the ore for the miners.^[13]

Swedish chemist Georg Brandt (1694–1768) is credited with discovering cobalt circa 1735, showing it to be a new previously unknown element different from bismuth and other traditional metals, and calling it a new "semi-metal."^[14][15] He was able to show that compounds of cobalt metal were the source of the blue color in glass, which previously had been attributed to the bismuth found with cobalt. Cobalt would thus become the first new metal to be discovered since the traditional metals known from antiquity, which have no named discoverer.

During the 19th century, cobalt blue and smalt were produced at the Norwegian Blaafarveværket (70–80% of world production), led by the Prussian industrialist Benjamin Wegner.^{[citation needed]} The first mining for the production of smalt in the 16th to 18th century were located in Norway, Sweden, Saxony and Hungary. With the discovery of cobalt ore in New Caledonia in 1864 the mining of cobalt in Europe declined. With the discovery of ore deposits in Ontario Canada in 1904 and the discovery of even larger deposits in the Katanga Province in the Kongo in 1914 the mining operations shifted again.^[13] With the Shaba conflict starting in the 1978 the main source for cobalt the copper mines of Katanga Province nearly stopped their production.^[16][17] The impact on the economy by was smaller than expected because industry established effective ways for recycling cobalt materials and changed to materials without cobalt.^[16][17]

In 1938, John Livingood and Glenn T. Seaborg discovered cobalt-60.^[18] This isotope was famously used at Columbia University in the 1950s to establish parity violation in beta decay.^[19][20]

[edit] Occurrence

The stable form of cobalt is created in supernovas via the r-process. It comprises 0.0029% of the Earth's crust and is one of the first transition metal series.

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.^[6] The mineral cattierite is similar to pyrite and occurs together vaesite in the copper deposits in the Katanga Province.^[21] Upon contact with the atmosphere weathering the sulfide minerals oxidatize to pink erythrite ("cobalt glance": Co₃(AsO₄)₂·8H₂O) and sphaerocobaltite (CoCO₃).^[22][23]

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.^[24][25]

[edit] Production

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.^[26] The political situation in the Congo influences the price of cobalt significantly.^[27]

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.^[28]

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.^[6]

[edit] Compounds

See also Category: Cobalt compounds Common oxidation states of cobalt include +2 and +3, although compounds with oxidation states ranging from −3 to +4 are also known. A common oxidation state for simple compounds is +2. Cobalt(II) salts form the red-pink [Co(H₂O)₆]²⁺ complex in aqueous solution. Addition of chloride gives the intensely blue [CoCl₄]²⁻.^[2]

Cobalt compounds release a blue-green flame when heated.

[edit] 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 600–700 °C, CoO oxidizes to the blue cobalt(II,III) oxide (Co₃O₄), which has spinel structure.^[2] Black cobalt(III) oxide (Co₂O₃) is also known.^[29] 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.^[30]

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.^[2]

[edit] Halides

The four dihalides of cobalt are known: cobalt(II) fluoride (CoF₂, pink), cobalt(II) chloride (CoCl₂, blue), cobalt(II) bromide (CoBr₂, green), cobalt(II) iodide (CoI₂, blue-black). These dihalides exist as anhydrous and hydrates. Whereas the anhydrous dichloride is blue, the hydrate is red.^[31]

The reduction potential for the reaction

Co³⁺ + e⁻ → Co²⁺

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.^[6]

[edit] 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).^[6]

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.^[32]

[edit] Organometallic compounds

Cobaltocene is a highly 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 and is the only vitamin to contain a metal atom.

[edit] Applications

The main application of cobalt is as the metal in alloys.

[edit] Alloys

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.^[33] Cobalt-based alloys are also corrosion and wear-resistant. The development of the wear resistant cobalt alloys started in the first decade of the 19th century with the development of the stellite alloys. The stellite alloys are cobalt chromium alloys with varying tungsten and carbon content. The formation of chromium and tungsten carbides makes the very hard and wear resistant.^[34] Special cobalt-chromium-molybdenum alloys are used for prosthetic parts such as hip and knee replacements.^[35] Cobalt alloys are also used for dental prosthetics, where they are useful to avoid allergies to nickel.^[36] 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.^[37]

[edit] Batteries

Lithium cobalt oxide (LiCoO₂) is widely used in lithium ion battery cathodes. The material is composed of cobalt oxide layers in which the lithium is intercalated. During discharging the lithium intercalated between the layers is set free as lithium ion.^[38] nickel-cadmium ^[39] (NiCd) and nickel metal hydride^[40] (NiMH) batteries also contain significant amounts of cobalt, the cobalt improves the oxidation capabilities of nickel in the battery.

[edit] Catalysis

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.^[38] 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 hydrogenation of carbon monoxide into liquid fuels.^[41] The hydroformylation of alkenes often rely on cobalt octacarbonyl as the catalyst,^[42] although such processes have been partially 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.^[38]

[edit] Pigments and coloring

Before the 19th century, the predominant use of cobalt was as pigment. Since the Middle Ages, it has been involved in the production of smalt, a blue colored glass. Smalt is produced by melting a mixture of the roasted mineral smaltite, quartz and potassium carbonate, yielding a dark blue silicate glass which is ground after the production.^[43] Smalt was widely used for the coloration of glass and as pigment for paintings.^[44] In 1780, Sven Rinman discovered cobalt green and in 1802 Louis Jacques Thénard discovered cobalt blue.^[45] 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.^[46][47]

[edit] Radioisotopes in medicine

Cobalt-60 (Co-60 or ⁶⁰Co) (Co-60 or ⁶⁰Co) is useful as a gamma ray source because it can be produced in predictable quantity and high activity by bombarding cobalt with neutrons. It produces two gamma rays with energies of 1.17 MeV and 1.33 MeV.

Its uses include radiotherapy, 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. The metal has the unfortunate habit of producing a fine dust, causing problems with radiation protection. 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.^[48][49]

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-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.^[50] Cobalt-57 is used as a source in Mössbauer spectroscopy and is one of several possible sources in XRF devices (Lead Paint Spectrum Analyzers).

[edit] 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 cobalt bomb.^[51]

[edit] Other uses

• Electroplating due to its appearance, hardness, and resistance to oxidation.^[52]
• Ground coats for porcelain enamels.^[53]

[edit] 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 anemia, a lethal disorder. Anemia secondary to cobalt deficiency is very rare, though, 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.^{[citation needed]}

The cobalamin-based proteins use corrin to hold the cobalt. Coenzyme B₁₂ features a reactive C-Co bond, which participates in its reactions.^[54] In humans, B₁₂ exists with two types of 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 B₁₂. These proteins include methionine aminopeptidase 2 and nitrile hydratase.^[55]

[edit] Precautions

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.^[56]

After nickel and chromium, cobalt is a major cause of contact dermatitis and is considered carcinogenic.^[57] In 1966, the addition of cobalt compounds to stabilize beer foam in Canada led to cardiomyopathy, which came to be known as beer drinker's cardiomyopathy.^[58]

[edit] References

[edit] External links

Wikimedia Commons has media related to: Cobalt

Look up cobalt in Wiktionary, the free dictionary.

• National Pollutant Inventory – Cobalt fact sheet
• WebElements.com – Cobalt
• London celebrates 50 years of Cobalt-60 Radiotherapy

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Large version

v • d • e   Cobalt compounds CoBr2 · Co(CN)2 · CoCO3 · CoCl2 · Co(ClO3)2 · CoF2 · CoF3 · CoI2 · Co(NO3)2 · CoO · Co(OH)2 · CoS · CoSO4 · Co2O3 · Co3O4