21.10A: The Metal

The origin of the name Cobalt is thought to stem from the German kobold for "evil spirits or goblins", who were superstitiously thought to cause trouble for miners, since the cobalt minerals contained arsenic that injured their health and the cobalt ores did not yield metals when treated using the normal methods. The name could also be derived from the Greek kobalos for "mine". Cobalt was discovered in 1735 by the Swedish chemist Georg Brandt.

Occurrence

The principal ores of Cobalt are cobaltite, [(Co,Fe)AsS], erythrite, [Co₃(AsO₄)₂.8(H₂O)], glaucodot, [(Co,Fe)AsS], and skutterudite, [CoAs₃]. World production of cobalt has steadily increased in recent years, almost trebling since 1993. The dominance of African copper-cobalt producers has been replaced by a more even spread of output between leading producing countries, with Canada, Norway and more recently Australia, together with exports from Russia, replacing lost production in the Democratic Republic of Congo (Zaire). The strongest growth in production of cobalt has come from Finland, where output grew at over 16% between 1990 and 2002.

Extraction

Not covered in this course.

Uses

• Alloys, such as:
  Superalloys, for parts in gas turbine aircraft engines.
  Corrosion- and wear-resistant alloys. Estimated as about 20% of production in 2003
• High-speed steels.
• Cemented carbides (also called hard metals) and diamond tools.
• Magnets and magnetic recording media.
• Catalysts for the petroleum and chemical industries.
• electroplating because of its appearance, hardness, and resistance to oxidation.
• Drying agents for paints, varnishes, and inks.
• Ground coats for porcelain enamels.
• Pigments (cobalt blue, known in ancient times, and Cobalt green).
• Battery sector (e.g. electrodes) estimated as about 11% of production in 2003.
• Steel-belted radial tires.
• Cobalt-60 has multiple uses as a gamma ray source:
  * It is used in radiotherapy.
  * It is used in radiation treatment of foods for sterilization (cold pasteurization).
  * It is used in industrial radiography to detect structural flaws in metal parts.

Cobalt compounds

Oxides

Table 1: Cobalt oxides

Formula	Color	Oxidation State	MP	Structure / comments
Co2O3		Co^3^+
Co3O4	black	Co^2+/3+	900-950decomp	normal spinel
CoO	olive green	Co^2^+	1795	NaCl -antiferromag. < 289 K

Preparations

• Co₂O₃ is formed from oxidation of Co(OH)₂.
• CoO when heated at 600-700°C converts to Co₃O₄
• Co₃O₄ when heated at 900-950°C reconverts back to CoO.

\[Co^{3+} + e^- \leftrightharpoons Co^{2+}\;\; 1.81\,V\]

\[Co^{2+} + 2e^- \leftrightharpoons Co \;\; -0.28\;V\]

no stable [Co(H₂O)₆]³⁺ or [Co(OH)₃ exist since these convert to CoO(OH).

[Co(H₂O)₆]²⁺ not acidic and a stable carbonate exists.

Cobalt Blue

One of the earliest uses of cobalt was in the coloring of glass by the addition of cobalt salts.

The cobalt blue pigment is based on the spinel CoAl₂O₄ and in the laboratory can be readily synthesized by pyrolysis of a mixture of AlCl₃ and CoCl₂.

Halides

Cobalt(II) halides

Formula	Color	MP	μ(BM)	Structure
CoF2	pink	1200	-	rutile
CoCl2	blue	724	5.47	CdCl2
CoBr2	green	678	-	CdI2
CoI2	blue-black	515	-	CdI2

Preparations:

Co or CoCO₃ + HX → CoX₂.aq → CoX₂

Cobalt complexes

The Cobalt(III) ion forms many stable complexes, which being inert, are capable of exhibiting various types of isomerism. The preparation and characterization of many of these complexes dates back to the pioneering work of Werner and his students. Coordination theory was developed on the basis of studies of complexes of the type:

Werner Complexes

[Co(NH3)6]Cl3	yellow
[CoCl(NH3)5]Cl2	red
trans-[CoCl2(NH3)4]Cl	green
cis-[CoCl2(NH3)4]Cl	purple

Another important complex in the history of coordination chemistry is hexol. This was the first complex that could be resolved into its optical isomers that did not contain carbon atoms. Since then, only three or four others have been found.

Recently a structure that Werner apparently misassigned has been determined to be related to the original hexol although in this case the complex contains 6 Co atoms, i.e. is hexanuclear. The dark green compound is not resolvable into optical isomers.

Highlight Co-O ON/OFF

Highlight Co-O ON/OFF

Werner's hexol and "2nd hexol"

A noticeable difference between chromium(III) and cobalt(III) chemistry is that cobalt complexes are much less susceptible to hydrolysis, though limited hydrolysis, leading to polynuclear cobaltammines with bridging OH- groups, is well known. Other commonly occurring bridging groups are NH₂^-, NH^2- and NO₂^-, which give rise to complexes such as the bright-blue amide bridged [(NH₃)₅Co-NH₂-Co(NH₃) ₅]⁵⁺.

In the preparation of cobalt(III) hexaammine salts by the oxidation in air of cobalt(II) in aqueous ammonia it is possible to isolate blue [(NH₃)₅Co-O₂-Co(NH₃) ₅]⁴⁺. This is moderately stable in concentrated aqueous ammonia and in the solid state but readily decomposes in acid solutions to Co(II) and O₂, while oxidizing agents such as (S₂O₈)^2- convert it to the green, paramagnetic [(NH₃)₅Co-O₂-Co(NH₃) ₅]⁵⁺ (μ₃₀₀ = 1.7 B.M.).

In the brown compound both cobalt atoms are Co(III) and are joined by a peroxo group, O₂^2-, this fits with the observed diamagnetism; in addition the stereochemistry of the central Co-O-O-Co group is similar to that of H₂O₂. The green compound is less straightforward. Werner thought that it too involved a peroxo group but in this instance bridging between Co(III) and Co(IV) atoms.

This could account for the paramagnetism, but EPR evidence shows that the 2 cobalt atoms are equivalent, and X-ray evidence shows the central Co-O-O-Co group to be planar with an O-O distance of 131 pm, which is very close to the 128 pm of the superoxide, O₂^-, ion.

A more satisfactory formulation therefore is that of 2 Co(III) atoms joined by a superoxide bridge.

A range of Co(II) dioxygen complexes are known, some of which are able to reversibly bind O₂ from the air. During WWII, some US aircraft carriers are reported to have used these complexes as a solid source for oxy-acetylene welding. By slightly warming the solid complex the oxygen was released and when cooled again oxygen would be coordinated again. Unlike an oxygen cylinder, the solid would not explode if hit by a stray bullet!

[CosalenO₂]

A laboratory experiment designed to measure the uptake of dioxygen by Cosalen is available online.

Co(acac)₃ is a green octahedral complex of Co(III). In the case of Co(II) a comparison can be made to the Ni(II) complexes.
Ni(acac)₂ is only found to be monomeric at temperatures around 200C in non-coordinating solvents such as n-decane. 6-coordinate monomeric species are formed at room temperature in solvents such as pyridine, but in the solid state Ni(acac)₂ is a trimer, where each Ni atom is 6-coordinate. Note that Co(acac)₂ actually exists as a tetramer.

[Ni(acac)₂]₃ [Co(acac)₂]₄

Cobalt(II) halide complexes with pyridine show structural isomerism. Addition of pyridine to cobalt(II) chloride in ethanol can produce blue, purple or pink complexes each having the composition "CoCl₂pyr₂". The structures are 4, 5 and 6 coordinate with either no bridging chlorides or mono- or di- bridged chlorides.

blue-[CoCl₂pyr₂] CN=4 pink-[CoCl₂pyr₂] CN=6

See the notes on isomerism for examples of Co(III) compounds that show linkage and structural isomerism.

Health

see the notes at The University of Bristol on Vitamin B12 and other Cobalt species essential for good health.

Contributors and Attributions

• Prof. Robert J. Lancashire (The Department of Chemistry, University of the West Indies)