Cobalt Chemistry

History

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.

An excellent site for finding the properties of the elements, including cobalt is at Start WebElements

Further information on Cobalt can be found at the Cobalt Institute Ltd and the page on Cobalt at Wikipedia.

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 The dominance of African copper-cobalt producers has been replaced by a more even spread of output between leading producing countries, with Canada, Norway and more recently Australia, together with exports from Russia, replacing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.

An article on The High Human Cost of Cobalt Mining may be of interest.

The International Centre for Environmental and Nuclear Sciences (ICENS) has an on-going programme of mapping the geochemical content of Jamaica. 'A Geochemical Atlas of Jamaica' was published in 1995 and is available from Amazon or ICENS.
The results found for Cobalt are shown below (courtesy of Prof G.C. Lalor).
ICENS Co data

Extraction
Not covered in this course.

Uses

Cobalt compounds

Oxides

Cobalt oxides
Formula Colour Oxidation State MP Structure / comments
Co2O3 Co3+
Co3O4 black Co2+/3+ 900-950decomp normal spinel
CoO olive green Co2+ 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.

Co3+ + e- ⇔ Co2+ 1.81V
Co2+ + 2e- ⇔ Co  -0.28V

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 colouring of glass by the addition of cobalt salts.
Murano blue miniature
The pigment is based on the spinel CoAl2O4 and in the laboratory can be readily synthesised by pyrolysis of a mixture of AlCl3 and CoCl2.

Halides
Cobalt(II) halides
Formula Colour 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 characterisation 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.


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-, NH2- 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, O22-, 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.
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The Department of Chemistry, University of the West Indies,
Mona Campus, Kingston 7, Jamaica.

Created July 2002. Links checked and/or last modified 28th September 2020.
URL http://wwwchem.uwimona.edu.jm/courses/cobalt.html