Many halides of cobalt(II) are known.e
cobalt(II) fluoride (CoF2) which is a pink solid,
cobalt(II) chloride (CoCl2) which is a blue solid,
cobalt(II) bromide (CoBr2) which is a green solid, and
cobalt(II) iodide (CoI2) which is a blue-black solid. In addition to the anhydrous forms, these cobalt halides also have hydrates. Anhydrous cobalt(II) chloride is blue, while the hexahydrate is red-purple.[1] Because the color change of cobalt(II) chloride in different hydrates, it can be used to manufacture color-changing
silica gel.
Anhydrous cobalt halides react with
nitric oxide at 70~120 °C to generate [Co(NO)2X]2 (X = Cl, Br or I). The complex of cobalt halides and
triethylphosphine ((C2H5)3P) can absorb nitric monoxide in
benzene to form the diamagnetic material Co(NO)X2(P(C2H5)3)[2]
In the reaction Co3+ + e− → Co2+ , the potential is +1.92 V, which is higher than that of
Cl2 to
Cl− (+1.36 V). Therefore, the interaction of Co3+ with Cl− produces Co2+ and releases chlorine gas. The potential from
F2 to
F− is as high as +2.87 V, and
cobalt(III) fluoride (CoF3) can exist stably. It is a fluorinated reagent and reacts violently with water.[3]
Oxides and hydroxides
Cobalt(II,III) oxide
Cobalt can form various oxides, such as
CoO,
Co2O3 and
Co3O4. Co3O4, at 950 °C, decomposes to CoO.[4]
Cobalt(II) hydroxide can be oxidized to the Co(III) compound CoO(OH) under
alkaline conditions.
Pnictogenides
Cobalt(II) nitrate hexahydrate
Cobalt powder reacts with
ammonia to form two kinds of nitrides, Co2N and Co3N. Cobalt reacts with
phosphorus or
arsenic to form Co2P, CoP,[2] CoP2,[6] CoAs2 and other substances.[2] The former three compounds are of interest as catalysts for
water electrolysis.[6][7][8]
Cobalt(II) azide (Co(N3)2) is another binary compound of cobalt and nitrogen that can explode when heated. Cobalt(II) and
azide can form Co(N 3)2− 4 complexes.[9] Cobalt pentazolide Co(N5)2 was discovered in 2017, and it exists in the form of the hydrate [Co(H2O)4(N5)2]·4H2O. It decomposes at 50~145 °C to form cobalt(II) azide, becoming anhydrous and releasing nitrogen, and exploding when heated further. This compound can be obtained by reacting (N5)6(H3O)3(NH4)4Cl[10] or Na(H2O)(N5)]·2H2O[11] and
[Co(H2O)6](NO3)2 at room temperature.
Hydrogen bonding of water stabilizes this molecule.[11]
Cobalt can easily react with
nitric acid to form
cobalt(II) nitrate Co(NO3)2. Cobalt(II) nitrate exists in the anhydrous form and the hydrate form, of which the hexahydrate is the most common. Cobalt nitrate hexahydrate (Co(NO3)2·6H2O) is a red deliquescence crystal that is easily soluble in water,[12] and its molecule contains cobalt(II) hydrated ions ([Co(H2O)62+) and free nitrate ions.[13] It can be obtained by precipitation from solution.
Coordination compounds
As for all metals, molecular compounds and polyatomic ions of cobalt are classified as
coordination complexes, that is, molecules or ions that contain cobalt linked to one or more
ligands. These can be combinations of a potentially infinite variety of molecules and ions, such as:
waterH 2O, as in the cation hexaaquocobalt(II) [Co(H 2O) 62+ . This pink-colored complex is the predominant cation in solid
cobalt sulfateCoSO 4·(H 2O)x, with x = 6 or 7, as well as in water solutions thereof.
ammoniaNH 3, as in cis-diaquotetraamminecobalt(III) [Co(NH 3) 4(H 2O) 23+ , in
hexol[Co(Co(NH 3) 4(HO) 2) 36− , in [Co(NO 2) 4(NH 3) 2− (the anion of Erdmann's salt),[14] and in [Co(NH 3) 5(CO 3)]− .[14]
carbonate[CO 32− , as in the green triscarbonatocobaltate(III) [Co(CO 3) 33− anion.[15][14][16]
chloride[Cl]− , as in tetrachloridocobaltate(II) CoCl 42− .
bicarbonate[HCO 3− , as in [Co(CO 3) 2(HCO 3)(H 2O)]3− .[14]
oxalate[C 2O 42− , as in trisoxalatocobaltate(III) [Co(C 2O 4)3− 3].[14]
These attached groups affect the stability of
oxidation states of the cobalt atoms, according to general principles of
electronegativity and of the
hardness–softness. For example, Co3+ complexes tend to have
ammine ligands. Because phosphorus is softer than nitrogen, phosphine ligands tend to feature the
softer Co2+ and Co+, an example being tris(triphenylphosphine)cobalt(I) chloride (P(C 6H 5) 3) 3CoCl). The more electronegative (and harder) oxide and fluoride can stabilize Co4+ and Co5+ derivatives, e.g.
caesium hexafluorocobaltate(IV) (Cs2CoF6) and potassium
percobaltate (K3CoO4).[17]
Vitamin B12 is a cobalt-centered organic biomolecule, soluble in water, and involved in the methylation and synthesis of
nucleic acid and
neurotransmitter.[20] The main source is the offal or meat of herbivorous animals.[21]
Cobaltocene (Co(C5H5)2) is a
cyclopentadiene complex of
cobalt. It has 19 valence electrons and is easily oxidized to Co(C 5H 5)+ 2 with a stable structure of 18 electrons by reaction.[24] It is a
structural analog to
ferrocene, with cobalt in place of iron. Cobaltocene is much more sensitive to
oxidation than ferrocene.[25]
^Holleman, A. F.; Wiberg, E.; Wiberg, N. (2007). "Cobalt". Lehrbuch der Anorganischen Chemie (in German) (102nd ed.). de Gruyter. pp. 1146–1152. ISBN 978-3-11-017770-1.
^US 4389339Archived 2019-07-01 at the
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^O. Glemser "Cobalt(II) Hydroxide" in Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 1521.
^
abJianmei Wang; Zhen Liu; Yiwei Zheng; Liang Cui; Wenrong Yang; Jingquan Liu (19 October 2017) [22 September 2017]. "Recent advances in cobalt phosphide based materials for energy-related applications". Journal of Materials Chemistry A. 5 (44). Royal Society of Chemistry: 22913–22932.
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^Popczun, Eric J.; Read, Carlos G.; Roske, Christopher W.; Lewis, Nathan S.; Schaak, Raymond E. (11 April 2014) [May 19, 2014]. "Highly Active Electrocatalysis of the Hydrogen Evolution Reaction by Cobalt Phosphide Nanoparticles". Angewandte Chemie International Edition. 53 (21): 5427–5430.
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^Senise, Paschoal (1959). "On the Reaction between Cobalt(II) and Azide Ions in Aqueous and Aqueous-organic Solutions1". Journal of the American Chemical Society. 81 (16): 4196–4199.
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^Zhang, Chong; Yang, Chen; Hu, Bingcheng; Yu, Chuanming; Zheng, Zhansheng; Sun, Chengguo (2017). "A Symmetric Co(N5 )2 (H2 O)4 ⋅4 H2 O High-Nitrogen Compound Formed by Cobalt(II) Cation Trapping of a Cyclo-N5− Anion". Angewandte Chemie International Edition. 56 (16): 4512–4514.
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^Prelesnik, P. V.; Gabela, F.; Ribar, B.; Krstanovic, I. (1973). "Hexaaquacobalt(II) nitrate". Cryst. Struct. Commun. 2 (4): 581–583.
^
abcdefMcCutcheon, Thomas P.; Schuele, William J. (1953). "Complex Acids of Cobalt and Chromium. The Green Carbonatocobalt(III) Anion*". Journal of the American Chemical Society. 75 (8): 1845–1846.
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^Bauer, H. F.; Drinkard, W. C. (1960). "A General Synthesis of Cobalt(III) Complexes; A New Intermediate, Na3[Co(CO3)3]·3H2O". Journal of the American Chemical Society. 82 (19): 5031–5032.
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^Tafesse, Fikru; Aphane, Elias; Mongadi, Elizabeth (2010). "Determination of the structural formula of sodium tris-carbonatocobaltate(III), Na3[Co(CO3)3]·3H2O by thermogravimetry". Journal of Thermal Analysis and Calorimetry. 102: 91–97.
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^Sweany, Ray L.; Brown, Theodore L. (1977). "Infrared spectra of matrix-isolated dicobalt octacarbonyl. Evidence for the third isomer". Inorganic Chemistry. 16 (2): 415–421.
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^Connelly, Neil G.; Geiger, William E. (1996). "Chemical Redox Agents for Organometallic Chemistry". Chemical Reviews. 96 (2): 877–910.
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