Cobalt(II) chloride

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Cobalt(II) chloride
Cobaltous chloride anhydrous.jpg
Anhydrous
Cobaltous chloride.jpg
Hexahydrate
Cobalt(II)-chloride-3D-balls.png
Structure of anhydrous compound
MCl2(aq)6forFeCoNi.png
Structure of hexahydrate
Names
IUPAC name
Cobalt(II) chloride
Other names
Cobaltous chloride
Cobalt dichloride
Muriate of cobalt[1]
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.028.718
EC Number 231-589-4
RTECS number GF9800000
UNII
UN number 3288
Properties
CoCl2
Molar mass 129.839 g/mol (anhydrous)
165.87 g/mol (dihydrate)
237.93 g/mol (hexahydrate)
Appearance blue crystals (anhydrous)
violet-blue (dihydrate)
rose red crystals (hexahydrate)
Density 3.356 g/cm3 (anhydrous)
2.477 g/cm3 (dihydrate)
1.924 g/cm3 (hexahydrate)
Melting point 726 °C (1,339 °F; 999 K) ±2 (anhydrous)[2]
140 °C (monohydrate)
100 °C (dihydrate)
86 °C (hexahydrate)
Boiling point 1,049 °C (1,920 °F; 1,322 K)
43.6 g/100 mL (0 °C)
45 g/100 mL (7 °C)
52.9 g/100 mL (20 °C)
105 g/100 mL (96 °C)
Solubility 38.5 g/100 mL (methanol)
8.6 g/100 mL (acetone)
soluble in ethanol, pyridine, glycerol
+12,660·10−6 cm3/mol
Structure
CdCl2 structure
hexagonal (anhydrous)
monoclinic (dihydrate)
Octahedral (hexahydrate)
Hazards
Safety data sheet ICSC 0783
GHS pictograms GHS06: Toxic GHS08: Health hazard GHS09: Environmental hazard
NFPA 704
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
80 mg/kg (rat, oral)
Related compounds
Other anions
Cobalt(II) fluoride
Cobalt(II) bromide
Cobalt(II) iodide
Other cations
Rhodium(III) chloride
Iridium(III) chloride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

Cobalt(II) chloride is an inorganic compound of cobalt and chlorine, with the formula CoCl
2
. It is a sky blue crystalline solid.

The compound forms several hydrates CoCl
2
nH
2
O
, for n = 1, 2, 6, and 9. Claims of the formation of tri- and tetrahydrates have not been confirmed.[4] The dihydrate is purple and hexahydrate is pink. It is usually supplied as the hexahydrate CoCl
2
·6H
2
O
, which is one of the most commonly used cobalt compounds in the lab.[5]

Because of the ease of the hydration/dehydration reaction, and the resulting color change, cobalt chloride is used as an indicator for water in desiccants.

Niche uses of cobalt chloride include its role in organic synthesis and electroplating objects with cobalt metal.

Cobalt chloride has been classified as a substance of very high concern by the European Chemicals Agency as it is a suspected carcinogen.

Properties[edit]

Anhydrous[edit]

At room temperature, anhydrous cobalt chloride has the CdCl
2
structure (R3m) in which the cobalt(II) ions are octahedrally coordinated. At about 706 C (20 degrees below the melting point), the coordination is believed to change to tetrahedral.[2] The vapor pressure has been reported as 7.6 mm Hg at the melting point.[6]

Solutions[edit]

Cobalt chloride is fairly soluble in water. Under atmospheric pressure, the mass concentration of a saturated solution of CoCl
2
in water is about 54% at the boiling point, 120.2 °C; 48% at 51.25 °C; 35% at 25 °C; 33% at 0 °C; and 29% at −27.8 °C.[4]

Diluted aqueous solutions of CoCl
2
contain the species [Co(H
2
O)
6
]2+
, besides chloride ions. Concentrated solutions are red at room temperature but become blue at higher temperatures.[7]

Hydrates[edit]

The crystal unit of the solid hexahydrate CoCl
2
•6H
2
O
contains the neutral molecule trans-CoCl
2
(H
2
O)
4
and two molecules of water of crystallization.[8] This species dissolves readily in water and alcohol.

The anhydrous salt is hygroscopic and the hexahydrate is deliquescent.[citation needed]

Preparation[edit]

Cobalt chloride can be prepared in aqueous solution from cobalt(II) hydroxide or cobalt(II) carbonate and hydrochloric acid:

CoCO
3
+ 2 HCl(aq)CoCl
2
(aq) + CO
2
Co(OH)
2
+ 2 HCl(aq)CoCl
2
(aq) + 2H
2
O

The solid dihydrate and hexahydrate can be obtained by evaporation. Cooling saturated aqueous solutions yields the dihydrate between 120.2 °C and 51.25 °C, and the hexahydrate below 51.25 °C. Water ice, rather than cobalt chloride, will crystallize from solutions with concentration below 29%. The monohydrate and the anhydrous forms can be obtained by cooling solutions only under high pressure, above 206 °C and 335 °C, respectively.[4]

The anhydrous compound can be prepared by heating the hydrates.[9] On rapid heating or in a closed container, each of the 6-, 2-, and 1- hydrates partially melts into a mixture of the next lower hydrate and a saturated solution -- at 51.25 °C, 206 °C, and 335 °C, respectively. On slow heating in an open container, water evaporates out of each of the solid 6-, 2-, and 1- hydrates, leaving the next lower hydrate -- at about 40 °C, 89 °C, and 126 °C, respectively.[4]

Dehydration can also be effected with trimethylsilyl chloride:[10]

CoCl
2
•6H
2
O
+ 12 (CH
3
)
3
SiCl
CoCl
2
+ 6[(CH
3
)
3
SiCl]
2
O
+ 12 HCl

The anhydrous compound can be purified by sublimation in vacuum.[2]

Reactions[edit]

In the laboratory, cobalt(II) chloride serves as a common precursor to other cobalt compounds. Generally, aqueous solutions of the salt behave like other cobalt(II) salts since these solutions consist of the [Co(H
2
O)
6
]2+
ion regardless of the anion. For example, such solutions give a precipitate of Cobalt sulfide CoS upon treatment with hydrogen sulfide H
2
S
.

Complexed chlorides[edit]

The hexahydrate and the anhydrous salt are weak Lewis acids. The adducts are usually either octahedral or tetrahedral. With pyridine (C
5
H
5
N
), one obtains an octahedral complex:

CoCl
2
·6H
2
O
+ 4 C
5
H
5
N
CoCl
2
(C
5
H
5
N)
4
+ 6 H
2
O

With triphenylphosphine (P(C
6
H
5
)
3
), a tetrahedral complex results:

CoCl
2
·6H
2
O
+ 2 P(C
6
H
5
)
3
CoCl
2
[P(C
6
H
5
)
3
]
2
+ 6 H
2
O

Salts of the anionic complex CoCl42− can be prepared using tetraethylammonium chloride:[11]

CoCl
2
+ 2 [(C2H5)4N]Cl → [(C2H5)4N)]2[CoCl4]

The [CoCl4]2− ion is the blue ion that forms upon addition of hydrochloric acid to aqueous solutions of hydrated cobalt chloride, which are pink.

Reduction[edit]

The structure of a cobalt(IV) coordination complex with the norbornyl anion

Reaction of the anhydrous compound with sodium cyclopentadienide gives cobaltocene Co(C
5
H
5
)
2
. This 19-electron species is a good reducing agent, being readily oxidised to the yellow 18-electron cobaltacenium cation [Co(C
5
H
5
)
2
]+
.

Oxidation to cobalt(III)[edit]

Compounds of cobalt in the +3 oxidation state exist, such as cobalt(III) fluoride CoF
3
, nitrate Co(NO
3
)
3
, and sulfate Co
2
(SO
4
)
3
; however, cobalt(III) chloride CoCl
3
is not stable in normal conditions, and would decompose immediately into CoCl
2
and chlorine.[12]

On the other hand, cobalt(III) chlorides can be obtained if the cobalt is bound also to other ligands of greater Lewis basicity than chloride, such as amines. For example, in the presence of ammonia, cobalt(II) chloride is readily oxidised by atmospheric oxygen to hexamminecobalt(III) chloride:

4 CoCl
2
·6H
2
O
+ 4 NH
4
Cl + 20 NH
3
+ O
2
→ 4 [Co(NH
3
)
6
]Cl
3
+ 26 H
2
O

Similar reactions occur with other amines. These reactions are often performed in the presence of charcoal as a catalyst, or with hydrogen peroxide H
2
O
2
substituted for atmospheric oxygen. Other highly basic ligands, including carbonate, acetylacetonate, and oxalate, induce the formation of Co(III) derivatives. Simple carboxylates and halides do not.

Unlike Co(II) complexes, Co(III) complexes are very slow to exchange ligands, so they are said to be kinetically inert. The German chemist Alfred Werner was awarded the Nobel prize in 1913 for his studies on a series of these cobalt(III) compounds, work that led to an understanding of the structures of such coordination compounds.

Oxidation to cobalt(IV)[edit]

Reaction of 1-norbonyllithium with the CoCl
2
·THF in pentane produces the brown, thermally stable cobalt(IV) tetralkyl[13][14] — a rare example of a stable transition metal/saturated alkane compound,[5] different products are obtained in other solvents.[15]

Safety[edit]

In 2005–06, cobalt chloride was the eighth-most-prevalent allergen in patch tests (8.4%).[16]

Other uses[edit]

  • Invisible ink: when suspended in solution, cobalt(II) chloride can be made to appear invisible on a surface; when that same surface is subsequently exposed to significant heat (such as from a handheld heat gun or lighter) the ink permanently/ irreversibly changes to blue.
  • Cobalt chloride is an established chemical inducer of hypoxia-like responses such as erythropoiesis.[17] Cobalt supplementation is not banned and therefore would not be detected by current anti-doping testing.[18] Cobalt chloride is a banned substance under the Australian Thoroughbred Racing Board.[19]

References[edit]

  1. ^ "Cobalt muriate, CAS Number: 7646-79-9". www.chemindustry.com. Retrieved 19 April 2018.
  2. ^ a b c Wojakowska, A., Krzyżak, E. and Plińska, S. (2007): "Melting and high-temperature solid state transitions in cobalt(II) halides". Journal of Thermal Analysis and Calorimetry, volume 88, issue 2, pages 525-530. doi:10.1007/s10973-006-8000-9
  3. ^ Santa Cruz Biotechnology: Cobalt(II) chloride
  4. ^ a b c d M. T. Saugier, M. Noailly, R. Cohen-Adad, F. Paulik, and J. Paulik (1977): "Equilibres solide ⇄ liquide ⇆ vapeur du systeme binaire CoCl
    2
    -H
    2
    O
    " Journal of Thermal Analysis, volume 11, issue 1, pages 87–100. doi:10.1007/BF02104087 Note: the lowest point of fig.6 is inconsistent with fig.7; probably should be at -27.8 C instead of 0 C.
  5. ^ a b Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  6. ^ Yuzo Saeki, Ryoko Matsuzaki, Naomi Aoyama (1977): "The vapor pressure of cobalt dichloride". Journal of the Less Common Metals, volume 55, issue 2, pages 289-291. doi:10.1016/0022-5088(77)90204-1
  7. ^ The Merck Index, 7th edition, Merck & Co, Rahway, New Jersey, USA, 1960.
  8. ^ Wells, A. F. (1984), Structural Inorganic Chemistry (5th ed.), Oxford: Clarendon Press, ISBN 0-19-855370-6
  9. ^ John Dallas Donaldson, Detmar Beyersmann, "Cobalt and Cobalt Compounds" in Ullmann's Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim, 2005. doi:10.1002/14356007.a07_281.pub2
  10. ^ Philip Boudjouk, Jeung-Ho So (1992). "Solvated and Unsolvated Anhydrous Metal Chlorides from Metal Chloride Hydrates". Inorg. Synth. 29: 108–111. doi:10.1002/9780470132609.ch26.CS1 maint: Uses authors parameter (link)
  11. ^ Gill, N. S. & Taylor, F. B. (1967). "Tetrahalo Complexes of Dipositive Metals in the First Transition Series". Inorg. Synth. 9: 136–142. doi:10.1002/9780470132401.ch37.
  12. ^ Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
  13. ^ Barton K. Bower & Howard G. Tennent (1972). "Transition metal bicyclo[2.2.1]hept-1-yls". J. Am. Chem. Soc. 94 (7): 2512–2514. doi:10.1021/ja00762a056.
  14. ^ Erin K. Byrne; Darrin S. Richeson & Klaus H. Theopold (1986). "Tetrakis(1-norbornyl)cobalt, a low spin tetrahedral complex of a first row transition metal". J. Chem. Soc., Chem. Commun. (19): 1491–1492. doi:10.1039/C39860001491.
  15. ^ Erin K. Byrne; Klaus H. Theopold (1989). "Synthesis, characterization, and electron-transfer reactivity of norbornyl complexes of cobalt in unusually high oxidation states". J. Am. Chem. Soc. 111 (11): 3887–3896. doi:10.1021/ja00193a021.
  16. ^ Zug KA, Warshaw EM, Fowler JF Jr, Maibach HI, Belsito DL, Pratt MD, Sasseville D, Storrs FJ, Taylor JS, Mathias CG, Deleo VA, Rietschel RL, Marks J. Patch-test results of the North American Contact Dermatitis Group 2005–2006. Dermatitis. 2009 May–Jun;20(3):149-60.
  17. ^ W. Jelkmann: The disparate roles of cobalt in erythropoiesis, and doping relevance. Open Journal of Hematology, 2012, 3–6. http://rossscience.org/ojhmt/2075-907X-3-6.php
  18. ^ Lippi G, Franchini M, Guidi GC (November 2005). "Cobalt chloride administration in athletes: a new perspective in blood doping?". Br J Sports Med. 39 (11): 872–3. doi:10.1136/bjsm.2005.019232. PMC 1725077. PMID 16244201.
  19. ^ Bartley, Patrick (6 February 2015). "Cobalt crisis turns the eyes of the world onto Australian racing". The Sydney Morning Herald.

External links[edit]