Cobalt(II) chloride hexahydrate (CoCl2·6H2O) is a cobalt salt composed of cobalt, chlorine and six molecules of water. It is a bright blue crystalline solid that readily absorbs water from the air and forms a deep blue solution. Anhydrous cobalt chloride (CoCl2) is a cobalt salt composed of cobalt and chlorine with no molecules of water. It is a yellow-orange crystalline solid that is hygroscopic and absorbs moisture from the air.

Cobalt(II) chloride

“CoCl2” redirects here. For the compound with molecular formula COCl2, see Phosgene.
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
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.028.718 Edit this at Wikidata
EC Number
  • 231-589-4
RTECS number
  • GF9800000
UNII
UN number 3288
  • InChI=1S/2ClH.Co/h2*1H;/q;;+2/p-2 checkY
    Key: GVPFVAHMJGGAJG-UHFFFAOYSA-L checkY
  • InChI=1/2ClH.Co/h2*1H;/q;;+2/p-2
    Key: GVPFVAHMJGGAJG-NUQVWONBAU
  • Cl[Co]Cl
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)
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
GHS labelling:
GHS06: Toxic GHS08: Health hazard GHS09: Environmental hazard
NFPA 704 (fire diamond)
Health 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gas Flammability 0: Will not burn. E.g. water Instability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogen Special hazards (white): no code

NFPA 704 four-colored diamond

3
0
0
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
80 mg/kg (rat, oral)
Safety data sheet (SDS) ICSC 0783
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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Chemical compound

Cobalt(II) chloride is an inorganic compound of cobalt and chlorine, with the formula CoCl
2
. 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. The anhydrous form is a blue crystalline solid; the dihydrate is purple and the hexahydrate is pink. Commercial samples are usually the hexahydrate, which is one of the most commonly used cobalt compounds in the lab.

Properties

Anhydrous

At room temperature, anhydrous cobalt chloride has the cadmium chloride structure (CdCl
2
) (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. The vapor pressure has been reported as 7.6 mmHg at the melting point.

Solutions

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.

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.

Hydrates

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

The anhydrous salt is hygroscopic and the hexahydrate is deliquescent.[citation needed]
The dihydrate, CoCl2(H2O)2, is a coordination polymer. Each Co center is coordinated to four doubly bridging chloride ligands. The octahedron is completed by a pair of mutually trans aquo ligands.

Subunit of CoCl2(H2O)2 lattice.

Preparation

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
+ H
2
O
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.

The anhydrous compound can be prepared by heating the hydrates.

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, so that the water vapor pressure over the solid is practically zero, 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 125°C, respectively. If the partial pressure of the water vapor is in equilibrium with the solid, as in a confined but not pressurized contained, the decomposition occurs at about 115°C, 145°C, and 195°C, respectively.

Dehydration can also be effected with trimethylsilyl chloride:

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.

Reactions

In the laboratory, cobalt(II) chloride serves as a common precursor to other cobalt compounds. Generally, diluted 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

The hexahydrate and the anhydrous salt are weak Lewis acids. The adducts are usually either octahedral or tetrahedral. It forms an octahedral complex with pyridine (C
5
H
5
N
):

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:

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

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

Reduction

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)

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.

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)

Reaction of 1-norbornyllithium with the CoCl
2
·THF in pentane produces the brown, thermally stable tetrakis(1-norbornyl)cobalt(IV) — a rare example of a stable transition metal/saturated alkane compound, different products are obtained in other solvents.

Moisture indication

The deep blue colour of this moisture indicating silica gel is due to cobalt chloride. When hydrated the colour changes to a light pink/purple.

Cobalt chloride is a common visual moisture indicator due to its distinct colour change when hydrated. The colour change is from some shade of blue when dry, to a pink when hydrated, although the shade of colour depends on the substrate and concentration. It is impregnated into paper to make test strips for detecting moisture in solutions, or more slowly, in air/gas. Desiccants such as silica gel can incorporate cobalt chloride to indicate when it is “spent” (i.e. hydrated).

Health issues

Cobalt is essential for most higher forms of life, but more than a few milligrams each day is harmful. Although poisonings have rarely resulted from cobalt compounds, their chronic ingestion has caused serious health problems at doses far less than the lethal dose. In 1966, the addition of cobalt compounds to stabilize beer foam in Canada led to a peculiar form of toxin-induced cardiomyopathy, which came to be known as beer drinker’s cardiomyopathy.

Furthermore, cobalt(II) chloride is suspected of causing cancer (i.e., possibly carcinogenic, IARC Group 2B) as per the International Agency for Research on Cancer (IARC) Monographs.

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

Other uses

  • 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.[citation needed] Cobalt supplementation is not banned and therefore would not be detected by current anti-doping testing. Cobalt chloride is a banned substance under the Australian Thoroughbred Racing Board.
  • Cobalt chloride is one method used to induce pulmonary arterial hypertension in animals for research and evaluation of treatment efficacy.

References

External links

Wikimedia Commons has media related to cobalt(II) chloride.
Salts and covalent derivatives of the chloride ion
HCl He
LiCl BeCl2 BCl3
B2Cl4
CCl4
+C
NCl3
ClN3
+N
Cl2O
ClO2
Cl2O7
+O
ClF
ClF3
ClF5
Ne
NaCl MgCl2 AlCl
AlCl3
SiCl4 P2Cl4
PCl3
PCl5
+P
S2Cl2
SCl2
SCl4
Cl2 Ar
KCl CaCl
CaCl2
ScCl3 TiCl2
TiCl3
TiCl4
VCl2
VCl3
VCl4
VCl5
CrCl2
CrCl3
CrCl4
MnCl2
MnCl3
FeCl2
FeCl3
CoCl2
CoCl3
NiCl2 CuCl
CuCl2
ZnCl2 GaCl3 GeCl2
GeCl4
AsCl3
AsCl5
+As
Se2Cl2
SeCl4
BrCl Kr
RbCl SrCl2 YCl3 ZrCl3
ZrCl4
NbCl3
NbCl4
NbCl5
MoCl2
MoCl3
MoCl4
MoCl5
MoCl6
TcCl3
TcCl4
RuCl3 RhCl3 PdCl2 AgCl CdCl2 InCl
InCl2
InCl3
SnCl2
SnCl4
SbCl3
SbCl5
Te3Cl2
TeCl4
ICl
ICl3
XeCl
XeCl2
XeCl4
CsCl BaCl2 * LuCl3 HfCl4 TaCl5 WCl2
WCl3
WCl4
WCl5
WCl6
ReCl3
ReCl4
ReCl5
ReCl6
OsCl4 IrCl2
IrCl3
IrCl4
PtCl2
PtCl4
AuCl
AuCl3
Hg2Cl2,
HgCl2
TlCl PbCl2,
PbCl4
BiCl3 PoCl2,
PoCl4
AtCl Rn
FrCl RaCl2 ** LrCl3 RfCl4 DbCl5 SgO2Cl2 BhO3Cl Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og
 
* LaCl3 CeCl3 PrCl3 NdCl2,
NdCl3
PmCl3 SmCl2,
SmCl3
EuCl2,
EuCl3
GdCl3 TbCl3 DyCl2,
DyCl3
HoCl3 ErCl3 TmCl2
TmCl3
YbCl2
YbCl3
** AcCl3 ThCl4 PaCl4
PaCl5
UCl3
UCl4
UCl5
UCl6
NpCl3 PuCl3 AmCl2
AmCl3
CmCl3 BkCl3 CfCl3 EsCl3 FmCl2 Md NoCl2


Source: Cobalt(II) chloride
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