Cobalt(II) chloride

Last updated
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 OOjs UI icon edit-ltr-progressive.svg
EC Number
  • 231-589-4
PubChem CID
RTECS number
  • GF9800000
UNII
UN number 3288
  • InChI=1S/2ClH.Co/h2*1H;/q;;+2/p-2 Yes check.svgY
    Key: GVPFVAHMJGGAJG-UHFFFAOYSA-L Yes check.svgY
  • 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)
Appearanceblue 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
GHS labelling:
GHS-pictogram-skull.svg GHS-pictogram-silhouette.svg GHS-pictogram-pollu.svg
NFPA 704 (fire diamond)
[3]
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).
X mark.svgN  verify  (what is  Yes check.svgYX mark.svgN ?)

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. [4] 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. [5]

Contents

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. [2] The vapor pressure has been reported as 7.6  mmHg at the melting point. [6]

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. [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

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 ] 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. [9]

Subunit of CoCl2(H2O)2 lattice. MX2(H2O)2.png
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. [4]

The anhydrous compound can be prepared by heating the hydrates. [10]

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. [4] 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. [4]

Dehydration can also be effected with trimethylsilyl chloride: [11]

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

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
): [12]

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: [13]

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 Co(norbornyl)4.png
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. [14]

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 cobalt(IV) tetralkyl [15] [16] — a rare example of a stable transition metal/saturated alkane compound, [5] different products are obtained in other solvents. [17]

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. Blue silica gel.jpg
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). [18]

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. [19] [20] [21]

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. [22]

In 2005–06, cobalt chloride was the eighth-most-prevalent allergen in patch tests (8.4%). [23]

Other uses

Related Research Articles

<span class="mw-page-title-main">Iron(III) chloride</span> Inorganic compound

Iron(III) chloride is the inorganic compound with the formula FeCl3. Also called ferric chloride, it is a common compound of iron in the +3 oxidation state. The anhydrous compound is a crystalline solid with a melting point of 307.6 °C. The color depends on the viewing angle: by reflected light the crystals appear dark green, but by transmitted light they appear purple-red.

<span class="mw-page-title-main">Copper(II) nitrate</span> Chemical compound

Copper(II) nitrate describes any member of the family of inorganic compounds with the formula Cu(NO3)2(H2O)x. The hydrates are blue solids. Anhydrous copper nitrate forms blue-green crystals and sublimes in a vacuum at 150-200 °C. Common hydrates are the hemipentahydrate and trihydrate.

<span class="mw-page-title-main">Zinc chloride</span> Chemical compound

Zinc chloride is the name of inorganic chemical compounds with the formula ZnCl2 and its hydrates. Zinc chlorides, of which nine crystalline forms are known, are colorless or white, and are highly soluble in water. This salt is hygroscopic and even deliquescent. Zinc chloride finds wide application in textile processing, metallurgical fluxes, and chemical synthesis. No mineral with this chemical composition is known aside from the very rare mineral simonkolleite, Zn5(OH)8Cl2·H2O.

In chemistry, water(s) of crystallization or water(s) of hydration are water molecules that are present inside crystals. Water is often incorporated in the formation of crystals from aqueous solutions. In some contexts, water of crystallization is the total mass of water in a substance at a given temperature and is mostly present in a definite (stoichiometric) ratio. Classically, "water of crystallization" refers to water that is found in the crystalline framework of a metal complex or a salt, which is not directly bonded to the metal cation.

<span class="mw-page-title-main">Europium(III) chloride</span> Chemical compound

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<span class="mw-page-title-main">Aluminium chloride</span> Chemical compound

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<span class="mw-page-title-main">Manganese(II) chloride</span> Chemical compound

Manganese(II) chloride is the dichloride salt of manganese, MnCl2. This inorganic chemical exists in the anhydrous form, as well as the dihydrate (MnCl2·2H2O) and tetrahydrate (MnCl2·4H2O), with the tetrahydrate being the most common form. Like many Mn(II) species, these salts are pink, with the paleness of the color being characteristic of transition metal complexes with high spin d5 configurations.

<span class="mw-page-title-main">Dysprosium(III) chloride</span> Chemical compound

Dysprosium(III) chloride (DyCl3), also known as dysprosium trichloride, is a compound of dysprosium and chlorine. It is a white to yellow solid which rapidly absorbs water on exposure to moist air to form a hexahydrate, DyCl3·6H2O. Simple rapid heating of the hydrate causes partial hydrolysis to an oxychloride, DyOCl.

<span class="mw-page-title-main">Copper(II) chloride</span> Chemical compound

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<span class="mw-page-title-main">Chromium(III) chloride</span> Chemical compound

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<span class="mw-page-title-main">Nickel(II) chloride</span> Chemical compound

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<span class="mw-page-title-main">Iron(II) chloride</span> Chemical compound

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<span class="mw-page-title-main">Gadolinium(III) chloride</span> Chemical compound

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<span class="mw-page-title-main">Cobalt(II) fluoride</span> Chemical compound

Cobalt(II) fluoride is a chemical compound with the formula (CoF2). It is a pink crystalline solid compound which is antiferromagnetic at low temperatures (TN=37.7 K) The formula is given for both the red tetragonal crystal, (CoF2), and the tetrahydrate red orthogonal crystal, (CoF2·4H2O). CoF2 is used in oxygen-sensitive fields, namely metal production. In low concentrations, it has public health uses. CoF2 is sparingly soluble in water. The compound can be dissolved in warm mineral acid, and will decompose in boiling water. Yet the hydrate is water-soluble, especially the di-hydrate CoF2·2H2 O and tri-hydrate CoF2·3H2O forms of the compound. The hydrate will also decompose with heat.

<span class="mw-page-title-main">Cobalt(II) nitrate</span> Chemical compound

Cobalt nitrate is the inorganic compound with the formula Co(NO3)2.xH2O. It is cobalt(II)'s salt. The most common form is the hexahydrate Co(NO3)2·6H2O, which is a red-brown deliquescent salt that is soluble in water and other polar solvents.

<span class="mw-page-title-main">Cobalt(II) bromide</span> Chemical compound

Cobalt(II) bromide (CoBr2) is an inorganic compound. In its anhydrous form, it is a green solid that is soluble in water, used primarily as a catalyst in some processes.

<span class="mw-page-title-main">Nickel(II) bromide</span> Chemical compound

Nickel(II) bromide is the name for the inorganic compounds with the chemical formula NiBr2(H2O)x. The value of x can be 0 for the anhydrous material, as well as 2, 3, or 6 for the three known hydrate forms. The anhydrous material is a yellow-brown solid which dissolves in water to give blue-green hexahydrate (see picture).

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Nickel(II) perchlorate is a inorganic compound with the chemical formula of Ni(ClO4)2, and it is a strong oxidizing agent. Its colours are different depending on water. For example, the hydrate forms cyan crystals, the pentahydrate forms green crystals, but the hexahydrate (Ni(ClO4)2·6H2O) forms blue crystals.

Cobalt compounds are chemical compounds formed by cobalt with other elements. In the compound, the most stable oxidation state of cobalt is the +2 oxidation state, and in the presence of specific ligands, there are also stable compounds with +3 valence. In addition, there are cobalt compounds in high oxidation states +4, +5 and low oxidation states -1, 0, +1.

References

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    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.
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  21. 11.1.5 The unusual type of myocardiopathy recognized in 1965 and 1966 in Quebec (Canada), Minneapolis (Minnesota), Leuven (Belgium), and Omaha (Nebraska) was associated with episodes of acute heart failure (e/g/, 50 deaths among 112 beer drinkers).
  22. [PDF
  23. 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.
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  25. Bartley, Patrick (6 February 2015). "Cobalt crisis turns the eyes of the world onto Australian racing". The Sydney Morning Herald.