Nickel(II) chloride

Nickel(II) chloride

Hexahydrate
Anhydrous
Names
IUPAC name Nickel(II) chloride
Other names Nickelous chloride, nickel(II) salt of hydrochloric acid
Identifiers
CAS Number	7718-54-9 7791-20-0 (hexahydrate)
ChEBI	CHEBI:34887
ChemSpider	22796
EC Number	231-743-0
Jmol 3D model	Interactive image
KEGG	C14711
PubChem	24385
RTECS number	QR6480000
InChI InChI=1S/2ClH.Ni/h2*1H;/q;;+2/p-2 Key: QMMRZOWCJAIUJA-UHFFFAOYSA-L InChI=1/2ClH.Ni/h2*1H;/q;;+2/p-2 Key: QMMRZOWCJAIUJA-NUQVWONBAR
SMILES Cl[Ni]Cl
Properties
Chemical formula	NiCl2
Molar mass	129.5994 g/mol (anhydrous) 237.69 g/mol (hexahydrate)
Appearance	yellow-brown crystals deliquescent (anhydrous) green crystals (hexahydrate)
Odor	odorless
Density	3.55 g/cm^3 (anhydrous) 1.92 g/cm^3 (hexahydrate)
Melting point	1,001 °C (1,834 °F; 1,274 K) (anhydrous) 140 °C (hexahydrate)
Solubility in water	anhydrous 67.525 g/100 mL (20 °C) [1] 87.6 g/100 mL (100 °C) hexahydrate 67.525 g/100 mL (20 °C) [1] 87.6 g/100 mL (100 °C)
Solubility	0.8 g/100 mL (hydrazine) soluble in ethylene glycol, ethanol, ammonium hydroxide insoluble in ammonia, nitric acid
Acidity (pKa)	4 (hexahydrate)
Structure
Crystal structure	Monoclinic
Coordination geometry	octahedral at Ni
Thermochemistry
Std molar entropy (S^o298)	107 J·mol^−1·K^−1[2]
Std enthalpy of formation (ΔfH^o298)	−316 kJ·mol^−1[2]
Hazards
Safety data sheet	Fischer Scientific
EU classification (DSD)	Carc. Cat. 1 Muta. Cat. 3 Repr. Cat. 2 Toxic (T) Irritant (Xi) Dangerous for the environment (N)
R-phrases	R49, R61, R23/25, R38, R42/43, R48/23, R68, R50/53
S-phrases	S53, S45, S60, S61
NFPA 704	0 2 0
Flash point	Non-flammable
Lethal dose or concentration (LD, LC):
LD50 (Median dose)	105 mg/kg (rat, oral)[3]
Related compounds
Other anions	Nickel(II) fluoride Nickel(II) bromide Nickel(II) iodide
Other cations	Palladium(II) chloride Platinum(II) chloride Platinum(II,IV) chloride Platinum(IV) chloride
Related compounds	Cobalt(II) chloride Copper(II) 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

Nickel(II) chloride (or just nickel chloride), is the chemical compound NiCl₂. The anhydrous salt is yellow, but the more familiar hydrate NiCl₂·6H₂O is green. In general nickel(II) chloride, in various forms, is the most important source of nickel for chemical synthesis. The nickel chlorides are deliquescent, absorbing moisture from the air to form a solution. Nickel salts are carcinogenic.

Production and syntheses

The largest scale production of nickel chloride involves the extraction with hydrochloric acid of nickel matte and residues obtained from roasting refining nickel-containing ores.

NiCl₂·6H₂O is prepared in the laboratory because it is inexpensive and has a long shelf-life. Heating the hexahydrate in the range 66-133.°C gives the yellowish dihydrate, NiCl₂·2H₂O.^[4] The hydrates convert to the anhydrous form upon heating in thionyl chloride or by heating under a stream of HCl gas. Simply heating the hydrates does not afford the anhydrous dichloride.

$\mathsf{NiCl_2 \cdot 6H_2O + 6SOCl_2 \ \xrightarrow \ NiCl_2 + 6SO_2 + 12HCl }$

The dehydration is accompanied by a color change from green to yellow.^[5]

In case one needs a pure compound without presence of cobalt, nickel chloride can be obtained cautiously heating hexammine nickel chloride:^[6]

$\mathsf{[Ni(NH_3)_6]Cl_2 \ \xrightarrow{175-200^oC}\ NiCl_2 + 6NH_3 }$

Structure of NiCl₂ and its hydrates

NiCl₂ adopts the CdCl₂ structure.^[7] In this motif, each Ni²⁺ center is coordinated to six Cl⁻ centers, and each chloride is bonded to three Ni(II) centers. In NiCl₂ the Ni-Cl bonds have "ionic character". Yellow NiBr₂ and black NiI₂ adopt similar structures, but with a different packing of the halides, adopting the CdI₂ motif.

In contrast, NiCl₂·6H₂O consists of separated trans-[NiCl₂(H₂O)₄] molecules linked more weakly to adjacent water molecules. Only four of the six water molecules in the formula are bound to the nickel, and the remaining two are water of crystallisation.^[7] Cobalt(II) chloride hexahydrate has a similar structure. The hexahydrate occurs in nature as the very rare mineral nickelbischofite.

The dihydrate NiCl₂·2H₂O adopts a structure intermediate between the hexahydrate and the anhydrous forms. It consists of infinite chains of NiCl₂, wherein both chloride centers are bridging ligands. The trans sites on the octahedral centers occupied by aquo ligands.^[8] A tetrahydrate NiCl₂·4H₂O is also known.

Reactions

Nickel(II) chloride solutions are acidic, with a pH of around 4 due to the hydrolysis of the Ni²⁺ ion.

Coordination complexes

Color of various Ni(II) complexes in aqueous solution. From left to right, [Ni(NH₃)₆]²⁺, [Ni(en)₃]²⁺, [NiCl₄]²⁻, [Ni(H₂O)₆]²⁺

Most of the reactions ascribed to "nickel chloride" involve the hexahydrate, although specialized reactions require the anhydrous form.

Reactions starting from NiCl₂·6H₂O can be used to form a variety of nickel coordination complexes because the H₂O ligands are rapidly displaced by ammonia, amines, thioethers, thiolates, and organophosphines. In some derivatives, the chloride remains within the coordination sphere, whereas chloride is displaced with highly basic ligands. Illustrative complexes include:

Complex	Color	Magnetism	Geometry
[Ni(NH3)6]Cl2	blue/violet	paramagnetic	octahedral
[Ni(en)3]^2+	violet	paramagnetic	octahedral
NiCl2(dppe)	orange	diamagnetic	square planar
[Ni(CN)4]^2−	colorless	diamagnetic	square planar
[NiCl4]^2−[9][10]	Yellowish-green	paramagnetic	tetrahedral

Crystals of hexammine nickel chloride

Some nickel chloride complexes exist as an equilibrium mixture of two geometries; these examples are some of the most dramatic illustrations of structural isomerism for a given coordination number. For example, NiCl₂(PPh₃)₂, containing four-coordinate Ni(II), exists in solution as a mixture of both the diamagnetic square planar and the paramagnetic tetrahedral isomers. Square planar complexes of nickel can often form five-coordinate adducts.

NiCl₂ is the precursor to acetylacetonate complexes Ni(acac)₂(H₂O)₂ and the benzene-soluble (Ni(acac)₂)₃, which is a precursor to Ni(1,5-cyclooctadiene)₂, an important reagent in organonickel chemistry.

In the presence of water scavengers, hydrated nickel(II) chloride reacts with dimethoxyethane (dme) to form the molecular complex NiCl₂(dme)₂. The dme ligands in this complex are labile. For example, this complex reacts with sodium cyclopentadienide to give the sandwich compound nickelocene.

Hexammine nickel chloride complex is soluble when respective cobalt complex is not, which allows for easy separating of these close-related metals in laboratory conditions.

Applications in organic synthesis

NiCl₂ and its hydrate are occasionally useful in organic synthesis.^[11]

• As a mild Lewis acid, e.g. for the regioselective isomerization of dienols:

• In combination with CrCl₂ for the coupling of an aldehyde and a vinylic iodide to give allylic alcohols.
• For selective reductions in the presence of LiAlH₄, e.g. for the conversion of alkenes to alkanes.
• As a precursor to nickel boride, prepared in situ from NiCl₂ and NaBH₄. This reagent behaves like Raney Nickel, comprising an efficient system for hydrogenation of unsaturated carbonyl compounds.
• As a precursor to finely divided Ni by reduction with Zn, for the reduction of aldehydes, alkenes, and nitro aromatic compounds. This reagent also promotes homo-coupling reactions, that is 2RX → R-R where R = aryl, vinyl.
• As a catalyst for making dialkyl arylphosphonates from phosphites and aryl iodide, ArI:

ArI + P(OEt)₃ → ArP(O)(OEt)₂ + EtI

Other uses

Nickel chloride solutions are used for electroplating nickel onto other metal items.

Safety

Nickel(II) chloride is irritating upon ingestion, inhalation, skin contact, and eye contact. Prolonged exposure to nickel and its compounds have been shown to produce cancer.

References

1. Lide, David S. (2003). CRC Handbook of Chemistry and Physics, 84th Edition. CRC Press. p. 4-71. ISBN 9780849304842.
2. Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22. ISBN 0-618-94690-X.
3. "Nickel metal and other compounds (as Ni)". Immediately Dangerous to Life and Health. National Institute for Occupational Safety and Health (NIOSH).
4. Laird G. L. Ward "Anhydrous Nickel (II) Halides and their Tetrakis(Ethanol) and 1,2-Dimethoxyethane Complexes Inorganic Syntheses, 1972 Volume 13, pages 154–164. doi:10.1002/9780470132449.ch30
5. Pray, A. P.; Tyree, S. Y.; Martin, Dean F.; Cook, James R. (1990). "Anhydrous Metal Chlorides". Inorganic Syntheses 28: 321–2. doi:10.1002/9780470132593.ch80.
6. Karyakin, Yu.V. (1947). Pure chemicals. Manual for laboratory preparation of inorganic substances (in Russian) (Moscow, Leningrad "State Scientific Technical Publishing of Chemical Literature" ed.). p. 416.
7. , Wells, A. F. Structural Inorganic Chemistry, Oxford Press, Oxford, United Kingdom, 1984.
8. B. Morosin "An X-ray diffraction study on nickel(II) chloride dihydrate" Acta Crystallogr. 1967. volume 23, pp. 630-634. doi:10.1107/S0365110X67003305
9. Gill, N. S. and Taylor, F. B. (1967). "Tetrahalo Complexes of Dipositive Metals in the First Transition Series". Inorganic Syntheses 9: 136–142. doi:10.1002/9780470132401.ch37.
10. G. D. Stucky, J. B. Folkers, T. J. Kistenmacher (1967). "The Crystal and Molecular Structure of Tetraethylammonium Tetrachloronickelate(II)". Acta Crystallographica 23 (6): 1064–1070. doi:10.1107/S0365110X67004268.
11. Tien-Yau Luh, Yu-Tsai Hsieh Nickel(II) Chloride" in Encyclopedia of Reagents for Organic Synthesis (L. A. Paquette, Ed.) 2001 J. Wiley & Sons, New York. doi:10.1002/047084289X.rn012. Article Online Posting Date: April 15, 2001.

External links

• NIOSH Pocket Guide to Chemical Hazards
• Linstrom, P.J.; Mallard, W.G. (eds.) NIST Chemistry WebBook, NIST Standard Reference Database Number 69. National Institute of Standards and Technology, Gaithersburg MD. http://webbook.nist.gov