Nickel(II) chloride

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. 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 have been shown to be carcinogenic to the lungs and nasal passages in cases of long-term inhalation exposure.^[4]

Nickel 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 Y 7791-20-0 (hexahydrate) Y
3D model (JSmol)	anhydrous: Interactive image hexahydrate: Interactive image
ChEBI	CHEBI:34887 Y
ChemSpider	22796 Y
ECHA InfoCard	100.028.858
EC Number	231-743-0
KEGG	C14711 Y
PubChem CID	24385
RTECS number	QR6480000
UNII	696BNE976J Y T8365BUD85 (hexahydrate) Y
UN number	3288 3077
CompTox Dashboard (EPA)	DTXSID7040316
InChI InChI=1S/2ClH.Ni/h2*1H;/q;;+2/p-2 Y Key: QMMRZOWCJAIUJA-UHFFFAOYSA-L Y InChI=1/2ClH.Ni/h2*1H;/q;;+2/p-2 Key: QMMRZOWCJAIUJA-NUQVWONBAR
SMILES anhydrous: [Ni+2].[Cl-].[Cl-] hexahydrate: Cl[Ni-4](Cl)([OH2+])([OH2+])([OH2+])[OH2+].O.O
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/cm3 (anhydrous) 1.92 g/cm3 (hexahydrate)
Melting point	1,001 °C (1,834 °F; 1,274 K) (anhydrous) 140 °C (hexahydrate)
Solubility in water	anhydrous 67.5 g/100 mL (25 °C) [1] 87.6 g/100 mL (100 °C) hexahydrate 282.5 g/100 mL (25 °C) [1] 578.5 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)
Magnetic susceptibility (χ)	+6145.0·10−6 cm3/mol
Structure
Crystal structure	Monoclinic
Coordination geometry	octahedral at Ni
Thermochemistry
Std molar entropy (S⦵298)	107 J·mol−1·K−1[2]
Std enthalpy of formation (ΔfH⦵298)	−316 kJ·mol−1[2]
Hazards
Occupational safety and health (OHS/OSH):
Main hazards	Very toxic (T+) Irritant (Xi) Dangerous for the environment (N) Carcinogen
GHS labelling:
Pictograms
Signal word	Danger
Hazard statements	H301, H315, H317, H331, H334, H341, H350i, H360D, H372, H410
Precautionary statements	P201, P202, P260, P261, P264, P270, P271, P272, P273, P280, P281, P285, P301+P310, P302+P352, P304+P340, P304+P341, P308+P313, P311, P314, P321, P330, P332+P313, P333+P313, P342+P311, P362, P363, P391, P403+P233, P405, P501
NFPA 704 (fire diamond)	3 0 0
Flash point	Non-flammable
Lethal dose or concentration (LD, LC):
LD50 (median dose)	105 mg/kg (rat, oral)[3]
Safety data sheet (SDS)	Fischer Scientific
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). N verify (what is YN ?) Infobox references

Production and syntheses

Large scale production and uses of nickel chloride are associated with the purification of nickel from its ores. It is generated upon extraction nickel matte and residues obtained from roasting refining nickel-containing ores using hydrochloric acid. Electrolysis of nickel chloride solutions are used in the production of nickel metal. Other significant routes to nickel chloride arise from processing of ore concentrates such as various reactions involving copper chlorides:^[5]

NiS + 2 CuCl₂ → NiCl₂ + 2 CuCl + S

NiO + 2 HCl → NiCl₂ + H₂O

Laboratory routes

Nickel chloride is not usually prepared in the laboratory because it is inexpensive and has a long shelf-life. The yellowish dihydrate, NiCl₂·2H₂O, is produced by heating the hexahydrate between 66 and 133 °C.^[6] 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.

NiCl₂·6H₂O + 6 SOCl₂ → NiCl₂ + 6SO₂ + 12HCl

The dehydration is accompanied by a color change from green to yellow.^[7]

In case one needs a pure compound without presence of cobalt, nickel chloride can be obtained by cautiously heating hexaamminenickel chloride:^[8]

${\displaystyle {\ce {{\overset {hexammine \atop nickel~chloride}{[Ni(NH3)6]Cl2}}->[175-200^{\circ }{\ce {C}}]NiCl2{}+6NH3}}}$

Structure of NiCl2 and its hydrates

Structure of hydrated nickel chloride based on X-ray crystallography. Color code: red = O, green = Cl

NiCl₂ adopts the CdCl₂ structure.^[9] 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 is bound to the nickel, and the remaining two are water of crystallization, so the formula of nickel(II) chloride hexahydrate is [NiCl₂(H₂O)₄]·2H₂O.^[9] 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.^[10] 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−[11][12]	Yellowish-green	paramagnetic	tetrahedral

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.^[6] The dme ligands in this complex are labile.

Applications in organic synthesis

NiCl₂ and its hydrate are occasionally useful in organic synthesis.^[13]

• 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 Brown's P-1 and P-2 nickel boride catalyst through reaction with NaBH₄.
• 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

NiCl₂-dme (or NiCl₂-glyme) is used due to its increased solubility in comparison to the hexahydrate.^[14]

Application of NiCl₂ precatalyst.

Safety

Nickel(II) chloride is irritating upon ingestion, inhalation, skin contact, and eye contact. Prolonged inhalation exposure to nickel and its compounds has been linked to increased cancer risk to the lungs and nasal passages.^[4]

References

1. Lide, David S. (2003). CRC Handbook of Chemistry and Physics, 84th Edition. CRC Press. pp. 4–71. ISBN 9780849304842.
2. Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22. ISBN 978-0-618-94690-7.
3. "Nickel metal and other compounds (as Ni)". Immediately Dangerous to Life or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
4. Grimsrud, Tom K; Andersen, Aage (2010). "Evidence of carcinogenicity in humans of water-soluble nickel salts". Journal of Occupational Medicine and Toxicology. 5 (1): 7. doi:10.1186/1745-6673-5-7. PMC 2868037. PMID 20377901.
5. Kerfoot, Derek G. E. (2000). "Nickel". Ullmann's Encyclopedia of Industrial Chemistry. doi:10.1002/14356007.a17_157. ISBN 978-3-527-30385-4.
6. Ward, Laird G. L. (1972). "Anhydrous Nickel(II) Halides and their Tetrakis(ethanol) and 1,2-Dimethoxyethane Complexes". Inorganic Syntheses. Vol. 13. pp. 154–164. doi:10.1002/9780470132449.ch30. ISBN 9780470132449.
7. Pray, A. P. (1990). "Anhydrous Metal Chlorides". Inorganic Syntheses. Inorganic Syntheses. Vol. 28. pp. 321–2. doi:10.1002/9780470132593.ch80. ISBN 9780470132593.
8. 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.
9. Wells, A. F. Structural Inorganic Chemistry, Oxford Press, Oxford, United Kingdom, 1984.
10. B. Morosin "An X-ray diffraction study on nickel(II) chloride dihydrate" Acta Crystallogr. 1967. volume 23, pp. 630-634. doi:10.1107/S0365110X67003305
11. Gill, N. S. & Taylor, F. B. (1967). Tetrahalo Complexes of Dipositive Metals in the First Transition Series. Inorganic Syntheses. Vol. 9. pp. 136–142. doi:10.1002/9780470132401.ch37. ISBN 9780470132401.
12. 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.
13. 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.
14. Cornella, Josep; Edwards, Jacob T.; Qin, Tian; Kawamura, Shuhei; Wang, Jie; Pan, Chung-Mao; Gianatassio, Ryan; Schmidt, Michael; Eastgate, Martin D. (2016-02-24). "Practical Ni-Catalyzed Aryl–Alkyl Cross-Coupling of Secondary Redox-Active Esters". Journal of the American Chemical Society. 138 (7): 2174–2177. doi:10.1021/jacs.6b00250. PMC 4768290. PMID 26835704.

External links

• NIOSH Pocket Guide to Chemical Hazards
• Linstrom, Peter J.; Mallard, William G. (eds.); NIST Chemistry WebBook, NIST Standard Reference Database Number 69, National Institute of Standards and Technology, Gaithersburg (MD)