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Nickel(II) chloride
IUPAC name Nickel(II) chloride
other names Nickelous chloride
Identifiers
CAS number	7718-54-9 Y,  7791-20-0 (hexahydrate)
PubChem	24385
EC number	231-743-0
RTECS number	QR6480000
Properties
Molecular formula	NiCl2
Molar mass	129.5994 g/mol (anhydrous) 237.69 g/mol (hexahydrate)
Appearance	yellow-green crystals deliquescent
Density	3.55 g/cm3 (anhydrous) 1.92 g/cm3 (hexahydrate)
Melting point	1001 °C (anhydrous) 140 °C (hexahydrate)
Solubility in water	anhydrous 64 g/100 mL hexahydrate 254 g/100 mL (20 °C) 600 g/100 mL (100 °C)
Solubility in ethanol	Soluble (hexahydrate)
Structure
Crystal structure	Monoclinic
Coordination geometry	octahedral at Ni
Thermochemistry
Std enthalpy of formation ΔfHo298	-304.93 kJ/mol
Standard molar entropy So298	98.11 J K−1 mol−1
Hazards
MSDS	Fischer Scientific
EU Index	028-011-00-6
EU classification	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
Flash point	Non-flammable
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
Y (what is this?)  (verify) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
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. It is very rarely found in nature as mineral nickelbischofite. A dihydrate is also known. In general nickel(II) chloride, in various forms, is the most important source of nickel for chemical synthesis. Nickel salts are carcinogenic.

Production and syntheses

Probably 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 rarely prepared in the laboratory because it is inexpensive and has a long shelf-life. The hydrate can be converted 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 + 6 SO₂ + 12 HCl

The dehydration is accompanied by a color change from green to yellow.^[1]

Structure and properties

NiCl₂ adopts the CdCl₂ structure.^[2] 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. Note that only four of the six water molecules in the formula are bound to the nickel, and the remaining two are water of crystallisation.^[2] Cobalt(II) chloride hexahydrate has a similar structure.

Many nickel(II) compounds are paramagnetic, due to the presence of two unpaired electrons on each metal center. Square planar nickel complexes are, however, diamagnetic.

Coordination chemistry

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 derivative, the chloride remains within the coordination sphere, whereas chloride is displaced with highly basic ligands. Illustrative complexes include:

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.

Applications in organic synthesis

NiCl₂ and its hydrate are occasionally useful in organic synthesis.^[5]

• 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.

References

1. ^ Pray, A. P. (1990). "Anhydrous Metal Chlorides". Inorganic Syntheses 28: 321–2. doi:10.1002/9780470132401.ch36. 
2. ^ ^{a b} , Wells, A. F. Structural Inorganic Chemistry, Oxford Press, Oxford, United Kingdom, 1984.
3. ^ 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. 
4. ^ G. D. Stucky, J. B. (1967). "The Crystal and Molecular Structure of Tetraethylammonium Tetrachloronickelate(II)". Acta Crystallographica 23: 1064–1070. doi:10.1107/S0365110X67004268. 
5. ^ 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

• National Pollutant Inventory: Nickel and compounds Fact Sheet(broken?)
• NIOSH Pocket Guide to Chemical Hazards
• IARC Monograph "Nickel and Nickel Compounds"(broken?)
• 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