Nickel(II) chloride

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Nickel(II) chloride

General
Systematic name	Nickel(II) chloride
Other names	Nickelous chloride
Molecular formula	NiCl₂
Molar mass g/mol	129.59 (anhydrous) 237.71 g/mol (hexahydrate)
Appearance	yellow crystals green crystals (hexahydrate)
CAS number	7718-54-9 (anhydr.) 7791-20-0 (6H₂O)
EINECS number	231-743-0
Properties
Density and phase	3.55 g/cm³, anhyd.
Solubility in water	254 g/100 mL (20 °C)
Solubility in ethanol	Soluble (hexahydrate)
Melting point	1001 °C hydrate loses water
Structure
Coordination geometry	octahedral at Ni
Crystal structure	Monoclinic
Thermodynamic data
Standard enthalpy of formation Δ_fH°_solid	-304.93 kJ/mol (anhyd form)
Standard molar entropy S°_solid	98.11 J.K⁻¹.mol⁻¹ for the hexahydrate?
Safety data
EU classification	not listed
PEL-TWA (OSHA)	1 mg/m³ (as Ni)
IDLH (NIOSH)	10 mg/m³ (as Ni)
RTECS number	QR6480000
Supplementary data page
Thermodynamic data	Phase behaviour Solid
Spectral data	UV
Related compounds
Other anions	Nickel(II) bromide Nickel(II) iodide
Other cations	Cadmium(II) chloride Palladium(II) chloride Platinum(II) chloride
Related compounds	Cobalt(II) chloride Copper(II) chloride
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox disclaimer and references

Nickel(II) chloride, or nickelous chloride or just nickel chloride, is the chemical compound NiCl₂. It is a high melting, paramagnetic yellow solid. The term nickel chloride often also refers to the yellow-green hexahydrate NiCl₂(H₂O)₆, which is commonly found in the laboratory. A dihydrate is also known, NiCl₂(H₂O)₂. In general nickel(II) chloride, in various forms, is the most important source of nickel for chemical synthesis. Like all nickel salts, this compound is carcinogenic.

1 Production and syntheses
2 Structure and properties
3 Coordination chemistry
4 Applications in organic synthesis
5 Other uses
6 References
7 External links

[edit] Production and syntheses

The largest scale process to a nickel chloride probably involves the extraction with hydrochloric acid of nickel matte and residues, obtained from roasting refining nickel-containing ores.

NiCl₂(H₂O)₆ 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. Simple heating the hydrates does not afford the anhydrous dichloride.

NiCl₂(H₂O)₆ + 6 SOCl₂ → NiCl₂ + 6 SO₂ + 12 HCl

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

[edit] 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, the CdI₂ motif.

In contrast, NiCl₂(H₂O)₆ 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 Ni.^[2]

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

[edit] Coordination chemistry

Most of the reactions ascribed to “nickel chloride” involve the hexahydrate, although specialized reactions require the anhydrous form. For example, NiCl₂ reacts with dimethoxyethane to form the molecular complex, NiCl₂(dme)₂. This complex reacts with sodium cyclopentadienide to give nickelocene.

NiCl₂(H₂O)₆ forms a vast array of coordination complexes because the H₂O ligands are rapidly displaced by ammonia, amines, thioethers, thiolates, and organophosphines. In some derivative, the chloride remains the coordination sphere whereas with highly basic ligands, chloride is displaced. Illustrative complexes include:

violet, paramagnetic, octahedral [Ni(NH₃)₆]Cl₂

orange, diamagnetic, square planar NiCl₂(Ph₂PCH₂CH₂PPh₂)

colorless, diamagnetic, square planar [Ni(CN)₄]^2-

blue, paramagnetic, tetrahedral [NiCl₄]^2-^[3]^[4]

Some nickel chlorides exist as an equilibrium mixture of two structures. In fact these examples are some of the most dramatic illustrations of structural isomerism for a given coordination number. For example, NiCl₂(PPh₃)₂ features four-coordinate Ni(II). In solution, this compound has been shown to exist as a mixture of both the diamagnetic square planar and the paramagnetic tetrahedral isomers. Square planar complexes can often form five-coordinate adducts.

NiCl₂ is the precursor to Ni(acac)₂ (with a trimeric structure), which is a key precursor to Ni(1,5-cyclooctadiene)₂, an important reagent in organoNi chemistry.

[edit] Applications in organic synthesis

NiCl₂ and its hydrate are useful in organic synthesis.

As a mild Lewis acid, e.g. for the regioselective isomerization of dienols, i.e. RCH=CH-CH=CH-CH₂OH to RCH(OH)-CH=CH-CH=CH₂.
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, especially.
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

[edit] Other uses

Nickel chloride solutions are used for electroplating.

[edit] References

^ Pray, A. P. “Anhydrous Metal Chlorides” "Inorganic Syntheses," vol. XXVIII, 321-2, 1990. Describes the formation of anhydrous LiCl, CuCl₂, ZnCl₂, CdCl₂, ThCl₄, CrCl₃, FeCl₃, CoCl₂, and NiCl₂ from the corresponding hydrates.
^ ^a ^b , Wells, A. F. Structural Inorganic Chemistry, Oxford Press, Oxford, United Kingdom, 1984.
^ Gill, N. S. and Taylor, F. B., "Tetrahalo Complexes of Dipositive Metals in the First Transition Series", Inorganic Syntheses, 1967, volume 9, pages 136-142.
^ G. D. Stucky, J. B. Folkers and T. J. Kistenmacher "The Crystal and Molecular Structure of Tetraethylammonium Tetrachloronickelate(II)" Acta Crystallographica, 1967. volume 23, pages 1064-1070. DOI:10.1107/S0365110X67004268]