Chemistry:Nickel(II) chloride

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Nickel chloride
Nickel(II) chloride hexahydrate
Hexahydrate
Anhydrous Nickel(II)-chloride.jpg
Anhydrous
MCl2(aq)6forFeCoNi.png
structure of hexahydrate
Names
IUPAC name
Nickel(II) chloride
Other names
Nickelous chloride, nickel(II) salt of hydrochloric acid
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
EC Number
  • 231-743-0
KEGG
RTECS number
  • QR6480000
UNII
UN number 3288 3077
Properties
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)
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)
+6145.0·10−6 cm3/mol
Structure
Monoclinic
octahedral at Ni
Thermochemistry
107 J·mol−1·K−1[2]
−316 kJ·mol−1[2]
Hazards
Main hazards Very toxic (T+)
Irritant (Xi)
Dangerous for the environment (N)
Carcinogen
Safety data sheet Fischer Scientific
GHS pictograms GHS06: ToxicGHS08: Health hazardGHS09: Environmental hazard
GHS Signal word Danger
H301, H315, H317, H331, H334, H341, H350i, H360D, H372, H410
P201, P202, P260, P261, P264, P270, P271, P272, P273, P280, P281, P285, P301+310, P302+352, P304+340, P304+341, P308+313, P311, P314, P321, P330, P332+313, P333+313, P342+311, P362
NFPA 704 (fire diamond)
Flammability code 0: Will not burn. E.g. waterHealth code 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasReactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no codeNFPA 704 four-colored diamond
0
3
0
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
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. 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]

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.

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.[5] 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.

[math]\ce{ NiCl2.6H2O + 6 SOCl2 -> NiCl2 + 6SO2 + 12HCl }[/math]

The dehydration is accompanied by a color change from green to yellow.[6]

In case one needs a pure compound without presence of cobalt, nickel chloride can be obtained by cautiously heating hexaamminenickel chloride:[7]

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

Structure of NiCl2 and its hydrates

NiCl2 adopts the CdCl2 structure.[8] 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.[8] 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.[9] 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−[10][11] 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.[5] 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.[12]

  • As a mild Lewis acid, e.g. for the regioselective isomerization of dienols:
General reaction scheme for the isomerisation 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.[13] File:NiCl2 schemes.tif

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 inhalation exposure to nickel and its compounds has been linked to increased cancer risk to the lungs and nasal passages.[4]

References

  1. ↑ 1.0 1.1 Lide, David S. (2003). CRC Handbook of Chemistry and Physics, 84th Edition. CRC Press. pp. 4–71. ISBN 9780849304842. 
  2. ↑ 2.0 2.1 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 and Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH). https://www.cdc.gov/niosh/idlh/7440020.html. 
  4. ↑ 4.0 4.1 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. PMID 20377901. 
  5. ↑ 5.0 5.1 Ward, Laird G. L. (1972). "Anhydrous Nickel(II) Halides and their Tetrakis(ethanol) and 1,2-Dimethoxyethane Complexes". Inorganic Syntheses. 13. pp. 154–164. doi:10.1002/9780470132449.ch30. ISBN 9780470132449. 
  6. ↑ Pray, A. P. (1990). "Anhydrous Metal Chlorides". Inorganic Syntheses. Inorganic Syntheses. 28. pp. 321–2. doi:10.1002/9780470132593.ch80. ISBN 9780470132593. 
  7. ↑ Karyakin, Yu.V. (1947) (in Russian). Pure chemicals. Manual for laboratory preparation of inorganic substances (Moscow, Leningrad "State Scientific Technical Publishing of Chemical Literature" ed.). pp. 416. 
  8. ↑ 8.0 8.1 Wells, A. F. Structural Inorganic Chemistry, Oxford Press, Oxford, United Kingdom , 1984.
  9. ↑ B. Morosin "An X-ray diffraction study on nickel(II) chloride dihydrate" Acta Crystallogr. 1967. volume 23, pp. 630-634. doi:10.1107/S0365110X67003305
  10. ↑ Gill, N. S.; Taylor, F. B. (1967). Tetrahalo Complexes of Dipositive Metals in the First Transition Series. Inorganic Syntheses. 9. pp. 136–142. doi:10.1002/9780470132401.ch37. ISBN 9780470132401. 
  11. ↑ 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. 
  12. ↑ 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.
  13. ↑ Cornella, Josep; Edwards, Jacob T.; Qin, Tian; Kawamura, Shuhei; Wang, Jie; Pan, Chung-Mao; Gianatassio, Ryan; Schmidt, Michael et al. (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. PMID 26835704. 

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