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| verifiedrevid = 443817436 |
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| Name = Nickel(II) chloride |
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|verifiedrevid = 445919159 |
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| ImageFile = Nickel(II)-chloride-hexahydrate-sample.jpg |
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| ImageName = Nickel(II) chloride hexahydrate |
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|Name = Nickel chloride |
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| IUPACName = Nickel(II) chloride |
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|ImageFile1 = Nickel(II)-chloride-hexahydrate-sample.jpg |
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|ImageName1 = Nickel(II) chloride hexahydrate |
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| OtherNames = Nickelous chloride, nickel(II) salt of hydrochloric acid |
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|ImageCaption1 = Hexahydrate |
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| Section1 = {{Chembox Identifiers |
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|ImageFile2 = Anhydrous Nickel(II)-chloride.jpg |
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| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}} |
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|ImageCaption2 = Anhydrous |
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| ChemSpiderID = 22796 |
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|IUPACName = Nickel(II) chloride |
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| KEGG_Ref = {{keggcite|correct|kegg}} |
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|OtherNames = Nickelous chloride, nickel(II) salt of hydrochloric acid |
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| KEGG = C14711 |
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|Section1 = {{Chembox Identifiers |
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| InChI = 1/2ClH.Ni/h2*1H;/q;;+2/p-2 |
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|CASNo = 7718-54-9 |
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| InChIKey = QMMRZOWCJAIUJA-NUQVWONBAR |
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| ChEBI_Ref = {{ebicite|correct|EBI}} |
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|CASNo_Ref = {{cascite|correct|CAS}} |
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|CASNo2_Ref = {{cascite|correct|CAS}} |
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| ChEBI = 34887 |
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|CASNo2 = 7791-20-0 |
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| SMILES = ClCl |
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|CASNo2_Comment = (hexahydrate) |
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| StdInChI_Ref = {{stdinchicite|correct|chemspider}} |
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|ChEBI_Ref = {{ebicite|correct|EBI}} |
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| StdInChI = 1S/2ClH.Ni/h2*1H;/q;;+2/p-2 |
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|ChEBI = 34887 |
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| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}} |
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|ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}} |
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| StdInChIKey = QMMRZOWCJAIUJA-UHFFFAOYSA-L |
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|ChemSpiderID = 22796 |
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| CASNo = 7718-54-9 |
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|EINECS = 231-743-0 |
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| CASNo_Ref = {{cascite|correct|CAS}} |
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|KEGG_Ref = {{keggcite|correct|kegg}} |
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| CASOther = <br/>7791-20-0 (hexahydrate) |
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|KEGG = C14711 |
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| PubChem = 24385 |
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|PubChem = 24385 |
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| EINECS = 231-743-0 |
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| RTECS = QR6480000 |
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|RTECS = QR6480000 |
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|UNNumber = 3288 3077 |
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}} |
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|UNII_Ref = {{fdacite|correct|FDA}} |
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| Section2 = {{Chembox Properties |
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|UNII = 696BNE976J |
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| Formula = NiCl<sub>2</sub> |
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|UNII2_Ref = {{fdacite|correct|FDA}} |
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| Appearance = yellow-green crystals <br> ] |
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|UNII2 = T8365BUD85 |
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| Density = 3.55 g/cm<sup>3</sup> (anhydrous) <br> 1.92 g/cm<sup>3</sup> (hexahydrate) |
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| MolarMass = 129.5994 g/mol (anhydrous) <br> 237.69 g/mol (hexahydrate) |
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|UNII2_Comment = (hexahydrate) |
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|SMILES = .. |
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| Solubility = ''anhydrous'' <br> 64 g/100 mL <hr> ''hexahydrate'' <br> 254 g/100 mL (20 °C) <br> 600 g/100 mL (100 °C) |
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|SMILES_Comment = anhydrous |
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| Solvent = ] |
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|SMILES1 = Cl(Cl)()()().O.O |
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| SolubleOther = Soluble (hexahydrate) |
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| MeltingPt = 1001 °C (anhydrous) <br> 140 °C (hexahydrate) |
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|SMILES1_Comment = hexahydrate |
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|InChI = 1/2ClH.Ni/h2*1H;/q;;+2/p-2 |
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}} |
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|InChIKey = QMMRZOWCJAIUJA-NUQVWONBAR |
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| Section3 = {{Chembox Structure |
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|StdInChI_Ref = {{stdinchicite|correct|chemspider}} |
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| Coordination = octahedral at Ni |
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|StdInChI = 1S/2ClH.Ni/h2*1H;/q;;+2/p-2 |
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| CrystalStruct = ] |
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|StdInChIKey_Ref = {{stdinchicite|correct|chemspider}} |
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}} |
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|StdInChIKey = QMMRZOWCJAIUJA-UHFFFAOYSA-L |
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| Section4 = {{Chembox Thermochemistry |
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}} |
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| DeltaHf = -304.93 kJ/mol |
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|Section2 = {{Chembox Properties |
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| Entropy = 98.11 J K<sup>−1</sup> mol<sup>−1</sup> |
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|Formula = NiCl<sub>2</sub> |
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|Appearance = yellow-brown crystals <br> ] (anhydrous) <br> green crystals (hexahydrate) |
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| Section7 = {{Chembox Hazards |
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|Odor = odorless |
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| ExternalMSDS = |
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|Density = 3.55 g/cm<sup>3</sup> (anhydrous) <br> 1.92 g/cm<sup>3</sup> (hexahydrate) |
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| EUClass = ]<br/>]<br/>Repr. Cat. 2<br/>Toxic ('''T''')<br/>Irritant ('''Xi''')<br/>Dangerous for the environment ('''N''') |
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|MolarMass = 129.5994 g/mol (anhydrous) <br> 237.69 g/mol (hexahydrate) |
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| EUIndex = 028-011-00-6 |
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|Solubility = ''anhydrous'' <br> 67.5 g/100 mL (25 °C) <ref name=b0>{{cite book| author = Lide, David S.|title =CRC Handbook of Chemistry and Physics, 84th Edition| publisher = CRC Press| year = 2003| isbn = 9780849304842|pages=4–71}}</ref><br> 87.6 g/100 mL (100 °C) <hr> ''hexahydrate'' <br> 282.5 g/100 mL (25 °C) <ref name=b0/><br> 578.5 g/100 mL (100 °C) |
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| NFPA-H = 2 |
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|SolubleOther = 0.8 g/100 mL (]) <br> soluble in ], ], ] <br> insoluble in ], ] |
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| NFPA-F = 0 |
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| NFPA-R = 0 |
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|MeltingPtC = 1001 |
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|MeltingPt_notes = (anhydrous) <br> 140 °C (hexahydrate) |
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| NFPA-O = |
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|pKa = 4 (hexahydrate) |
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| RPhrases = {{R49}}, {{R61}}, {{R23/25}}, {{R38}}, {{R42/43}}, {{R48/23}}, {{R68}}, {{R50/53}} |
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|MagSus = +6145.0·10<sup>−6</sup> cm<sup>3</sup>/mol |
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| SPhrases = {{S53}}, {{S45}}, {{S60}}, {{S61}} |
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}} |
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| FlashPt = Non-flammable |
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|Section3 = {{Chembox Structure |
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| PEL = |
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|Coordination = octahedral at Ni |
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|CrystalStruct = ] |
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| Section8 = {{Chembox Related |
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}} |
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| OtherAnions = ]<br/>]<br/>] |
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|Section4 = {{Chembox Thermochemistry |
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| OtherCations = ]<br/>]<br/>]<br/>] |
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|DeltaHf = −316 kJ·mol<sup>−1</sup><ref name=b1>{{cite book| author = Zumdahl, Steven S.|title =Chemical Principles 6th Ed.| publisher = Houghton Mifflin Company| year = 2009| isbn = 978-0-618-94690-7|page=A22}}</ref> |
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| OtherCpds = ]<br/>]}} |
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|Entropy = 107 J·mol<sup>−1</sup>·K<sup>−1</sup><ref name=b1/> |
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}} |
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|Section5 = {{Chembox Hazards |
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|MainHazards = Very toxic ('''T+''')<br/>Irritant ('''Xi''')<br/>Dangerous for the environment ('''N''')<br/>Carcinogen |
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|ExternalSDS = |
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|NFPA-H = 3 |
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|NFPA-F = 0 |
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|NFPA-R = 0 |
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|GHSPictograms = {{GHS06}}{{GHS08}}{{GHS09}} |
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|GHSSignalWord = Danger |
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|HPhrases = {{H-phrases|301|315|317|331|334|341|350i|360D|372|410}} |
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|PPhrases = {{P-phrases|201|202|260|261|264|270|271|272|273|280|281|285|301+310|302+352|304+340|304+341|308+313|311|314|321|330|332+313|333+313|342+311|362|363|391|403+233|405|501}} |
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|FlashPt = Non-flammable |
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|LD50 = 105 mg/kg (rat, oral)<ref>{{IDLH|7440020|Nickel metal and other compounds (as Ni)}}</ref> |
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}} |
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|Section6 = {{Chembox Related |
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|OtherAnions = ]<br/>]<br/>] |
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|OtherCations = ]<br/>]<br/>Platinum(II,IV) chloride<br/>] |
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|OtherCompounds = ]<br/>]}} |
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'''Nickel(II) chloride''' (or just '''nickel chloride'''), is the ] NiCl<sub>2</sub>. The ] ] is yellow, but the more familiar ] NiCl<sub>2</sub>·6H<sub>2</sub>O is green. It is very rarely found in nature as mineral ]. A dihydrate is also known. In general nickel(II) chloride, in various forms, is the most important source of nickel for chemical synthesis. ] salts are ]ic. They are also ], absorbing moisture from the air to form a solution. |
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'''Nickel(II) chloride''' (or just '''nickel chloride''') is the ] NiCl<sub>2</sub>. The ] ] is yellow, but the more familiar ] NiCl<sub>2</sub>·6H<sub>2</sub>O is green. Nickel(II) chloride, in various forms, is the most important source of nickel for chemical synthesis. The nickel chlorides are ], absorbing moisture from the air to form a solution. ] salts have been shown to be ]ic to the lungs and nasal passages in cases of long-term ].<ref name="Grimsrud">{{cite journal|last1=Grimsrud|first1=Tom K|last2=Andersen|first2=Aage|title=Evidence of carcinogenicity in humans of water-soluble nickel salts|journal=Journal of Occupational Medicine and Toxicology|date=2010|volume=5|issue=1|pages=7|doi=10.1186/1745-6673-5-7|pmid=20377901|pmc=2868037 |doi-access=free }}</ref> |
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==Production and syntheses== |
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==Production and syntheses== |
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Probably the largest scale production of nickel chloride involves the extraction with hydrochloric acid of nickel ] and residues obtained from roasting refining nickel-containing ores. |
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Large scale production and uses of nickel chloride are associated with the purification of nickel from its ores. It is generated upon extraction nickel ] 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:<ref>{{cite book |doi=10.1002/14356007.a17_157 |chapter=Nickel |title=Ullmann's Encyclopedia of Industrial Chemistry |date=2000 |last1=Kerfoot |first1=Derek G. E. |isbn=978-3-527-30385-4 }}</ref> |
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:{{chem2|NiS + 2 CuCl2 -> NiCl2 + 2 CuCl + S}} |
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:{{chem2|NiO + 2 HCl -> NiCl2 + H2O}} |
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===Laboratory routes=== |
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NiCl<sub>2</sub>·6H<sub>2</sub>O 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 ] or by heating under a stream of HCl gas. Simply heating the hydrates does not afford the anhydrous dichloride. |
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Nickel chloride is not usually prepared in the laboratory because it is inexpensive and has a long shelf-life. The yellowish dihydrate, NiCl<sub>2</sub>·2H<sub>2</sub>O, is produced by heating the hexahydrate between 66 and 133 °C.<ref name=dme>{{cite book|doi=10.1002/9780470132449.ch30|chapter=Anhydrous Nickel(II) Halides and their Tetrakis(ethanol) and 1,2-Dimethoxyethane Complexes|year=1972|last1=Ward|first1=Laird G. L.|title=Inorganic Syntheses|volume=13|pages=154–164|isbn=9780470132449}}</ref> The hydrates convert to the anhydrous form upon heating in ] or by heating under a stream of HCl gas. Simply heating the hydrates does not afford the anhydrous dichloride. |
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:NiCl<sub>2</sub>·6H<sub>2</sub>O + 6 SOCl<sub>2</sub> → NiCl<sub>2</sub> + 6 SO<sub>2</sub> + 12 ] |
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:{{chem2|NiCl2*6H2O + 6 SOCl2 -> NiCl2 + 6SO2 + 12HCl}} |
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The dehydration is accompanied by a color change from green to yellow.<ref>{{cite journal | author = Pray, A. P. | title = Anhydrous Metal Chlorides | journal = ] | volume = 28 | pages = 321–2 | year = 1990 | doi = 10.1002/9780470132401.ch36 | last2 = Tyree | first2 = S. Y. | last3 = Martin | first3 = Dean F. | last4 = Cook | first4 = James R.}}</ref> |
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The dehydration is accompanied by a color change from green to yellow.<ref>{{cite book | author = Pray, A. P. | title = Inorganic Syntheses | chapter = Anhydrous Metal Chlorides | series = ] | volume = 28 | pages = 321–2 | year = 1990 | doi = 10.1002/9780470132593.ch80| isbn = 9780470132593 }}</ref> |
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==Structure and properties== |
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NiCl<sub>2</sub> adopts the CdCl<sub>2</sub> structure.<ref name=Wells>, Wells, A. F. ''Structural Inorganic Chemistry'', Oxford Press, ], ], 1984.</ref> In this motif, each Ni<sup>2+</sup> center is coordinated to six Cl<sup>-</sup> centers, and each chloride is bonded to three Ni(II) centers. In NiCl<sub>2</sub> the Ni-Cl bonds have “ionic character”. Yellow NiBr<sub>2</sub> and black NiI<sub>2</sub> adopt similar structures, but with a different packing of the halides, adopting the CdI<sub>2</sub> motif. |
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In case one needs a pure compound without presence of cobalt, nickel chloride can be obtained by cautiously heating ]:<ref name="karyakin">{{cite book|language = Russian | author = Karyakin, Yu.V. |title = Pure chemicals. Manual for laboratory preparation of inorganic substances |edition = Moscow, Leningrad "State Scientific Technical Publishing of Chemical Literature" |year = 1947 |pages=416}}</ref> |
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In contrast, NiCl<sub>2</sub>·6H<sub>2</sub>O consists of separated ''trans''- 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 ].<ref name=Wells/> ] hexahydrate has a similar structure. |
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:<chem>\overset{hexammine\atop nickel~chloride}{Cl2} -> NiCl2{} + 6NH3</chem> |
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==Structure of NiCl<sub>2</sub> and its hydrates== |
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Many nickel(II) compounds are ], due to the presence of two unpaired electrons on each metal center. Square planar nickel complexes are, however, ]. |
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]. Color code: red = O, green = Cl|thumb|left]] |
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NiCl<sub>2</sub> adopts the ].<ref name=Wells>Wells, A. F. ''Structural Inorganic Chemistry'', Oxford Press, ], ], 1984.</ref> In this motif, each Ni<sup>2+</sup> center is coordinated to six Cl<sup>−</sup> centers, and each chloride is bonded to three Ni(II) centers. In NiCl<sub>2</sub> the Ni-Cl bonds have "ionic character". Yellow NiBr<sub>2</sub> and black NiI<sub>2</sub> adopt similar structures, but with a different packing of the halides, adopting the CdI<sub>2</sub> motif. |
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In contrast, NiCl<sub>2</sub>·6H<sub>2</sub>O consists of separated ''trans''- 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 ], so the formula of nickel(II) chloride hexahydrate is ·2H<sub>2</sub>O.<ref name=Wells/> ] hexahydrate has a similar structure. The hexahydrate occurs in nature as the very rare mineral nickelbischofite. |
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The dihydrate NiCl<sub>2</sub>·2H<sub>2</sub>O adopts a structure intermediate between the hexahydrate and the anhydrous forms. It consists of infinite chains of NiCl<sub>2</sub>, wherein both chloride centers are ]s. The trans sites on the octahedral centers occupied by ]s.<ref>B. Morosin "An X-ray diffraction study on nickel(II) chloride dihydrate" Acta Crystallogr. 1967. volume 23, pp. 630-634. {{doi|10.1107/S0365110X67003305}}</ref> A tetrahydrate NiCl<sub>2</sub>·4H<sub>2</sub>O is also known. |
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==Reactions== |
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Nickel(II) chloride solutions are acidic, with a pH of around 4 due to the hydrolysis of the Ni<sup>2+</sup> ion. |
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Nickel(II) chloride solutions are acidic, with a pH of around 4 due to the hydrolysis of the Ni<sup>2+</sup> ion. |
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==Coordination chemistry== |
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===Coordination complexes=== |
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<sup>2+</sup>, ])<sub>3</sub>]<sup>2+</sup>, <sup>2−</sup>, <sup>2+</sup>]] |
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] |
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Most of the reactions ascribed to “nickel chloride” involve the hexahydrate, although specialized reactions require the anhydrous form. |
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Most of the reactions ascribed to "nickel chloride" involve the hexahydrate, although specialized reactions require the ] form. |
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Reactions starting from NiCl<sub>2</sub>·6H<sub>2</sub>O can be used to form a variety of nickel ] complexes because the H<sub>2</sub>O ligands are rapidly displaced by ], ]s, ]s, ]s, and organo]s. In some derivatives, the chloride remains within the ], whereas chloride is displaced with highly basic ligands. Illustrative complexes include: |
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Reactions starting from NiCl<sub>2</sub>·6H<sub>2</sub>O can be used to form a variety of nickel ] complexes because the H<sub>2</sub>O ligands are rapidly displaced by ], ]s, ]s, ]s, and organo]s. In some derivatives, the chloride remains within the ], whereas chloride is displaced with highly basic ligands. Illustrative complexes include: |
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| Cl<sub>2</sub> |
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| Cl<sub>2</sub>]] |
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| blue/violet |
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| paramagnetic |
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| octahedral |
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| ])<sub>3</sub>]<sup>2+</sup> |
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| violet |
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| violet |
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| paramagnetic |
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| paramagnetic |
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| octahedral |
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| octahedral |
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| NiCl<sub>2</sub>(]) |
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| ] |
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| orange |
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| orange |
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| diamagnetic |
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| diamagnetic |
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| square planar |
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| square planar |
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| ])<sub>4</sub>]<sup>2-</sup> |
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| ])<sub>4</sub>]<sup>2−</sup> |
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| colorless |
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| colorless |
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| diamagnetic |
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| diamagnetic |
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| square planar |
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| square planar |
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|<sup>2-</sup><ref>{{cite journal | author = Gill, N. S. and Taylor, F. B. | title = Tetrahalo Complexes of Dipositive Metals in the First Transition Series | journal = ] | year = 1967 | volume = 9 | pages = 136–142 | doi = 10.1002/9780470132401.ch37}}</ref><ref>{{cite journal | unused_data = Folkers and T. J. Kistenmacher | author = G. D. Stucky, J. B. | title = The Crystal and Molecular Structure of Tetraethylammonium Tetrachloronickelate(II) | journal = ] | year = 1967 | volume = 23 | issue = 6 | pages = 1064–1070 | doi =10.1107/S0365110X67004268}}</ref> |
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| <sup>2−</sup><ref>{{cite book |author1=Gill, N. S. |author2=Taylor, F. B. |name-list-style=amp | title = Tetrahalo Complexes of Dipositive Metals in the First Transition Series | series = ] | year = 1967 | volume = 9 | pages = 136–142 | doi = 10.1002/9780470132401.ch37 |isbn=9780470132401 }}</ref><ref>{{cite journal |author1=G. D. Stucky |author2=J. B. Folkers |author3=T. J. Kistenmacher | title = The Crystal and Molecular Structure of Tetraethylammonium Tetrachloronickelate(II) | journal = ] | year = 1967 | volume = 23 | issue = 6 | pages = 1064–1070 | doi =10.1107/S0365110X67004268}}</ref> |
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| Yellowish-Brown |
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| Yellowish-green |
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| paramagnetic |
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| paramagnetic |
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| tetrahedral |
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| tetrahedral |
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NiCl<sub>2</sub> is the precursor to ] complexes Ni(acac)<sub>2</sub>(H<sub>2</sub>O)<sub>2</sub> and the benzene-soluble (Ni(acac)<sub>2</sub>)<sub>3</sub>, which is a precursor to ], an important reagent in organonickel chemistry. |
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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 ]. For example, NiCl<sub>2</sub>(PPh<sub>3</sub>)<sub>2</sub>, 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. |
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In the presence of water scavengers, hydrated nickel(II) chloride reacts with ] (dme) to form the molecular complex NiCl<sub>2</sub>(dme)<sub>2</sub>.<ref name=dme/> The dme ligands in this complex are labile. |
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NiCl<sub>2</sub> is the precursor to ] complexes Ni(acac)<sub>2</sub>(H<sub>2</sub>O)<sub>2</sub> and the benzene-soluble (Ni(acac)<sub>2</sub>)<sub>3</sub>, which is a precursor to Ni(1,5-])<sub>2</sub>, an important reagent in organonickel chemistry. |
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===Applications in organic synthesis=== |
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In the presence of water scavengers, hydrated nickel(II) chloride reacts with ] (dme) to form the molecular complex NiCl<sub>2</sub>(dme)<sub>2</sub>. The dme ligands in this complex are labile. For example, this complex reacts with ] to give the ] ]. |
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NiCl<sub>2</sub> and its hydrate are occasionally useful in ].<ref>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.</ref> |
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==Applications in organic synthesis== |
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NiCl<sub>2</sub> and its hydrate are occasionally useful in ].<ref>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.</ref> |
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*As a mild Lewis acid, e.g. for the regioselective isomerization of dienols: |
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*As a mild Lewis acid, e.g. for the regioselective isomerization of dienols: |
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::] |
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::] |
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*In combination with ] for the coupling of an aldehyde and a vinylic iodide to give allylic alcohols. |
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*In combination with ] for the coupling of an ] and a vinylic iodide to give ]. |
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*For selective reductions in the presence of ], e.g. for the conversion of alkenes to alkanes. |
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*For selective reductions in the presence of ], e.g. for the conversion of alkenes to alkanes. |
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*As a precursor to “nickel boride”, prepared in situ from NiCl<sub>2</sub> and ]. This reagent behaves like Raney Nickel, comprising an efficient system for hydrogenation of unsaturated carbonyl compounds. |
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*As a precursor to ]'s ] through reaction with ]. |
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*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. |
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*As a precursor to finely divided Ni by reduction with Zn, for the reduction of aldehydes, ], and nitro ]. This reagent also promotes homo-coupling reactions, that is 2RX → R-R where R = aryl, vinyl. |
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*As a catalyst for making dialkyl arylphosphonates from phosphites and ] iodide, ArI: |
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*As a catalyst for making dialkyl arylphosphonates from phosphites and ] iodide, ArI: |
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::ArI + P(OEt)<sub>3</sub> → ArP(O)(OEt)<sub>2</sub> + EtI |
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::ArI + P(OEt)<sub>3</sub> → ArP(O)(OEt)<sub>2</sub> + EtI |
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NiCl<sub>2</sub>-dme (or NiCl<sub>2</sub>-glyme) is used due to its increased solubility in comparison to the hexahydrate.<ref>{{Cite journal |last1=Cornella |first1=Josep |last2=Edwards |first2=Jacob T. |last3=Qin |first3=Tian |last4=Kawamura |first4=Shuhei |last5=Wang |first5=Jie |last6=Pan |first6=Chung-Mao |last7=Gianatassio |first7=Ryan |last8=Schmidt |first8=Michael |last9=Eastgate |first9=Martin D. |date=2016-02-24 |title=Practical Ni-Catalyzed Aryl–Alkyl Cross-Coupling of Secondary Redox-Active Esters |journal=Journal of the American Chemical Society |volume=138 |issue=7 |pages=2174–2177 |doi=10.1021/jacs.6b00250 |doi-access=free |pmc=4768290 |pmid=26835704}}</ref> |
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==Other uses== |
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Nickel chloride solutions are used for ] nickel onto other metal items. |
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==Safety== |
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==Safety== |
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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. |
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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.<ref name="Grimsrud" /> |
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==References== |
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==References== |
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==External links== |
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==External links== |
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{{Commons cat|Nickel(II) chloride}} |
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{{Commons category|Nickel(II) chloride}} |
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{{Nickel compounds}} |
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