Nickel(II) chloride: Difference between revisions

|Verifiedfields = changed

|Watchedfields = changed

|verifiedrevid = 445919159

|Name = Nickel chloride

|ImageFile1 = Nickel(II)-chloride-hexahydrate-sample.jpg

|ImageName1 = Nickel(II) chloride hexahydrate

|ImageCaption1 = Hexahydrate

|ImageFile2 = Anhydrous Nickel(II)-chloride.jpg

|ImageCaption2 = Anhydrous

|ImageFile3 = MCl2(aq)6forFeCoNi.png

|ImageCaption3 = structure of hexahydrate

|IUPACName = Nickel(II) chloride

|OtherNames = Nickelous chloride, nickel(II) salt of hydrochloric acid

|Section1 = {{Chembox Identifiers

|CASNo = 7718-54-9

|CASNo_Ref = {{cascite|correct|CAS}}

|CASNo2_Ref = {{cascite|correct|CAS}}

|CASNo2 = 7791-20-0

|CASNo2_Comment = (hexahydrate)

|ChEBI_Ref = {{ebicite|correct|EBI}}

|ChEBI = 34887

|ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}

|ChemSpiderID = 22796

|EINECS = 231-743-0

|KEGG_Ref = {{keggcite|correct|kegg}}

|KEGG = C14711

|PubChem = 24385

|RTECS = QR6480000

|UNNumber = 3288 3077

|UNII_Ref = {{fdacite|correct|FDA}}

|UNII = 696BNE976J

|UNII2_Ref = {{fdacite|correct|FDA}}

|UNII2 = T8365BUD85

|UNII2_Comment = (hexahydrate)

|SMILES = Cl[Ni]Cl

|InChI = 1/2ClH.Ni/h2*1H;/q;;+2/p-2

|InChIKey = QMMRZOWCJAIUJA-NUQVWONBAR

|StdInChI_Ref = {{stdinchicite|correct|chemspider}}

|StdInChI = 1S/2ClH.Ni/h2*1H;/q;;+2/p-2

|StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}

|StdInChIKey = QMMRZOWCJAIUJA-UHFFFAOYSA-L

}}

|Section2 = {{Chembox Properties

|Formula = NiCl₂

|Appearance = yellow-brown crystals <br> [[deliquescent]] (anhydrous) <br> green crystals (hexahydrate)

|Odor = odorless

|Density = 3.55 g/cm³ (anhydrous) <br> 1.92 g/cm³ (hexahydrate)

|MolarMass = 129.5994 g/mol (anhydrous) <br> 237.69 g/mol (hexahydrate)

|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)

|SolubleOther = 0.8 g/100 mL ([[hydrazine]]) <br> soluble in [[ethylene glycol]], [[ethanol]], [[ammonium hydroxide]] <br> insoluble in [[ammonia]], [[nitric acid]]

|MeltingPtC = 1001

|MeltingPt_notes = (anhydrous) <br> 140 °C (hexahydrate)

|pKa = 4 (hexahydrate)

|MagSus = +6145.0·10⁻⁶ cm³/mol

}}

|Section3 = {{Chembox Structure

|Coordination = octahedral at Ni

|CrystalStruct = [[Monoclinic]]

}}

|Section4 = {{Chembox Thermochemistry

|DeltaHf = −316&nbsp;kJ·mol⁻¹<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>

|Entropy = 107&nbsp;J·mol⁻¹·K⁻¹<ref name=b1/>

}}

|Section5 = {{Chembox Hazards

|MainHazards = Very toxic ('''T+''')<br/>Irritant ('''Xi''')<br/>Dangerous for the environment ('''N''')<br/>Carcinogen

|ExternalSDS = [https://fscimage.fishersci.com/msds/16310.htm Fischer Scientific]

|NFPA-H = 3

|NFPA-F = 0

|NFPA-R = 0

|GHSPictograms = {{GHS06}}{{GHS08}}{{GHS09}}

|GHSSignalWord = Danger

|HPhrases = {{H-phrases|301|315|317|331|334|341|350i|360D|372|410}}

|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}}

|FlashPt = Non-flammable

|LD50 = 105 mg/kg (rat, oral)<ref>{{IDLH|7440020|Nickel metal and other compounds (as Ni)}}</ref>

}}

|Section6 = {{Chembox Related

|OtherAnions = [[Nickel(II) fluoride]]<br/>[[Nickel(II) bromide]]<br/>[[Nickel(II) iodide]]

|OtherCations = [[Palladium(II) chloride]]<br/>[[Platinum(II) chloride]]<br/>Platinum(II,IV) chloride<br/>[[Platinum(IV) chloride]]

|OtherCompounds = [[Cobalt(II) chloride]]<br/>[[Copper(II) chloride]]}}

'''Nickel(II) chloride''' (or just '''nickel chloride''') is the [[chemical compound]] NiCl₂. The [[anhydrous]] [[salt (chemistry)|salt]] is yellow, but the more familiar [[Water of crystallization|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 [[carcinogen]]ic to the lungs and nasal passages in cases of long-term [[inhalation exposure]].<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>

Large scale production and uses of nickel chloride are associated with extraction and purification of nickel from its ores. The extraction with hydrochloric acid of nickel [[matte (metallurgy)|matte]] and residues obtained from roasting refining nickel-containing ores. 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>

:{{chem2|NiS + 2 CuCl2 -> NiCl2 + 2 CuCl + S}}

:{{chem2|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.<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 [[thionyl chloride]] or by heating under a stream of HCl gas. Simply heating the hydrates does not afford the anhydrous dichloride.

:{{chem2|NiCl2*6H2O + 6 SOCl2 -> NiCl2 + 6SO2 + 12HCl}}

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 = [[Inorganic Syntheses]] | volume = 28 | pages = 321–2 | year = 1990 | doi = 10.1002/9780470132593.ch80| isbn = 9780470132593 }}</ref>

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

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

==Structure of NiCl₂ and its hydrates==

NiCl₂ adopts the [[Cadmium chloride#Structure|CdCl₂ structure]].<ref name=Wells>Wells, A. F. ''Structural Inorganic Chemistry'', Oxford Press, [[Oxford]], [[United Kingdom]], 1984.</ref> 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.<ref name=Wells/> [[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 ligand]]s. The trans sites on the octahedral centers occupied by [[aquo ligand]]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 NiCl₂·4H₂O is also known.

==Reactions==

===Coordination complexes===

[[File:Color of various Ni(II) complexes in aqueous solution.jpg|left|thumb|400px|Color of various Ni(II) complexes in aqueous solution. From left to right, [Ni(NH₃)₆]²⁺, [Ni([[ethylenediamine|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.

| [[Hexaamminenickel chloride|[Ni(NH₃)₆]Cl₂]]

| blue/violet

| paramagnetic

| octahedral

|-

| [Ni([[ethylenediamine|en]])₃]²⁺

| [[Dichloro(1,2-bis(diphenylphosphino)ethane)nickel|NiCl₂(dppe)]]

| [Ni([[cyanide|CN]])₄]²⁻

| [NiCl₄]²⁻<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 = [[Inorganic Syntheses]] | 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 = [[Acta Crystallographica]] | year = 1967 | volume = 23 | issue = 6 | pages = 1064–1070 | doi =10.1107/S0365110X67004268}}</ref>

| Yellowish-green

[[File:Hexammine nickel chloride, bulk.jpg|left|thumb|200px|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 [[acetylacetone|acetylacetonate]] complexes Ni(acac)₂(H₂O)₂ and the benzene-soluble (Ni(acac)₂)₃, which is a precursor to [[Bis(cyclooctadiene)nickel(0)|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.<ref name=dme/> The dme ligands in this complex are labile. For example, this complex reacts with [[sodium cyclopentadienide]] to give the [[sandwich compound]] [[nickelocene]].

[[Hexaamminenickel chloride|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]].<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>

::[[File:Dienol-isomerisation-2D-skeletal.png|300px|General reaction scheme for the isomerisation of dienols]]

*In combination with [[Chromium(II) chloride|CrCl₂]] for the coupling of an [[aldehyde]] and a vinylic iodide to give [[Allyl alcohol|allylic alcohols]].

*As a precursor to [[Charles Allan Brown|Brown]]'s [[nickel boride catalyst|P-1 and P-2 nickel boride catalyst]] through reaction with [[sodium borohydride|NaBH₄]].

*As a precursor to finely divided Ni by reduction with Zn, for the reduction of aldehydes, [[Alkene|alkenes]], and nitro [[Aromatic compound|aromatic compounds]]. This reagent also promotes homo-coupling reactions, that is 2RX → R-R where R = aryl, vinyl.

NiCl₂-dme (or NiCl₂-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>

[[File:NiCl2 schemes.tif|center|344x344px|thumb|Application of NiCl₂ precatalyst.]]

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" />

{{Commons category|Nickel(II) chloride}}

*[https://www.cdc.gov/niosh/npg/npgd0445.html NIOSH Pocket Guide to Chemical Hazards]

*{{nist}}

{{Chlorides}}

[[Category:Nickel compounds]]

[[Category:Coordination complexes]]