Nickel(II) chloride

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Nickel(II) chloride
Nickel(II) chloride hexahydrate
Anhydrous Nickel(II)-chloride.jpg
IUPAC name
Nickel(II) chloride
Other names
Nickelous chloride, nickel(II) salt of hydrochloric acid
7718-54-9 YesY
7791-20-0 (hexahydrate) N
ChEBI CHEBI:34887 YesY
ChemSpider 22796 YesY
EC Number 231-743-0
Jmol interactive 3D Image
KEGG C14711 YesY
PubChem 24385
RTECS number QR6480000
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)
64.2 g/100 mL (20 °C)
87.6 g/100 mL (100 °C)
254 g/100 mL (20 °C)
600 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)
octahedral at Ni
107 J·mol−1·K−1[1]
−316 kJ·mol−1[1]
Safety data sheet Fischer Scientific
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
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g., chloroform Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no codeNFPA 704 four-colored diamond
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
105 mg/kg (rat, oral)[2]
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. In general 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 are carcinogenic.

Production and syntheses[edit]

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. Heating the hexahydrate in the range 66-133.°C gives the yellowish dihydrate, NiCl2·2H2O.[3] 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.

\mathsf{NiCl_2 \cdot 6H_2O  +  6SOCl_2 \ \xrightarrow \ NiCl_2  +  6SO_2  +  12HCl }

The dehydration is accompanied by a color change from green to yellow.[4]

In case one needs a pure compound without presence of cobalt, nickel chloride can be obtained cautiously heating hexammine nickel chloride:[5]

\mathsf{[Ni(NH_3)_6]Cl_2 \ \xrightarrow{175-200^oC}\ NiCl_2 + 6NH_3 }

Structure of NiCl2 and its hydrates[edit]

NiCl2 adopts the CdCl2 structure.[6] 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 are bound to the nickel, and the remaining two are water of crystallisation.[6] 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.[7] A tetrahydrate NiCl2·4H2O is also known.


Nickel(II) chloride solutions are acidic, with a pH of around 4 due to the hydrolysis of the Ni2+ ion.

Coordination complexes[edit]

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−[8][9] 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. 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[edit]

NiCl2 and its hydrate are occasionally useful in organic synthesis.[10]

  • 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 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[edit]

Nickel chloride solutions are used for electroplating nickel onto other metal items.


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.


  1. ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22. ISBN 0-618-94690-X. 
  2. ^ "Nickel metal and other compounds (as Ni)". Immediately Dangerous to Life and Health. National Institute for Occupational Safety and Health (NIOSH). 
  3. ^ Laird G. L. Ward "Anhydrous Nickel (II) Halides and their Tetrakis(Ethanol) and 1,2-Dimethoxyethane Complexes Inorganic Syntheses, 1972 Volume 13, pages 154–164. doi:10.1002/9780470132449.ch30
  4. ^ Pray, A. P.; Tyree, S. Y.; Martin, Dean F.; Cook, James R. (1990). "Anhydrous Metal Chlorides". Inorganic Syntheses 28: 321–2. doi:10.1002/9780470132593.ch80. 
  5. ^ Karyakin, Yu.V. (1947). Pure chemicals. Manual for laboratory preparation of inorganic substances (in Russian) (Moscow, Leningrad "State Scientific Technical Publishing of Chemical Literature" ed.). p. 416. 
  6. ^ a b , Wells, A. F. Structural Inorganic Chemistry, Oxford Press, Oxford, United Kingdom, 1984.
  7. ^ B. Morosin "An X-ray diffraction study on nickel(II) chloride dihydrate" Acta Cryst. 1967. volume 23, pp. 630-634. doi:10.1107/S0365110X67003305
  8. ^ 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. 
  9. ^ 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. 
  10. ^ 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[edit]