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.[1]
Large scale production and uses of nickel chloride are associated with the purification of nickel from its ores. It is generated upon extraction nickel matte 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:[2]
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.[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.
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 by cautiously heating hexaamminenickel chloride:[5]
Structure of hydrated nickel chloride based on X-ray crystallography. Color code: red = O, green = Cl|thumb|leftNiCl2 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-[NiCl<sub>2</sub>(H<sub>2</sub>O)<sub>4</sub>] 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 [NiCl<sub>2</sub>(H<sub>2</sub>O)<sub>4</sub>]·2H2O.[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.
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(NH<sub>3</sub>)<sub>6</sub>]Cl2 | blue/violet | paramagnetic | octahedral | |
| [Ni([[ethylenediamine|en]])3]2+ | violet | paramagnetic | octahedral | |
| NiCl2(dppe) | orange | diamagnetic | square planar | |
| [Ni([[cyanide|CN]])4]2− | colorless | diamagnetic | square planar | |
| [NiCl<sub>4</sub>]2−[8] [9] | Yellowish-green | paramagnetic | tetrahedral |
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.[3] The dme ligands in this complex are labile.
NiCl2 and its hydrate are occasionally useful in organic synthesis.[10]
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.[11]
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.