Nickel(II) sulfate

Nickel(II) sulfate
Nickel(II) sulfate


CAS number 7786-81-4 YesY, (anhydrous)
10101-97-0 (hexahydrate)
10101-98-1 (heptahydrate)
PubChem 24586
ChemSpider 22989 YesY
EC number 232-104-9
ChEBI CHEBI:53001 YesY
RTECS number QR9600000
Jmol-3D images Image 1
Molecular formula NiSO4
Molar mass 154.75 g/mol (anhydrous)
262.85 g/mol (hexahydrate)
280.86 g/mol (heptahydrate)
Appearance yellow solid (anhydrous)
green crystals (hexahydrate)
Density 3.68 g/cm3 (anhydrous)
2.07 g/cm3 (hexahydrate)
1.948 g/cm3 (heptahydrate)
Melting point

> 100 °C (anhydrous)
53 °C (hexahydrate)

Boiling point

840 °C (anhydrous, decomp)

Solubility in water 650 g/L (20°C)
775 g/L (30°C) (heptahydrate)
Solubility anhydrous
insoluble in ethanol, ether, acetone
very soluble in ethanol, ammonia
soluble in alcohol
Refractive index (nD) 1.511 (hexahydrate)
1.467 (heptahydrate)
Crystal structure cubic (anhydrous)
tetragonal (hexahydrate)
rhombohedral (hexahydrate)
MSDS External MSDS
EU Index 028-009-00-5
EU classification Carc. Cat. 1
Muta. Cat. 3
Repr. Cat. 2
Toxic (T)
Harmful (Xn)
Irritant (Xi)
Dangerous for the environment (N)
R-phrases R49, R61, R20/22, R38, R42/43, R48/23, R68, R50/53
S-phrases S53, S45, S60, S61
Flash point Non-flammable
LD50 264 mg/kg
Related compounds
Other cations Cobalt(II) sulfate
Copper(II) sulfate
 YesY sulfate (verify) (what is: YesY/N?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Nickel(II) sulfate, or just nickel sulfate, usually refers to the inorganic compound with the formula NiSO4(H2O)6. This highly soluble blue-coloured salt is a common source of the Ni2+ ion for electroplating. Nickel sulfate is paramagnetic.

At least seven sulfate salts of nickel(II) are known. These salts differ, sometimes subtly, in terms of their hydration or structural details. The common tetragonal hexahydrate crystallizes from aqueous solution between 30.7 and 53.8 °C. Below these temperatures, a heptahydrate crystallises and above these temperatures an orthorhombic hexahydrate forms. The yellow anhydrous form, NiSO4, is a high melting solid. This material produced by heating the hydrates above 330 °C. It decomposes at still higher temperatures to nickel oxide.[1]

X-ray crystallography measurements show that NiSO4·6H2O consists of the octahedral [Ni(H2O)6]2+ ions. These ions in turn are hydrogen bonded to sulfate ions.[2] Dissolution of the salt in water gives solutions containing the aquo complex [Ni(H2O)6]2+.


Production, applications, and coordination chemistry

The salt is usually obtained as a by-product of copper refining. It is also produced by dissolution of nickel metal or nickel oxides in sulfuric acid. Approximately 40,000 tonnes were produced in 2005. It is mainly used for electroplating of nickel.[1]

Aqueous solutions of nickel sulfate reacts with sodium carbonate to precipitate nickel carbonate, a precursor to nickel-based catalysts and pigments. Addition of ammonium sulfate to concentrated aqueous solutions of nickel sulfate precipitates Ni(NH4)2SO4·6H2O. This blue-coloured solid is analogous to Mohr's salt, Fe(NH4)2SO4·6H2O.[1]

Nickel sulfate is used in the laboratory. Columns used in polyhistidine-tagging, useful in biochemistry and molecular biology, are regenerated with nickel sulfate. Aqueous solutions of NiSO4·6H2O and related hydrates react with ammonia to give [Ni(NH3)6]SO4 and with ethylenediamine to give [Ni(H2NCH2CH2NH2)3]SO4. The latter is occasionally used as a calibrant for magnetic susceptibility measurements because it has no tendency to hydrate.

Natural occurrence

Nickel sulfate occurs as the rare mineral retgersite, which is a hexahydrate. The second hexahydrate is known as nickelhexahydrite (Ni,Mg,Fe)SO4·6H2O. The heptahydrate, which is relatively unstable in air, occurs as morenosite. The monohydrate occurs as very rare mineral dwornikite (Ni,Fe)SO4·H2O.


Nickel salts are carcinogenic and irritate the skin.


  1. ^ a b c K. Lascelles, L. G. Morgan, D. Nicholls, D. Beyersmann “Nickel Compounds” in Ullmann's Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim, 2005. doi:10.1002/14356007.a17_235.pub2.
  2. ^ Wells, A. F. (1984). Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.

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