Chemistry:Mercury(II) sulfate

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Mercury(II) sulfate
Mercury(II) sulfate
Other names
Mercuric sulfate, Mercurypersulfate, Mercury Bisulfate[1]
Molar mass 296.653 g/mol
Appearance white monoclinic crystals
Odor odorless
Density 6.47 g/cm3, solid
450 °C (dec.)[2]
Decomposes in water to yellow mercuric subsulfate and sulfuric acid
Solubility soluble in hot H2SO4, NaCl solution
insoluble in alcohol, acetone, ammonia
−78.1·10−6 cm3/mol
−707.5 kJ mol−1[3]
NFPA 704 (fire diamond)
Flammability code 0: Will not burn. E.g. waterHealth code 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasReactivity code 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g. calciumSpecial hazards (white): no codeNFPA 704 four-colored diamond
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

Mercury(II) sulfate, commonly called mercuric sulfate, is the chemical compound HgSO4. It is an odorless solid that forms white granules or crystalline powder. In water, it separates into an insoluble sulfate with a yellow color and sulfuric acid.[3]


Portion of structure of HgSO4 illustrating the distorted tetrahedral geometry at Hg (dark blue spheres).

The anhydrous compound features Hg2+ in a highly distorted tetrahedral HgO4 environment. Two Hg-O distances are 2.22 Å and the others are 2.28 and 2.42 Å.[4] In the monohydrate, Hg2+ adopts a linear coordination geometry with Hg-O (sulfate) and Hg-O (water) bond lengths of 2.179 and 2.228 Å, respectively. Four weaker bonds are also observed with Hg---O distances >2.5 Å.[5]


In 1932, the Japanese chemical company Chisso Corporation began using mercury sulfate as the catalyst for the production of acetaldehyde from acetylene and water. Though it was unknown at the time, methylmercury is formed as side product of this reaction. Exposure and consumption of the mercury waste products, including methylmercury, that were dumped into Minamata Bay by Chisso are believed to be the cause of Minamata disease in Minamata, Japan .[6]


Mercury sulfate can be produced by treating mercury with hot concentrated sulfuric acid:[7]

Hg + 2 H2SO4 → HgSO4 + SO2 + 2 H2O

Alternatively yellow mercuric oxide reacts also with concentrated sulfuric acid.[8]


Denigés' reagent

An acidic solution of mercury sulfate is known as Denigés' reagent. It was commonly used throughout the 20th century as a qualitative analysis reagent. If Denigés' reagent is added to a solution containing compounds that have tertiary alcohols, a yellow or red precipitate will form.[9]

Production of acetaldehyde

As previously mentioned, Hg S O4 was used as the catalyst for the production of acetaldehyde from acetylene and water.[10]

Oxymercuration-demercuration of alkenes

Mercury Compounds such as mercury sulfate and mercury(II) acetate are commonly used as catalysts in the oxymercuration-demercuration, a type of Electrophilic Addition reaction. The hydration of an alkene results in an alcohol that follows regioselectivity that is predicted by Markovnikov's Rule.

Hydration of alkynes

The reaction scheme is provided below. The conversion of 2,5-dimethyhexyn-2,5-diol to 2,2,5,5-tetramethylte-trahydrofuran using aqueous mercury sulfate without the addition of acid.[11]

Conversion of 2,5-dimethyhexyn-2,5-diol to 2,2,5,5-tetramethylte-trahydrofuran-3-one

Health issues

Inhalation of HgSO4 can result in acute poisoning: causing tightness in the chest, difficulties breathing, coughing and pain. Exposure of HgSO4 to the eyes can cause ulceration of conjunctiva and cornea. If mercury sulfate is exposed to the skin it may cause sensitization dermatitis. Lastly, ingestion of mercury sulfate will cause necrosis, pain, vomiting, and severe purging. Ingestion can result in death within a few hours due to peripheral vascular collapse.[1]

It was used in the late 19th century to induce vomiting for medical reasons. [1]


  1. 1.0 1.1 "Chemicalbook". 
  2. Wu, Shengji; Uddin, Md. Azhar; Nagano, Saori; Ozaki, Masaki; Sasaoka, Eiji (2011). "Fundamental Study on Decomposition Characteristics of Mercury Compounds over Solid Powder by Temperature-Programmed Decomposition Desorption Mass Spectrometry". Energy & Fuels 25 (1): 144–153. doi:10.1021/ef1009499. 
  3. 3.0 3.1 Lide, David R. (1998). Handbook of Chemistry and Physics (87 ed.). Boca Raton, Florida: CRC Press. pp. 5–19. ISBN 0-8493-0594-2. 
  4. Aurivillius, Karin; Stålhandske, Claes (1980). "A Reinvestigation of the Crystal Structures of HgSO4 and CdSO4". Zeitschrift für Kristallographie - Crystalline Materials 153 (1–2): 121–129. doi:10.1524/zkri.1980.0011. Bibcode1980ZK....153..121A. 
  5. Stålhandske, C. (1980). "An X-ray and Neutron Diffraction Study of Mercury(II) Sulphate Monohydrate". Acta Crystallographica Section B Structural Crystallography and Crystal Chemistry 36: 23–26. doi:10.1107/s0567740880002361. 
  6. Minamata Disease. Boston University. Retrieved 2016-11-10.
  7. Simon, Matthias; Jönk, Peter; Wühl-Couturier, Gabriele; Halbach, Stefan (2006). "Mercury, Mercury Alloys, and Mercury Compounds". Ullmann's Encyclopedia of Industrial Chemistry. doi:10.1002/14356007.a16_269.pub2. ISBN 3527306730. 
  8. Robey, R. F.; Robertson, N. C. (May 1947). "Test for tert-Butyl and Isopropyl Alcohols with Deniges Reagent". Analytical Chemistry 19 (5): 310–311. doi:10.1021/ac60005a007. 
  9. Marks, E. M.; Lipkin, D. (1939). "Reaction of Aliphatic Ethers with Denigés' Reagent". J. Org. Chem. 3 (6): 598–602. doi:10.1021/jo01223a008. 
  10. Vogt, R; Nieuwland, J (September 1921). "The role of mercury salts in the catalytic transformation of acetylene into acetaldehyde, and a new commercial process for manufacture of paraaldehyde". J. Am. Chem. Soc. 43 (9): 2071–81. doi:10.1021/ja01442a010. 
  11. Wasacz, J. P.; Badding, V. G. (1982). "A hydration of an alkyne illustrating steam and vacuum distillation". Journal of Chemical Education 59 (8): 694. doi:10.1021/ed059p694. Bibcode1982JChEd..59..694W. 

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