Mercury(II) sulfate
Mercury(II) sulfate
Names
Other names
Mercuric sulfate, Mercurypersulfate, Mercury Bisulfate[1]
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.029.083
EC Number
  • 231-992-5
RTECS number
  • OX0500000
UNII
UN number 1645
  • InChI=1S/Hg.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2
    Key: DOBUSJIVSSJEDA-UHFFFAOYSA-L
  • [O-]S(=O)(=O)[O-].[Hg+2]
Properties
HgSO4
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
Structure
rhombic
Thermochemistry
−707.5 kJ mol−1[3]
Hazards
GHS labelling:[4]
GHS06: ToxicGHS08: Health hazardGHS09: Environmental hazard
Danger
H300, H310, H330, H373, H410
P260, P262, P264, P270, P271, P273, P280, P284, P301+P316, P302+P352, P304+P340, P316, P319, P320, P321, P330, P361+P364, P391, P403+P233, P405, P501
NFPA 704 (fire diamond)
NFPA 704 four-colored diamond
3
0
1
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 salt that forms white granules or crystalline powder. In water, it separates into an insoluble basic sulfate with a yellow color and sulfuric acid.[3]

Structure

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.[5] 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.[6]

History

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.[7]

Production

Mercury sulfate can be produced by treating mercury with hot concentrated sulfuric acid:[8]

Hg + 2 H2SO4 → HgSO4 + SO2 + 2 H2O

Alternatively yellow mercuric oxide reacts also with concentrated sulfuric acid.[9]

Uses

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.[10]

Hydration reactions

Mercury compounds such as mercury sulfate and mercury(II) acetate are commonly used as catalysts in oxymercuration-demercuration, a type of electrophilic addition reaction which results in hydration of an unsaturated compound. The hydration of an alkene results in an alcohol that follows the regioselectivity predicted by Markovnikov's rule. For an alkyne, the result is an enol, which tautomerizes to give a ketone.[11] An example is the conversion of 2,5-dimethyhexyn-2,5-diol to 2,2,5,5-tetramethyltetrahydrofuran using aqueous mercury sulfate without the addition of acid.[12]

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

As previously mentioned, HgSO4 was used as the catalyst for the production of acetaldehyde from acetylene and water.[13]

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.[14]

References

  1. 1 2 "Chemicalbook". Retrieved 2 May 2011.
  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. 1 2 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. ↑ "Mercuric sulfate". pubchem.ncbi.nlm.nih.gov.
  5. ↑ 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. Bibcode:1980ZK....153..121A. doi:10.1524/zkri.1980.0011.
  6. ↑ 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.
  7. ↑ Minamata Disease Archived 13 November 2019 at the Wayback Machine. Boston University. Retrieved 2016-11-10.
  8. ↑ 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.
  9. ↑ 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.
  10. ↑ Marks, E. M.; Lipkin, D. (1939). "Reaction of Aliphatic Ethers with Denigés' Reagent". J. Org. Chem. 3 (6): 598–602. doi:10.1021/jo01223a008.
  11. ↑ Clayden, Jonathan; Greeves, Nick; Warren, Stuart (2012). Organic chemistry (2nd ed.). Oxford: Oxford university press. pp. 444–445. ISBN 978-0-19-927029-3.
  12. ↑ Wasacz, J. P.; Badding, V. G. (1982). "A hydration of an alkyne illustrating steam and vacuum distillation". Journal of Chemical Education. 59 (8): 694. Bibcode:1982JChEd..59..694W. doi:10.1021/ed059p694.
  13. ↑ 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.
  14. ↑ Hubbard (24 June 1846). "On the Effects of the Sub-Sulphate of Mercury, (Turpeth Mineral,) As An Emetic". BMJ. s1-10 (25): 288–289. doi:10.1136/bmj.s1-10.25.288. ISSN 0959-8138. PMC 2559525. PMID 20794001.
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