Mercury(I) sulfate

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Mercury(I) sulfate
Names
IUPAC name
Mercury(I) sulfate
Other names
Mercurous sulfate
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.029.084 Edit this at Wikidata
EC Number
  • 231-993-0
UNII
  • InChI=1S/2Hg.H2O4S/c;;1-5(2,3)4/h;;(H2,1,2,3,4)/q2*+1;/p-2 ☒N
    Key: MINVSWONZWKMDC-UHFFFAOYSA-L ☒N
  • [O-]S(=O)(=O)[O-].[Hg+].[Hg+]
Properties
Hg2SO4
Molar mass 497.24 g/mol
Appearance whitish-yellow crystals
Density 7.56 g/cm3
0.051 g/100 mL (25 °C)
0.09 g/100 mL (100 °C)
6.5×10−7[1]
Solubility soluble in dilute nitric acid, Insoluble in water, Soluble in hot sulfuric acid.
−123.0·10−6 cm3/mol
Structure
monoclinic
Thermochemistry
132 J·mol−1·K−1[2]
200.7 J·mol−1·K−1
-743.1 kJ·mol−1
Related compounds
Other anions
 Mercury(I) fluoride
Mercury(I) chloride
Mercury(I) bromide
Mercury(I) iodide
Other cations
 Mercury(II) sulfate
Cadmium sulfate
Thallium(I) sulfate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Mercury(I) sulfate, commonly called mercurous sulphate (UK) or mercurous sulfate (US) is the chemical compound Hg2SO4.[3] Mercury(I) sulfate is a metallic compound that is a white, pale yellow or beige powder.[4] It is a metallic salt of sulfuric acid formed by replacing both hydrogen atoms with mercury(I). It is highly toxic; it could be fatal if inhaled, ingested, or absorbed by skin.

Structure

Simplified depiction of the structure of mercurous sulfate.

In the crystal, mercurous sulfate is made up of Hg22+ center with an Hg-Hg distance of about 2.50 Å. The SO42− anions form both long and short Hg-O bonds ranging from 2.23 to 2.93 Å.[5]

Focusing on the shorter Hg-O bonds, the Hg – Hg – O bond angle is 165°±1°.[6][7]

Preparation

One way to prepare mercury(I) sulfate is to mix the acidic solution of mercury(I) nitrate with 1 to 6 sulfuric acid solution:,[8][9]

Hg2(NO3)2 + H2SO4 → Hg2SO4 + 2 HNO3

It can also be prepared by reacting an excess of mercury with concentrated sulfuric acid:[8]

2 Hg + 2 H2SO4 → Hg2SO4 + 2 H2O + SO2

Use in electrochemical cells

Mercury(I) sulfate is often used in electrochemical cells.[10][11][12] It was first introduced in electrochemical cells by Latimer Clark in 1872,[13] It was then alternatively[clarification needed] used in Weston cells made by George Augustus Hulett in 1911.[13] It has been found to be a good electrode at high temperatures above 100 °C along with silver sulfate.[14]

Mercury(I) sulfate has been found to decompose at high temperatures. The decomposition process is endothermic, and it occurs between 335 °C and 500 °C.

Mercury(I) sulfate has unique properties that make the standard cells possible. It has a rather low solubility (about one gram per liter); diffusion from the cathode system is not excessive; and it is sufficient to give a large potential at a mercury electrode.[15]

References

  1. ^ John Rumble (June 18, 2018). CRC Handbook of Chemistry and Physics (99 ed.). CRC Press. pp. 5–189. ISBN 978-1-138-56163-2.
  2. ^ Lide, David R. (1998), Handbook of Chemistry and Physics (87 ed.), Boca Raton, FL: CRC Press, pp. 5–19, ISBN 0-8493-0594-2
  3. ^ Intermediate Inorganic Chemistry by J. W. Mellor, published by Longmans, Green and Company, London, 1941, page 388
  4. ^ "Mercurous Sulfate | 7783-36-0".
  5. ^ Matthias Weil; Michael Puchberger; Enrique J. Baran (2004). "Preparation and Characterization of Dimercury(I)Monofluorophosphate(V), Hg2PO3F: Crystal Structure, Thermal Behavior, Vibrational Spectra, and Solid-State 31P and 19F NMR Spectra". Inorg. Chem. 43 (26): 8330–8335. doi:10.1021/ic048741e. PMID 15606179.
  6. ^ Dorm, E. (1969). "Structural Studies on Mercury(I) Compounds. VI. Crystal Structure of Mercury(I) Sulfate and Selenate". Acta Chemica Scandinavica. 23: 1607–15. doi:10.3891/acta.chem.scand.23-1607.
  7. ^ Weil, Matthias (2014). "Crystal structure of Hg2SO4– a redetermination". Acta Crystallographica Section E. 70 (9): i44. doi:10.1107/S1600536814011155. PMC 4186147. PMID 25309168.
  8. ^ a b Google Books result, accessed 11 December 2010
  9. ^ Mercurous Sulphate, cadmium sulphate, and the cadmium cell. by Hulett G. A. The physical review.1907. p.19.
  10. ^ "Influence of Microstucture on the Charge Storage Properties of Chemically Synthesized Manganese Dioxide" by Mathieu Toupin, Thiery Brousse, and Daniel Belanger. Chem. Mater. 2002, 14, 3945–3952
  11. ^ "Electromotive Force Studies of Cell, CdxHgy | CdSO4,(m) I Hg2SO4, Hg, in Dioxane-Water Media" by Somesh Chakrabarti and Sukumar Aditya. Journal of Chemical and Engineering Data, Vol.17, No. 1, 1972
  12. ^ "Characterization of Lithium Sulfate as an Unsymmetrical-Valence Salt Bridge for the Minimization of Liquid Junction Potentials in Aqueous – Organic Solvent Mixtures" by Cristiana L. Faverio, Patrizia R. Mussini, and Torquato Mussini. Anal. Chem. 1998, 70, 2589–2595
  13. ^ a b "George Augustus Hulett: from Liquid Crystals to Standard Cell" by John T. Stock. Bull. Hist. Chem. Volume 25, Number 2, 2000, p.91-98
  14. ^ Lietzke, M. H.; Stoughton, R. W. (November 1953). "The Behavior of the Silver—Silver Sulfate and the Mercury—Mercurous Sulfate Electrodes at High Temperatures 1". Journal of the American Chemical Society. 75 (21): 5226–5227. doi:10.1021/ja01117a024.(subscription required)
  15. ^ "Sulphates of Mercury and Standard Cells." by Elliott, R. B. and Hulett, G. A. The Journal of Physical Chemistry 36.7 (1932): 2083–2086.
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