Mercury(I) sulfate

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Mercury(I) sulfate
IUPAC name
Mercury(I) sulfate
Other names
Mercurous sulfate
3D model (JSmol)
ECHA InfoCard 100.029.084
EC Number 231-993-0
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)
Solubility soluble in dilute nitric acid, Insoluble in water, Soluble in hot sulfuric acid.
−123.0·10−6 cm3/mol
132 J·mol−1·K−1[1]
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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Infobox references

Mercury(I) sulfate, commonly called mercurous sulphate (UK) or mercurous sulfate (US) is the chemical compound Hg2SO4.[2] Mercury(I) sulfate is a metallic compound that is white, pale yellow or beige powder.[3] 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.


The crystal structure of mercurous sulfate is made up of Hg22+ dumbbells and SO42− anions as main building units. Hg22+ dumbbell is surrounded by four Oxygen atoms with Hg₋O distance ranging from 2.23 to 2.93 Å, whereas Hg-Hg distance is approximately 2.500Å.[4] Studies have shown mercury(I) sulfate to have the mercury atoms arranged in doublets with a bond distance of 2.500Å. The metal atom doublets are oriented parallel to the an axis in a unit cell. Mercury doublets form part of infinite chain SO4 - Hg - Hg - SO4 - Hg - Hg - … The Hg - Hg - O bond angle is 165°±1 . The chain crosses the unit cell diagonally. The mercury sulfate structure is held together by weak Hg-O interactions. The SO4 does not act as a single anion, but rather coordinated to the mercury metal.[5]


One way to prepare mercury(I) sulfate is to mix the acidic solution of mercury(I) nitrate with 1 to 6 sulfuric acid solution:,[6][7]

Hg2(NO3)2 + H2SO4 → Hg2SO4 + 2 HNO3

It can also be prepared by reacting an excess of mercury with concentrated sulfuric acid:[6]

2 Hg + 2 H2SO4 → Hg2SO4 + 2 H2O + SO2

Use in Electrochemical Cells[edit]

Mercury(I) Sulfate is often used in electrochemical cells,[8][9][10] It was first introduced in electrochemical cells by Latimer Clark in 1872,[11] It was then alternatively used in Weston Cells mady by George Augustus Hulett in 1911.[11] It has been found to be a good electrode at high temperatures above 100 °C along with Silver Sulfate.[12] Mercury(I) Sulfate was found to decompose at high temperatures. The decomposition process is endothermic and occurs between 335 and 500. Mercury(I) sulfate has unique properties that make the standard cells possible. It has a rather small solubility (about a gram per liter) that diffusion from the cathode system is not excessive, and it is sufficient to give a large potential at a mercury electrode.[13]


  1. ^ Lide, David R. (1998), Handbook of Chemistry and Physics (87 ed.), Boca Raton, FL: CRC Press, pp. 5–19, ISBN 0-8493-0594-2 
  2. ^ Intermediate Inorganic Chemistry by J. W. Mellor, published by Longmans, Green and Company, London, 1941, page 388
  3. ^
  4. ^ Preparation and Characterization of Dimercury(I)Monofluorophosphate(V), Hg2PO3F: Crystal Structure, Thermal Behavior, Vibrational Spectra, and Solid-State 31P and 19F NMR Spectra by Matthias Weil, Michael Puchberger, and Enrique J. Baran, published by Inorg. Chem. 2004, 43. pages 8330-8335
  5. ^ Dorm, E. 1969. Structural studies on mercury(I) compounds. VI. Crystal structure of mercury(I) sulfate and selenate. Acta Chemica Scandinavica (1947-1973) 23:1607–15.
  6. ^ a b Google Books result, accessed 11 December 2010
  7. ^ Mercurous Sulphate, cadmium sulphate, and the cadmium cell. by Hulett G. A. The physical review.1907. p.19.
  8. ^ 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
  9. ^ 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
  10. ^ 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
  11. ^ 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
  12. ^ The Behavior of the Silver—Silver Sulfate and the Mercury—Mercurous Sulfate Electrodes at High Temperatures by M. H. Lietzke and R. W. Stoughton. J. Am. Chem. Soc., 1953, 75 (21), pp 5226–5227 DOI: 10.1021/ja01117a024
  13. ^ Sulphates of Mercury and Standard Cells. by Elliott, R. B. and Hulett, G. A. The Journal of Physical Chemistry 36.7 (1932): 2083-2086.
Salts and esters of the sulfate ion
H2SO4 He
Li2SO4 BeSO4 B esters
MgSO4 Al2(SO4)3
Si P SO42−
Cl Ar
CaSO4 Sc2(SO4)3 Ti(SO4)2
ZnSO4 Ga2(SO4)3 Ge As Se Br Kr
SrSO4 Y2(SO4)3 Zr(SO4)2 Nb Mo Tc Ru Rh PdSO4 Ag2SO4 CdSO4 In2(SO4)3 SnSO4 Sb2(SO4)3 Te I Xe
Cs2SO4 BaSO4   Hf Ta W Re Os Ir Pt Au Hg2SO4
PbSO4 Bi2(SO4)3 Po At Rn
Fr Ra   Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og
La Ce2(SO4)3
Pr2(SO4)3 Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb2(SO4)3 Lu
Ac Th Pa U(SO4)2
Np Pu Am Cm Bk Cf Es Fm Md No Lr