Mercury(I) sulfate: Difference between revisions

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| ImageFile1 = Mercury(I)sulfate.svg

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|Section1={{Chembox Identifiers

| CASNo_Ref = {{cascite|correct|??}}

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| UNII = PI950N9DYS

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| PubChem = 24545

| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}

| ChemSpiderID = 22951

| SMILES = [O-]S(=O)(=O)[O-].[Hg+][Hg+]

| StdInChI_Ref = {{stdinchicite|changed|chemspider}}

| StdInChI = 1S/2Hg.H2O4S/c;;1-5(2,3)4/h;;(H2,1,2,3,4)/q2*+1;/p-2

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| StdInChIKey = MINVSWONZWKMDC-UHFFFAOYSA-L

|Section2={{Chembox Properties

| Solubility = 0.051 g/100 mL

| SolubleOther = soluble in dilute [[nitric acid]]

| SolubilityProduct = 6.5{{e|&minus;7}}<ref name="crc">{{cite book |author1=John Rumble |title=CRC Handbook of Chemistry and Physics |date=June 18, 2018 |publisher=CRC Press |isbn=978-1-138-56163-2 |pages=5–189|edition=99 |language=English}}</ref>

| MagSus = −123.0·10⁻⁶ cm³/mol

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|Section4={{Chembox Thermochemistry

| DeltaHf = -743.1 [[kJ/mol|kJ·mol]]

| Entropy = 200.7 J·mol·K

| HeatCapacity = 132 J·mol·K<ref name="hand2">

| last = Lide

| first = David R.

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| title = Handbook of Chemistry and Physics

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| location = Boca Raton, FL

| publisher = CRC Press

| isbn = 0-8493-0594-2

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'''Mercury(I) sulfate''', commonly called sulfate ([[]]) or mercurous ([[]]) is the [[chemical compound]] Hg₂SO₄.<ref>''Intermediate Inorganic Chemistry'' by J. W. Mellor, published by Longmans, Green and Company, London, 1941, page 388</ref> a a <ref =http://../=& ,

==Structure==

[[File:Hg2SO4chemdraw.svg|thumb|left|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.<ref>{{cite journal|title=Preparation and Characterization of Dimercury(I)Monofluorophosphate(V), Hg₂PO₃F: Crystal Structure, Thermal Behavior, Vibrational Spectra, and Solid-State ³¹P and ¹⁹F NMR Spectra|author=Matthias Weil |author2=Michael Puchberger |author3=Enrique J. Baran |journal= Inorg. Chem.|year=2004|volume=43|issue=26|pages=8330–8335|doi=10.1021/ic048741e|pmid=15606179}}</ref>

Focusing on the shorter Hg-O bonds, the Hg – Hg – O bond angle is 165°±1°.<ref>{{cite journal|author=Dorm, E.|year=1969|title=Structural Studies on Mercury(I) Compounds. VI. Crystal Structure of Mercury(I) Sulfate and Selenate|journal=Acta Chemica Scandinavica|volume=23|pages=1607–15|doi=10.3891/acta.chem.scand.23-1607|doi-access=free}}</ref><ref>{{cite journal |doi=10.1107/S1600536814011155|title=Crystal structure of Hg2SO4– a redetermination|year=2014|last1=Weil|first1=Matthias|journal=Acta Crystallographica Section E|volume=70|issue=9|pages=i44|pmid=25309168|pmc=4186147}}</ref>

==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:,<ref name="google">[https://books.google.com/books?id=VrTVAAAAMAAJ&q=%22mercury%28I%29+sulfate%22+prepared Google Books result], accessed 11 December 2010</ref><ref>''Mercurous Sulphate, cadmium sulphate, and the cadmium cell.'' by Hulett G. A. The physical review.1907. p.19.

</ref>

:{{chem2 | Hg2(NO3)2 + H2SO4 -> Hg2SO4 + 2 HNO3 }}

It can also be prepared by reacting an excess of [[mercury (element)|mercury]] with concentrated [[sulfuric acid]]:<ref name="google"/>

:{{chem2 | 2 Hg + 2 H2SO4 -> Hg2SO4 + 2 H2O + SO2}}<!--typical reaction of sulfuric acid as oxidizing agent-->

== Use in electrochemical cells==

Mercury(I) sulfate is often used in [[electrochemical cell]]s.<ref>"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</ref><ref>"Electromotive Force Studies of Cell, Cd_xHg_y | 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</ref><ref>"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</ref> It was first introduced in electrochemical cells by Latimer Clark in 1872,<ref name="GEORGE AUGUSTUS HULETT 2000, p.91-98">"George Augustus Hulett: from Liquid Crystals to Standard Cell" by John T. Stock. ''Bull. Hist. Chem.'' Volume 25, Number 2, 2000, p.91-98</ref> It was then alternatively{{clarify|date=December 2019}} used in [[Weston cell]]s made by George Augustus Hulett in 1911.<ref name="GEORGE AUGUSTUS HULETT 2000, p.91-98"/> It has been found to be a good electrode at high temperatures above 100&nbsp;°C along with silver sulfate.<ref>{{cite journal |last1=Lietzke |first1=M. H. |last2=Stoughton |first2=R. W. |title=The Behavior of the Silver—Silver Sulfate and the Mercury—Mercurous Sulfate Electrodes at High Temperatures 1 |journal=Journal of the American Chemical Society |date=November 1953 |volume=75 |issue=21 |pages=5226–5227 |doi=10.1021/ja01117a024}}{{subscription required}}</ref>

Mercury(I) sulfate has been found to decompose at high temperatures. The decomposition process is [[endothermic]], and it occurs between 335&nbsp;°C and 500&nbsp;°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.<ref>"Sulphates of Mercury and Standard Cells." by Elliott, R. B. and Hulett, G. A. ''The Journal of Physical Chemistry'' 36.7 (1932): 2083–2086.

</ref>

{{Sulfates}}

[[Category:Mercury compounds]]