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Mercury(II) oxide, also called mercuric oxide or simply mercury oxide, has a formula of HgO. It has a red or orange color. Mercury(II) oxide is a solid at room temperature and pressure. The mineral form montroydite is very rarely found.

Mercury(II) oxide
Mercury(II) oxide
Mercury(II) oxide
Names
IUPAC name
Mercury(II) oxide
Other names
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.040.580
KEGG
RTECS number OW8750000
UN number 1641
Properties
HgO
Molar mass 216.59 g·mol−1
Appearance Yellow or red solid
Odor odorless
Density 11.14 g/cm3
Melting point 500 °C (932 °F; 773 K) (decomposes)
0.0053 g/100 mL (25 °C)
0.0395 g/100 mL (100 °C)
Solubility insoluble in alcohol, ether, acetone, ammonia
Band gap 2.2 eV[1]
−44.0·10−6 cm3/mol
2.5 (550 nm)[1]
Structure
orthorhombic
Thermochemistry
70 J·mol−1·K−1[2]
−90 kJ·mol−1[2]
Hazards
Safety data sheet ICSC 0981
Very toxic (T+)
Dangerous for the environment (N)
R-phrases (outdated) R26/27/28, R33, R50/53
S-phrases (outdated) (S1/2), S13, S28, S45, S60, S61
NFPA 704
Flammability code 0: Will not burn. E.g., waterHealth code 4: Very short exposure could cause death or major residual injury. E.g., VX gasReactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogenSpecial hazards (white): no codeNFPA 704 four-colored diamond
0
4
0
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
18 mg/kg (oral, rat)[3]
Related compounds
Other anions
Mercury sulfide
Mercury selenide
Mercury telluride
Other cations
Zinc oxide
Cadmium oxide
Related compounds
Mercury(I) oxide
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

Contents

HistoryEdit

In 1774, Joseph Priestley discovered that oxygen was released by heating mercuric oxide, although he did not identify the gas as oxygen (rather, Priestley called it "dephlogisticated air," as that was the paradigm that he was working under at the time).[4]

SynthesisEdit

 
Montroydite structure (red atoms are oxygens)
 
Cinnabar structure

The red form of HgO can be made by heating Hg in oxygen at roughly 350 °C, or by pyrolysis of Hg(NO3)2.[5] The yellow form can be obtained by precipitation of aqueous Hg2+ with alkali.[5] The difference in color is due to particle size, both forms have the same structure consisting of near linear O-Hg-O units linked in zigzag chains with an Hg-O-Hg angle of 108°.[5]

StructureEdit

Under atmospheric pressure mercuric oxide has two crystalline forms: one is called montroydite (orthorhombic, 2/m 2/m 2/m, Pnma), and the second is analogous to the sulfide mineral cinnabar (hexagonal, hP6, P3221); both are characterized by Hg-O chains.[6] At pressures above 10 GPa both structures convert to a tetragonal form.[1]

UsesEdit

HgO is sometimes used in the production of mercury as it decomposes quite easily. When it decomposes, oxygen gas is generated.

It is also used as a material for cathodes for mercury batteries.[7]

Health issuesEdit

 
The label on an HgO powder bottle.

Mercury oxide is a highly toxic substance which can be absorbed into the body by inhalation of its aerosol, through the skin and by ingestion. The substance is irritating to the eyes, the skin and the respiratory tract and may have effects on the kidneys, resulting in kidney impairment. In the food chain important to humans, bioaccumulation takes place, specifically in aquatic organisms. The substance is banned as a pesticide in the EU.[8]

Evaporation at 20 °C is negligible. HgO decomposes on exposure to light or on heating above 500 °C. Heating produces highly toxic mercury fumes and oxygen, which increases the fire hazard. Mercury(II) oxide reacts violently with reducing agents, chlorine, hydrogen peroxide, magnesium (when heated), disulfur dichloride and hydrogen trisulfide. Shock-sensitive compounds are formed with metals and elements such as sulfur and phosphorus.[9]

ReferencesEdit

  1. ^ a b c "Mercury oxide (HgO) crystal structure, physical properties". Semiconductors · II-VI and I-VII Compounds; Semimagnetic Compounds. Landolt-Börnstein – Group III Condensed Matter. 41B. Springer-Verlag. 1999. pp. 1–7. doi:10.1007/b71137. ISBN 978-3-540-64964-9. 
  2. ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22. ISBN 0-618-94690-X. 
  3. ^ Chambers, Michael. "ChemIDplus - 21908-53-2 - UKWHYYKOEPRTIC-UHFFFAOYSA-N - Mercuric oxide [ISO] - Similar structures search, synonyms, formulas, resource links, and other chemical information". chem.sis.nlm.nih.gov. 
  4. ^ Almqvist, Ebbe (2003). History of Industrial Gases. Springer. p. 23. ISBN 0-306-47277-5. 
  5. ^ a b c Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9. 
  6. ^ Aurivillius, Karin; Carlsson, Inga-Britt; Pedersen, Christian; Hartiala, K.; Veige, S.; Diczfalusy, E. (1958). "The Structure of Hexagonal Mercury(II)oxide". Acta Chemica Scandinavica. 12: 1297–1304. doi:10.3891/acta.chem.scand.12-1297. Retrieved November 17, 2010. 
  7. ^ Moore, John W.; Conrad L. Stanitski; Peter C. Jurs (2005). Chemistry: The Molecular Science. Thomson Brooks/Cole. p. 941. ISBN 0-534-42201-2. 
  8. ^ Chemicals Regulation Directorate. "Banned and Non-Authorised Pesticides in the United Kingdom". Retrieved 1 December 2009. 
  9. ^ "Mercury (II) oxide". International Occupational Safety and Health Information Centre. Retrieved 2009-06-06. 

External linksEdit