Mercury(II) oxide

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
Names
IUPAC name
Mercury(II) oxide
Other names
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.040.580
KEGG
RTECS number
  • OW8750000
UNII
UN number 1641
Properties
HgO
Molar mass 216.591 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]
Std enthalpy of
formation (ΔfH⦵298)
 −90 kJ·mol−1[2]
Hazards
Main hazards Highly toxic
Safety data sheet ICSC 0981
GHS pictograms
NFPA 704 (fire diamond)
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 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g. white phosphorusSpecial hazards (white): no code
0
4
2
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).
N verify (what is YN ?)
Infobox references

History
See also: Maslama al-Majriti

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]

Synthesis
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]

Structure

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]

Uses

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 issues
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]

References
  1. "Mercury oxide (HgO) crystal structure, physical properties". Semiconductors · II-VI and I-VII Compounds; Semimagnetic Compounds. Landolt-Börnstein – Group III Condensed Matter. Landolt-Börnstein - Group III Condensed Matter. 41B. Springer-Verlag. 1999. pp. 1–7. doi:10.1007/b71137. ISBN 978-3-540-64964-9.
  2. Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22. ISBN 978-0-618-94690-7.
  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 978-0-306-47277-0.
  5. Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  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 978-0-534-42201-1. Mercury(II) oxide anode mercury battery.
  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.
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