|Jmol-3D images||Image 1|
|Molar mass||216.59 g mol−1|
|Appearance||Yellow or red solid|
|Melting point||500 °C (932 °F; 773 K) (decomposes)|
|Solubility in water||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|
|Refractive index (nD)||2.5 (550 nm)|
|Std enthalpy of
|EU classification||Very toxic (T+)
Dangerous for the environment (N)
|R-phrases||R26/27/28, R33, R50/53|
|S-phrases||(S1/2), S13, S28, S45, S60, S61|
|LD50||18 mg/kg (oral, rat)|
|Other anions||Mercury sulfide
|Other cations||Zinc oxide
|Related compounds||Mercury(I) oxide|
|Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)|
|(what is: / ?)|
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.
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.)
The red form of HgO can be made by heating Hg in oxygen at roughly 350 °C, or by pyrolysis of Hg(NO3)2. The yellow form can be obtained by precipitation of aqueous Hg2+ with alkali. 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 O-Hg-O angle of 108°.
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. At pressures above 10 GPa both of those structures convert to a tetragonal form.
Mercury oxide is a 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.
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.
- "Mercury oxide (HgO) crystal structure, physical properties". Landolt-Börnstein – Group III Condensed Matter 41B (Springer-Verlag). 1999. pp. 1–7. doi:10.1007/b71137. ISBN 978-3-540-64964-9.
- Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22. ISBN 0-618-94690-X.
- Almqvist, Ebbe (2003). History of Industrial Gases. Springer. p. 23. ISBN 0-306-47277-5.
- Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0080379419.
- 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.
- Moore, John W.; Conrad L. Stanitski; Peter C. Jurs (2005). Chemistry: The Molecular Science. Thomson Brooks/Cole. p. 941. ISBN 0-534-42201-2.
- Chemicals Regulation Directorate. "Banned and Non-Authorised Pesticides in the United Kingdom". Retrieved 1 December 2009.
- "Mercury (II) oxide". International Occupational Safety and Health Information Centre. Retrieved 2009-06-06.
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