Magnesium peroxide
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| IUPAC name
Magnesium peroxide
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| Other names
Magnesium dioxide, magnesium bioxide, UN 1476
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3D model (JSmol)
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PubChem CID
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| Properties | |
| MgO2 | |
| Molar mass | 56.3038 g/mol |
| Hazards | |
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 | |
Editor-In-Chief: C. Michael Gibson, M.S., M.D. [1]
Overview
Magnesium peroxide (MgO2) is an odorless fine powder peroxide with a white to off-white color. It is similar to calcium peroxide because magnesium peroxide also releases oxygen by breaking down at a controlled rate with water. Commercially, magnesium peroxide often exists as a compound of magnesium peroxide and magnesium hydroxide.
Structure
O2, similarly to N2, has the ability to bind either side-on or end-on. The structure of MgO2 has been calculated as a triangular shape with the O2 molecule binding side-on to the magnesium. This arrangement is a result of the Mg+ donating charge to the oxygen and creating a Mg2+O22−. The bond between to O2 and the magnesium atom has an approximate dissociation energy of 90 kJ mol−1.[1]
In the solid state, MgO2 has a cubic pyrite-type crystal structure with 6-coordinate Mg2+ ions and O22− peroxide-groups, according to experimental data [2] and evolutionary crystal structure prediction,[3] the latter predicting a phase transition at the pressure of 53 GPa to a tetragonal structure with 8-coordinate Mg2+ ions. While at normal conditions MgO2 is a metastable compound (less stable than MgO+1/2O2), at pressures above 116 GPa it is predicted to become thermodynamically stable in the tetragonal phase.
Synthesis
MgO2 can be produced by mixing MgO with hydrogen peroxide to create magnesium peroxide and water. This being an exothermic reaction should be cooled and kept around 30–40 degrees Celsius. It is also important to remove as much iron from the reaction environment as possible due to iron's ability to catalyze the degradation of the peroxide. The addition of oxygen stabilizers such as sodium silicate can also be used to help prevent the premature degradation of the peroxide. Regardless, a good yield from this reaction is only about 35%.[4]
- MgO + H2O2 --> MgO2 + H2O
High yields are further complicated by the fact that MgO2 reacts with water to degrade the peroxide into magnesium hydroxide, also known as milk of magnesia.
Applications
Magnesium peroxide is a stable oxygen releasing compound, which is used in agricultural and environmental industries. It is used to reduce contaminant levels in groundwater. Magnesium peroxide is used in the bioremediation of contaminated soil and can improve the soil quality for plant growth and metabolism. It also used in the aquaculture industry for bioremediation.
For sanitation purposes magnesium peroxide is often used as a source of oxygen for aerobic organisms in the treatment and disposal of biological waste. Since the breakdown of hydrocarbons in soil is usually quicker in aerobic conditions, MgO2 can also added to compost piles or in soil to speed up the microbe activities and to reduce the odors produced in the process.[5][6]
In certain circumstances MgO2 has also been shown to inhibit growth of bacteria. In particular, the growth of sulfate-reducing bacteria can be inhibited in an environment containing magnesium peroxide. While the oxygen slowly dissociates, it is theorized that it may then act to displace the sulfate that normally acts as the terminal electron acceptor in their electron transport chain.[7]
Toxicity
Magnesium peroxide is an irritant that can cause redness, itching, swelling, and may burn the skin and eyes on contact. Inhalation can also cause irritation to the lungs, nose, and throat, as well as causing coughing. Long term exposure may lead to lung damage, shortness of breath, and tightening of the chest. Ingestion of MgO2 can cause numerous adverse effects including: bloating, belching, abdominal pain, irritation of the mouth and throat, nausea, vomiting, and diarrhea.[8][9]
Environmentally, Magnesium peroxide is not a naturally occurring compound and is not known to persist in the environment for prolonged times, in its complete state, or to bio-accumulate. The natural degradation of MgO2 leads to Magnesium hydroxide, O2, and H2O. If spilled MgO2 should be contained and isolated from any waterways, sewer drains, and it should be isolated from combustible materials or chemicals including paper, cloth, and wood[6].
Common Environmental Reactions
Magnesium exists in the upper atmosphere in a variety of different molecular forms. Due to its ability to react with common oxygen and simple carbon-oxygen compounds the magnesium may exist in oxidized compounds including MgO2, OMgO2, MgO, and O2MgO2.[10]
- MgCO3 + O → MgO2 + CO2
- OMgO2 + O → MgO2 + O2
- MgO + O3 → MgO2 + O2
- MgO2 + O2 → O2MgO2
- MgO2 + O → MgO + O2
In contact with water it decomposes by the reactions:
- MgO2 + 2 H2O → Mg(OH)2 + H2O2
- 2 H2O2 → 2 H2O + O2
References
- ↑ Plowright, Richard J.; Thomas J. McDonnell; Timothy G. Wright; John M. C. Plane (28 July 2009). "Theoretical Study of Mg+−X and [X−Mg−Y]+Complexes Important in the Chemistry of Ionospheric Magnesium (X, Y = H2O, CO2, N2, O2, and O)". Journal of Physical Chemistry. 113 (33): 9354–9364. doi:10.1021/jp905642h. PMID 19637880.
- ↑ Vannerberg N. (1959). Ark. Kemi. 14: 99–105. Missing or empty
|title=(help) - ↑ Zhu, Qiang; Oganov A.R.; Lyakhov A.O. (2013). "Novel stable compounds in the Mg-O system under high pressure" (PDF). Phys. Chem. Chem. Phys. 15: 7696–7700. doi:10.1039/c3cp50678a.
- ↑ Shand, Mark A. (2006). THE CHEMISTRY AND TECHNOLOGY OF MAGNESIA (PDF). Hoboken, NJ: John Wiley & Sons. ISBN 978-0-471-65603-6.
- ↑ Vidali, M. (2001). "Bioremediation. An overview" (PDF). Pure Appl. Chem. 73 (7): 1163–1172. doi:10.1351/pac200173071163. Retrieved 2012-04-18.
- ↑ Rockne, Karl J.; Krishna R. Reddy (October 2003). "BIOREMEDIATION OF CONTAMINATED SITES" (PDF). International e-Conference on Modern Trends in Foundation Engineering: Geotechnical Challenges and Solutions. Retrieved 2012-04-18.
- ↑ Chang, Yu-Jie (2008). "The use of magnesium peroxide for the inhibition of sulfate-reducing bacteria under anoxic conditions". J Ind Microbiol Biotechnol: 1481–1491. Unknown parameter
|coauthors=ignored (help) - ↑ "Product Safety Summary: Magnesium Peroxide" (PDF). Solvay America Inc. Retrieved 25 April 2012.
- ↑ Pohanish, Richard P. (2011). Sittig's Handbook of Toxic and Hazardous Chemicals and Carcinogens. William Andrew. pp. 1645–1646. ISBN 1437778704.
- ↑ Plane, John M. C.; Charlotte L. Whalley. "A New Model for Magnesium Chemistry in the Upper Atmosphere". Journal of Physical Chemistry A.
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