Sodium peroxide
 | |
 | |
 | |
| Names | |
|---|---|
| IUPAC name
Sodium peroxide | |
| Other names
Flocool Solozone Disodium peroxide | |
| Identifiers | |
3D model (JSmol) |
|
| ChemSpider | |
| ECHA InfoCard | 100.013.828 |
| EC Number |
|
PubChem CID |
|
| RTECS number |
|
| UNII | |
| UN number | 1504 |
CompTox Dashboard (EPA) |
|
| |
| |
| Properties | |
| Na2O2 | |
| Molar mass | 77.98 g/mol |
| Appearance | white solid (pure) yellowish solid (commercial samples)[1] |
| Density | 2.805 g/cm3 |
| Melting point | 460 °C (860 °F; 733 K) (decomposes) |
| Boiling point | 657 °C (1,215 °F; 930 K) (decomposes) |
| Reacts | |
| Solubility | Soluble in acid Insoluble in base Reacts with ethanol |
| −28.10·10−6 cm3/mol | |
| Structure | |
| hexagonal | |
| Thermochemistry | |
Heat capacity (C) |
89.37 J/(mol·K) |
Std molar entropy (S⦵298) |
95 J/(mol·K)[2] |
Std enthalpy of formation (ΔfH⦵298) |
−515 kJ·mol−1[2] |
Gibbs free energy (ΔfG⦵) |
−446.9 kJ/mol |
| Hazards | |
| Occupational safety and health (OHS/OSH): | |
Main hazards |
Caustic, reacts with water and ethanol, oxidizer |
| GHS labelling: | |
  | |
| Danger | |
| H271, H314 | |
| P210, P220, P221, P260, P264, P280, P283, P301+P330+P331, P303+P361+P353, P304+P340, P305+P351+P338, P306+P360, P310, P321, P363, P370+P378, P371+P380+P375, P405, P501 | |
| NFPA 704 (fire diamond) | |
| Flash point | Non-flammable |
| Safety data sheet (SDS) | External MSDS |
| Related compounds | |
Other cations |
Lithium peroxide Potassium peroxide Rubidium peroxide Caesium peroxide |
| Sodium oxide Sodium superoxide Sodium ozonide | |
Related compounds |
Sodium hydroxide Hydrogen peroxide |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
| |
Sodium peroxide is an inorganic compound with the formula Na2O2. It is the product of sodium ignited in excess oxygen.[4] The pure substance is a white solid, but commercial and laboratory samples often appear as a yellowish solid due to sodium superoxide impurities.[1] It is a strong base. This metal peroxide exists in several hydrates and peroxyhydrates including Na2O2·2H2O2·4H2O, Na2O2·2H2O, Na2O2·2H2O2, and Na2O2·8H2O.[5] The octahydrate, which is simple to prepare, is white, like the pure anhydrous material.[6]
Properties
[edit]Sodium peroxide crystallizes with hexagonal symmetry.[7] Upon heating, the hexagonal form undergoes a transition into a phase of unknown symmetry at 512 °C.[8] With further heating above the 657 °C boiling point, the compound decomposes to Na2O, releasing O2.[9]
- 2 Na2O2 → 2 Na2O + O2
Preparation
[edit]Commercially, sodium peroxide is produced from the elements in a two-stage process. First sodium is oxidized to sodium oxide:[8][10]
- 4Na + O2 → 2 Na2O
Subsequently, this oxide is treated with more oxygen:
- 2 Na2O + O2 → 2 Na2O2
This was the method by which the substance was discovered in 1810 by Joseph Louis Gay-Lussac and Louis Jacques Thénard, as well as how it was for the first time commercially made by Hamilton Castner in the 1890s.[11]
It may also be produced by passing ozone gas over solid sodium iodide inside a platinum or palladium tube. The ozone oxidizes the sodium to form sodium peroxide. The iodine can be sublimed by mild heating. The platinum or palladium catalyzes the reaction and is not attacked by the sodium peroxide.
The octahydrate can be produced by treating sodium hydroxide with hydrogen peroxide.[6]
Uses
[edit]Sodium peroxide hydrolyzes to give sodium hydroxide and hydrogen peroxide according to the reaction[10]
- Na2O2 + 2 H2O → 2 NaOH + H2O2
Sodium peroxide was used to bleach wood pulp for the production of paper and textiles. Presently it is mainly used for specialized laboratory operations, e.g., the extraction of minerals from various ores. Sodium peroxide may go by the commercial names of Solozone[8] and Flocool.[9] In chemistry preparations, sodium peroxide is used as an oxidizing agent. It is also used as an oxygen source by reacting it with carbon dioxide to produce oxygen and sodium carbonate:
- Na2O2 + CO2 → Na2CO3 + 1⁄2 O2
- Na2O2 + H2O + 2 CO2 → 2 NaHCO3 + 1⁄2 O2
It is thus particularly useful in scuba gear, submarines, etc. Lithium peroxide and potassium superoxide have similar uses.
Sodium peroxide was once used on a large scale for the production of sodium perborate, but alternative routes to that cleaning agent have been developed.[5]
References
[edit]- 1 2 Creighton, J. A.; Lippincott, E. R. (1964). "Vibrational Frequency and Dissociation Energy of the Superoxide Ion". The Journal of Chemical Physics. 40 (6): 1779–1780. doi:10.1063/1.1725405.
- 1 2 Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A23. ISBN 978-0-618-94690-7.
- "Hazard Rating Information for NFPA Fire Diamonds". Archived from the original on 2004-09-04.
- ↑ Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. p. 98. ISBN 978-0-08-022057-4.
- 1 2 Jakob, Harald; Leininger, Stefan; Lehmann, Thomas; Jacobi, Sylvia; Gutewort, Sven (2007). "Peroxo Compounds, Inorganic". Ullmann's Encyclopedia of Industrial Chemistry. doi:10.1002/14356007.a19_177.pub2. ISBN 978-3-527-30673-2.
- 1 2 Penneman, R. A.; Toy, A. D. F. (1950). "Sodium Peroxide 8-Hydrate". Inorganic Syntheses. Vol. 3. pp. 1–4. doi:10.1002/9780470132340.ch1.
- ↑ Tallman, R. L.; Margrave, J. L.; Bailey, S. W. (1957). "The Crystal Structure Of Sodium Peroxide". J. Am. Chem. Soc. 79 (11): 2979–80. Bibcode:1957JAChS..79.2979T. doi:10.1021/ja01568a087.
- 1 2 3 Macintyre, J. E., ed. Dictionary of Inorganic Compounds, Chapman & Hall: 1992.
- 1 2 Lewis, R. J. Sax's Dangerous Properties of Industrial Materials, 10th ed., John Wiley & Sons, Inc.: 2000.
- 1 2 E. Dönges "Lithium and Sodium Peroxides" in Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 979.
- ↑ Bulletin of Pharmacy. 1895.