Sodium peroxide
| Sodium peroxide | |
|---|---|
 |
|
 |
|
 |
|
|
Other names
Sodium dioxide
Flocool Solozone Disodium peroxide |
|
| Identifiers | |
| CAS number | 1313-60-6  |
| PubChem | 14803 |
| EC number | 215-209-4 |
| UN number | 1504 |
| RTECS number | WD3450000 |
| Properties | |
| Molecular formula | Na2O2 |
| Molar mass | 77.98 g/mol |
| Appearance | yellow to white powder |
| Density | 2.805 g/cm3 |
| Melting point |
675 °C |
| Boiling point |
decomp. |
| Solubility in water | reacts violently |
| Structure | |
| Crystal structure | Hexagonal |
| Thermochemistry | |
| Std enthalpy of formation ΔfH |
−513 kJ/mol |
| Standard molar entropy S |
95 J K−1 mol−1 |
| Hazards | |
| MSDS | External MSDS |
| EU Index | 011-003-00-1 |
| EU classification | Oxidant (O) Corrosive (C) |
| R-phrases | R8, R35 |
| S-phrases | (S1/2), S8, S27, S39, S45 |
| NFPA 704 |
 0
2
1
OX
|
| Flash point | Non-flammable |
| Related compounds | |
| Other cations | Lithium peroxide Potassium peroxide Rubidium peroxide Caesium peroxide |
| Related sodium oxides | Sodium oxide Sodium superoxide |
| Related compounds | Sodium hydroxide Hydrogen peroxide |
| (what is this?) (verify) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) |
|
| Infobox references | |
Sodium peroxide is the inorganic compound with the formula Na2O2. This solid is the product when sodium is burned.[1] It is a strong base and a potent oxidizing agent. It exists in several hydrates and peroxyhydrates including Na2O2·2H2O2·4H2O, Na2O2·2H2O, Na2O2·2H2O2, and Na2O2·8H2O.[2]
Contents |
Properties
Sodium peroxide crystallizes with hexagonal symmetry.[3] Upon heating, the hexagonal form undergoes a transition into a phase of unknown symmetry at 512 °C.[4] With further heating above the 675 °C melting point, the compound decomposes, releasing O2, before reaching a boiling point.[5]
Sodium peroxide is hydrolyzed to give sodium hydroxide and hydrogen peroxide according to the reaction:
- Na2O2 + 2 H2O → 2 NaOH + H2O2
Preparation
The synthesis is no longer of commercial significance since the development of efficient routes hydrogen peroxide.[2] Sodium peroxide formerly was prepared on a large scale by the reaction with sodium with oxygen at 130–200 °C, a process that generates sodium oxide, which in a separate stage absorbs oxygen:[4]
- 4 Na + O2 → 2 Na2O
- 2 Na2O + O2 → 2 Na2O2
More specialized routes have been developed. At ambient temperatures (0–20 °C), O2 reacts with a dilute (0.1–5.0 mole percent) sodium amalgam. 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 is freed into iodine crystals, which can be sublimed by mild heating. The platinum or palladium catalyzes the reaction and is not attacked by the sodium peroxide.
Uses
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[4] and Flocool.[5] In chemistry preparations, sodium peroxide is used as an oxidising agent. It is also used as an oxygen source by reacting it with carbon dioxide to produce oxygen and sodium carbonate; it is thus particularly useful in scuba gear, submarines, etc. Lithium peroxide has similar uses.
It is also used for the preparation of samples by peroxide fusion and later analysis by AA or ICP.
References
- ↑ Greenwood, Norman N.; Earnshaw, A. (1984), Chemistry of the Elements, Oxford: Pergamon, p. 98, ISBN 0-08-022057-6
- ↑ 2.0 2.1 Harald Jakob, Stefan Leininger, Thomas Lehmann, Sylvia Jacobi, Sven Gutewort “Peroxo Compounds, Inorganic” Ullmann's Encyclopedia of Industrial Chemistry, 2007, Wiley-VCH, Weinheim. doi:10.1002/14356007.a19_177.pub2
- ↑ Tallman, R. L.; Margrave, J. L.; Bailey, S. W. (1957). "The Crystal Structure Of Sodium Peroxide". J. Am. Chem. Soc. 79: 2979–80. doi:10.1021/ja01568a087.
- ↑ 4.0 4.1 4.2 Macintyre, J. E., ed. Dictionary of Inorganic Compounds, Chapman & Hall: 1992.
- ↑ 5.0 5.1 Lewis, R. J. Sax's Dangerous Properties of Industrial Materials, 10th ed., John Wiley & Sons, Inc.: 2000.
External links
|
|||||