Sodium peroxide

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Sodium peroxide
Sodium peroxide
Sodium peroxide 2grams.jpg
Identifiers
CAS number 1313-60-6 YesY
PubChem 14803
EC number 215-209-4
UN number 1504
RTECS number WD3450000
Jmol-3D images Image 1
Properties
Molecular formula Na2O2
Molar mass 77.98 g/mol
Appearance yellow to white powder
Density 2.805 g/cm3
Melting point 460 °C (860 °F; 733 K) (decomposes)
Boiling point 657 °C (1,215 °F; 930 K) (decomposes)
Solubility in water reacts violently
Solubility soluble in acid
insoluble in alkali
Structure
Crystal structure Hexagonal
Thermochemistry
Std molar
entropy
So298
95 J·mol−1·K−1[1]
Std enthalpy of
formation
ΔfHo298
−515 kJ·mol−1[1]
Hazards
MSDS External MSDS
EU Index 011-003-00-1
EU classification Oxidising agent O Corrosive C
R-phrases R8, R35
S-phrases (S1/2), S8, S27, S39, S45
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g., chloroform Reactivity code 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g., calcium Special hazard OX: Oxidizer. E.g., potassium perchlorateNFPA 704 four-colored diamond
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
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
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Infobox references

Sodium peroxide is the inorganic compound with the formula Na2O2. This solid is the product when sodium is burned with oxygen.[2] 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.[3]

Properties[edit]

Sodium peroxide crystallizes with hexagonal symmetry.[4] Upon heating, the hexagonal form undergoes a transition into a phase of unknown symmetry at 512 °C.[5] With further heating above the 675 °C melting point, the compound decomposes to NaO, releasing O2, before reaching a boiling point.[6]

2Na2O2 → 2Na2 + O2

Sodium peroxide is hydrolyzed to give sodium hydroxide and hydrogen peroxide according to the reaction:

Na2O2 + 2 H2O → 2 NaOH +H2O2

Preparation[edit]

The synthesis is no longer of commercial significance since the development of efficient routes to hydrogen peroxide.[dubious ][3] 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:[5]

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[edit]

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[5] and Flocool.[6] 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; it is thus particularly useful in scuba gear, submarines, etc. Lithium peroxide has similar uses.

References[edit]

  1. ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A23. ISBN 0-618-94690-X. 
  2. ^ Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. p. 98. ISBN 0-08-022057-6. 
  3. ^ a b 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
  4. ^ Tallman, R. L.; Margrave, J. L.; Bailey, S. W. (1957). "The Crystal Structure Of Sodium Peroxide". J. Am. Chem. Soc. 79 (11): 2979–80. doi:10.1021/ja01568a087. 
  5. ^ a b c Macintyre, J. E., ed. Dictionary of Inorganic Compounds, Chapman & Hall: 1992.
  6. ^ a b Lewis, R. J. Sax's Dangerous Properties of Industrial Materials, 10th ed., John Wiley & Sons, Inc.: 2000.

External links[edit]