Sodium peroxide

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Sodium peroxide
Sodium peroxide
Sodium-peroxide-3D-vdW.png
Sodium peroxide 2grams.jpg
Names
Other names
Sodium dioxide
Flocool
Solozone
Disodium peroxide
Identifiers
1313-60-6 YesY
EC number 215-209-4
Jmol-3D images Image
PubChem 14803
RTECS number WD3450000
UN number 1504
Properties
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)
reacts violently
Solubility soluble in acid
insoluble in alkali
reacts with ethanol
Structure
hexagonal
Thermochemistry
89.37 J/mol K
95 J·mol−1·K−1[1]
−515 kJ·mol−1[1]
-446.9 kJ/mol
Hazards
Safety data sheet External MSDS
EU classification Oxidizing 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 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas Reactivity code 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g., phosphorus Special hazard W+OX: Reacts with water in an unusual or dangerous manner AND is oxidizer.NFPA 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 otherwise noted, 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 yellowish solid is the product of sodium ignited in oxygen.[3] It is a strong base. It exists in several hydrates and peroxyhydrates including Na2O2·2H2O2·4H2O, Na2O2·2H2O, Na2O2·2H2O2, and Na2O2·8H2O.[4]

Properties[edit]

Sodium peroxide crystallizes with hexagonal symmetry.[5] Upon heating, the hexagonal form undergoes a transition into a phase of unknown symmetry at 512 °C.[6] With further heating above the 675 °C melting point, the compound decomposes to Na2O, releasing O2, before reaching a boiling point.[7]

2 Na2O2 → 2 Na2O + O2

Preparation[edit]

Sodium peroxide can be prepared on a large scale by the reaction of metallic sodium with oxygen at 130–200 °C, a process that generates sodium oxide, which in a separate stage absorbs oxygen:[6]

4 Na + O2 → 2 Na2O
2 Na2O + O2 → 2 Na2O2

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]

On contact with water sodium peroxide is hydrolyzed to give sodium hydroxide and hydrogen peroxide according to the reaction:

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[6] and Flocool.[7] 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. ^ http://www.nmsu.edu/safety/programs/chem_safety/NFPA-ratingS-Z.htm
  3. ^ Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. p. 98. ISBN 0-08-022057-6. 
  4. ^ 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
  5. ^ 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. 
  6. ^ a b c Macintyre, J. E., ed. Dictionary of Inorganic Compounds, Chapman & Hall: 1992.
  7. ^ a b Lewis, R. J. Sax's Dangerous Properties of Industrial Materials, 10th ed., John Wiley & Sons, Inc.: 2000.

External links[edit]