Sodium peroxide

Sodium peroxide
Sodium peroxide
Names
IUPAC name
Sodium peroxide
Other names
Flocool
Solozone
Disodium peroxide
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.013.828 Edit this at Wikidata
EC Number
  • 215-209-4
RTECS number
  • WD3450000
UNII
UN number 1504
  • InChI=1S/2Na.O2/c;;1-2/q2*+1;-2 ☒N
    Key: PFUVRDFDKPNGAV-UHFFFAOYSA-N ☒N
  • [O-][O-].[Na+].[Na+]
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
Solubility Soluble in acid
Insoluble in base
Reacts with ethanol
−28.10·10−6 cm3/mol
Structure
hexagonal
Thermochemistry
89.37 J/(mol·K)
95 J/(mol·K)[1]
−515 kJ·mol−1[1]
−446.9 kJ/mol
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
 caustic
GHS labelling:
GHS03: OxidizingGHS05: Corrosive
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)
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 0: Will not burn. E.g. waterInstability 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g. white phosphorusSpecial hazard W+OX: Reacts with water in an unusual or dangerous manner AND is oxidizer
3
0
2
W
OX
Flash point Non-flammable
Safety data sheet (SDS) External MSDS
Related compounds
Other cations
 Lithium peroxide
Potassium peroxide
Rubidium peroxide
Caesium peroxide
Related sodium oxides
 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).
☒N verify (what is checkY☒N ?)

Sodium peroxide is an inorganic compound with the formula Na2O2. This yellowish solid is the product of sodium ignited in excess oxygen.[3] 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.[4] The octahydrate, which is simple to prepare, is white, in contrast to the anhydrous material.[5]

Properties

Sodium peroxide crystallizes with hexagonal symmetry.[6] Upon heating, the hexagonal form undergoes a transition into a phase of unknown symmetry at 512 °C.[7] With further heating above the 657 °C boiling point, the compound decomposes to Na2O, releasing O2.[8]

2 Na2O2 → 2 Na2O + O2

Preparation

Commercially, sodium peroxide is produced from the elements in a two-stage process. First sodium is oxidized to sodium oxide:[7][9]

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.[10]

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.[5]

Uses

Sodium peroxide hydrolyzes to give sodium hydroxide and hydrogen peroxide according to the reaction[9]

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[7] and Flocool.[8] 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.[4]

References

  1. ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A23. ISBN 978-0-618-94690-7.
  2. ^ "Hazard Rating Information for NFPA Fire Diamonds". Archived from the original on 2004-09-04.
  3. ^ Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. p. 98. ISBN 978-0-08-022057-4.
  4. ^ 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.
  5. ^ a b R. A. Penneman (1950). "Potassium Sodium Peroxide 8-Hydrate". Inorg. Synth. 3: 1–4. doi:10.1002/9780470132340.ch1.
  6. ^ 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.
  7. ^ a b c Macintyre, J. E., ed. Dictionary of Inorganic Compounds, Chapman & Hall: 1992.
  8. ^ a b Lewis, R. J. Sax's Dangerous Properties of Industrial Materials, 10th ed., John Wiley & Sons, Inc.: 2000.
  9. ^ a b 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.
  10. ^ Bulletin of Pharmacy. 1895.