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Sodium peroxide



Sodium peroxide


Sodium peroxide formations on burning sodium metal
Systematic name Sodium peroxide
Other names Sodium dioxide, Natrium peroxide,
Flocool, Solozone, Disodium peroxide
Molecular formula Na2O2
Molar mass 78 g/mol
Density 2.8 g/cm3
Solubility (water) reacts with water
Melting point 675 °C
Boiling point N/A, decomposes
Appearance yellow to white powder
CAS number [1313-60-6]
Crystal structure hexagonal[1]
Hazards
MSDS External MSDS
Main hazards Strong oxidizer, reacts with water
NFPA 704
0
3
1
OX
R/S statement R: R7, R14, R26/27/28, R29, R41
S: S7/8, S37/39
Supplementary data
Thermodynamic
properties[2]
ΔH °f = -513 kJ mol-1

ΔG °f = -449 kJ mol-1
S ° = 95 J K-1 mol-1

Related compounds
Other cations hydrogen peroxide, lithium peroxide,
calcium peroxide
Related oxides sodium oxide, sodium superoxide
Except where noted otherwise, data are given for
materials in their standard state (at 25 °C, 100 kPa)
Infobox disclaimer and references

Sodium peroxide, Na2O2, is the normal product when sodium is burned. It is a strong oxidizer.

Additional recommended knowledge

Contents

Chemical Properties

Sodium peroxide is hydrolyzed by water to form sodium hydroxide plus hydrogen peroxide according to the reaction:

Na2O2 + 2 H2O → 2 NaOH + H2O2

The hydrogen peroxide thus formed decomposes rapidly in the ensuing basic solution, producing water and oxygen. The reaction is substantially exothermic and can set fire to combustible materials.

Sodium peroxide will also set fire to many organic liquids on contact (particularly alcohols and glycols), and reacts violently with powdered metals and numerous other compounds after minimal initiation.

Structural Transitions

The hexagonal crystal structure of sodium peroxide was discovered by Tallman et al.[1]. Upon heating, the structure undergoes a transition into a phase of unknown symmetry at 512 °C.[2] With further heating above the 675 °C melting point, the compound decomposes, releasing O2, before reaching a boiling point.[3]

Preparation

Sodium peroxide can be synthesized by direct reaction with sodium and oxygen at 130 - 200 °C.[2] Lower temperature (0 - 20 °C) synthesis can be achieved by passing O2 over a dilute (0.1 - 5.0 mole percent) sodium metal amalgam, thus oxidizing the sodium.[4] 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

Given its strong oxidation properties, sodium peroxide is used to bleach wood pulp for the production of paper. It has also been used for the extraction of minerals from various ores. Sodium peroxide may go by the commercial names of Solozone[2] and Flocool.[3] In chemistry preparations, sodium peroxide is used as an oxidative reagent.

References

  1. ^ a b Tallman, R. L.; Margrave, J. L.; Bailey, S. W. J. Am. Chem. Soc. 1957, 79, 2979-80.
  2. ^ a b c d Macintyre, J. E., ed. Dictionary of Inorganic Compounds, Chapman & Hall: 1992.
  3. ^ a b Lewis, R. J. Sax's Dangerous Properties of Industrial Materials, 10th ed., John Wiley & Sons, Inc.: 2000.
  4. ^ Schechter, D. L.; Bon, C. K.; Leddy, J. J.; Process for the Preparation of Sodium Peroxide by the Oxidation of a Sodium Amalgam, U.S. Patent 3,141,736. Filed 7 May 1962.
 
This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Sodium_peroxide". A list of authors is available in Wikipedia.
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