Sodium peroxide

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

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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.