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.
The hexagonal crystal structure of sodium peroxide was discovered by Tallman et al.. Upon heating, the structure undergoes a transition into a phase of unknown symmetry at 512 °C. With further heating above the 675 °C melting point, the compound decomposes, releasing O2, before reaching a boiling point.
Sodium peroxide can be synthesized by direct reaction with sodium and oxygen at 130 - 200 °C. Lower temperature (0 - 20 °C) synthesis can be achieved by passing O2 over a dilute (0.1 - 5.0 mole percent) sodium amalgam, thus oxidizing the sodium. 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.
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 and Flocool. In chemistry preparations, sodium peroxide is used as an oxidising 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. See also lithium peroxide (similar use).