Sodium peroxide

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Sodium peroxide
IUPAC name
Sodium peroxide
Other names
Disodium dioxide
Disodium peroxide
Molar mass 77.98 g/mol
Appearance White-yellow solid
Odor Odorless
Density 2.805 g/cm3
Melting point 460 °C (860 °F; 733 K) (decomposes)
Boiling point 657 °C (1,215 °F; 930 K) (decomposition)
Reacts exothermically
Solubility Reacts with acids, alcohols
Insoluble in bases
95 J·mol−1·K−1
−515 kJ/mol
Safety data sheet Sigma-Aldrich
Flash point Non-flammable
Related compounds
Related compounds
Sodium oxide
Hydrogen peroxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Infobox references

Sodium peroxide is a sodium salt of hydrogen peroxide. Its formula is Na2O2. This compound is somewhat unstable and a strong oxidizer.



Sodium peroxide exists in two main forms: anhydrous and octahydrate (Na2O2·8H2O). Other lesser known hydrates also exist, but they have little importance.

Anhydrous sodium peroxide is a yellow powder (white if very high-grade: the yellow color is caused by contamination with sodium superoxide). The octahydrate appears as colorless crystals with a low melting point (30 °C).


Anhydrous sodium peroxide reacts exothermically with water. Part of it is hydrolyzed irreversibly into oxygen, sodium hydroxide and water. Part is converted into the hydrate, which undergoes reversible hydrolysis and can be crystallized out of the solution by adding ethanol. The irreversible reaction with water is actually caused by heat generated by common dissolution of sodium peroxide, and low temperatures suppress it: it is possible to dissolve anhydrous sodium peroxide in ice-cold water without much loss to decomposition. The octahydrate is soluble in water without decomposition.

It reacts with carbon dioxide to produce oxygen and sodium carbonate.

Attempts to calcine the octahydrate result in its decomposition, the crystals turn into sodium hydroxide. There is no easily available way to turn the octahydrate into the anhydrous salt, but there were reports that putting it in a desiccator with very concentrated (98%) sulfuric acid for a long time does the trick.


Sodium peroxide, usually anhydrous, is commonly sold by various chemical suppliers.


Anhydrous sodium peroxide is usually prepared by burning sodium metal in air or in a gentle stream of oxygen. The octahydrate can be easily prepared by a simple neutralization reaction with hydrogen peroxide (a weak acid) and sodium hydroxide. It is recommended to use dilute solutions of hydrogen peroxide, because the side reaction of H2O2 decomposition becomes prevalent in concentrated solutions. The crystals can be displaced out of solution with ethanol.




Avoid contact of anhydrous sodium peroxide with organic solvents. This may result in fire or explosions.

Sodium peroxide will cause burns on contact with bare skin.


Anhydrous sodium peroxide should be kept in a tightly closed jar, to avoid contact with water vapor. The octahydrate should be kept in a cold place, also in a closed jar.


Any form of sodium peroxide can be neutralized by adding any acid and manganese dioxide. This reduces the compound to a sodium salt of that acid and oxygen.


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