Sodium peroxide

Sodium peroxide
Names
	Other names Disodium dioxide; Flocool; Solozone; Disodium peroxide
Identifiers
CAS Number	1313-60-6 ;
3D model (JSmol)	Interactive image;
ChemSpider	14119 ;
ECHA InfoCard	100.013.828
EC Number	215-209-4;
PubChem CID	14803;
RTECS number	WD3450000;
UNII	3GSN3JCJ5K ;
UN number	1504
CompTox Dashboard (EPA)	DTXSID0061660 ;
	InChI InChI=1S/2Na.O2/c;;1-2/q2*+1;-2 Key: PFUVRDFDKPNGAV-UHFFFAOYSA-N ;
	SMILES [O-][O-].[Na+].[Na+];
Properties
Chemical formula	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)
Solubility in water	Reacts
Solubility	Soluble in acid ; Insoluble in base ; Reacts with ethanol
Magnetic susceptibility (χ)	−28.10·10−6 cm3/mol
Structure
Crystal structure	hexagonal
Thermochemistry
Heat capacity (C)	89.37 J/(mol·K)
Std molar; entropy (S⦵298)	95 J/(mol·K)[1]
Std enthalpy of; formation (ΔfH⦵298)	−515 kJ·mol−1[1]
Gibbs free energy (ΔfG⦵)	−446.9 kJ/mol
Hazards
	Occupational safety and health (OHS/OSH):
Main hazards	caustic
	GHS labelling:
Pictograms
Signal word	Danger
Hazard statements	H271, H314
Precautionary statements	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)	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). verify (what is ?) Infobox references

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

The octahydrate is produced by treating sodium hydroxide with hydrogen peroxide.[5]

Sodium peroxide can be prepared on a large scale by the reaction of metallic sodium with oxygen at 130–200 °C, a process that generates sodium oxide, which in a separate stage absorbs oxygen:[7][9]

4 Na + O₂ → 2 Na₂O

2 Na₂O + O₂ → 2 Na₂O₂

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.

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:

2 Na₂O₂ + 2 CO₂ → 2 Na₂CO₃ + O₂

It is thus particularly useful in scuba gear, submarines, etc. Lithium peroxide and potassium superoxide have similar uses.

References

Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A23. ISBN 978-0-618-94690-7.
"Hazard Rating Information for NFPA Fire Diamonds". Archived from the original on 2004-09-04.
Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. p. 98. ISBN 978-0-08-022057-4.
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.
R. A. Penneman (1950). "Potassium Sodium Peroxide 8-Hydrate". Inorg. Synth. 3: 1–4. doi:10.1002/9780470132340.ch1.{{cite journal}}: CS1 maint: uses authors parameter (link)
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.
Macintyre, J. E., ed. Dictionary of Inorganic Compounds, Chapman & Hall: 1992.
Lewis, R. J. Sax's Dangerous Properties of Industrial Materials, 10th ed., John Wiley & Sons, Inc.: 2000.
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.

External links

International Chemical Safety Card 1606

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[b1-1] 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.

[Ullmann-4] 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.

[IS-5] R. A. Penneman (1950). "Potassium Sodium Peroxide 8-Hydrate". Inorg. Synth. 3: 1–4. doi:10.1002/9780470132340.ch1.{{cite journal}}: CS1 maint: uses authors parameter (link)

[Tallman-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.

[dic-7] Macintyre, J. E., ed. Dictionary of Inorganic Compounds, Chapman & Hall: 1992.

[Lewis-8] Lewis, R. J. Sax's Dangerous Properties of Industrial Materials, 10th ed., John Wiley & Sons, Inc.: 2000.

[Brauser-9] 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.