Sodium peroxide

Crystal structure
__ Na +      __ O -
General
Surname	Sodium peroxide
Ratio formula	Na 2 O 2
Brief description	white to yellowish powder
External identifiers / databases
CAS number 1313-60-6 EC number 215-209-4 ECHA InfoCard 100,013,828 PubChem 14803 ChemSpider 14119 Wikidata Q205459
properties
Molar mass	77.98 g · mol -1
Physical state	firmly
density	2.805 g cm −3
Melting point	675 ° C 660 ° C
solubility	Decomposes in water
safety instructions
GHS hazard labeling from  Regulation (EC) No. 1272/2008 (CLP) , expanded if necessary danger H and P phrases H: 271-314 P: 210-221-280-301 + 330 + 331-305 + 351 + 338-308 + 310
As far as possible and customary, SI units are used. Unless otherwise noted, the data given apply to standard conditions .

Sodium peroxide is the peroxide of sodium and salt of hydrogen peroxide . It has the molecular formula Na ₂ O ₂ .

history

Sodium peroxide was first made in 1811 by JL Gay-Lussac and LJ Thénard by burning sodium. The first industrial manufacturing process was developed in 1891 by H. Castner as a so-called block process and carried out at Degussa in Rheinfelden from 1899 . From 1928 a continuous process in a rotary kiln was introduced.

Presentation and extraction

The technical production of sodium peroxide takes place by burning sodium in an air atmosphere .

${\displaystyle \mathrm {2\ Na+O_{2}\longrightarrow \ Na_{2}O_{2}} }$

The reaction is strongly exothermic at −511.2 kJ mol ⁻¹ (Na ₂ O ₂ ). To avoid local overheating, which would cause the product formed to break down, the process is carried out using the countercurrent principle. Fresh sodium impinges on an oxygen-poor, stale air, initially at temperatures between 150 ° C and 200 ° C, oxidation to sodium oxide is carried out. In this step, −431.6 kJ mol ⁻¹ (Na ₂ O), most of the total heat of reaction is released.

${\displaystyle \mathrm {4\ Na+O_{2}\longrightarrow 2\ Na_{2}O} }$

This is then further converted into the target product with fresh air at 350 ° C. Here, −79.6 kJ mol ⁻¹ (Na ₂ O ₂ ) are released.

${\displaystyle \mathrm {2\ Na_{2}O+O_{2}\longrightarrow 2\ Na_{2}O_{2}} }$

In the laboratory, the peroxide can be represented as octahydrate by precipitating a concentrated sodium hydroxide solution with hydrogen peroxide :

${\displaystyle {\ce {2 NaOH + H2O2 ->[H_2O][] Na2O2 * 8 H2O}}}$

properties

Sodium peroxide is a colorless powder which can also be colored yellow due to contamination with yellow sodium oxide. The compound is thermally stable up to 500 ° C and can be melted almost undecomposed at 675 ° C. It is very hygroscopic. It reacts with water to form sodium hydroxide and hydrogen peroxide .

${\displaystyle \mathrm {Na_{2}O_{2}+2\ H_{2}O\longrightarrow 2\ NaOH+\ H_{2}O_{2}} }$

This reaction is highly exothermic, so that if the cooling is inadequate, the temperature may rise to a temperature range in which the product hydrogen peroxide decomposes. A vigorous evolution of oxygen is observed.

${\displaystyle \mathrm {Na_{2}O_{2}+H_{2}O\longrightarrow 2\ NaOH+{1/2}\ O_{2}} }$

The compound is a strong oxidizing agent , which can react explosively with oxidizable substances such as sulfur , carbon or aluminum powder .

The compound has a hexagonal crystal structure with the space group P 6 2 m (space group no. 189) . The standard enthalpy of formation of sodium peroxide is ΔH _f ⁰  = −513 kJ / mol. There further comprising the dihydrate Na ₂ O ₂ · 2H ₂ O and the octahydrate Na ₂ O ₂ · 8H ₂ O two hydrates , a Diperhydrat Na ₂ O ₂ · 2H ₂ O ₂ and a mixed hydrate perhydrate Na ₂ O ₂ · 2H ₂ O ₂ · 4H ₂ O.

use

In industry, sodium peroxide is used as an oxidizing and bleaching agent. The compound is used extensively in textile and paper bleaching. The sodium hydroxide solution present in aqueous solution is neutralized by the addition of sulfuric acid or precipitated using magnesium sulfate . Among other things, it is used for CO ₂ absorption and oxygen production (e.g. in breathing apparatus, submarines or space stations). In space travel, however, the lighter lithium peroxide is preferred.

${\displaystyle \mathrm {Na_{2}O_{2}+CO_{2}\longrightarrow Na_{2}CO_{3}+1/2\ O_{2}} }$

In organic synthesis it is used as a reagent for the epoxidation of α, β-unsaturated carbonyl compounds and in the direct sulfonation of methane using sulfur trioxide in the production of methanesulfonic acid .

${\displaystyle \mathrm {CH_{4}+SO_{3}\longrightarrow CH_{3}SO_{3}H} }$

In the cation separation process , sodium peroxide can be used instead of the sodium hydroxide / hydrogen peroxide mixture for alkaline fall .

See also

• Sodium oxide
• Sodium hyperoxide

Individual evidence

1. ↑ ^{a b c d e f g} Entry on sodium peroxide in the GESTIS substance database of the IFA , accessed on January 30, 2017(JavaScript required) .
2. ↑ ^{a b c d e f g} Entry on sodium peroxide. In: Römpp Online . Georg Thieme Verlag, accessed on January 30, 2017.
3. ^ ^{A b c d e f g} E. Wiberg , N. Wiberg , AF Holleman : Inorganische Chemie . 103rd edition. Walter de Gruyter, Berlin / Boston 2017, ISBN 978-3-11-026932-1 , p. 602, (accessed via De Gruyter Online).
4. ↑ Entry on Disodium peroxide in the Classification and Labeling Inventory of the European Chemicals Agency (ECHA), accessed on February 1, 2016. Manufacturers or distributors can expand the harmonized classification and labeling .
5. ↑ ^{a b c d e f g} H. Jakob, S. Leininger, T. Lehmann, S. Jacobi, S. Gutewort: Inorganic Peroxo Compounds. In: Ullmann's Encyclopedia of Technical Chemistry . Wiley-VCH, Weinheim 2012. doi : 10.1002 / 14356007.a19_177.pub2 .
6. ↑ Patent US 1,685,520 (Roessler & Hasslacher Chemical, 1928); U.S. Patent 7,796,241 (Roessler & Hasslacher Chemical, 1931).
7. ↑ RA Penneman: Sodium proxide 8-hydrate . In: Ludwig F. Audrieth (Ed.): Inorganic Syntheses . tape 3 . McGraw-Hill, Inc., 1950, pp. 1-3 (English).
8. ↑ Roger Blachnik (Ed.): Paperback for chemists and physicists . Volume III: Elements, Inorganic Compounds and Materials, Minerals . founded by Jean d'Ans, Ellen Lax. 4th, revised and revised edition. Springer, Berlin 1998, ISBN 3-540-60035-3 , pp. 614 ( limited preview in Google Book search).
9. ^ E. Wiberg , N. Wiberg , AF Holleman : Inorganic Chemistry . 103rd edition. Walter de Gruyter, Berlin / Boston 2017, ISBN 978-3-11-026932-1 , p. 1515, (accessed via De Gruyter Online).
10. ↑ VK Reddy, B. Haritha, M. Yamashita: Highly Diastereoselective Epoxidation of α, β-Unsaturated Carbonyl Compounds Using Sodium Peroxide. In: Lett. Org. Chem. 2, 2005, pp. 128-131. doi: 10.2174 / 1570178053202964 .
11. ↑ S. Mukhopadhyay, AT Bell: Direct Liquid-Phase Sulfonation of Methane to Methanesulfonic Acid by SO ₃ in the Presence of a Metal Peroxide. In: Angew. Chem. 115, 2005, pp. 1049-1051. doi: 10.1002 / anie.200390235 .
12. ↑ Blasius Jander: Textbook of analytical and preparative inorganic chemistry. 14th edition. S. Hirzel Verlag, Stuttgart 1995, ISBN 3-7776-0612-X , pp. 561-567.