Sodium oxide

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Sodium oxide
Sodium oxide Sodium oxide
Identifiers
CAS number 1313-59-3 YesY
PubChem 73971
EC number 215-208-9
UN number 1825
Jmol-3D images Image 1
Properties
Molecular formula Na2O
Molar mass 61.98 g mol−1
Appearance white solid
Density 2.27 g/cm3
Melting point 1,132 °C (2,070 °F; 1,405 K)
Boiling point 1,950 °C (3,540 °F; 2,220 K) sublimates
Sublimation conditions sublimates at 1275 °C
Solubility in water reacts violently to form NaOH
Solubility reacts with ethanol
Structure
Crystal structure Antifluorite (face centered cubic), cF12
Space group Fm3m, No. 225
Coordination
geometry
Tetrahedral (Na+); cubic (O2–)
Thermochemistry
Specific
heat capacity
C
72.95 J/mol·K
Std molar
entropy
So298
73 J/mol·K[1]
Std enthalpy of
formation
ΔfHo298
-416 kJ/mol[1]
Gibbs free energy ΔG -377.1 kJ/mol
Hazards
MSDS ICSC 1653
GHS pictograms The corrosion pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)[2]
GHS hazard statements H314[2]
GHS precautionary statements P280[2]
EU Index not listed
EU classification Corrosive C
R-phrases R14, R8, R34, R35
S-phrases (S1/2), S8, S26, S27, S36/37/39, S43, S45
Main hazards corrosive, reacts violently with water
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g., chloroform Reactivity code 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g., calcium Special hazard OX: Oxidizer. E.g., potassium perchlorateNFPA 704 four-colored diamond
Flash point non-flammable
Related compounds
Other anions Sodium sulfide
Sodium selenide
Sodium telluride
Other cations Lithium oxide
Potassium oxide
Rubidium oxide
Caesium oxide
Related sodium oxides Sodium peroxide
Sodium superoxide
Related compounds Sodium hydroxide
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
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Infobox references

Sodium oxide (SOX) is a chemical compound with the formula Na2O. It is used in ceramics and glasses, though not in a raw form. It is the base anhydride of sodium hydroxide, so when water is added to sodium oxide NaOH is produced.

Na2O + H2O → 2 NaOH

The alkali metal oxides M2O (M = Li, Na, K, Rb) crystallise in the antifluorite structure. In this motif the positions of the anions and cations are reversed relative to their positions in CaF2, with sodium ions tetrahedrally coordinated to 4 oxide ions and oxide cubically coordinated to 8 sodium ions.[3][4]

Preparation[edit]

Sodium oxide is produced by the reaction of sodium with sodium hydroxide, sodium peroxide, or sodium nitrite:[5]

2 NaOH + 2 Na → 2 Na2O + H2
Na2O2 + 2 Na → 2 Na2O
2 NaNO2 + 6 Na → 4 Na2O + N2

Most of these reactions rely on the reduction of something by sodium, whether it is hydroxide, peroxide, or nitrite.

Burning sodium in air will produce Na2O and about 20% sodium peroxide Na2O2.

6 Na + 2 O2 → 2 Na2O + Na2O2

Alternatively, sodium carbonate can be heated to 851 °C, producing carbon dioxide and sodium oxide.

Na2CO3 → Na2O + CO2

Applications[edit]

Glass making[edit]

Sodium oxide is a significant component of glasses and windows although it is added in the form of "soda" (sodium carbonate). Sodium oxide does not explicitly exist in glasses, since glasses are complex cross-linked polymers. Typically, manufactured glass contains around 15% sodium oxide, 70% silica (silicon dioxide) and 9% lime (calcium oxide). The sodium carbonate "soda" serves as a flux to lower the temperature at which the silica melts. Soda glass has a much lower melting temperature than pure silica, and has slightly higher elasticity. These changes arise because the silicon dioxide and soda react to form sodium silicates of the general formula Na2[SiO2]x[SiO3].

Na2CO3 → Na2O + CO2
Na2O + SiO2 → Na2SiO3

References[edit]

  1. ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A23. ISBN 0-618-94690-X. 
  2. ^ a b c Sigma-Aldrich Co., Sodium oxide. Retrieved on 2014-05-25.
  3. ^ Zintl, E.; Harder, A.; Dauth B. (1934). "Gitterstruktur der oxyde, sulfide, selenide und telluride des lithiums, natriums und kaliums". Z. Elektrochem. Angew. Phys. Chem. 40: 588–93. 
  4. ^ Wells, A.F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.
  5. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0080379419. 

External links[edit]