Silver oxide

Not to be confused with silver(I,III) oxide, used in silver-oxide batteries.
Silver oxide
Names
IUPAC name
Silver(I) oxide
Other names
Silver rust, Argentous oxide, Silver monoxide
Identifiers
20667-12-3 YesY
3D model (Jmol) Interactive image
ChemSpider 7970393 N
ECHA InfoCard 100.039.946
EC Number 243-957-1
MeSH silver+oxide
PubChem 9794626
RTECS number VW4900000
Properties
Ag2O
Molar mass 231.74 g·mol−1
Appearance Black/ brown cubic crystals
Odor Odorless[1]
Density 7.14 g/cm3
Melting point 300 °C (572 °F; 573 K) decomposes from ≥200 °C[2][3]
0.013 g/L (20 °C)
0.025 g/L (25 °C)[4]
0.053 g/L (80 °C)[2]
Solubility product (Ksp) of AgOH
 1.52·10−8 (20 °C)
Solubility Soluble in acid, alkali
Insoluble in ethanol[4]
Structure
Cubic
Pn3m, 224
Thermochemistry
65.9 J/mol·K[4]
122 J/mol·K[5]
−31 kJ/mol[5]
−11.3 kJ/mol[3]
Hazards
GHS pictograms [6]
GHS signal word Danger
H272, H315, H319, H335[6]
P220, P261, P305+351+338[6]
O Xi
R-phrases R36/37/38
S-phrases S17, S26, S36
NFPA 704
Lethal dose or concentration (LD, LC):
2.82 g/kg (rats, oral)[1]
Related compounds
Related compounds
 Silver(I,III) oxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

Silver(I) oxide is the chemical compound with the formula Ag2O. It is a fine black or dark brown powder that is used to prepare other silver compounds.

Preparation

Silver(I) oxide produced by reacting lithium hydroxide with a very dilute silver nitrate solution

Silver oxide can be prepared by combining aqueous solutions of silver nitrate and an alkali hydroxide.[7][8] This reaction does not afford appreciable amounts of silver hydroxide due to the favorable energetics for the following reaction:[9]

2 AgOH → Ag2O + H2O (pK = 2.875[10])

US patent 20050050990 describes the preparation of Ag2O with properties suitable for use as a fine grained conductive paste filler.

Structure and properties

Ag2O features linear, two-coordinate Ag centers linked by tetrahedral oxides. It is isostructural with Cu2O. It "dissolves" in solvents that degrade it. It is slightly soluble in water due to the formation of the ion Ag(OH)2− and possibly related hydrolysis products.[11] It dissolves in ammonia solution to give soluble derivatives. A slurry of Ag2O is readily attacked by acids:

Ag2O + 2 HX → 2 AgX + H2O

where HX = HF, HCl, HBr, or HI, HO2CCF3. It will also react with solutions of alkali chlorides to precipitate silver chloride, leaving a solution of the corresponding alkali hydroxide.[11][12]

Like many silver compounds, silver oxide is photosensitive. It also decomposes at temperatures above 280 °C.[13]

Applications

This oxide is used in some silver-oxide batteries, as is the silver(I,III)oxide, Ag4O4. In organic chemistry, silver oxide is used as a mild oxidizing agent. For example, it oxidizes aldehydes to carboxylic acids. Such reactions often work best when the silver oxide is prepared in situ from silver nitrate and alkali hydroxide.

References

  1. 1 2 3 "Silver Oxide MSDS". SaltLakeMetals.com. Salt Lake Metals. Retrieved 2014-06-08.
  2. 1 2 Perry, Dale L. (1995). Handbook of Inorganic Compounds (illustrated ed.). CRC Press. p. 354. ISBN 0849386713.
  3. 1 2 http://chemister.ru/Database/properties-en.php?dbid=1&id=4098
  4. 1 2 3 Lide, David R. (1998). Handbook of Chemistry and Physics (87 ed.). Boca Raton, FL: CRC Press. pp. 4–83. ISBN 0-8493-0594-2.
  5. 1 2 Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A23. ISBN 0-618-94690-X.
  6. 1 2 3 Sigma-Aldrich Co., Silver(I) oxide. Retrieved on 2014-06-07.
  7. ↑ O. Glemser and H. Sauer "Silver Oxide" in Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 1037.
  8. ↑ Janssen, D. E.; Wilson, C. V. (1963). "4-Iodoveratrole". Org. Synth.; Coll. Vol., 4, p. 547
  9. ↑ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  10. ↑ Biedermann, George; Sillén, Lars Gunnar (1960). "Studies on the Hydrolysis of Metal Ions. Part 30. A Critical Survey of the Solubility Equilibria of Ag2O". Acta Chemica Scandinavica. 14: 717. doi:10.3891/acta.chem.scand.14-0717.
  11. 1 2 Cotton, F. Albert; Wilkinson, Geoffrey (1966). Advanced Inorganic Chemistry (2nd Ed.). New York:Interscience. p. 1042.
  12. ↑ General Chemistry by Linus Pauling, 1970 Dover ed. p703-704
  13. ↑ Merck Index of Chemicals and Drugs, 14th ed. monograph 8521
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