Silver oxide

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Silver oxide
IUPAC name
Silver(I) oxide
Other names
Argentous oxide
Disilver monoxide
Disilver oxide
Silver monoxide
Silver rust
Molar mass 231.735 g/mol
Appearance Black or dark brown powder
Odor Odorless
Density 7.14 g/cm3 (cubic)[1]
9.53 g/cm3 (hexagonal)[1]
Melting point 200 °C (392 °F; 473 K) (decomposes)
Boiling point Decomposes
0.00070 g/100 ml (10 °C)
0.00116 g/100 ml (15 °C)
0.00174 g/100 ml (20 °C)
0.0022 g/100 ml (25 °C)
0.0027 g/100 ml (30 °C)
0.0035 g/100 ml (40 °C)
0.0042 g/100 ml (50 °C)
0.0049 g/100 ml (60 °C)
0.0052 g/100 ml (70 °C)
0.0055 g/100 ml (80 °C)[2]
Solubility Reacts with acids
Insoluble in alcohols
Vapor pressure ~0 mmHg
122 J·mol-1·K-1
−31 kJ/mol[3]
Safety data sheet Sigma-Aldrich
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
2820 mg/kg (rat, oral)
Related compounds
Related compounds
Copper(II) oxide
Silver(I,III) oxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Infobox references

Silver oxide is the chemical compound with the formula Ag2O.



Silver oxide readily reacts with acids, forming silver salts. Thus, it is the most important precursor in the synthesis of Ag salts.

Ag2O + 2 HX → 2 AgX + H2O

Unlike other silver compounds, silver oxide does not break down when exposed to light, but will decompose to silver metal if heated above 200 °C:

Ag2O → 2 Ag + ½ O2

Silver oxide can be used to synthesize epoxides from halocarbons. For example, reaction with 1,2-diiodoethane at high temperatures, yields ethylene oxide.

Ag2O + (CH2)2I2 → 2 AgI + (CH2CH2)O

Silver oxide is a mild oxidizing agent, capable of oxidizing aldehydes to carboxylic acids.


Silver oxide is a dense black powder, odorless. It is slightly soluble in water due to the formation of the ion Ag+ or Ag(OH)2, and the solubility increases at high and low pH (amphoteric). A saturated solution in water has a pH of about 10, and the solubility is lowest at pH 12.[1] It will also dissolve in aq. ammonia, producing active compound of Tollens' reagent.


Silver oxide can be bought from chemical suppliers, but it's not cheap.

Silver oxide can be found in silver oxide batteries, but it's impractical to extract it from them, as you need a large amount of batteries.


The most common way of preparing silver oxide is by adding an alkali base, like NaOH, to an aq. solution containing soluble silver salts, like silver nitrate.

2 AgNO3 + 2 NaOH → Ag2O + 2 NaNO3 + H2O

Other salts, like silver perchlorate can be used, but this one is a bit more expensive than the nitrate.[4]

It's possible to use silver salts that are poorly soluble in water, if silver nitrate is not available, like silver sulfate, and the reaction will work, provided that the silver sulfate is extremely fine. The trick is to stir the alkali suspension of Ag2 for a while, until no more white precipitate is seen in the suspension. If the silver sulfate is not fine enough, any lumps of it will only react at the surface with the alkali hydroxide, and thus the final product will be contaminated with unreacted Ag2SO4. To check for any unreacted silver sulfate, take the dry crude Ag2 and grind it using a mortar and pestle, an see if crushing any black lumps produces white powder. The crude silver oxide can be further purified, by dissolving it in an acid, like formic acid (if nitric acid is not available), then precipitated again with NaOH to silver oxide.


  • Make silver salts
  • Synthesis of ethylene oxide
  • Compound collecting
  • Make expensive thermite
  • Make silver-oxide battery



Silver oxide is harmful if ingested or inhaled, and will stain the skin. Proper protection must be worn when handling the compound.


In closed bottles, plastic or glass, away from any acids or reducing agents. Unlike most silver compounds, it is not light sensitive, thus it's not necessary to be kept in dark.


Silver oxide can be reduced to silver metal, which is recycled.


  1. 1.0 1.1 1.2 Gmelins Handbuch der Anorganischen Chemie, Silber Teil B1, Verlag Chemie GmbH, Weinheim/Bergstraße, 8th edition 1971, p. 68, 96-98
  2. A. Seidell, W. F. Linke, Solubilities of Inorganic and Metal-Organic Compounds, 4th edition, volume 1, Toronto, New York, London 1958, p. 124
  3. Barin I. Thermochemical Data of Pure Substances. - VCH, 1995 pp. 12
  4. An EXPENSIVE Metal Oxide Silver(I) Oxide

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