Silver sulfide

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Silver sulfide
Ag2S-bas.png
Sulfid stribrny.PNG
Names
IUPAC name
Silver(I) sulfide
Other names
Silver sulfide
Argentous sulfide
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.040.384 OOjs UI icon edit-ltr-progressive.svg
EC Number
  • 244-438-2
PubChem CID
UNII
  • InChI=1S/2Ag.S/q2*+1;-2 X mark.svgN
    Key: XUARKZBEFFVFRG-UHFFFAOYSA-N X mark.svgN
  • [S-2].[Ag+].[Ag+]
Properties
Ag2S
Molar mass 247.80 g·mol−1
AppearanceGrayish-blackish crystal
Odor Odorless
Density 7.234 g/cm3 (25 °C) [1] [2]
7.12 g/cm3 (117 °C) [3]
Melting point 836 °C (1,537 °F; 1,109 K) [1]
6.21·10−15 g/L (25 °C)
6.31·10−50
Solubility Soluble in aq. HCN, aq. citric acid with KNO3
Insoluble in acids, alkalies, aqueous ammoniums [4]
Structure
Cubic, cI8 (α-form)
Monoclinic, mP12 (β-form)
Cubic, cF12 (γ-form) [3] [5]
Im3m, No. 229 (α-form) [5]
P21/n, No. 14 (β-form)
Fm3m, No. 225 (γ-form) [3]
2/m (α-form) [5]
4/m 3 2/m (β-form, γ-form) [3]
a = 4.23 Å, b = 6.91 Å, c = 7.87 Å (α-form) [5]
α = 90°, β = 99.583°, γ = 90°
Thermochemistry
76.57 J/mol·K [6]
Std molar
entropy (S⦵298)
143.93 J/mol·K [6]
−32.59 kJ/mol [6]
−40.71 kJ/mol [6]
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
May cause irritation
GHS labelling:
GHS-pictogram-exclam.svg [2]
Warning
H315, H319, H335 [2]
P261, P305+P351+P338 [2]
NFPA 704 (fire diamond)
[7]
NFPA 704.svgHealth 0: Exposure under fire conditions would offer no hazard beyond that of ordinary combustible material. E.g. sodium chlorideFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
0
0
0
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
X mark.svgN  verify  (what is  Yes check.svgYX mark.svgN ?)

Silver sulfide is an inorganic compound with the formula Ag
2S. A dense black solid, it is the only sulfide of silver. It is useful as a photosensitizer in photography. It constitutes the tarnish that forms over time on silverware and other silver objects. Silver sulfide is insoluble in most solvents, but is degraded by strong acids. Silver sulfide is a network solid made up of silver (electronegativity of 1.98) and sulfur (electronegativity of 2.58) where the bonds have low ionic character (approximately 10%).

Contents

Formation

Silver sulfide naturally occurs as the tarnish on silverware. When combined with silver, hydrogen sulfide gas creates a layer of black silver sulfide patina on the silver, protecting the inner silver from further conversion to silver sulfide. [8] Silver whiskers can form when silver sulfide forms on the surface of silver electrical contacts operating in an atmosphere rich in hydrogen sulfide and high humidity. [9] Such atmospheres can exist in sewage treatment and paper mills. [10] [11]

Structure and properties

Three forms are known: monoclinic acanthite (β-form), stable below 179 °C, body centered cubic so-called argentite (α-form), stable above 180 °C, and a high temperature face-centred cubic (γ-form) stable above 586 °C. [5] The higher temperature forms are electrical conductors. It is found in nature as relatively low temperature mineral acanthite. Acanthite is an important ore of silver. The acanthite, monoclinic, form features two kinds of silver centers, one with two and the other with three near neighbour sulfur atoms. [12] Argentite refers to a cubic form, which, due to instability in "normal" temperatures, is found in form of the pseudomorphosis of acanthite after argentite.

History

In 1833 Michael Faraday noticed that the resistance of silver sulfide decreased dramatically as temperature increased. This constituted the first report of a semiconducting material. [13]

Silver sulfide is a component of classical qualitative inorganic analysis. [14]

Related Research Articles

<span class="mw-page-title-main">Chalcogen</span> Group of chemical elements

The chalcogens are the chemical elements in group 16 of the periodic table. This group is also known as the oxygen family. Group 16 consists of the elements oxygen (O), sulfur (S), selenium (Se), tellurium (Te), and the radioactive elements polonium (Po) and livermorium (Lv). Often, oxygen is treated separately from the other chalcogens, sometimes even excluded from the scope of the term "chalcogen" altogether, due to its very different chemical behavior from sulfur, selenium, tellurium, and polonium. The word "chalcogen" is derived from a combination of the Greek word khalkόs (χαλκός) principally meaning copper, and the Latinized Greek word genēs, meaning born or produced.

<span class="mw-page-title-main">Sulfur</span> Chemical element, symbol S and atomic number 16

Sulfur (also spelled sulphur in British English) is a chemical element; it has symbol S and atomic number 16. It is abundant, multivalent and nonmetallic. Under normal conditions, sulfur atoms form cyclic octatomic molecules with the chemical formula S8. Elemental sulfur is a bright yellow, crystalline solid at room temperature.

<span class="mw-page-title-main">Hydrogen sulfide</span> Poisonous, corrosive and flammable gas

Hydrogen sulfide is a chemical compound with the formula H2S. It is a colorless chalcogen-hydride gas, and is poisonous, corrosive, and flammable, with trace amounts in ambient atmosphere having a characteristic foul odor of rotten eggs. Swedish chemist Carl Wilhelm Scheele is credited with having discovered the chemical composition of purified hydrogen sulfide in 1777.

<span class="mw-page-title-main">Noble metal</span> Metallic elements that are nearly chemically inert

A noble metal is ordinarily regarded as a metallic chemical element that is generally resistant to corrosion and is usually found in nature in its raw form. Gold, platinum, and the other platinum group metals are most often so classified. Silver, copper, and mercury are sometimes included as noble metals, but each of these usually occurs in nature combined with sulfur.

Sulfide (British English also sulphide) is an inorganic anion of sulfur with the chemical formula S2− or a compound containing one or more S2− ions. Solutions of sulfide salts are corrosive. Sulfide also refers to large families of inorganic and organic compounds, e.g. lead sulfide and dimethyl sulfide. Hydrogen sulfide (H2S) and bisulfide (SH−) are the conjugate acids of sulfide.

<span class="mw-page-title-main">Argentite</span>

In mineralogy, argentite (from Latin argentum ' silver') is cubic silver sulfide (Ag2S), which can only exist at temperatures above 173 °C (343 °F), 177 °C (351 °F), or 179 °C (354 °F). When it cools to ordinary temperatures it turns into its monoclinic polymorph, acanthite. The International Mineralogical Association has decided to reject argentite as a proper mineral.

In chemistry, a reactivity series (or activity series) is an empirical, calculated, and structurally analytical progression of a series of metals, arranged by their "reactivity" from highest to lowest. It is used to summarize information about the reactions of metals with acids and water, single displacement reactions and the extraction of metals from their ores.

<span class="mw-page-title-main">Acanthite</span> Mineral, silver sulfide

Acanthite is a form of silver sulfide with the chemical formula Ag2S. It crystallizes in the monoclinic system and is the stable form of silver sulfide below 173 °C (343 °F). Argentite is the stable form above that temperature. As argentite cools below that temperature its cubic form is distorted to the monoclinic form of acanthite. Below 173 °C acanthite forms directly. Acanthite is the only stable form in normal air temperature.

<span class="mw-page-title-main">Lead(II) sulfide</span> Chemical compound

Lead(II) sulfide is an inorganic compound with the formula PbS. Galena is the principal ore and the most important compound of lead. It is a semiconducting material with niche uses.

<span class="mw-page-title-main">Ammonium hydrosulfide</span> Chemical compound

Ammonium hydrosulfide is the chemical compound with the formula [NH4]SH.

Sulfur compounds are chemical compounds formed the element sulfur (S). Common oxidation states of sulfur range from −2 to +6. Sulfur forms stable compounds with all elements except the noble gases.

<span class="mw-page-title-main">Silver(I) selenide</span> Chemical compound

Silver selenide (Ag2Se) is the reaction product formed when selenium toning analog silver gelatine photo papers in photographic print toning. The selenium toner contains sodium selenite (Na2SeO3) as one of its active ingredients, which is the source of the selenide (Se2−) anion combining with the silver in the toning process.

<span class="mw-page-title-main">Copper(I) sulfide</span> Chemical compound

Copper(I) sulfide is a copper sulfide, a chemical compound of copper and sulfur. It has the chemical compound Cu2S. It is found in nature as the mineral chalcocite. It has a narrow range of stoichiometry ranging from Cu1.997S to Cu2.000S. Samples are typically black.

<span class="mw-page-title-main">Gold(I) sulfide</span> Chemical compound

Gold(I) sulfide is the inorganic compound with the formula Au2S. It is the principal sulfide of gold. It decomposes to gold metal and elemental sulfur, illustrating the "nobility" of gold.

<span class="mw-page-title-main">Californium compounds</span>

Few compounds of californium have been made and studied. The only californium ion that is stable in aqueous solutions is the californium(III) cation. The other two oxidation states are IV (strong oxidizing agents) and II (strong reducing agents). The element forms a water-soluble chloride, nitrate, perchlorate, and sulfate and is precipitated as a fluoride, oxalate or hydroxide. If problems of availability of the element could be overcome, then CfBr2 and CfI2 would likely be stable.

<span class="mw-page-title-main">Berkelium compounds</span> Any chemical compound having at least one berkelium atom

Berkelium forms a number of chemical compounds, where it normally exists in an oxidation state of +3 or +4, and behaves similarly to its lanthanide analogue, terbium. Like all actinides, berkelium easily dissolves in various aqueous inorganic acids, liberating gaseous hydrogen and converting into the trivalent oxidation state. This trivalent state is the most stable, especially in aqueous solutions, but tetravalent berkelium compounds are also known. The existence of divalent berkelium salts is uncertain and has only been reported in mixed lanthanum chloride-strontium chloride melts. Aqueous solutions of Bk3+ ions are green in most acids. The color of the Bk4+ ions is yellow in hydrochloric acid and orange-yellow in sulfuric acid. Berkelium does not react rapidly with oxygen at room temperature, possibly due to the formation of a protective oxide surface layer; however, it reacts with molten metals, hydrogen, halogens, chalcogens and pnictogens to form various binary compounds. Berkelium can also form several organometallic compounds.

<span class="mw-page-title-main">Polonium dioxide</span> Chemical compound

Polonium dioxide (also known as polonium(IV) oxide) is a chemical compound with the formula PoO2. It is one of three oxides of polonium, the other two being polonium monoxide (PoO) and polonium trioxide (PoO3). It is a pale yellow crystalline solid at room temperature. Under lowered pressure (such as a vacuum), it decomposes into elemental polonium and oxygen at 500 °C. It is the most stable oxide of polonium and is an interchalcogen.

<span class="mw-page-title-main">Gallium(III) sulfide</span> Chemical compound

Gallium(III) sulfide, Ga2S3, is a compound of sulfur and gallium, that is a semiconductor that has applications in electronics and photonics.

<span class="mw-page-title-main">Chromium(II) sulfide</span> Chemical compound

Chromium(II) sulfide is an inorganic compound of chromium and sulfur with the chemical formula CrS. The compound forms black hexagonal crystals, insoluble in water.

<span class="mw-page-title-main">Holmium(III) sulfide</span> Chemical compound

Holmium(III) sulfide is the sulfide of holmium, with the chemical formula of Ho2S3. Like other rare earth sulfides, it is used as a high-performance inorganic pigment.

References

  1. 1 2 Lide, David R., ed. (2009). CRC Handbook of Chemistry and Physics (90th ed.). Boca Raton, Florida: CRC Press. ISBN   978-1-4200-9084-0.
  2. 1 2 3 4 Sigma-Aldrich Co., Silver sulfide. Retrieved on 2014-07-13.
  3. 1 2 3 4 Tonkov, E. Yu (1992). High Pressure Phase Transformations: A Handbook. Vol. 1. Gordon and Breach Science Publishers. p. 13. ISBN   978-2-88124-761-3.
  4. ↑ Comey, Arthur Messinger; Hahn, Dorothy A. (February 1921). A Dictionary of Chemical Solubilities: Inorganic (2nd ed.). New York: The MacMillan Company. p.  835.
  5. 1 2 3 4 5 "Silver sulfide (Ag2S) crystal structure". Non-Tetrahedrally Bonded Elements and Binary Compounds I. Landolt-Börnstein - Group III Condensed Matter. Vol. 41C. Springer Berlin Heidelberg. 1998. pp. 1–4. doi:10.1007/10681727_86. ISBN   978-3-540-31360-1.
  6. 1 2 3 4 Pradyot, Patnaik (2003). Handbook of Inorganic Chemicals. The McGraw-Hill Companies, Inc. p. 845. ISBN   978-0-07-049439-8.
  7. ↑ "MSDS of Silver Sulfide". saltlakemetals.com. Utah, USA: Salt Lake Metals. Archived from the original on 2014-08-10. Retrieved 2014-07-13.
  8. ↑ Zumdahl, Steven S.; DeCoste, Donald J. (2013). Chemical Principles (7th ed.). Cengage Learning. p. 505. ISBN   978-1-111-58065-0.
  9. ↑ "Degradation of Power Contacts in Industrial Atmosphere: Silver Corrosion and Whiskers" (PDF). 2002.
  10. ↑ Dutta, Paritam K.; Rabaey, Korneel; Yuan, Zhiguo; Rozendal, René A.; Keller, Jürg (2010). "Electrochemical sulfide removal and recovery from paper mill anaerobic treatment effluent". Water Research. 44 (8): 2563–2571. Bibcode:2010WatRe..44.2563D. doi:10.1016/j.watres.2010.01.008. ISSN   0043-1354. PMID   20163816.
  11. ↑ "Control of Hydrogen Sulfide Generation | Water & Wastes Digest". www.wwdmag.com. 5 March 2012. Retrieved 2018-07-05.
  12. ↑ Frueh, A. J. (1958). The crystallography of silver sulfide, Ag2S. Zeitschrift für Kristallographie-Crystalline Materials, 110(1-6), 136-144.
  13. ↑ "1833 - First Semiconductor Effect is Recorded". Computer History Museum. Retrieved 24 June 2014.
  14. ↑ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN   978-0-08-037941-8.
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