NATIVE SILVER Natural Crystal Cluster Mineral Specimen PRECIOUS METAL AUSTRALIA

Silver

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Silver, 47Ag

Silver

Appearance lustrous white metal

Standard atomic weight Ar°(Ag)

107.8682±0.0002[1]

107.87±0.01 (abridged)[2]

Atomic number (Z) 47

Group group 11

Period period 5

Block   d-block

Electron configuration [Kr] 4d10 5s1

Electrons per shell 2, 8, 18, 18, 1

Physical properties

Phase at STP solid

Melting point 1234.93 K ​(961.78 °C, ​1763.2 °F)

Boiling point 2435 K ​(2162 °C, ​3924 °F)

Density (at 20° C) 10.503 g/cm3[3]

when liquid (at m.p.) 9.320 g/cm3

Heat of fusion 11.28 kJ/mol

Heat of vaporisation 254 kJ/mol

Molar heat capacity 25.350 J/(mol·K)

Vapour pressure

P (Pa) 1 10 100 1 k 10 k 100 k

at T (K) 1283 1413 1575 1782 2055 2433

Atomic properties

Oxidation states −2, −1, 0,[4] +1, +2, +3 (an amphoteric oxide)

Electronegativity Pauling scale: 1.93

Ionisation energies

1st: 731.0 kJ/mol

2nd: 2070 kJ/mol

3rd: 3361 kJ/mol

Atomic radius empirical: 144 pm

Covalent radius 145±5 pm

Van der Waals radius 172 pm

Color lines in a spectral range

Spectral lines of silver

Other properties

Natural occurrence primordial

Crystal structure ​face-centred cubic (fcc) (cF4)

Lattice constant Face-centered cubic crystal structure for silvera = 408.60 pm (at 20 °C)[3]

Thermal expansion 18.92×10−6/K (at 20 °C)[3]

Thermal conductivity 429 W/(m⋅K)

Thermal diffusivity 174 mm2/s (at 300 K)

Electrical resistivity 15.87 nΩ⋅m (at 20 °C)

Magnetic ordering diamagnetic[5]

Molar magnetic susceptibility −19.5×10−6 cm3/mol (296 K)[6]

Young's modulus 83 GPa

Shear modulus 30 GPa

Bulk modulus 100 GPa

Speed of sound thin rod 2680 m/s (at r.t.)

Poisson ratio 0.37

Mohs hardness 2.5

Vickers hardness 251 MPa

Brinell hardness 206–250 MPa

CAS Number 7440-22-4

History

Discovery before 5000 BC

Symbol "Ag": from Latin argentum

Isotopes of silverve

Main isotopes[7] Decay

abun­dance half-life (t1/2) mode pro­duct

105Ag synth 41.3 d ε 105Pd

γ –

106mAg synth 8.28 d ε 106Pd

γ –

107Ag 51.8% stable

108mAg synth 439 y ε 108Pd

IT 108Ag

γ –

109Ag 48.2% stable

110m2Ag synth 249.86 d β− 110Cd

γ –

111Ag synth 7.43 d β− 111Cd

γ –

 Category: Silver

Silver is a chemical element; it has symbol Ag (from Latin argentum 'silver', derived from Proto-Indo-European *h₂erǵ 'shiny, white') and atomic number 47. A soft, white, lustrous transition metal, it exhibits the highest electrical conductivity, thermal conductivity, and reflectivity of any metal.[8] Silver is found in the Earth's crust in the pure, free elemental form ("native silver"), as an alloy with gold and other metals, and in minerals such as argentite and chlorargyrite. Most silver is produced as a byproduct of copper, gold, lead, and zinc refining.

Silver has long been valued as a precious metal. Silver metal is used in many bullion coins, sometimes alongside gold:[9] while it is more abundant than gold, it is much less abundant as a native metal.[10] Its purity is typically measured on a per-mille basis; a 94%-pure alloy is described as "0.940 fine". As one of the seven metals of antiquity, silver has had an enduring role in most human cultures.

Other than in currency and as an investment medium (coins and bullion), silver is used in solar panels, water filtration, jewellery, ornaments, high-value tableware and utensils (hence the term "silverware"), in electrical contacts and conductors, in specialized mirrors, window coatings, in catalysis of chemical reactions, as a colorant in stained glass, and in specialized confectionery. Its compounds are used in photographic and X-ray film. Dilute solutions of silver nitrate and other silver compounds are used as disinfectants and microbiocides (oligodynamic effect), added to bandages, wound-dressings, catheters, and other medical instruments.

Characteristics

Silver is extremely ductile, and can be drawn into a wire one atom wide.[11]

Silver is similar in its physical and chemical properties to its two vertical neighbours in group 11 of the periodic table: copper, and gold. Its 47 electrons are arranged in the configuration [Kr]4d105s1, similarly to copper ([Ar]3d104s1) and gold ([Xe]4f145d106s1); group 11 is one of the few groups in the d-block which has a completely consistent set of electron configurations.[12] This distinctive electron configuration, with a single electron in the highest occupied s subshell over a filled d subshell, accounts for many of the singular properties of metallic silver.[13]

Silver is a relatively soft and extremely ductile and malleable transition metal, though it is slightly less malleable than gold. Silver crystallizes in a face-centered cubic lattice with bulk coordination number 12, where only the single 5s electron is delocalized, similarly to copper and gold.[14] Unlike metals with incomplete d-shells, metallic bonds in silver are lacking a covalent character and are relatively weak. This observation explains the low hardness and high ductility of single crystals of silver.[15]

Silver has a brilliant, white, metallic luster that can take a high polish,[16] and which is so characteristic that the name of the metal itself has become a color name.[13] Protected silver has greater optical reflectivity than aluminium at all wavelengths longer than ~450 nm.[17] At wavelengths shorter than 450 nm, silver's reflectivity is inferior to that of aluminium and drops to zero near 310 nm.[18]

Very high electrical and thermal conductivity are common to the elements in group 11, because their single s electron is free and does not interact with the filled d subshell, as such interactions (which occur in the preceding transition metals) lower electron mobility.[19] The thermal conductivity of silver is among the highest of all materials, although the thermal conductivity of carbon (in the diamond allotrope) and superfluid helium-4 are higher.[12] The electrical conductivity of silver is the highest of all metals, greater even than copper. Silver also has the lowest contact resistance of any metal.[12] Silver is rarely used for its electrical conductivity, due to its high cost, although an exception is in radio-frequency engineering, particularly at VHF and higher frequencies where silver plating improves electrical conductivity because those currents tend to flow on the surface of conductors rather than through the interior. During World War II in the US, 13540 tons of silver were used for the electromagnets in calutrons for enriching uranium, mainly because of the wartime shortage of copper.[20][21][22]

Silver readily forms alloys with copper, gold, and zinc. Zinc-silver alloys with low zinc concentration may be considered as face-centred cubic solid solutions of zinc in silver, as the structure of the silver is largely unchanged while the electron concentration rises as more zinc is added. Increasing the electron concentration further leads to body-centred cubic (electron concentration 1.5), complex cubic (1.615), and hexagonal close-packed phases (1.75).[14]

Isotopes

Main article: Isotopes of silver

Naturally occurring silver is composed of two stable isotopes, 107Ag and 109Ag, with 107Ag being slightly more abundant (51.839% natural abundance). This almost equal abundance is rare in the periodic table. The atomic weight is 107.8682(2) u;[23][24] this value is very important because of the importance of silver compounds, particularly halides, in gravimetric analysis.[23] Both isotopes of silver are produced in stars via the s-process (slow neutron capture), as well as in supernovas via the r-process (rapid neutron capture).[25]

Twenty-eight radioisotopes have been characterized, the most stable being 105Ag with a half-life of 41.29 days, 111Ag with a half-life of 7.45 days, and 112Ag with a half-life of 3.13 hours. Silver has numerous nuclear isomers, the most stable being 108mAg (t1/2 = 418 years), 110mAg (t1/2 = 249.79 days) and 106mAg (t1/2 = 8.28 days). All of the remaining radioactive isotopes have half-lives of less than an hour, and the majority of these have half-lives of less than three minutes.[26]

Isotopes of silver range in relative atomic mass from 92.950 u (93Ag) to 129.950 u (130Ag);[27] the primary decay mode before the most abundant stable isotope, 107Ag, is electron capture and the primary mode after is beta decay. The primary decay products before 107Ag are palladium (element 46) isotopes, and the primary products after are cadmium (element 48) isotopes.[26]

The palladium isotope 107Pd decays by beta emission to 107Ag with a half-life of 6.5 million years. Iron meteorites are the only objects with a high-enough palladium-to-silver ratio to yield measurable variations in 107Ag abundance. Radiogenic 107Ag was first discovered in the Santa Clara meteorite in 1978.[28] 107Pd–107Ag correlations observed in bodies that have clearly been melted since the accretion of the Solar System must reflect the presence of unstable nuclides in the early Solar System.[29]

Chemistry

Oxidation states and stereochemistries of silver[30]

Oxidation

state Coordination

number Stereochemistry Representative

compound

0 (d10s1) 3 Planar Ag(CO)3

1 (d10) 2 Linear [Ag(CN)2]−

3 Trigonal planar AgI(PEt2Ar)2

4 Tetrahedral [Ag(diars)2]+

6 Octahedral AgF, AgCl, AgBr

2 (d9) 4 Square planar [Ag(py)4]2+

3 (d8) 4 Square planar [AgF4]−

6 Octahedral [AgF6]3−

Silver is a rather unreactive metal. This is because its filled 4d shell is not very effective in shielding the electrostatic forces of attraction from the nucleus to the outermost 5s electron, and hence silver is near the bottom of the electrochemical series (E0(Ag+/Ag) = +0.799 V).[13] In group 11, silver has the lowest first ionization energy (showing the instability of the 5s orbital), but has higher second and third ionization energies than copper and gold (showing the stability of the 4d orbitals), so that the chemistry of silver is predominantly that of the +1 oxidation state, reflecting the increasingly limited range of oxidation states along the transition series as the d-orbitals fill and stabilize.[31] Unlike copper, for which the larger hydration energy of Cu2+ as compared to Cu+ is the reason why the former is the more stable in aqueous solution and solids despite lacking the stable filled d-subshell of the latter, with silver this effect is swamped by its larger second ionisation energy. Hence, Ag+ is the stable species in aqueous solution and solids, with Ag2+ being much less stable as it oxidizes water.[31]

Most silver compounds have significant covalent character due to the small size and high first ionization energy (730.8 kJ/mol) of silver.[13] Furthermore, silver's Pauling electronegativity of 1.93 is higher than that of lead (1.87), and its electron affinity of 125.6 kJ/mol is much higher than that of hydrogen (72.8 kJ/mol) and not much less than that of oxygen (141.0 kJ/mol).[32] Due to its full d-subshell, silver in its main +1 oxidation state exhibits relatively few properties of the transition metals proper from groups 4 to 10, forming rather unstable organometallic compounds, forming linear complexes showing very low coordination numbers like 2, and forming an amphoteric oxide[33] as well as Zintl phases like the post-transition metals.[34] Unlike the preceding transition metals, the +1 oxidation state of silver is stable even in the absence of π-acceptor ligands.[31]

Silver does not react with air, even at red heat, and thus was considered by alchemists as a noble metal, along with gold. Its reactivity is intermediate between that of copper (which forms copper(I) oxide when heated in air to red heat) and gold. Like copper, silver reacts with sulfur and its compounds; in their presence, silver tarnishes in air to form the black silver sulfide (copper forms the green sulfate instead, while gold does not react). While silver is not attacked by non-oxidizing acids, the metal dissolves readily in hot concentrated sulfuric acid, as well as dilute or concentrated nitric acid. In the presence of air, and especially in the presence of hydrogen peroxide, silver dissolves readily in aqueous solutions of cyanide.[30]

The three main forms of deterioration in historical silver artifacts are tarnishing, formation of silver chloride due to long-term immersion in salt water, as well as reaction with nitrate ions or oxygen. Fresh silver chloride is pale yellow, becoming purplish on exposure to light; it projects slightly from the surface of the artifact or coin. The precipitation of copper in ancient silver can be used to date artifacts, as copper is nearly always a constituent of silver alloys.[35]

Silver metal is attacked by strong oxidizers such as potassium permanganate (KMnO

4) and potassium dichromate (K

2Cr

2O

7), and in the presence of potassium bromide (KBr). These compounds are used in photography to bleach silver images, converting them to silver bromide that can either be fixed with thiosulfate or redeveloped to intensify the original image. Silver forms cyanide complexes (silver cyanide) that are soluble in water in the presence of an excess of cyanide ions. Silver cyanide solutions are used in electroplating of silver.[36]

The common oxidation states of silver are (in order of commonness): +1 (the most stable state; for example, silver nitrate, AgNO3); +2 (highly oxidising; for example, silver(II) fluoride, AgF2); and even very rarely +3 (extreme oxidising; for example, potassium tetrafluoroargentate(III), KAgF4).[37] The +3 state requires very strong oxidising agents to attain, such as fluorine or peroxodisulfate, and some silver(III) compounds react with atmospheric moisture and attack glass.[38] Indeed, silver(III) fluoride is usually obtained by reacting silver or silver monofluoride with the strongest known oxidizing agent, krypton difluoride.[39]