Re: Silver Dat Ii Keygen Torrent
Tacio Allaire
Silver has long been valued as a precious metal. Silver metal is used in many bullion coins, sometimes alongside gold:[7] while it is more abundant than gold, it is much less abundant as a native metal.[8] 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.
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.[10] 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.[11]
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.[12] 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.[13]
Silver has a brilliant, white, metallic luster that can take a high polish,[14] and which is so characteristic that the name of the metal itself has become a color name.[11] Protected silver has greater optical reflectivity than aluminium at all wavelengths longer than 450 nm.[15] At wavelengths shorter than 450 nm, silver's reflectivity is inferior to that of aluminium and drops to zero near 310 nm.[16]
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.[17] 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.[10] 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.[10] 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.[18][19][20]
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).[12]
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;[21][22] this value is very important because of the importance of silver compounds, particularly halides, in gravimetric analysis.[21] 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).[23]
Isotopes of silver range in relative atomic mass from 92.950 u (93Ag) to 129.950 u (130Ag);[25] 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.[24]
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).[11] 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.[29] 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.[29]
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.[28]
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.[33]
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.[34]
Silver and gold have rather low chemical affinities for oxygen, lower than copper, and it is therefore expected that silver oxides are thermally quite unstable. Soluble silver(I) salts precipitate dark-brown silver(I) oxide, Ag2O, upon the addition of alkali. (The hydroxide AgOH exists only in solution; otherwise it spontaneously decomposes to the oxide.) Silver(I) oxide is very easily reduced to metallic silver, and decomposes to silver and oxygen above 160 C.[38] This and other silver(I) compounds may be oxidized by the strong oxidizing agent peroxodisulfate to black AgO, a mixed silver(I,III) oxide of formula AgIAgIIIO2. Some other mixed oxides with silver in non-integral oxidation states, namely Ag2O3 and Ag3O4, are also known, as is Ag3O which behaves as a metallic conductor.[38]
Silver(I) sulfide, Ag2S, is very readily formed from its constituent elements and is the cause of the black tarnish on some old silver objects. It may also be formed from the reaction of hydrogen sulfide with silver metal or aqueous Ag+ ions. Many non-stoichiometric selenides and tellurides are known; in particular, AgTe3 is a low-temperature superconductor.[38]
The only known dihalide of silver is the difluoride, AgF2, which can be obtained from the elements under heat. A strong yet thermally stable and therefore safe fluorinating agent, silver(II) fluoride is often used to synthesize hydrofluorocarbons.[39]
White silver nitrate, AgNO3, is a versatile precursor to many other silver compounds, especially the halides, and is much less sensitive to light. It was once called lunar caustic because silver was called luna by the ancient alchemists, who believed that silver was associated with the Moon.[41][42] It is often used for gravimetric analysis, exploiting the insolubility of the heavier silver halides which it is a common precursor to.[21] Silver nitrate is used in many ways in organic synthesis, e.g. for deprotection and oxidations. Ag+ binds alkenes reversibly, and silver nitrate has been used to separate mixtures of alkenes by selective absorption. The resulting adduct can be decomposed with ammonia to release the free alkene.[43]
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