Cobalt is a chemical element with symbol Co and atomic number 27. Like nickel, cobalt is found in the Earth's crust only in chemically combined form, save for small deposits found in alloys of natural meteoric iron; the free element, produced by reductive smelting, is a hard, silver-gray metal. Cobalt-based blue pigments have been used since ancient times for jewelry and paints, to impart a distinctive blue tint to glass, but the color was thought by alchemists to be due to the known metal bismuth. Miners had long used the name kobold ore for some of the blue-pigment producing minerals. In 1735, such ores were found to be reducible to a new metal, this was named for the kobold. Today, some cobalt is produced from one of a number of metallic-lustered ores, such as for example cobaltite; the element is however more produced as a by-product of copper and nickel mining. The copper belt in the Democratic Republic of the Congo and Zambia yields most of the global cobalt production; the DRC alone accounted for more than 50% of world production in 2016, according to Natural Resources Canada.
Cobalt is used in the manufacture of magnetic, wear-resistant and high-strength alloys. The compounds cobalt silicate and cobalt aluminate give a distinctive deep blue color to glass, inks and varnishes. Cobalt occurs as only one stable isotope, cobalt-59. Cobalt-60 is a commercially important radioisotope, used as a radioactive tracer and for the production of high energy gamma rays. Cobalt is the active center of a group of coenzymes called cobalamins. Vitamin B12, the best-known example of the type, is an essential vitamin for all animals. Cobalt in inorganic form is a micronutrient for bacteria and fungi. Cobalt is a ferromagnetic metal with a specific gravity of 8.9. The Curie temperature is 1,115 °C and the magnetic moment is 1.6–1.7 Bohr magnetons per atom. Cobalt has a relative permeability two-thirds. Metallic cobalt occurs as two crystallographic structures: fcc; the ideal transition temperature between the hcp and fcc structures is 450 °C, but in practice the energy difference between them is so small that random intergrowth of the two is common.
Cobalt is a weakly reducing metal, protected from oxidation by a passivating oxide film. It is attacked by halogens and sulfur. Heating in oxygen produces Co3O4 which loses oxygen at 900 °C to give the monoxide CoO; the metal reacts with fluorine at 520 K to give CoF3. It does not react with hydrogen gas or nitrogen gas when heated, but it does react with boron, phosphorus and sulfur. At ordinary temperatures, it reacts with mineral acids, slowly with moist, but not with dry, air. Common oxidation states of cobalt include +2 and +3, although compounds with oxidation states ranging from −3 to +5 are known. A common oxidation state for simple compounds is +2; these salts form the pink-colored metal aquo complex 2+ in water. Addition of chloride gives the intensely blue 2−. In a borax bead flame test, cobalt shows deep blue in both reducing flames. Several oxides of cobalt are known. Green cobalt oxide has rocksalt structure, it is oxidized with water and oxygen to brown cobalt hydroxide. At temperatures of 600 -- 700 °C, CoO oxidizes to the blue cobalt oxide.
Black cobalt oxide is known. Cobalt oxides are antiferromagnetic at low temperature: CoO and Co3O4, analogous to magnetite, with a mixture of +2 and +3 oxidation states; the principal chalcogenides of cobalt include the black cobalt sulfides, CoS2, which adopts a pyrite-like structure, cobalt sulfide. Four dihalides of cobalt are known: cobalt fluoride, cobalt chloride, cobalt bromide, cobalt iodide; these halides exist in hydrated forms. Whereas the anhydrous dichloride is blue, the hydrate is red; the reduction potential for the reaction Co3+ + e− → Co2+ is +1.92 V, beyond that for chlorine to chloride, +1.36 V. Consequently and chloride would result in the cobalt being reduced to cobalt; because the reduction potential for fluorine to fluoride is so high, +2.87 V, cobalt fluoride is one of the few simple stable cobalt compounds. Cobalt fluoride, used in some fluorination reactions, reacts vigorously with water; as for all metals, molecular compounds and polyatomic ions of cobalt are classified as coordination complexes, that is, molecules or ions that contain cobalt linked to several ligands.
The principles of electronegativity and hardness–softness of a series of ligands can be used to explain the usual oxidation state of cobalt. For example, Co+3 complexes tend to have ammine ligands; because phosphorus is softer than nitrogen, phosphine ligands tend to feature the softer Co2+ and Co+, an example being triscobalt chloride. The more electronegative oxide and fluoride can stabilize Co4+ and Co5+ derivatives, e.g. caesium hexafluorocobaltate and potassium percobaltate. Alfred Werner, a Nobel-prize winning pioneer in coordination chemistry, worked with compounds of empirical formula 3+. One of the isomers
Gold is a chemical element with symbol Au and atomic number 79, making it one of the higher atomic number elements that occur naturally. In its purest form, it is a bright reddish yellow, soft and ductile metal. Chemically, gold is a group 11 element, it is solid under standard conditions. Gold occurs in free elemental form, as nuggets or grains, in rocks, in veins, in alluvial deposits, it occurs in a solid solution series with the native element silver and naturally alloyed with copper and palladium. Less it occurs in minerals as gold compounds with tellurium. Gold is resistant to most acids, though it does dissolve in aqua regia, a mixture of nitric acid and hydrochloric acid, which forms a soluble tetrachloroaurate anion. Gold is insoluble in nitric acid, which dissolves silver and base metals, a property that has long been used to refine gold and to confirm the presence of gold in metallic objects, giving rise to the term acid test. Gold dissolves in alkaline solutions of cyanide, which are used in mining and electroplating.
Gold dissolves in mercury, forming amalgam alloys. A rare element, gold is a precious metal, used for coinage and other arts throughout recorded history. In the past, a gold standard was implemented as a monetary policy, but gold coins ceased to be minted as a circulating currency in the 1930s, the world gold standard was abandoned for a fiat currency system after 1971. A total of 186,700 tonnes of gold exists above ground, as of 2015; the world consumption of new gold produced is about 50% in jewelry, 40% in investments, 10% in industry. Gold's high malleability, resistance to corrosion and most other chemical reactions, conductivity of electricity have led to its continued use in corrosion resistant electrical connectors in all types of computerized devices. Gold is used in infrared shielding, colored-glass production, gold leafing, tooth restoration. Certain gold salts are still used as anti-inflammatories in medicine; as of 2017, the world's largest gold producer by far was China with 440 tonnes per year.
Gold is the most malleable of all metals. It can be drawn into a monoatomic wire, stretched about twice before it breaks; such nanowires distort via formation and migration of dislocations and crystal twins without noticeable hardening. A single gram of gold can be beaten into a sheet of 1 square meter, an avoirdupois ounce into 300 square feet. Gold leaf can be beaten thin enough to become semi-transparent; the transmitted light appears greenish blue, because gold reflects yellow and red. Such semi-transparent sheets strongly reflect infrared light, making them useful as infrared shields in visors of heat-resistant suits, in sun-visors for spacesuits. Gold is a good conductor of electricity. Gold has a density of 19.3 g/cm3 identical to that of tungsten at 19.25 g/cm3. By comparison, the density of lead is 11.34 g/cm3, that of the densest element, osmium, is 22.588±0.015 g/cm3. Whereas most metals are gray or silvery white, gold is reddish-yellow; this color is determined by the frequency of plasma oscillations among the metal's valence electrons, in the ultraviolet range for most metals but in the visible range for gold due to relativistic effects affecting the orbitals around gold atoms.
Similar effects impart a golden hue to metallic caesium. Common colored gold alloys include the distinctive eighteen-karat rose gold created by the addition of copper. Alloys containing palladium or nickel are important in commercial jewelry as these produce white gold alloys. Fourteen-karat gold-copper alloy is nearly identical in color to certain bronze alloys, both may be used to produce police and other badges. White gold alloys can be made with nickel. Fourteen- and eighteen-karat gold alloys with silver alone appear greenish-yellow and are referred to as green gold. Blue gold can be made by alloying with iron, purple gold can be made by alloying with aluminium. Less addition of manganese, aluminium and other elements can produce more unusual colors of gold for various applications. Colloidal gold, used by electron-microscopists, is red. Gold has only one stable isotope, 197Au, its only occurring isotope, so gold is both a mononuclidic and monoisotopic element. Thirty-six radioisotopes have been synthesized, ranging in atomic mass from 169 to 205.
The most stable of these is 195Au with a half-life of 186.1 days. The least stable is 171Au. Most of gold's radioisotopes with atomic masses below 197 decay by some combination of proton emission, α decay, β+ decay; the exceptions are 195Au, which decays by electron capture, 196Au, which decays most by electron capture with a minor β− decay path. All of gold's radioisotopes with atomic masses above 197 decay by β− decay. At least 32 nuclear isomers have been characterized, ranging in atomic mass from 170 to 200. Within that range, only 178Au, 180Au, 181Au, 182Au, 188Au do not have isomers. Gold's most stable isomer is 198m2Au with a half-life of 2.27 days. Gold's least stable isomer is 177m2Au with a half-life of only 7 ns. 184m1Au has three decay paths: β+ decay, isomeric
A metal is a material that, when freshly prepared, polished, or fractured, shows a lustrous appearance, conducts electricity and heat well. Metals are malleable or ductile. A metal may be an alloy such as stainless steel. In physics, a metal is regarded as any substance capable of conducting electricity at a temperature of absolute zero. Many elements and compounds that are not classified as metals become metallic under high pressures. For example, the nonmetal iodine becomes a metal at a pressure of between 40 and 170 thousand times atmospheric pressure; some materials regarded as metals can become nonmetals. Sodium, for example, becomes a nonmetal at pressure of just under two million times atmospheric pressure. In chemistry, two elements that would otherwise qualify as brittle metals—arsenic and antimony—are instead recognised as metalloids, on account of their predominately non-metallic chemistry. Around 95 of the 118 elements in the periodic table are metals; the number is inexact as the boundaries between metals and metalloids fluctuate due to a lack of universally accepted definitions of the categories involved.
In astrophysics the term "metal" is cast more to refer to all chemical elements in a star that are heavier than the lightest two and helium, not just traditional metals. A star fuses lighter atoms hydrogen and helium, into heavier atoms over its lifetime. Used in that sense, the metallicity of an astronomical object is the proportion of its matter made up of the heavier chemical elements. Metals are present in many aspects of modern life; the strength and resilience of some metals has led to their frequent use in, for example, high-rise building and bridge construction, as well as most vehicles, many home appliances, tools and railroad tracks. Precious metals were used as coinage, but in the modern era, coinage metals have extended to at least 23 of the chemical elements; the history of metals is thought to begin with the use of copper about 11,000 years ago. Gold, iron and brass were in use before the first known appearance of bronze in the 5th millennium BCE. Subsequent developments include the production of early forms of steel.
Metals are lustrous, at least when freshly prepared, polished, or fractured. Sheets of metal thicker than a few micrometres appear opaque; the solid or liquid state of metals originates in the capacity of the metal atoms involved to lose their outer shell electrons. Broadly, the forces holding an individual atom’s outer shell electrons in place are weaker than the attractive forces on the same electrons arising from interactions between the atoms in the solid or liquid metal; the electrons involved become delocalised and the atomic structure of a metal can be visualised as a collection of atoms embedded in a cloud of mobile electrons. This type of interaction is called a metallic bond; the strength of metallic bonds for different elemental metals reaches a maximum around the center of the transition metal series, as these elements have large numbers of delocalized electrons. Although most elemental metals have higher densities than most nonmetals, there is a wide variation in their densities, lithium being the least dense and osmium the most dense.
Magnesium and titanium are light metals of significant commercial importance. Their respective densities of 1.7, 2.7 and 4.5 g/cm3 can be compared to those of the older structural metals, like iron at 7.9 and copper at 8.9 g/cm3. An iron ball would thus weigh about as much as three aluminium balls. Metals are malleable and ductile, deforming under stress without cleaving; the nondirectional nature of metallic bonding is thought to contribute to the ductility of most metallic solids. In contrast, in an ionic compound like table salt, when the planes of an ionic bond slide past one another, the resultant change in location shifts ions of the same charge into close proximity, resulting in the cleavage of the crystal; such a shift is not observed in a covalently bonded crystal, such as a diamond, where fracture and crystal fragmentation occurs. Reversible elastic deformation in metals can be described by Hooke's Law for restoring forces, where the stress is linearly proportional to the strain. Heat or forces larger than a metal's elastic limit may cause a permanent deformation, known as plastic deformation or plasticity.
An applied force may be a compressive force, or a shear, bending or torsion force. A temperature change may affect the movement or displacement of structural defects in the metal such as grain boundaries, point vacancies and screw dislocations, stacking faults and twins in both crystalline and non-crystalline metals. Internal slip and metal fatigue may ensue; the atoms of metallic substances are arranged in one of three common crystal structures, namely body-centered cubic, face-centered cubic, hexagonal close-packed. In bcc, each atom is positioned at the center of a cube of eight others. In fcc and hcp, each atom is surrounded by twelve others; some metals adopt different structures depending on the temperature. The
Chromium is a chemical element with symbol Cr and atomic number 24. It is the first element in group 6, it is a steely-grey, lustrous and brittle transition metal. Chromium boasts a high usage rate as a metal, able to be polished while resisting tarnishing. Chromium is the main additive in stainless steel, a popular steel alloy due to its uncommonly high specular reflection. Simple polished chromium reflects 70% of the visible spectrum, with 90% of infrared light being reflected; the name of the element is derived from the Greek word χρῶμα, chrōma, meaning color, because many chromium compounds are intensely colored. Ferrochromium alloy is commercially produced from chromite by silicothermic or aluminothermic reactions and chromium metal by roasting and leaching processes followed by reduction with carbon and aluminium. Chromium metal is of high value for hardness. A major development in steel production was the discovery that steel could be made resistant to corrosion and discoloration by adding metallic chromium to form stainless steel.
Stainless steel and chrome plating together comprise 85% of the commercial use. In the United States, trivalent chromium ion is considered an essential nutrient in humans for insulin and lipid metabolism. However, in 2014, the European Food Safety Authority, acting for the European Union, concluded that there was not sufficient evidence for chromium to be recognized as essential. While chromium metal and Cr ions are not considered toxic, hexavalent chromium is both toxic and carcinogenic. Abandoned chromium production sites require environmental cleanup. Chromium is the fourth transition metal found on the periodic table, has an electron configuration of 3d5 4s1, it is the first element in the periodic table whose ground-state electron configuration violates the Aufbau principle. This occurs again in the periodic table with other elements and their electron configurations, such as copper and molybdenum; this occurs. In the previous elements, the energetic cost of promoting an electron to the next higher energy level is too great to compensate for that released by lessening inter-electronic repulsion.
However, in the 3d transition metals, the energy gap between the 3d and the next-higher 4s subshell is small, because the 3d subshell is more compact than the 4s subshell, inter-electron repulsion is smaller between 4s electrons than between 3d electrons. This lowers the energetic cost of promotion and increases the energy released by it, so that the promotion becomes energetically feasible and one or two electrons are always promoted to the 4s subshell. Chromium is the first element in the 3d series where the 3d electrons start to sink into the inert core. Chromium is a strong oxidising agent in contrast to the tungsten oxides. Chromium is hard, is the third hardest element behind carbon and boron, its Mohs hardness is 8.5, which means that it can scratch samples of quartz and topaz, but can be scratched by corundum. Chromium is resistant to tarnishing, which makes it useful as a metal that preserves its outermost layer from corroding, unlike other metals such as copper and aluminium. Chromium has a melting point of 1907 °C, low compared to the majority of transition metals.
However, it still has the second highest melting point out of all the Period 4 elements, being topped by vanadium by 3 °C at 1910 °C. The boiling point of 2671 °C, however, is comparatively lower, having the third lowest boiling point out of the Period 4 transition metals alone behind manganese and zinc. Chromium has an unusually high specular reflection in comparison to that of other transition metals. At 425 μm, chromium was found to have a relative maximum reflection of about 72% reflectance, before entering a depression in reflectivity, reaching a minimum of 62% reflectance at 750 μm before rising again to reflecting 90% of 4000 μm of infrared waves.. When chromium is formed into a stainless steel alloy and polished, the specular reflection decreases with the inclusion of additional metals, yet is still rather high in comparison with other alloys. Between 40% and 60% of the visible spectrum is reflected from polished stainless steel; the explanation on why chromium displays such a high turnout of reflected photon waves in general the 90% of infrared waves that were reflected, can be attributed to chromium's magnetic properties.
Chromium has unique magnetic properties in the sense that chromium is the only elemental solid which shows antiferromagnetic ordering at room temperature. Above 38 °C, its magnetic ordering changes to paramagnetic.. The antiferromagnetic properties, which cause the chromium atoms to temporarily ionize and bond with themselves, are present because the body-centric cubic's magnetic properties are disproportionate to the lattice periodicity; this is due to the fact that the magnetic moments at the cube's corners and the cube centers are not equal, but are still antiparallel. From here, the frequency-dependent relative permittivity of chromium, deriving from Maxwell's equations in conjunction with chromium's antiferromagnetivity, leaves chromium with a high infrared and visible light reflectance. Chromium metal left standing in air is passivated by oxidation, forming a th
Tin is a chemical element with the symbol Sn and atomic number 50. It is a post-transition metal in group 14 of the periodic table of elements, it is obtained chiefly from the mineral cassiterite, which contains stannic oxide, SnO2. Tin shows a chemical similarity to both of its neighbors in group 14, germanium and lead, has two main oxidation states, +2 and the more stable +4. Tin is the 49th most abundant element and has, with 10 stable isotopes, the largest number of stable isotopes in the periodic table, thanks to its magic number of protons, it has two main allotropes: at room temperature, the stable allotrope is β-tin, a silvery-white, malleable metal, but at low temperatures it transforms into the less dense grey α-tin, which has the diamond cubic structure. Metallic tin does not oxidize in air; the first tin alloy used on a large scale was bronze, made of 1/8 tin and 7/8 copper, from as early as 3000 BC. After 600 BC, pure metallic tin was produced. Pewter, an alloy of 85–90% tin with the remainder consisting of copper and lead, was used for flatware from the Bronze Age until the 20th century.
In modern times, tin is used in many alloys, most notably tin/lead soft solders, which are 60% or more tin, in the manufacture of transparent, electrically conducting films of indium tin oxide in optoelectronic applications. Another large application for tin is corrosion-resistant tin plating of steel; because of the low toxicity of inorganic tin, tin-plated steel is used for food packaging as tin cans. However, some organotin compounds can be as toxic as cyanide. Tin is a soft, malleable and crystalline silvery-white metal; when a bar of tin is bent, a crackling sound known as the "tin cry" can be heard from the twinning of the crystals. Tin melts at low temperatures of about 232 °C, the lowest in group 14; the melting point is further lowered to 177.3 °C for 11 nm particles. Β-tin, stable at and above room temperature, is malleable. In contrast, α-tin, stable below 13.2 °C, is brittle. Α-tin has a diamond cubic crystal structure, similar to silicon or germanium. Α-tin has no metallic properties at all because its atoms form a covalent structure in which electrons cannot move freely.
It is a dull-gray powdery material with no common uses other than a few specialized semiconductor applications. These two allotropes, α-tin and β-tin, are more known as gray tin and white tin, respectively. Two more allotropes, γ and σ, exist at temperatures above 161 pressures above several GPa. In cold conditions, β-tin tends to transform spontaneously into α-tin, a phenomenon known as "tin pest". Although the α-β transformation temperature is nominally 13.2 °C, impurities lower the transition temperature well below 0 °C and, on the addition of antimony or bismuth, the transformation might not occur at all, increasing the durability of the tin. Commercial grades of tin resist transformation because of the inhibiting effect of the small amounts of bismuth, antimony and silver present as impurities. Alloying elements such as copper, bismuth and silver increase its hardness. Tin tends rather to form hard, brittle intermetallic phases, which are undesirable, it does not form wide solid solution ranges in other metals in general, few elements have appreciable solid solubility in tin.
Simple eutectic systems, occur with bismuth, lead and zinc. Tin was one of the first superconductors to be studied. Tin can be attacked by acids and alkalis. Tin can be polished and is used as a protective coat for other metals. A protective oxide layer prevents further oxidation, the same that forms on pewter and other tin alloys. Tin helps to accelerate the chemical reaction. Tin has ten stable isotopes, with atomic masses of 112, 114 through 120, 122 and 124, the greatest number of any element. Of these, the most abundant are 120Sn, 118Sn, 116Sn, while the least abundant is 115Sn; the isotopes with mass numbers have no nuclear spin, while those with odd have a spin of +1/2. Tin, with its three common isotopes 116Sn, 118Sn and 120Sn, is among the easiest elements to detect and analyze by NMR spectroscopy, its chemical shifts are referenced against SnMe4; this large number of stable isotopes is thought to be a direct result of the atomic number 50, a "magic number" in nuclear physics. Tin occurs in 29 unstable isotopes, encompassing all the remaining atomic masses from 99 to 137.
Apart from 126Sn, with a half-life of 230,000 years, all the radioisotopes have a half-life of less than a year. The radioactive 100Sn, discovered in 1994, 132Sn are one of the few nuclides with a "doubly magic" nucleus: despite being unstable, having lopsided proton–neutron ratios, they represent endpoints beyond which stability drops off rapidly. Another 30 metastable isomers have been characterized for isotopes between 111 and 131, the most stable being 121mSn with a half-life of 43.9 years. The relative differences in the abundances of tin's stable isotopes can be explained by their different modes of formation in stellar nucleosynthesis. 116Sn through 120Sn inclusive are formed in the s-process in most stars and hence they are the most common isotopes, while 122Sn and 124Sn are only formed in the r-process (rapid neutr
Molybdenum is a chemical element with symbol Mo and atomic number 42. The name is from Neo-Latin molybdaenum, from Ancient Greek Μόλυβδος molybdos, meaning lead, since its ores were confused with lead ores. Molybdenum minerals have been known throughout history, but the element was discovered in 1778 by Carl Wilhelm Scheele; the metal was first isolated in 1781 by Peter Jacob Hjelm. Molybdenum does not occur as a free metal on Earth; the free element, a silvery metal with a gray cast, has the sixth-highest melting point of any element. It forms hard, stable carbides in alloys, for this reason most of world production of the element is used in steel alloys, including high-strength alloys and superalloys. Most molybdenum compounds have low solubility in water, but when molybdenum-bearing minerals contact oxygen and water, the resulting molybdate ion MoO2−4 is quite soluble. Industrially, molybdenum compounds are used in high-pressure and high-temperature applications as pigments and catalysts. Molybdenum-bearing enzymes are by far the most common bacterial catalysts for breaking the chemical bond in atmospheric molecular nitrogen in the process of biological nitrogen fixation.
At least 50 molybdenum enzymes are now known in bacteria and animals, although only bacterial and cyanobacterial enzymes are involved in nitrogen fixation. These nitrogenases contain molybdenum in a form different from other molybdenum enzymes, which all contain oxidized molybdenum in a molybdenum cofactor; these various molybdenum cofactor enzymes are vital to the organisms, molybdenum is an essential element for life in all higher eukaryote organisms, though not in all bacteria. In its pure form, molybdenum is a silvery-grey metal with a Mohs hardness of 5.5, a standard atomic weight of 95.95 g/mol. It has a melting point of 2,623 °C, it has one of the lowest coefficients of thermal expansion among commercially used metals. The tensile strength of molybdenum wires increases about 3 times, from about 10 to 30 GPa, when their diameter decreases from ~50–100 nm to 10 nm. Molybdenum is a transition metal with an electronegativity of 2.16 on the Pauling scale. It does not visibly react with water at room temperature.
Weak oxidation of molybdenum starts at 300 °C. Like many heavier transition metals, molybdenum shows little inclination to form a cation in aqueous solution, although the Mo3+ cation is known under controlled conditions. There are 35 known isotopes of molybdenum, ranging in atomic mass from 83 to 117, as well as four metastable nuclear isomers. Seven isotopes occur with atomic masses of 92, 94, 95, 96, 97, 98, 100. Of these occurring isotopes, only molybdenum-100 is unstable. Molybdenum-98 is the most abundant isotope, comprising 24.14% of all molybdenum. Molybdenum-100 has a half-life of about 1019 y and undergoes double beta decay into ruthenium-100. Molybdenum isotopes with mass numbers from 111 to 117 all have half-lives of 150 ns. All unstable isotopes of molybdenum decay into isotopes of niobium and ruthenium; as noted below, the most common isotopic molybdenum application involves molybdenum-99, a fission product. It is a parent radioisotope to the short-lived gamma-emitting daughter radioisotope technetium-99m, a nuclear isomer used in various imaging applications in medicine.
In 2008, the Delft University of Technology applied for a patent on the molybdenum-98-based production of molybdenum-99. Molybdenum forms chemical compounds in oxidation states from -II to +VI. Higher oxidation states are more relevant to its terrestrial occurrence and its biological roles, mid-level oxidation states are associated with metal clusters, low oxidation states are associated with organomolybdenum compounds. Mo and W chemistry shows strong similarities; the relative rarity of molybdenum, for example, contrasts with the pervasiveness of the chromium compounds. The highest oxidation state is seen in molybdenum oxide, whereas the normal sulfur compound is molybdenum disulfide MoS2. From the perspective of commerce, the most important compounds are molybdenum disulfide and molybdenum trioxide; the black disulfide is the main mineral. It is roasted in air to give the trioxide: 2 MoS2 + 7 O2 → 2 MoO3 + 4 SO2The trioxide, volatile at high temperatures, is the precursor to all other Mo compounds as well as alloys.
Molybdenum has several oxidation states, the most stable being +4 and +6. Molybdenum oxide is soluble in strong alkaline water, forming molybdates. Molybdates are weaker oxidants than chromates, they tend to form structurally complex oxyanions by condensation at lower pH values, such as 6− and 4−. Polymolybdates can incorporate other ions; the dark-blue phosphorus-containing heteropolymolybdate P3− is used for the spectroscopic detection of phosphorus. The broad range of oxidation states of molybdenum is reflected in various molybdenum chlorides: Molybdenum chloride MoCl2, which exists as the hexamer Mo6Cl12 and the related dianion 2-. Molybdenum chloride MoCl3, a dark red solid, which converts to the anion trianionic complex 3-. Molybdenum chloride MoCl4, a black solid, which adopts a polymeric structure. Molybdenum chloride MoCl5 dark green solid that
Borax known as sodium borate, sodium tetraborate, or disodium tetraborate, is an important boron compound, a mineral, a salt of boric acid. Powdered borax is white, consisting of soft colorless crystals that dissolve in water. A number of related minerals or chemical compounds that differ in their crystal water content are referred to as borax, but the word is used to refer to the octahydrate. Commercially sold borax is dehydrated. Borax is a component of many detergents and enamel glazes, it is used to make buffer solutions in biochemistry, as a fire retardant, as an anti-fungal compound, in the manufacture of fiberglass, as a flux in metallurgy, neutron-capture shields for radioactive sources, a texturing agent in cooking, as a precursor for other boron compounds, along with its inverse, boric acid, is useful as an insecticide. In artisanal gold mining, borax is sometimes used as part of a process meant to eliminate the need for toxic mercury in the gold extraction process, although it cannot directly replace mercury.
Borax was used by gold miners in parts of the Philippines in the 1900s. Borax was first discovered in dry lake beds in Tibet and was imported via the Silk Road to the Arabian Peninsula in the 8th century AD. Borax first came into common use in the late 19th century when Francis Marion Smith's Pacific Coast Borax Company began to market and popularize a large variety of applications under the 20 Mule Team Borax trademark, named for the method by which borax was hauled out of the California and Nevada deserts; the term borax is used for a number of related minerals or chemical compounds that differ in their crystal water content: anhydrous sodium tetraborate, Na2B4O7 sodium tetraborate pentahydrate, Na2B4O7·5H2O sodium tetraborate decahydrate, Na2B4O7·10H2O or equivalently the octahydrate, Na2B4O54·8H2OFrom the chemical perspective, borax contains the 2− ion. In this structure, there are two four-coordinate boron centers and two three-coordinate boron centers. Borax is easily converted to boric acid and other borates, which have many applications.
Its reaction with hydrochloric acid to form boric acid is: Na2B4O7·10H2O + 2 HCl → 4 B3 + 2 NaCl + 5 H2OThe "decahydrate" is sufficiently stable to find use as a primary standard for acid base titrimetry. When borax is added to a flame, it produces a yellow green color. Borax is not used for this purpose in fireworks due to the overwhelming yellow color of sodium. Boric acid is used to color methanol flames a transparent green. Borax is soluble in ethylene glycol, moderately soluble in diethylene glycol and methanol soluble in acetone, it is poorly soluble in cold water, but its solubility increases with temperature. The English word borax is Latinized: the Middle English form was boras, from Old French boras, bourras; that may have been from medieval Latin baurach, borax, along with Spanish borrax and Italian borrace, in the 9th century. Another name for borax is tincal, from Sanskrit; the word tincal "tinkle", or tincar "tinker", refers to crude borax, before it is purified, as mined from lake deposits in Tibet and other parts of Asia.
The word was adopted in the 17th century from Malay tingkal and from Urdu/Persian/Arabic تنکار tinkār/tankār. These all appear to be related to the Sanskrit टांकण ṭānkaṇa. Borax occurs in evaporite deposits produced by the repeated evaporation of seasonal lakes; the most commercially important deposits are found in: Turkey. Borax has been found at many other locations in the Southwestern United States, the Atacama desert in Chile, newly discovered deposits in Bolivia, in Tibet and Romania. Borax can be produced synthetically from other boron compounds. Occurring borax is refined by a process of recrystallization. Borax is used in various household laundry and cleaning products, including the "20 Mule Team Borax" laundry booster, "Boraxo" powdered hand soap, some tooth bleaching formulas. Borate ions are used in biochemical and chemical laboratories to make buffers, e.g. for polyacrylamide gel electrophoresis of DNA and RNA, such as TBE buffer or the newer SB buffer or BBS buffer in coating procedures.
Borate buffers are used as preferential equilibration solution in dimethyl pimelimidate based crosslinking reactions. Borax as a source of borate has been used to take advantage of the co-complexing ability of borate with other agents in water to form complex ions with various substances. Borate and a suitable polymer bed are used to chromatograph non-glycosylated hemoglobin differentially from glycosylated hemoglobin, an indicator of long term hyperglycemia in diabetes mellitus. Borax alone does not have a high affinity for the hardness cations, although it has been used for water-softening, its chemical equation for water-softening is given below: Ca2+ + Na2B4O7 → CaB4O7 ↓ + 2 Na+ Mg2+ + Na2B4O7 → MgB4O7 ↓ + 2 Na+ The sodium ions introduced do not make water ‘hard’. This method is suitable for removing both permanent types of hardness. A mixture of borax and ammonium chloride is used as a flux when welding steel, it lowers the melting point of the unwanted iron oxide. Borax is used mixed with water as a flux when soldering jewelry metals such as gold or silver, where it allows the molten solder to wet the metal and flow evenly int