Lead is a chemical element with symbol Pb and atomic number 82. It is a heavy metal, denser than most common materials. Lead is soft and malleable, has a low melting point; when freshly cut, lead is silvery with a hint of blue. Lead has the highest atomic number of any stable element and three of its isotopes each include a major decay chain of heavier elements. Lead is a unreactive post-transition metal, its weak metallic character is illustrated by its amphoteric nature. Compounds of lead are found in the +2 oxidation state rather than the +4 state common with lighter members of the carbon group. Exceptions are limited to organolead compounds. Like the lighter members of the group, lead tends to bond with itself. Lead is extracted from its ores. Galena, a principal ore of lead bears silver, interest in which helped initiate widespread extraction and use of lead in ancient Rome. Lead production declined after the fall of Rome and did not reach comparable levels until the Industrial Revolution. In 2014, the annual global production of lead was about ten million tonnes, over half of, from recycling.
Lead's high density, low melting point and relative inertness to oxidation make it useful. These properties, combined with its relative abundance and low cost, resulted in its extensive use in construction, batteries and shot, solders, fusible alloys, white paints, leaded gasoline, radiation shielding. In the late 19th century, lead's toxicity was recognized, its use has since been phased out of many applications. However, many countries still allow the sale of products that expose humans to lead, including some types of paints and bullets. Lead is a toxin that accumulates in soft tissues and bones, it acts as a neurotoxin damaging the nervous system and interfering with the function of biological enzymes, causing neurological disorders, such as brain damage and behavioral problems. A lead atom has 82 electrons, arranged in an electron configuration of 4f145d106s26p2; the sum of lead's first and second ionization energies—the total energy required to remove the two 6p electrons—is close to that of tin, lead's upper neighbor in the carbon group.
This is unusual. The similarity of ionization energies is caused by the lanthanide contraction—the decrease in element radii from lanthanum to lutetium, the small radii of the elements from hafnium onwards; this is due to poor shielding of the nucleus by the lanthanide 4f electrons. The sum of the first four ionization energies of lead exceeds that of tin, contrary to what periodic trends would predict. Relativistic effects, which become significant in heavier atoms, contribute to this behavior. One such effect is the inert pair effect: the 6s electrons of lead become reluctant to participate in bonding, making the distance between nearest atoms in crystalline lead unusually long. Lead's lighter carbon group congeners form stable or metastable allotropes with the tetrahedrally coordinated and covalently bonded diamond cubic structure; the energy levels of their outer s- and p-orbitals are close enough to allow mixing into four hybrid sp3 orbitals. In lead, the inert pair effect increases the separation between its s- and p-orbitals, the gap cannot be overcome by the energy that would be released by extra bonds following hybridization.
Rather than having a diamond cubic structure, lead forms metallic bonds in which only the p-electrons are delocalized and shared between the Pb2+ ions. Lead has a face-centered cubic structure like the sized divalent metals calcium and strontium. Pure lead has a silvery appearance with a hint of blue, it tarnishes on contact with moist air and takes on a dull appearance, the hue of which depends on the prevailing conditions. Characteristic properties of lead include high density, malleability and high resistance to corrosion due to passivation. Lead's close-packed face-centered cubic structure and high atomic weight result in a density of 11.34 g/cm3, greater than that of common metals such as iron and zinc. This density is the origin of the idiom to go over like a lead balloon; some rarer metals are denser: tungsten and gold are both at 19.3 g/cm3, osmium—the densest metal known—has a density of 22.59 g/cm3 twice that of lead. Lead is a soft metal with a Mohs hardness of 1.5. It is somewhat ductile.
The bulk modulus of lead—a measure of its ease of compressibility—is 45.8 GPa. In comparison, that of aluminium is 75.2 GPa. Lead's tensile strength, at 12–17 MPa, is low; the melting point of lead—at 327.5 °C —is low compared to most metals. Its boiling point of 1749 °C is the lowest among the carbon group elements; the electrical resistivity of lead at 20 °C is 192 nanoohm-meters an order of magnitude higher than those of other industrial metals. Lead is a superconductor at temperatures lower than 7.19 K.
Iron is a chemical element with symbol Fe and atomic number 26. It is a metal, that belongs to group 8 of the periodic table, it is by mass the most common element on Earth, forming much of Earth's inner core. It is the fourth most common element in the Earth's crust. Pure iron is rare on the Earth's crust being limited to meteorites. Iron ores are quite abundant, but extracting usable metal from them requires kilns or furnaces capable of reaching 1500 °C or higher, about 500 °C higher than what is enough to smelt copper. Humans started to dominate that process in Eurasia only about 2000 BCE, iron began to displace copper alloys for tools and weapons, in some regions, only around 1200 BCE; that event is considered the transition from the Bronze Age to the Iron Age. Iron alloys, such as steel and special steels are now by far the most common industrial metals, because of their mechanical properties and their low cost. Pristine and smooth pure iron surfaces are mirror-like silvery-gray. However, iron reacts with oxygen and water to give brown to black hydrated iron oxides known as rust.
Unlike the oxides of some other metals, that form passivating layers, rust occupies more volume than the metal and thus flakes off, exposing fresh surfaces for corrosion. The body of an adult human contains about 3 to 5 grams of elemental iron in hemoglobin and myoglobin; these two proteins play essential roles in vertebrate metabolism oxygen transport by blood and oxygen storage in muscles. To maintain the necessary levels, human iron metabolism requires a minimum of iron in the diet. Iron is the metal at the active site of many important redox enzymes dealing with cellular respiration and oxidation and reduction in plants and animals. Chemically, the most common oxidation states of iron are +2 and +3. Iron shares many properties of other transition metals, including the other group 8 elements and osmium. Iron forms compounds in a wide range of oxidation states, −2 to +7. Iron forms many coordination compounds. At least four allotropes of iron are known, conventionally denoted α, γ, δ, ε; the first three forms are observed at ordinary pressures.
As molten iron cools past its freezing point of 1538 °C, it crystallizes into its δ allotrope, which has a body-centered cubic crystal structure. As it cools further to 1394 °C, it changes to its γ-iron allotrope, a face-centered cubic crystal structure, or austenite. At 912 °C and below, the crystal structure again becomes the bcc α-iron allotrope; the physical properties of iron at high pressures and temperatures have been studied extensively, because of their relevance to theories about the cores of the Earth and other planets. Above 10 GPa and temperatures of a few hundred kelvin or less, α-iron changes into another hexagonal close-packed structure, known as ε-iron; the higher-temperature γ-phase changes into ε-iron, but does so at higher pressure. Some controversial experimental evidence exists for a stable β phase at pressures above 50 GPa and temperatures of at least 1500 K, it is supposed to have a double hcp structure. The inner core of the Earth is presumed to consist of an iron-nickel alloy with ε structure.
The melting and boiling points of iron, along with its enthalpy of atomization, are lower than those of the earlier 3d elements from scandium to chromium, showing the lessened contribution of the 3d electrons to metallic bonding as they are attracted more and more into the inert core by the nucleus. This same trend appears for ruthenium but not osmium; the melting point of iron is experimentally well defined for pressures less than 50 GPa. For greater pressures, published data still varies by tens of gigapascals and over a thousand kelvin. Below its Curie point of 770 °C, α-iron changes from paramagnetic to ferromagnetic: the spins of the two unpaired electrons in each atom align with the spins of its neighbors, creating an overall magnetic field; this happens because the orbitals of those two electrons do not point toward neighboring atoms in the lattice, therefore are not involved in metallic bonding. In the absence of an external source of magnetic field, the atoms get spontaneously partitioned into magnetic domains, about 10 micrometres across, such that the atoms in each domain have parallel spins, but different domains have other orientations.
Thus a macroscopic piece of iron will have a nearly zero overall magnetic field. Application of an external magnetic field causes the domains that are magnetized in the same general direction to grow at the expense of adjacent ones that point in other directions, reinforcing the external field; this effect is exploited in devices that needs to channel magnetic fields, such as electrical transformers, magnetic recording heads, electric motors. Impurities, lattice defects, or grain and particle boundaries can "pin" the domains in the new positions, so that the effect persists after the external field is removed -- thus turning the iron object into a magnet. Similar behavior is exhibited by some iron compounds, such as the fer
Boron is a chemical element with symbol B and atomic number 5. Produced by cosmic ray spallation and supernovae and not by stellar nucleosynthesis, it is a low-abundance element in the Solar system and in the Earth's crust. Boron is concentrated on Earth by the water-solubility of its more common occurring compounds, the borate minerals; these are mined industrially as evaporites, such as kernite. The largest known boron deposits are in the largest producer of boron minerals. Elemental boron is a metalloid, found in small amounts in meteoroids but chemically uncombined boron is not otherwise found on Earth. Industrially pure boron is produced with difficulty because of refractory contamination by carbon or other elements. Several allotropes of boron exist: amorphous boron is a brown powder; the primary use of elemental boron is as boron filaments with applications similar to carbon fibers in some high-strength materials. Boron is used in chemical compounds. About half of all boron consumed globally is an additive in fiberglass for insulation and structural materials.
The next leading use is in polymers and ceramics in high-strength, lightweight structural and refractory materials. Borosilicate glass is desired for its greater strength and thermal shock resistance than ordinary soda lime glass. Boron as sodium perborate is used as a bleach. A small amount of boron is used as a dopant in semiconductors, reagent intermediates in the synthesis of organic fine chemicals. A few boron-containing organic pharmaceuticals are in study. Natural boron is composed of two stable isotopes, one of which has a number of uses as a neutron-capturing agent. In biology, borates have low toxicity in mammals, but are more toxic to arthropods and are used as insecticides. Boric acid is mildly antimicrobial, several natural boron-containing organic antibiotics are known. Boron is an essential plant nutrient and boron compounds such as borax and boric acid are used as fertilizers in agriculture, although it's only required in small amounts, with excess being toxic. Boron compounds play a strengthening role in the cell walls of all plants.
There is no consensus on whether boron is an essential nutrient for mammals, including humans, although there is some evidence it supports bone health. The word boron was coined from borax, the mineral from which it was isolated, by analogy with carbon, which boron resembles chemically. Borax, its mineral form known as tincal, glazes were used in China from AD 300, some crude borax reached the West, where the Perso-Arab alchemist Jābir ibn Hayyān mentioned it in AD 700. Marco Polo brought some glazes back to Italy in the 13th century. Agricola, around 1600, reports the use of borax as a flux in metallurgy. In 1777, boric acid was recognized in the hot springs near Florence and became known as sal sedativum, with medical uses; the rare mineral is called sassolite, found at Sasso, Italy. Sasso was the main source of European borax from 1827 to 1872. Boron compounds were rarely used until the late 1800s when Francis Marion Smith's Pacific Coast Borax Company first popularized and produced them in volume at low cost.
Boron was not recognized as an element until it was isolated by Sir Humphry Davy and by Joseph Louis Gay-Lussac and Louis Jacques Thénard. In 1808 Davy observed that electric current sent through a solution of borates produced a brown precipitate on one of the electrodes. In his subsequent experiments, he used potassium to reduce boric acid instead of electrolysis, he named the element boracium. Gay-Lussac and Thénard used iron to reduce boric acid at high temperatures. By oxidizing boron with air, they showed. Jöns Jakob Berzelius identified boron as an element in 1824. Pure boron was arguably first produced by the American chemist Ezekiel Weintraub in 1909; the earliest routes to elemental boron involved the reduction of boric oxide with metals such as magnesium or aluminium. However, the product is always contaminated with borides of those metals. Pure boron can be prepared by reducing volatile boron halides with hydrogen at high temperatures. Ultrapure boron for use in the semiconductor industry is produced by the decomposition of diborane at high temperatures and further purified by the zone melting or Czochralski processes.
The production of boron compounds does not involve the formation of elemental boron, but exploits the convenient availability of borates. Boron is similar to carbon in its capability to form stable covalently bonded molecular networks. Nominally disordered boron contains regular boron icosahedra which are, bonded randomly to each other without long-range order. Crystalline boron is a hard, black material with a melting point of above 2000 °C, it forms four major polymorphs: β-rhombohedral, γ and β-tetragonal. Most of the phases are based on B12 icosahedra, but the γ-phase can be described as a rocksalt-type arrangement of the icosahedra and B2 atomic pairs, it can be produced by compressing other boron phases to 12–20 GPa and heating to 1500–1800 °C. The T phase is produced at similar pressures, but higher temperatures of 1800–2200 °C; as to the α and β phases, they might both coexist at ambient conditions with the β phase being more stable
Aluminium or aluminum is a chemical element with symbol Al and atomic number 13. It is a silvery-white, soft and ductile metal in the boron group. By mass, aluminium makes up about 8% of the Earth's crust; the chief ore of aluminium is bauxite. Aluminium metal is so chemically reactive that native specimens are rare and limited to extreme reducing environments. Instead, it is found combined in over 270 different minerals. Aluminium is remarkable for its low density and its ability to resist corrosion through the phenomenon of passivation. Aluminium and its alloys are vital to the aerospace industry and important in transportation and building industries, such as building facades and window frames; the oxides and sulfates are the most useful compounds of aluminium. Despite its prevalence in the environment, no known form of life uses aluminium salts metabolically, but aluminium is well tolerated by plants and animals; because of these salts' abundance, the potential for a biological role for them is of continuing interest, studies continue.
Of aluminium isotopes, only 27Al is stable. This is consistent with aluminium having an odd atomic number, it is the only aluminium isotope that has existed on Earth in its current form since the creation of the planet. Nearly all the element on Earth is present as this isotope, which makes aluminium a mononuclidic element and means that its standard atomic weight equates to that of the isotope; the standard atomic weight of aluminium is low in comparison with many other metals, which has consequences for the element's properties. All other isotopes of aluminium are radioactive; the most stable of these is 26Al and therefore could not have survived since the formation of the planet. However, 26Al is produced from argon in the atmosphere by spallation caused by cosmic ray protons; the ratio of 26Al to 10Be has been used for radiodating of geological processes over 105 to 106 year time scales, in particular transport, sediment storage, burial times, erosion. Most meteorite scientists believe that the energy released by the decay of 26Al was responsible for the melting and differentiation of some asteroids after their formation 4.55 billion years ago.
The remaining isotopes of aluminium, with mass numbers ranging from 21 to 43, all have half-lives well under an hour. Three metastable states are known, all with half-lives under a minute. An aluminium atom has 13 electrons, arranged in an electron configuration of 3s23p1, with three electrons beyond a stable noble gas configuration. Accordingly, the combined first three ionization energies of aluminium are far lower than the fourth ionization energy alone. Aluminium can easily surrender its three outermost electrons in many chemical reactions; the electronegativity of aluminium is 1.61. A free aluminium atom has a radius of 143 pm. With the three outermost electrons removed, the radius shrinks to 39 pm for a 4-coordinated atom or 53.5 pm for a 6-coordinated atom. At standard temperature and pressure, aluminium atoms form a face-centered cubic crystal system bound by metallic bonding provided by atoms' outermost electrons; this crystal system is shared by some other metals, such as copper. Aluminium metal, when in quantity, is shiny and resembles silver because it preferentially absorbs far ultraviolet radiation while reflecting all visible light so it does not impart any color to reflected light, unlike the reflectance spectra of copper and gold.
Another important characteristic of aluminium is its low density, 2.70 g/cm3. Aluminium is a soft, lightweight and malleable with appearance ranging from silvery to dull gray, depending on the surface roughness, it is nonmagnetic and does not ignite. A fresh film of aluminium serves as a good reflector of visible light and an excellent reflector of medium and far infrared radiation; the yield strength of pure aluminium is 7–11 MPa, while aluminium alloys have yield strengths ranging from 200 MPa to 600 MPa. Aluminium has stiffness of steel, it is machined, cast and extruded. Aluminium atoms are arranged in a face-centered cubic structure. Aluminium has a stacking-fault energy of 200 mJ/m2. Aluminium is a good thermal and electrical conductor, having 59% the conductivity of copper, both thermal and electrical, while having only 30% of copper's density. Aluminium is capable of superconductivity, with a superconducting critical temperature of 1.2 kelvin and a critical magnetic field of about 100 gauss.
Aluminium is the most common material for the fabrication of superconducting qubits. Aluminium's corrosion resistance can be excellent due to a thin surface layer of aluminium oxide that forms when the bare metal is exposed to air preventing further oxidation, in a process termed passivation; the strongest aluminium alloys are less corrosion resistant due to galvanic reactions with alloyed copper. This corrosion resistance is reduced by aqueous salts in the presence of dissimilar metals. In acidic solutions, aluminium reacts with water to form hydrogen, in alkaline ones to form aluminates—protective passivation under these conditions is negligible; because it is corroded by dissolved chlorides, such as common sodium chloride, household plumbing is never made from aluminium. However, because
Nickel is a chemical element with symbol Ni and atomic number 28. It is a silvery-white lustrous metal with a slight golden tinge. Nickel is hard and ductile. Pure nickel, powdered to maximize the reactive surface area, shows a significant chemical activity, but larger pieces are slow to react with air under standard conditions because an oxide layer forms on the surface and prevents further corrosion. So, pure native nickel is found in Earth's crust only in tiny amounts in ultramafic rocks, in the interiors of larger nickel–iron meteorites that were not exposed to oxygen when outside Earth's atmosphere. Meteoric nickel is found in combination with iron, a reflection of the origin of those elements as major end products of supernova nucleosynthesis. An iron–nickel mixture is thought to compose Earth's outer and inner cores. Use of nickel has been traced as far back as 3500 BCE. Nickel was first isolated and classified as a chemical element in 1751 by Axel Fredrik Cronstedt, who mistook the ore for a copper mineral, in the cobalt mines of Los, Hälsingland, Sweden.
The element's name comes from a mischievous sprite of German miner mythology, who personified the fact that copper-nickel ores resisted refinement into copper. An economically important source of nickel is the iron ore limonite, which contains 1–2% nickel. Nickel's other important ore minerals include pentlandite and a mixture of Ni-rich natural silicates known as garnierite. Major production sites include the Sudbury region in Canada, New Caledonia in the Pacific, Norilsk in Russia. Nickel is oxidized by air at room temperature and is considered corrosion-resistant, it has been used for plating iron and brass, coating chemistry equipment, manufacturing certain alloys that retain a high silvery polish, such as German silver. About 9% of world nickel production is still used for corrosion-resistant nickel plating. Nickel-plated objects sometimes provoke nickel allergy. Nickel has been used in coins, though its rising price has led to some replacement with cheaper metals in recent years. Nickel is one of four elements that are ferromagnetic at room temperature.
Alnico permanent magnets based on nickel are of intermediate strength between iron-based permanent magnets and rare-earth magnets. The metal is valuable in modern times chiefly in alloys. A further 10% is used for nickel-based and copper-based alloys, 7% for alloy steels, 3% in foundries, 9% in plating and 4% in other applications, including the fast-growing battery sector; as a compound, nickel has a number of niche chemical manufacturing uses, such as a catalyst for hydrogenation, cathodes for batteries and metal surface treatments. Nickel is an essential nutrient for some microorganisms and plants that have enzymes with nickel as an active site. Nickel is a silvery-white metal with a slight golden tinge, it is one of only four elements that are magnetic at or near room temperature, the others being iron and gadolinium. Its Curie temperature is 355 °C; the unit cell of nickel is a face-centered cube with the lattice parameter of 0.352 nm, giving an atomic radius of 0.124 nm. This crystal structure is stable to pressures of at least 70 GPa.
Nickel belongs to the transition metals. It is hard and ductile, has a high for transition metals electrical and thermal conductivity; the high compressive strength of 34 GPa, predicted for ideal crystals, is never obtained in the real bulk material due to the formation and movement of dislocations. The nickel atom has two electron configurations, 3d8 4s2 and 3d9 4s1, which are close in energy – the symbol refers to the argon-like core structure. There is some disagreement. Chemistry textbooks quote the electron configuration of nickel as 4s2 3d8, which can be written 3d8 4s2; this configuration agrees with the Madelung energy ordering rule, which predicts that 4s is filled before 3d. It is supported by the experimental fact that the lowest energy state of the nickel atom is a 3d8 4s2 energy level the 3d8 4s2 3F, J = 4 level. However, each of these two configurations splits into several energy levels due to fine structure, the two sets of energy levels overlap; the average energy of states with configuration 3d9 4s1 is lower than the average energy of states with configuration 3d8 4s2.
For this reason, the research literature on atomic calculations quotes the ground state configuration of nickel as 3d9 4s1. The isotopes of nickel range in atomic weight from 48 u to 78 u. Occurring nickel is composed of five stable isotopes. Isotopes heavier than 62Ni cannot be formed by nuclear fusion without losing energy. Nickel-62 has the highest mean nuclear binding energy per nucleon of any nuclide, at 8.7946 MeV/nucleon. Its binding energy is greater than both 56Fe and 58Fe, more abundant elements incorrectly cited as having the most tightly-bound nuclides. Although this would seem to predict nickel-62 as the most abundant heavy element in the universe, the high rate of photodisintegration of nickel in stellar interiors causes iron to be by far the most abundant. Stable isotope nickel-60 is the daughter product of the extinct radionuclide 60Fe, whi
A chemical element is a species of atom having the same number of protons in their atomic nuclei. For example, the atomic number of oxygen is 8, so the element oxygen consists of all atoms which have 8 protons. 118 elements have been identified, of which the first 94 occur on Earth with the remaining 24 being synthetic elements. There are 80 elements that have at least one stable isotope and 38 that have radionuclides, which decay over time into other elements. Iron is the most abundant element making up Earth, while oxygen is the most common element in the Earth's crust. Chemical elements constitute all of the ordinary matter of the universe; however astronomical observations suggest that ordinary observable matter makes up only about 15% of the matter in the universe: the remainder is dark matter. The two lightest elements and helium, were formed in the Big Bang and are the most common elements in the universe; the next three elements were formed by cosmic ray spallation, are thus rarer than heavier elements.
Formation of elements with from 6 to 26 protons occurred and continues to occur in main sequence stars via stellar nucleosynthesis. The high abundance of oxygen and iron on Earth reflects their common production in such stars. Elements with greater than 26 protons are formed by supernova nucleosynthesis in supernovae, when they explode, blast these elements as supernova remnants far into space, where they may become incorporated into planets when they are formed; the term "element" is used for atoms with a given number of protons as well as for a pure chemical substance consisting of a single element. For the second meaning, the terms "elementary substance" and "simple substance" have been suggested, but they have not gained much acceptance in English chemical literature, whereas in some other languages their equivalent is used. A single element can form multiple substances differing in their structure; when different elements are chemically combined, with the atoms held together by chemical bonds, they form chemical compounds.
Only a minority of elements are found uncombined as pure minerals. Among the more common of such native elements are copper, gold and sulfur. All but a few of the most inert elements, such as noble gases and noble metals, are found on Earth in chemically combined form, as chemical compounds. While about 32 of the chemical elements occur on Earth in native uncombined forms, most of these occur as mixtures. For example, atmospheric air is a mixture of nitrogen and argon, native solid elements occur in alloys, such as that of iron and nickel; the history of the discovery and use of the elements began with primitive human societies that found native elements like carbon, sulfur and gold. Civilizations extracted elemental copper, tin and iron from their ores by smelting, using charcoal. Alchemists and chemists subsequently identified many more; the properties of the chemical elements are summarized in the periodic table, which organizes the elements by increasing atomic number into rows in which the columns share recurring physical and chemical properties.
Save for unstable radioactive elements with short half-lives, all of the elements are available industrially, most of them in low degrees of impurities. The lightest chemical elements are hydrogen and helium, both created by Big Bang nucleosynthesis during the first 20 minutes of the universe in a ratio of around 3:1 by mass, along with tiny traces of the next two elements and beryllium. All other elements found in nature were made by various natural methods of nucleosynthesis. On Earth, small amounts of new atoms are produced in nucleogenic reactions, or in cosmogenic processes, such as cosmic ray spallation. New atoms are naturally produced on Earth as radiogenic daughter isotopes of ongoing radioactive decay processes such as alpha decay, beta decay, spontaneous fission, cluster decay, other rarer modes of decay. Of the 94 occurring elements, those with atomic numbers 1 through 82 each have at least one stable isotope. Isotopes considered stable are those. Elements with atomic numbers 83 through 94 are unstable to the point that radioactive decay of all isotopes can be detected.
Some of these elements, notably bismuth and uranium, have one or more isotopes with half-lives long enough to survive as remnants of the explosive stellar nucleosynthesis that produced the heavy metals before the formation of our Solar System. At over 1.9×1019 years, over a billion times longer than the current estimated age of the universe, bismuth-209 has the longest known alpha decay half-life of any occurring element, is always considered on par with the 80 stable elements. The heaviest elements undergo radioactive decay with half-lives so short that they are not found in nature and must be synthesized; as of 2010, there are 118 known elements (in this context, "known" means observed well enough from just a few de
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