Titanium is a chemical element with symbol Ti and atomic number 22. It is a lustrous transition metal with a silver color, low density, high strength. Titanium is resistant to corrosion in sea water, aqua regia, chlorine. Titanium was discovered in Cornwall, Great Britain, by William Gregor in 1791, was named by Martin Heinrich Klaproth after the Titans of Greek mythology; the element occurs within a number of mineral deposits, principally rutile and ilmenite, which are distributed in the Earth's crust and lithosphere, it is found in all living things, water bodies and soils. The metal is extracted from its principal mineral ores by the Hunter processes; the most common compound, titanium dioxide, is a popular photocatalyst and is used in the manufacture of white pigments. Other compounds include a component of smoke screens and catalysts. Titanium can be alloyed with iron, aluminium and molybdenum, among other elements, to produce strong, lightweight alloys for aerospace, industrial processes, agri-food, medical prostheses, orthopedic implants and endodontic instruments and files, dental implants, sporting goods, mobile phones, other applications.
The two most useful properties of the metal are corrosion resistance and strength-to-density ratio, the highest of any metallic element. In its unalloyed condition, titanium is less dense. There are two allotropic forms and five occurring isotopes of this element, 46Ti through 50Ti, with 48Ti being the most abundant. Although they have the same number of valence electrons and are in the same group in the periodic table and zirconium differ in many chemical and physical properties; as a metal, titanium is recognized for its high strength-to-weight ratio. It is a strong metal with low density, quite ductile and metallic-white in color; the high melting point makes it useful as a refractory metal. It is paramagnetic and has low electrical and thermal conductivity. Commercially pure grades of titanium have ultimate tensile strength of about 434 MPa, equal to that of common, low-grade steel alloys, but are less dense. Titanium is 60% denser than aluminium, but more than twice as strong as the most used 6061-T6 aluminium alloy.
Certain titanium alloys achieve tensile strengths of over 1,400 MPa. However, titanium loses strength when heated above 430 °C. Titanium is not as hard as some grades of heat-treated steel. Machining requires precautions, because the material can gall unless sharp tools and proper cooling methods are used. Like steel structures, those made from titanium have a fatigue limit that guarantees longevity in some applications; the metal is a dimorphic allotrope of an hexagonal α form that changes into a body-centered cubic β form at 882 °C. The specific heat of the α form increases as it is heated to this transition temperature but falls and remains constant for the β form regardless of temperature. Like aluminium and magnesium, titanium metal and its alloys oxidize upon exposure to air. Titanium reacts with oxygen at 1,200 °C in air, at 610 °C in pure oxygen, forming titanium dioxide, it is, slow to react with water and air at ambient temperatures because it forms a passive oxide coating that protects the bulk metal from further oxidation.
When it first forms, this protective layer continues to grow slowly. Atmospheric passivation gives titanium excellent resistance to corrosion equivalent to platinum. Titanium is capable of withstanding attack by dilute sulfuric and hydrochloric acids, chloride solutions, most organic acids. However, titanium is corroded by concentrated acids; as indicated by its negative redox potential, titanium is thermodynamically a reactive metal that burns in normal atmosphere at lower temperatures than the melting point. Melting is possible only in a vacuum. At 550 °C, it combines with chlorine, it reacts with the other halogens and absorbs hydrogen. Titanium is one of the few elements that burns in pure nitrogen gas, reacting at 800 °C to form titanium nitride, which causes embrittlement; because of its high reactivity with oxygen and some other gases, titanium filaments are applied in titanium sublimation pumps as scavengers for these gases. Such pumps inexpensively and reliably produce low pressures in ultra-high vacuum systems.
Titanium is the ninth-most abundant element in the seventh-most abundant metal. It is present as oxides in most igneous rocks, in sediments derived from them, in living things, natural bodies of water. Of the 801 types of igneous rocks analyzed by the United States Geological Survey, 784 contained titanium, its proportion in soils is 0.5 to 1.5%. Common titanium-containing minerals are anatase, ilmenite, perovskite and titanite. Akaogiite is an rare mineral consisting of titanium dioxide. Of these minerals, only rutile and ilmenite have economic importance, yet they are difficult to find in high concentrations. About 6.0 and 0.7 million tonnes of those minerals were mined in 2011, respectively. Signi
Arsenic is a chemical element with symbol As and atomic number 33. Arsenic occurs in many minerals in combination with sulfur and metals, but as a pure elemental crystal. Arsenic is a metalloid, it has various allotropes, but only the gray form, which has a metallic appearance, is important to industry. The primary use of arsenic is in alloys of lead. Arsenic is a common n-type dopant in semiconductor electronic devices, the optoelectronic compound gallium arsenide is the second most used semiconductor after doped silicon. Arsenic and its compounds the trioxide, are used in the production of pesticides, treated wood products and insecticides; these applications are declining due to the toxicity of its compounds. A few species of bacteria are able to use arsenic compounds as respiratory metabolites. Trace quantities of arsenic are an essential dietary element in rats, goats and other species. A role in human metabolism is not known. However, arsenic poisoning occurs in multicellular life. Arsenic contamination of groundwater is a problem.
The United States' Environmental Protection Agency states that all forms of arsenic are a serious risk to human health. The United States' Agency for Toxic Substances and Disease Registry ranked arsenic as number 1 in its 2001 Priority List of Hazardous Substances at Superfund sites. Arsenic is classified as a Group-A carcinogen; the three most common arsenic allotropes are gray and black arsenic, with gray being the most common. Gray arsenic adopts a double-layered structure consisting of many interlocked, six-membered rings; because of weak bonding between the layers, gray arsenic is brittle and has a low Mohs hardness of 3.5. Nearest and next-nearest neighbors form a distorted octahedral complex, with the three atoms in the same double-layer being closer than the three atoms in the next; this close packing leads to a high density of 5.73 g/cm3. Gray arsenic becomes a semiconductor with a bandgap of 1.2 -- 1.4 eV if amorphized. Gray arsenic is the most stable form. Yellow arsenic is soft and waxy, somewhat similar to tetraphosphorus.
Both have four atoms arranged in a tetrahedral structure in which each atom is bound to each of the other three atoms by a single bond. This unstable allotrope, being molecular, is the most volatile, least dense, most toxic. Solid yellow arsenic is produced by rapid cooling of arsenic vapor, As4, it is transformed into gray arsenic by light. The yellow form has a density of 1.97 g/cm3. Black arsenic is similar in structure to black phosphorus. Black arsenic can be formed by cooling vapor at around 100–220 °C, it is brittle. It is a poor electrical conductor. Arsenic occurs in nature as a monoisotopic element, composed of 75As; as of 2003, at least 33 radioisotopes have been synthesized, ranging in atomic mass from 60 to 92. The most stable of these is 73As with a half-life of 80.30 days. All other isotopes have half-lives of under one day, with the exception of 71As, 72As, 74As, 76As, 77As. Isotopes that are lighter than the stable 75As tend to decay by β+ decay, those that are heavier tend to decay by β− decay, with some exceptions.
At least 10 nuclear isomers have been described, ranging in atomic mass from 66 to 84. The most stable of arsenic's isomers is 68mAs with a half-life of 111 seconds. Arsenic has a similar electronegativity and ionization energies to its lighter congener phosphorus and as such forms covalent molecules with most of the nonmetals. Though stable in dry air, arsenic forms a golden-bronze tarnish upon exposure to humidity which becomes a black surface layer; when heated in air, arsenic oxidizes to arsenic trioxide. This odor can be detected on striking arsenide minerals such as arsenopyrite with a hammer, it burns in oxygen to form arsenic trioxide and arsenic pentoxide, which have the same structure as the more well-known phosphorus compounds, in fluorine to give arsenic pentafluoride. Arsenic sublimes upon heating at atmospheric pressure, converting directly to a gaseous form without an intervening liquid state at 887 K; the triple point is 3.63 MPa and 1,090 K. Arsenic makes arsenic acid with concentrated nitric acid, arsenous acid with dilute nitric acid, arsenic trioxide with concentrated sulfuric acid.
Arsenic reacts with metals to form arsenides, though these are not ionic compounds containing the As3− ion as the formation of such an anion would be endothermic and the group 1 arsenides have properties of intermetallic compounds. Like germanium and bromine, which like arsenic succeed the 3d transition series, arsenic is much less stable in the group oxidation state of +5 than its vertical neighbors phosphorus and antimony, hence arsenic pentoxide and arsenic acid are potent oxidizers. Compounds of arsenic resemble in some respects those of phosphorus which occupies the same group of the periodic table; the most common oxidation states for arsenic are: −3 in the arsenides, which are alloy-like intermetallic compounds, +3 in the arsenites, +5 in the arsenates and most organoarsenic compounds. Arsenic bonds to itself as seen in the square As3−4 ions in the mineral skutterudite. In the +3 oxidation state, arsenic is pyramidal owing to the i
Carbon is a chemical element with symbol C and atomic number 6. It is nonmetallic and tetravalent—making four electrons available to form covalent chemical bonds, it belongs to group 14 of the periodic table. Three isotopes occur 12C and 13C being stable, while 14C is a radionuclide, decaying with a half-life of about 5,730 years. Carbon is one of the few elements known since antiquity. Carbon is the 15th most abundant element in the Earth's crust, the fourth most abundant element in the universe by mass after hydrogen and oxygen. Carbon's abundance, its unique diversity of organic compounds, its unusual ability to form polymers at the temperatures encountered on Earth enables this element to serve as a common element of all known life, it is the second most abundant element in the human body by mass after oxygen. The atoms of carbon can bond together in different ways, termed allotropes of carbon; the best known are graphite and amorphous carbon. The physical properties of carbon vary with the allotropic form.
For example, graphite is opaque and black while diamond is transparent. Graphite is soft enough to form a streak on paper, while diamond is the hardest occurring material known. Graphite is a good electrical conductor. Under normal conditions, carbon nanotubes, graphene have the highest thermal conductivities of all known materials. All carbon allotropes are solids under normal conditions, with graphite being the most thermodynamically stable form at standard temperature and pressure, they are chemically resistant and require high temperature to react with oxygen. The most common oxidation state of carbon in inorganic compounds is +4, while +2 is found in carbon monoxide and transition metal carbonyl complexes; the largest sources of inorganic carbon are limestones and carbon dioxide, but significant quantities occur in organic deposits of coal, peat and methane clathrates. Carbon forms a vast number of compounds, more than any other element, with ten million compounds described to date, yet that number is but a fraction of the number of theoretically possible compounds under standard conditions.
For this reason, carbon has been referred to as the "king of the elements". The allotropes of carbon include graphite, one of the softest known substances, diamond, the hardest occurring substance, it bonds with other small atoms, including other carbon atoms, is capable of forming multiple stable covalent bonds with suitable multivalent atoms. Carbon is known to form ten million different compounds, a large majority of all chemical compounds. Carbon has the highest sublimation point of all elements. At atmospheric pressure it has no melting point, as its triple point is at 10.8±0.2 MPa and 4,600 ± 300 K, so it sublimes at about 3,900 K. Graphite is much more reactive than diamond at standard conditions, despite being more thermodynamically stable, as its delocalised pi system is much more vulnerable to attack. For example, graphite can be oxidised by hot concentrated nitric acid at standard conditions to mellitic acid, C66, which preserves the hexagonal units of graphite while breaking up the larger structure.
Carbon sublimes in a carbon arc, which has a temperature of about 5800 K. Thus, irrespective of its allotropic form, carbon remains solid at higher temperatures than the highest-melting-point metals such as tungsten or rhenium. Although thermodynamically prone to oxidation, carbon resists oxidation more than elements such as iron and copper, which are weaker reducing agents at room temperature. Carbon is the sixth element, with a ground-state electron configuration of 1s22s22p2, of which the four outer electrons are valence electrons, its first four ionisation energies, 1086.5, 2352.6, 4620.5 and 6222.7 kJ/mol, are much higher than those of the heavier group-14 elements. The electronegativity of carbon is 2.5 higher than the heavier group-14 elements, but close to most of the nearby nonmetals, as well as some of the second- and third-row transition metals. Carbon's covalent radii are taken as 77.2 pm, 66.7 pm and 60.3 pm, although these may vary depending on coordination number and what the carbon is bonded to.
In general, covalent radius decreases with higher bond order. Carbon compounds form the basis of all known life on Earth, the carbon–nitrogen cycle provides some of the energy produced by the Sun and other stars. Although it forms an extraordinary variety of compounds, most forms of carbon are comparatively unreactive under normal conditions. At standard temperature and pressure, it resists all but the strongest oxidizers, it does not react with hydrochloric acid, chlorine or any alkalis. At elevated temperatures, carbon reacts with oxygen to form carbon oxides and will rob oxygen from metal oxides to leave the elemental metal; this exothermic reaction is used in the iron and steel industry to smelt iron and to control the carbon content of steel: Fe3O4 + 4 C → 3 Fe + 4 COCarbon monoxide can be recycled to smelt more iron: Fe3O4 + 4 CO → 3 Fe + 4 CO2with sulfur to form carbon disulfide and with steam in the coal-gas reaction: C + H2O → CO + H2. Carbon combines with some metals at high temperatures to form metallic carbides, such as the iron carbide cementite in steel and tungsten carbide used as an abrasive and for making hard tips for cutting tools.
The system of carbon allotropes spans a range of extremes: Atomic carbon is a ver
Fluorine is a chemical element with symbol F and atomic number 9. It is the lightest halogen and exists as a toxic pale yellow diatomic gas at standard conditions; as the most electronegative element, it is reactive, as it reacts with all other elements, except for helium and neon. Among the elements, fluorine ranks 24th in universal 13th in terrestrial abundance. Fluorite, the primary mineral source of fluorine which gave the element its name, was first described in 1529. Proposed as an element in 1810, fluorine proved difficult and dangerous to separate from its compounds, several early experimenters died or sustained injuries from their attempts. Only in 1886 did French chemist Henri Moissan isolate elemental fluorine using low-temperature electrolysis, a process still employed for modern production. Industrial production of fluorine gas for uranium enrichment, its largest application, began during the Manhattan Project in World War II. Owing to the expense of refining pure fluorine, most commercial applications use fluorine compounds, with about half of mined fluorite used in steelmaking.
The rest of the fluorite is converted into corrosive hydrogen fluoride en route to various organic fluorides, or into cryolite, which plays a key role in aluminium refining. Molecules containing a Carbon–fluorine bond have high chemical and thermal stability. Pharmaceuticals such as atorvastatin and fluoxetine contain C-F bonds, the fluoride ion inhibits dental cavities, so finds use in toothpaste and water fluoridation. Global fluorochemical sales amount to more than US$15 billion a year. Fluorocarbon gases are greenhouse gases with global-warming potentials 100 to 20,000 times that of carbon dioxide. Organofluorine compounds persist in the environment due to the strength of the carbon–fluorine bond. Fluorine has no known metabolic role in mammals. Fluorine atoms have nine electrons, one fewer than neon, electron configuration 1s22s22p5: two electrons in a filled inner shell and seven in an outer shell requiring one more to be filled; the outer electrons are ineffective at nuclear shielding, experience a high effective nuclear charge of 9 − 2 = 7.
Fluorine's first ionization energy is third-highest among all elements, behind helium and neon, which complicates the removal of electrons from neutral fluorine atoms. It has a high electron affinity, second only to chlorine, tends to capture an electron to become isoelectronic with the noble gas neon. Fluorine atoms have a small covalent radius of around 60 picometers, similar to those of its period neighbors oxygen and neon; the bond energy of difluorine is much lower than that of either Cl2 or Br2 and similar to the cleaved peroxide bond. Conversely, bonds to other atoms are strong because of fluorine's high electronegativity. Unreactive substances like powdered steel, glass fragments, asbestos fibers react with cold fluorine gas. Reactions of elemental fluorine with metals require varying conditions. Alkali metals cause; some solid nonmetals react vigorously in liquid air temperature fluorine. Hydrogen sulfide and sulfur dioxide combine with fluorine, the latter sometimes explosively. Hydrogen, like some of the alkali metals, reacts explosively with fluorine.
Carbon, as lamp black, reacts at room temperature to yield fluoromethane. Graphite combines with fluorine above 400 °C to produce non-stoichiometric carbon monofluoride. Carbon dioxide and carbon monoxide react at or just above room temperature, whereas paraffins and other organic chemicals generate strong reactions: fully substituted haloalkanes such as carbon tetrachloride incombustible, may explode. Although nitrogen trifluoride is stable, nitrogen requires an electric discharge at elevated temperatures for reaction with fluorine to occur, due to the strong triple bond in elemental nitrogen. Oxygen does not combine with fluorine under ambient conditions, but can be made to react using electric discharge at low temperatures and pressures. Heavier halogens react with fluorine as does the noble gas radon. At room temperature, fluorine is a gas of diatomic molecules, pale yellow, it has a characteristic halogen-like biting odor detectable at 20 ppb. Fluorine condenses into a bright yellow liquid at −188 °C, a transition temperature similar to those of oxygen and nitrogen.
Fluorine has two solid forms, α- and β-fluorine. The latter crystallizes at −220 °C and is transparent and sof
Copper is a chemical element with symbol Cu and atomic number 29. It is a soft and ductile metal with high thermal and electrical conductivity. A freshly exposed surface of pure copper has a pinkish-orange color. Copper is used as a conductor of heat and electricity, as a building material, as a constituent of various metal alloys, such as sterling silver used in jewelry, cupronickel used to make marine hardware and coins, constantan used in strain gauges and thermocouples for temperature measurement. Copper is one of the few metals; this led to early human use in several regions, from c. 8000 BC. Thousands of years it was the first metal to be smelted from sulfide ores, c. 5000 BC, the first metal to be cast into a shape in a mold, c. 4000 BC and the first metal to be purposefully alloyed with another metal, tin, to create bronze, c. 3500 BC. In the Roman era, copper was principally mined on Cyprus, the origin of the name of the metal, from aes сyprium corrupted to сuprum, from which the words derived and copper, first used around 1530.
The encountered compounds are copper salts, which impart blue or green colors to such minerals as azurite and turquoise, have been used and as pigments. Copper used in buildings for roofing, oxidizes to form a green verdigris. Copper is sometimes used in decorative art, both in its elemental metal form and in compounds as pigments. Copper compounds are used as bacteriostatic agents and wood preservatives. Copper is essential to all living organisms as a trace dietary mineral because it is a key constituent of the respiratory enzyme complex cytochrome c oxidase. In molluscs and crustaceans, copper is a constituent of the blood pigment hemocyanin, replaced by the iron-complexed hemoglobin in fish and other vertebrates. In humans, copper is found in the liver and bone; the adult body contains between 2.1 mg of copper per kilogram of body weight. Copper and gold are in group 11 of the periodic table; the filled d-shells in these elements contribute little to interatomic interactions, which are dominated by the s-electrons through metallic bonds.
Unlike metals with incomplete d-shells, metallic bonds in copper are lacking a covalent character and are weak. This observation explains the low high ductility of single crystals of copper. At the macroscopic scale, introduction of extended defects to the crystal lattice, such as grain boundaries, hinders flow of the material under applied stress, thereby increasing its hardness. For this reason, copper is supplied in a fine-grained polycrystalline form, which has greater strength than monocrystalline forms; the softness of copper explains its high electrical conductivity and high thermal conductivity, second highest among pure metals at room temperature. This is because the resistivity to electron transport in metals at room temperature originates from scattering of electrons on thermal vibrations of the lattice, which are weak in a soft metal; the maximum permissible current density of copper in open air is 3.1×106 A/m2 of cross-sectional area, above which it begins to heat excessively. Copper is one of a few metallic elements with a natural color other than silver.
Pure copper acquires a reddish tarnish when exposed to air. The characteristic color of copper results from the electronic transitions between the filled 3d and half-empty 4s atomic shells – the energy difference between these shells corresponds to orange light; as with other metals, if copper is put in contact with another metal, galvanic corrosion will occur. Copper does not react with water, but it does react with atmospheric oxygen to form a layer of brown-black copper oxide which, unlike the rust that forms on iron in moist air, protects the underlying metal from further corrosion. A green layer of verdigris can be seen on old copper structures, such as the roofing of many older buildings and the Statue of Liberty. Copper tarnishes when exposed to some sulfur compounds, with which it reacts to form various copper sulfides. There are 29 isotopes of copper. 63Cu and 65Cu are stable, with 63Cu comprising 69% of occurring copper. The other isotopes are radioactive, with the most stable being 67Cu with a half-life of 61.83 hours.
Seven metastable isotopes have been characterized. Isotopes with a mass number above 64 decay by β−, whereas those with a mass number below 64 decay by β+. 64Cu, which has a half-life of 12.7 hours, decays both ways.62Cu and 64Cu have significant applications. 62Cu is used in 62Cu-PTSM as a radioactive tracer for positron emission tomography. Copper is produced in massive stars and is present in the Earth's crust in a proportion of about 50 parts per million. In nature, copper occurs in a variety of minerals, including native copper, copper sulfides such as chalcopyrite, digenite and chalcocite, copper sulfosalts such as tetrahedite-tennantite, enargite, copper carbonates such as azurite and malachite, as copper or copper oxides such as cuprite and tenorite, respectively; the largest mass of elemental copper discovered weighed 420 tonnes and was found in 1857 on the Keweenaw Peninsula in Michigan, US. Native copper is a polycrystal
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
Oxygen is the chemical element with the symbol O and atomic number 8. It is a member of the chalcogen group on the periodic table, a reactive nonmetal, an oxidizing agent that forms oxides with most elements as well as with other compounds. By mass, oxygen is the third-most abundant element in the universe, after helium. At standard temperature and pressure, two atoms of the element bind to form dioxygen, a colorless and odorless diatomic gas with the formula O2. Diatomic oxygen gas constitutes 20.8% of the Earth's atmosphere. As compounds including oxides, the element makes up half of the Earth's crust. Dioxygen is used in cellular respiration and many major classes of organic molecules in living organisms contain oxygen, such as proteins, nucleic acids and fats, as do the major constituent inorganic compounds of animal shells and bone. Most of the mass of living organisms is oxygen as a component of water, the major constituent of lifeforms. Oxygen is continuously replenished in Earth's atmosphere by photosynthesis, which uses the energy of sunlight to produce oxygen from water and carbon dioxide.
Oxygen is too chemically reactive to remain a free element in air without being continuously replenished by the photosynthetic action of living organisms. Another form of oxygen, ozone absorbs ultraviolet UVB radiation and the high-altitude ozone layer helps protect the biosphere from ultraviolet radiation. However, ozone present at the surface is a byproduct of thus a pollutant. Oxygen was isolated by Michael Sendivogius before 1604, but it is believed that the element was discovered independently by Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, Joseph Priestley in Wiltshire, in 1774. Priority is given for Priestley because his work was published first. Priestley, called oxygen "dephlogisticated air", did not recognize it as a chemical element; the name oxygen was coined in 1777 by Antoine Lavoisier, who first recognized oxygen as a chemical element and characterized the role it plays in combustion. Common uses of oxygen include production of steel and textiles, brazing and cutting of steels and other metals, rocket propellant, oxygen therapy, life support systems in aircraft, submarines and diving.
One of the first known experiments on the relationship between combustion and air was conducted by the 2nd century BCE Greek writer on mechanics, Philo of Byzantium. In his work Pneumatica, Philo observed that inverting a vessel over a burning candle and surrounding the vessel's neck with water resulted in some water rising into the neck. Philo incorrectly surmised that parts of the air in the vessel were converted into the classical element fire and thus were able to escape through pores in the glass. Many centuries Leonardo da Vinci built on Philo's work by observing that a portion of air is consumed during combustion and respiration. In the late 17th century, Robert Boyle proved. English chemist John Mayow refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus. In one experiment, he found that placing either a mouse or a lit candle in a closed container over water caused the water to rise and replace one-fourteenth of the air's volume before extinguishing the subjects.
From this he surmised that nitroaereus is consumed in both combustion. Mayow observed that antimony increased in weight when heated, inferred that the nitroaereus must have combined with it, he thought that the lungs separate nitroaereus from air and pass it into the blood and that animal heat and muscle movement result from the reaction of nitroaereus with certain substances in the body. Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "De respiratione". Robert Hooke, Ole Borch, Mikhail Lomonosov, Pierre Bayen all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as a chemical element; this may have been in part due to the prevalence of the philosophy of combustion and corrosion called the phlogiston theory, the favored explanation of those processes. Established in 1667 by the German alchemist J. J. Becher, modified by the chemist Georg Ernst Stahl by 1731, phlogiston theory stated that all combustible materials were made of two parts.
One part, called phlogiston, was given off when the substance containing it was burned, while the dephlogisticated part was thought to be its true form, or calx. Combustible materials that leave little residue, such as wood or coal, were thought to be made of phlogiston. Air did not play a role in phlogiston theory, nor were any initial quantitative experiments conducted to test the idea. Polish alchemist and physician Michael Sendivogius in his work De Lapide Philosophorum Tractatus duodecim e naturae fonte et manuali experientia depromti described a substance contained in air, referring to it as'cibus vitae', this substance is identical with oxygen. Sendivogius, during his experiments performed between 1598 and 1604, properly recognized that the substance is equivalent to the gaseous byproduct released by the thermal decomposition of potassium nitrate. In Bugaj’s view, the isolation of oxygen and the proper association of the substance to that part of air, required for life, lends sufficient weight to the discovery of oxygen by Sendivogius.