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.
Beryllium is a chemical element with symbol Be and atomic number 4. It is a rare element in the universe occurring as a product of the spallation of larger atomic nuclei that have collided with cosmic rays. Within the cores of stars beryllium is depleted as it creates larger elements, it is a divalent element which occurs only in combination with other elements in minerals. Notable gemstones which contain beryllium include chrysoberyl; as a free element it is a steel-gray, strong and brittle alkaline earth metal. Beryllium improves many physical properties when added as an alloying element to aluminium, copper and nickel. Beryllium does not form oxides until it reaches high temperatures. Tools made of beryllium copper alloys are strong and hard and do not create sparks when they strike a steel surface. In structural applications, the combination of high flexural rigidity, thermal stability, thermal conductivity and low density make beryllium metal a desirable aerospace material for aircraft components, missiles and satellites.
Because of its low density and atomic mass, beryllium is transparent to X-rays and other forms of ionizing radiation. The high thermal conductivities of beryllium and beryllium oxide have led to their use in thermal management applications; the commercial use of beryllium requires the use of appropriate dust control equipment and industrial controls at all times because of the toxicity of inhaled beryllium-containing dusts that can cause a chronic life-threatening allergic disease in some people called berylliosis. Beryllium is a steel gray and hard metal, brittle at room temperature and has a close-packed hexagonal crystal structure, it has a reasonably high melting point. The modulus of elasticity of beryllium is 50% greater than that of steel; the combination of this modulus and a low density results in an unusually fast sound conduction speed in beryllium – about 12.9 km/s at ambient conditions. Other significant properties are high specific heat and thermal conductivity, which make beryllium the metal with the best heat dissipation characteristics per unit weight.
In combination with the low coefficient of linear thermal expansion, these characteristics result in a unique stability under conditions of thermal loading. Occurring beryllium, save for slight contamination by the cosmogenic radioisotopes, is isotopically pure beryllium-9, which has a nuclear spin of 3/2. Beryllium has a large scattering cross section for high-energy neutrons, about 6 barns for energies above 10 keV. Therefore, it works as a neutron reflector and neutron moderator slowing the neutrons to the thermal energy range of below 0.03 eV, where the total cross section is at least an order of magnitude lower – exact value depends on the purity and size of the crystallites in the material. The single primordial beryllium isotope 9Be undergoes a neutron reaction with neutron energies over about 1.9 MeV, to produce 8Be, which immediately breaks into two alpha particles. Thus, for high-energy neutrons, beryllium is a neutron multiplier, releasing more neutrons than it absorbs; this nuclear reaction is: 94Be + n → 2 42He + 2 nNeutrons are liberated when beryllium nuclei are struck by energetic alpha particles producing the nuclear reaction 94Be + 42He → 126C + n, where 42He is an alpha particle and 126C is a carbon-12 nucleus.
Beryllium releases neutrons under bombardment by gamma rays. Thus, natural beryllium bombarded either by alphas or gammas from a suitable radioisotope is a key component of most radioisotope-powered nuclear reaction neutron sources for the laboratory production of free neutrons. Small amounts of tritium are liberated when 94Be nuclei absorb low energy neutrons in the three-step nuclear reaction 94Be + n → 42He + 62He, 62He → 63Li + β−, 63Li + n → 42He + 31HNote that 62He has a half-life of only 0.8 seconds, β− is an electron, 63Li has a high neutron absorption cross-section. Tritium is a radioisotope of concern in nuclear reactor waste streams; as a metal, beryllium is transparent to most wavelengths of X-rays and gamma rays, making it useful for the output windows of X-ray tubes and other such apparatus. Both stable and unstable isotopes of beryllium are created in stars, but the radioisotopes do not last long, it is believed that most of the stable beryllium in the universe was created in the interstellar medium when cosmic rays induced fission in heavier elements found in interstellar gas and dust.
Primordial beryllium contains only one stable isotope, 9Be, therefore beryllium is a monoisotopic element. Radioactive cosmogenic 10Be is produced in the atmosphere of the Earth by the cosmic ray spallation of oxygen. 10Be accumulates at the soil surface, where its long half-life permits a long residence time before decaying to boron-10. Thus, 10Be and its daughter products are used to examine natural soil erosion, soil formation and the development of lateritic soils, as a proxy for measurement of the variations in solar activity and the age of ice cores; the production of 10Be is inversely proportional to solar activity, because increased solar wind during periods of high solar activity decreases the flux of galactic cosmic rays that reach the Earth. Nuclear explosions form 10Be by the reaction of fast neutrons with 13C in the carbon dioxide in air; this is one of the indicators of past activity at nuclear weapon
Potassium is a chemical element with symbol K and atomic number 19. It was first isolated from the ashes of plants, from which its name derives. In the periodic table, potassium is one of the alkali metals. All of the alkali metals have a single valence electron in the outer electron shell, removed to create an ion with a positive charge – a cation, which combines with anions to form salts. Potassium in nature occurs only in ionic salts. Elemental potassium is a soft silvery-white alkali metal that oxidizes in air and reacts vigorously with water, generating sufficient heat to ignite hydrogen emitted in the reaction, burning with a lilac-colored flame, it is found dissolved in sea water, is part of many minerals. Potassium is chemically similar to sodium, the previous element in group 1 of the periodic table, they have a similar first ionization energy, which allows for each atom to give up its sole outer electron. That they are different elements that combine with the same anions to make similar salts was suspected in 1702, was proven in 1807 using electrolysis.
Occurring potassium is composed of three isotopes, of which 40K is radioactive. Traces of 40K are found in all potassium, it is the most common radioisotope in the human body. Potassium ions are vital for the functioning of all living cells; the transfer of potassium ions across nerve cell membranes is necessary for normal nerve transmission. Fresh fruits and vegetables are good dietary sources of potassium; the body responds to the influx of dietary potassium, which raises serum potassium levels, with a shift of potassium from outside to inside cells and an increase in potassium excretion by the kidneys. Most industrial applications of potassium exploit the high solubility in water of potassium compounds, such as potassium soaps. Heavy crop production depletes the soil of potassium, this can be remedied with agricultural fertilizers containing potassium, accounting for 95% of global potassium chemical production; the English name for the element potassium comes from the word "potash", which refers to an early method of extracting various potassium salts: placing in a pot the ash of burnt wood or tree leaves, adding water and evaporating the solution.
When Humphry Davy first isolated the pure element using electrolysis in 1807, he named it potassium, which he derived from the word potash. The symbol "K" stems from kali, itself from the root word alkali, which in turn comes from Arabic: القَلْيَه al-qalyah "plant ashes". In 1797, the German chemist Martin Klaproth discovered "potash" in the minerals leucite and lepidolite, realized that "potash" was not a product of plant growth but contained a new element, which he proposed to call kali. In 1807, Humphry Davy produced the element via electrolysis: in 1809, Ludwig Wilhelm Gilbert proposed the name Kalium for Davy's "potassium". In 1814, the Swedish chemist Berzelius advocated the name kalium for potassium, with the chemical symbol "K"; the English and French speaking countries adopted Davy and Gay-Lussac/Thénard's name Potassium, while the Germanic countries adopted Gilbert/Klaproth's name Kalium. The "Gold Book" of the International Union of Physical and Applied Chemistry has designated the official chemical symbol as K.
Potassium is the second least dense metal after lithium. It is a soft solid with a low melting point, can be cut with a knife. Freshly cut potassium is silvery in appearance, but it begins to tarnish toward gray on exposure to air. In a flame test and its compounds emit a lilac color with a peak emission wavelength of 766.5 nanometers. Neutral potassium atoms have 19 electrons, one more than the stable configuration of the noble gas argon; because of this and its low first ionization energy of 418.8 kJ/mol, the potassium atom is much more to lose the last electron and acquire a positive charge than to gain one and acquire a negative charge. This process requires so little energy that potassium is oxidized by atmospheric oxygen. In contrast, the second ionization energy is high, because removal of two electrons breaks the stable noble gas electronic configuration. Potassium therefore does not form compounds with the oxidation state of higher. Potassium is an active metal that reacts violently with oxygen in water and air.
With oxygen it forms potassium peroxide, with water potassium forms potassium hydroxide. The reaction of potassium with water is dangerous because of its violent exothermic character and the production of hydrogen gas. Hydrogen reacts again with atmospheric oxygen, producing water, which reacts with the remaining potassium; this reaction requires only traces of water. Because of the sensitivity of potassium to water and air, reactions with other elements are possible only in an inert atmosphere such as argon gas using air-free techniques. Potassium does not react with most hydrocarbons such as mineral kerosene, it dissolves in liquid ammonia, up to 480 g per 1000 g of ammonia at 0 °C. Depending on the concentration, the ammonia solutions are blue to yellow, their electrical conductivity is similar to that of liquid metals. In a pure solution, potassium reacts with ammonia to form KNH2, but this reaction is accelerated by minute amounts of transition metal s
Xenon is a chemical element with symbol Xe and atomic number 54. It is a colorless, odorless noble gas found in the Earth's atmosphere in trace amounts. Although unreactive, xenon can undergo a few chemical reactions such as the formation of xenon hexafluoroplatinate, the first noble gas compound to be synthesized. Xenon is used in flash lamps and arc lamps, as a general anesthetic; the first excimer laser design used a xenon dimer molecule as the lasing medium, the earliest laser designs used xenon flash lamps as pumps. Xenon is used to search for hypothetical weakly interacting massive particles and as the propellant for ion thrusters in spacecraft. Occurring xenon consists of eight stable isotopes. More than 40 unstable xenon isotopes undergo radioactive decay, the isotope ratios of xenon are an important tool for studying the early history of the Solar System. Radioactive xenon-135 is produced by beta decay from iodine-135, is the most significant neutron absorber in nuclear reactors. Xenon was discovered in England by the Scottish chemist William Ramsay and English chemist Morris Travers in September 1898, shortly after their discovery of the elements krypton and neon.
They found xenon in the residue left over from evaporating components of liquid air. Ramsay suggested the name xenon for this gas from the Greek word ξένον, neuter singular form of ξένος, meaning'foreign','strange', or'guest'. In 1902, Ramsay estimated the proportion of xenon in the Earth's atmosphere to be one part in 20 million. During the 1930s, American engineer Harold Edgerton began exploring strobe light technology for high speed photography; this led him to the invention of the xenon flash lamp in which light is generated by passing brief electric current through a tube filled with xenon gas. In 1934, Edgerton was able to generate flashes as brief as one microsecond with this method. In 1939, American physician Albert R. Behnke Jr. began exploring the causes of "drunkenness" in deep-sea divers. He tested the effects of varying the breathing mixtures on his subjects, discovered that this caused the divers to perceive a change in depth. From his results, he deduced. Although Russian toxicologist Nikolay V. Lazarev studied xenon anesthesia in 1941, the first published report confirming xenon anesthesia was in 1946 by American medical researcher John H. Lawrence, who experimented on mice.
Xenon was first used as a surgical anesthetic in 1951 by American anesthesiologist Stuart C. Cullen, who used it with two patients. Xenon and the other noble gases were for a long time considered to be chemically inert and not able to form compounds. However, while teaching at the University of British Columbia, Neil Bartlett discovered that the gas platinum hexafluoride was a powerful oxidizing agent that could oxidize oxygen gas to form dioxygenyl hexafluoroplatinate. Since O2 and xenon have the same first ionization potential, Bartlett realized that platinum hexafluoride might be able to oxidize xenon. On March 23, 1962, he mixed the two gases and produced the first known compound of a noble gas, xenon hexafluoroplatinate. Bartlett thought its composition to be Xe+−, but work revealed that it was a mixture of various xenon-containing salts. Since many other xenon compounds have been discovered, in addition to some compounds of the noble gases argon and radon, including argon fluorohydride, krypton difluoride, radon fluoride.
By 1971, more than 80 xenon compounds were known. In November 1989, IBM scientists demonstrated a technology capable of manipulating individual atoms; the program, called IBM in atoms, used a scanning tunneling microscope to arrange 35 individual xenon atoms on a substrate of chilled crystal of nickel to spell out the three letter company initialism. It was the first time atoms had been positioned on a flat surface. Xenon has atomic number 54. At standard temperature and pressure, pure xenon gas has a density of 5.761 kg/m3, about 4.5 times the density of the Earth's atmosphere at sea level, 1.217 kg/m3. As a liquid, xenon has a density of up to 3.100 g/mL, with the density maximum occurring at the triple point. Liquid xenon has a high polarizability due to its large atomic volume, thus is an excellent solvent, it can dissolve hydrocarbons, biological molecules, water. Under the same conditions, the density of solid xenon, 3.640 g/cm3, is greater than the average density of granite, 2.75 g/cm3.
Under gigapascals of pressure, xenon forms a metallic phase. Solid xenon changes from face-centered cubic to hexagonal close packed crystal phase under pressure and begins to turn metallic at about 140 GPa, with no noticeable volume change in the hcp phase, it is metallic at 155 GPa. When metallized, xenon appears sky blue because it absorbs red light and transmits other visible frequencies; such behavior is unusual for a metal and is explained by the small width of the electron bands in that state. Liquid or solid xenon nanoparticles can be formed at room temperature by implanting Xe+ ions into a solid matrix. Many solids have lattice constants smaller than solid Xe; this results in compression of the implanted Xe to pressures that may be sufficient for its liquefaction or solidification. Xenon is a member of the zero-valence elements that are called inert gases, it is inert to most common chemical reactions because the outer valence shell contains eight electrons. This produces a stable, minimum energy configuration in which the outer electrons are bound.
In a gas-filled tube, xenon em
Rhodium is a chemical element with symbol Rh and atomic number 45. It is a rare, silvery-white, corrosion-resistant, chemically inert transition metal, it is a member of the platinum group. It has only one occurring isotope, 103Rh. Occurring rhodium is found as the free metal, alloyed with similar metals, as a chemical compound in minerals such as bowieite and rhodplumsite, it is one of most valuable precious metals. Rhodium is found in platinum or nickel ores together with the other members of the platinum group metals, it was discovered in 1803 by William Hyde Wollaston in one such ore, named for the rose color of one of its chlorine compounds, produced after it reacted with the powerful acid mixture aqua regia. The element's major use is as one of the catalysts in the three-way catalytic converters in automobiles; because rhodium metal is inert against corrosion and most aggressive chemicals, because of its rarity, rhodium is alloyed with platinum or palladium and applied in high-temperature and corrosion-resistive coatings.
White gold is plated with a thin rhodium layer to improve its appearance while sterling silver is rhodium-plated for tarnish resistance. Rhodium detectors are used in nuclear reactors to measure the neutron flux level. Rhodium was discovered in 1803 by William Hyde Wollaston, soon after his discovery of palladium, he used crude platinum ore obtained from South America. His procedure involved dissolving the ore in aqua regia and neutralizing the acid with sodium hydroxide, he precipitated the platinum as ammonium chloroplatinate by adding ammonium chloride. Most other metals like copper, lead and rhodium were precipitated with zinc. Diluted nitric acid dissolved all but rhodium. Of these, palladium dissolved in aqua regia but rhodium did not, the rhodium was precipitated by the addition of sodium chloride as Na3·nH2O. After being washed with ethanol, the rose-red precipitate was reacted with zinc, which displaced the rhodium in the ionic compound and thereby released the rhodium as free metal. After the discovery, the rare element had only minor applications.
The first major application was electroplating for decorative uses and as corrosion-resistant coating. The introduction of the three-way catalytic converter by Volvo in 1976 increased the demand for rhodium; the previous catalytic converters used platinum or palladium, while the three-way catalytic converter used rhodium to reduce the amount of NOx in the exhaust. Rhodium is a hard, durable metal that has a high reflectance. Rhodium metal does not form an oxide when heated. Oxygen is absorbed from the atmosphere only at the melting point of rhodium, but is released on solidification. Rhodium has both lower density than platinum, it is not attacked by most acids: it is insoluble in nitric acid and dissolves in aqua regia. Rhodium belongs to group 9 of the periodic table, but the configuration of electrons in the outermost shells is atypical for the group; this anomaly is observed in the neighboring elements, niobium and palladium. The common oxidation state of rhodium is +3, but oxidation states from 0 to +6 are observed.
Unlike ruthenium and osmium, rhodium forms no volatile oxygen compounds. The known stable oxides include Rh2O3, RhO2, RhO2·xH2O, Na2RhO3, Sr3LiRhO6 and Sr3NaRhO6. Halogen compounds are known in nearly the full range of possible oxidation states. Rhodium chloride, rhodium fluoride, rhodium fluoride and rhodium fluoride are examples; the lower oxidation states are stable only in the presence of ligands. The best-known rhodium-halogen compound is the Wilkinson's catalyst chlorotrisrhodium; this catalyst is used in the hydrogenation of alkenes. Occurring rhodium is composed of only one isotope, 103Rh; the most stable radioisotopes are 101Rh with a half-life of 3.3 years, 102Rh with a half-life of 207 days, 102mRh with a half-life of 2.9 years, 99Rh with a half-life of 16.1 days. Twenty other radioisotopes have been characterized with atomic weights ranging from 92.926 u to 116.925 u. Most of these have half-lives shorter except 100Rh and 105Rh. Rhodium has numerous meta states, the most stable being 102mRh with a half-life of about 2.9 years and 101mRh with a half-life of 4.34 days.
In isotopes weighing less than 103, the primary decay mode is electron capture and the primary decay product is ruthenium. In isotopes greater than 103, the primary decay mode is beta emission and the primary product is palladium. Rhodium is one of the rarest elements in the Earth's crust, comprising an estimated 0.0002 parts per million. Its rarity affects its use in commercial applications; the industrial extraction of rhodium is complex because the ores are mixed with other metals such as palladium, silver and gold and there are few rhodium-bearing minerals. It is found in platinum ores and extracted as a white inert metal, difficult to fuse. Principal sources are located in South Africa. Although the quantity at Sudbury is small, the large amount of processed nickel ore makes rhodium recovery cost-effective; the main exporter of rhodium is South
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
Vanadium is a chemical element with symbol V and atomic number 23. It is a hard, silvery-grey, malleable transition metal; the elemental metal is found in nature, but once isolated artificially, the formation of an oxide layer somewhat stabilizes the free metal against further oxidation. Andrés Manuel del Río discovered compounds of vanadium in 1801 in Mexico by analyzing a new lead-bearing mineral he called "brown lead", presumed its qualities were due to the presence of a new element, which he named erythronium since, upon heating, most of the salts turned red. Four years however, he was convinced by other scientists that erythronium was identical to chromium. Chlorides of vanadium were generated in 1830 by Nils Gabriel Sefström who thereby proved that a new element was involved, which he named "vanadium" after the Scandinavian goddess of beauty and fertility, Vanadís. Both names were attributed to the wide range of colors found in vanadium compounds. Del Rio's lead mineral was renamed vanadinite for its vanadium content.
In 1867 Henry Enfield Roscoe obtained the pure element. Vanadium occurs in about 65 minerals and in fossil fuel deposits, it is produced in Russia from steel smelter slag. It is used to produce specialty steel alloys such as high-speed tool steels; the most important industrial vanadium compound, vanadium pentoxide, is used as a catalyst for the production of sulfuric acid. The vanadium redox battery for energy storage may be an important application in the future. Large amounts of vanadium ions are found in a few organisms as a toxin; the oxide and some other salts of vanadium have moderate toxicity. In the ocean, vanadium is used by some life forms as an active center of enzymes, such as the vanadium bromoperoxidase of some ocean algae. Vanadium was discovered by Andrés Manuel del Río, a Spanish-Mexican mineralogist, in 1801. Del Río extracted the element from a sample of Mexican "brown lead" ore named vanadinite, he found that its salts exhibit a wide variety of colors, as a result he named the element panchromium.
Del Río renamed the element erythronium because most of the salts turned red upon heating. In 1805, the French chemist Hippolyte Victor Collet-Descotils, backed by del Río's friend Baron Alexander von Humboldt, incorrectly declared that del Río's new element was only an impure sample of chromium. Del Río retracted his claim. In 1831, the Swedish chemist Nils Gabriel Sefström rediscovered the element in a new oxide he found while working with iron ores; that same year, Friedrich Wöhler confirmed del Río's earlier work. Sefström chose a name beginning with V, he called the element vanadium after Old Norse Vanadís, because of the many beautifully colored chemical compounds it produces. In 1831, the geologist George William Featherstonhaugh suggested that vanadium should be renamed "rionium" after del Río, but this suggestion was not followed; the isolation of vanadium metal proved difficult. In 1831, Berzelius reported the production of the metal, but Henry Enfield Roscoe showed that Berzelius had in fact produced the nitride, vanadium nitride.
Roscoe produced the metal in 1867 by reduction of vanadium chloride, VCl2, with hydrogen. In 1927, pure vanadium was produced by reducing vanadium pentoxide with calcium; the first large-scale industrial use of vanadium was in the steel alloy chassis of the Ford Model T, inspired by French race cars. Vanadium steel allowed for reduced weight while increasing tensile strength. For the first decade of the 20th century, most vanadium ore was mined by American Vanadium Company from the Minas Ragra in Peru; the demand for uranium rose, leading to increased mining of that metal's ores. One major uranium ore was carnotite, which contains vanadium. Uranium mining began to supply a large share of the demand for vanadium. German chemist Martin Henze discovered vanadium in the hemovanadin proteins found in blood cells of Ascidiacea in 1911. Vanadium is a medium-hard, steel-blue metal, it is electrically thermally insulating. Some sources describe vanadium as "soft" because it is ductile and not brittle. Vanadium steels.
It has good resistance to corrosion and it is stable against alkalis and sulfuric and hydrochloric acids. It is oxidized in air at about 933 K, although an oxide passivation layer forms at room temperature. Occurring vanadium is composed of one stable isotope, 51V, one radioactive isotope, 50V; the latter has a half-life of 1.5×1017 years and a natural abundance of 0.25%. 51V has a nuclear spin of 7⁄2, useful for NMR spectroscopy. Twenty-four artificial radioisotopes have been characterized, ranging in mass number from 40 to 65; the most stable of these isotopes are 49V with a half-life of 330 days, 48V with a half-life of 16.0 days. The remaining radioactive isotopes have half-lives shorter than an hour, most below 10 seconds. At least four isotopes have metastable excited states. Electron capture is the main decay mode for isotopes lighter than 51V. For the heavier ones, the most common mode is beta decay; the electron capture reactions lead