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
Radium is a chemical element with symbol Ra and atomic number 88. It is the sixth element in group 2 of the periodic table known as the alkaline earth metals. Pure radium is silvery-white, but it reacts with nitrogen on exposure to air, forming a black surface layer of radium nitride. All isotopes of radium are radioactive, with the most stable isotope being radium-226, which has a half-life of 1600 years and decays into radon gas; when radium decays, ionizing radiation is a product, which can excite fluorescent chemicals and cause radioluminescence. Radium, in the form of radium chloride, was discovered by Marie and Pierre Curie in 1898, they extracted the radium compound from uraninite and published the discovery at the French Academy of Sciences five days later. Radium was isolated in its metallic state by Marie Curie and André-Louis Debierne through the electrolysis of radium chloride in 1911. In nature, radium is found in uranium and thorium ores in trace amounts as small as a seventh of a gram per ton of uraninite.
Radium is not necessary for living organisms, adverse health effects are when it is incorporated into biochemical processes because of its radioactivity and chemical reactivity. Other than its use in nuclear medicine, radium has no commercial applications. Today, these former applications are no longer in vogue because radium's toxicity has since become known, less dangerous isotopes are used instead in radioluminescent devices. Radium is the only radioactive member of its group, its physical and chemical properties most resemble its lighter congener barium. Pure radium is a volatile silvery-white metal, although its lighter congeners calcium and barium have a slight yellow tint; this tint vanishes on exposure to air, yielding a black layer of radium nitride. Its melting point is either 700 °C or 960 °C and its boiling point is 1,737 °C. Both of these values are lower than those of barium, confirming periodic trends down the group 2 elements. Like barium and the alkali metals, radium crystallizes in the body-centered cubic structure at standard temperature and pressure: the radium–radium bond distance is 514.8 picometers.
Radium has a density of 5.5 g/cm3, higher than that of barium, again confirming periodic trends. Radium has 33 known isotopes, with mass numbers from 202 to 234: all of them are radioactive. Four of these – 223Ra, 224Ra, 226Ra, 228Ra – occur in the decay chains of primordial thorium-232, uranium-235, uranium-238; these isotopes still have half-lives too short to be primordial radionuclides and only exist in nature from these decay chains. Together with the artificial 225Ra, these are the five most stable isotopes of radium. All other known radium isotopes have half-lives under two hours, the majority have half-lives under a minute. At least 12 nuclear isomers have been reported. In the early history of the study of radioactivity, the different natural isotopes of radium were given different names. In this scheme, 223Ra was named actinium X, 224Ra thorium X, 226Ra radium, 228Ra mesothorium 1; when it was realized that all of these are isotopes of the same element, many of these names fell out of use, "radium" came to refer to all isotopes, not just 226Ra.
Some of radium-226's decay products received historical names including "radium", ranging from radium A to radium G, with the letter indicating how far they were down the chain from their parent 226Ra.226Ra is the most stable isotope of radium and is the last isotope in the decay chain of uranium-238 with a half-life of over a millennium: it makes up all of natural radium. Its immediate decay product is the dense radioactive noble gas radon, responsible for much of the danger of environmental radium, it is 2.7 million times more radioactive than the same molar amount of natural uranium, due to its proportionally shorter half-life. A sample of radium metal maintains itself at a higher temperature than its surroundings because of the radiation it emits – alpha particles, beta particles, gamma rays. More natural radium emits alpha particles, but other steps in its decay chain emit alpha or beta particles, all particle emissions are accompanied by gamma rays. In 2013 it was discovered; this was the first discovery of an asymmetric nucleus.
Radium, like barium, is a reactive metal and always exhibits its group oxidation state of +2. It forms the colorless Ra2+ cation in aqueous solution, basic and does not form complexes readily. Most radium compounds are therefore simple ionic compounds, though participation from the 6s and 6p electrons is expected due to relativistic effects and would enhance the covalent character of radium compounds such as RaF2 and RaAt2. For this reason, the standard electrode potential for the half-reaction Ra2+ + 2e− →
The pascal is the SI derived unit of pressure used to quantify internal pressure, Young's modulus and ultimate tensile strength. It is defined as one newton per square metre, it is named after the French polymath Blaise Pascal. Common multiple units of the pascal are the hectopascal, equal to one millibar, the kilopascal, equal to one centibar; the unit of measurement called. Meteorological reports in the United States state atmospheric pressure in millibars. In Canada these reports are given in kilopascals; the unit is named after Blaise Pascal, noted for his contributions to hydrodynamics and hydrostatics, experiments with a barometer. The name pascal was adopted for the SI unit newton per square metre by the 14th General Conference on Weights and Measures in 1971; the pascal can be expressed using SI derived units, or alternatively SI base units, as: 1 P a = 1 N m 2 = 1 k g m ⋅ s 2 = 1 J m 3 where N is the newton, m is the metre, kg is the kilogram, s is the second, J is the joule. One pascal is the pressure exerted by a force of magnitude one newton perpendicularly upon an area of one square metre.
The unit of measurement called a standard atmosphere is 101325 Pa.. This value is used as a reference pressure and specified as such in some national and international standards, such as the International Organization for Standardization's ISO 2787, ISO 2533 and ISO 5024. In contrast, International Union of Pure and Applied Chemistry recommends the use of 100 kPa as a standard pressure when reporting the properties of substances. Unicode has dedicated code-points U+33A9 ㎩ SQUARE PA and U+33AA ㎪ SQUARE KPA in the CJK Compatibility block, but these exist only for backward-compatibility with some older ideographic character-sets and are therefore deprecated; the pascal or kilopascal as a unit of pressure measurement is used throughout the world and has replaced the pounds per square inch unit, except in some countries that still use the imperial measurement system or the US customary system, including the United States. Geophysicists use the gigapascal in measuring or calculating tectonic stresses and pressures within the Earth.
Medical elastography measures tissue stiffness non-invasively with ultrasound or magnetic resonance imaging, displays the Young's modulus or shear modulus of tissue in kilopascals. In materials science and engineering, the pascal measures the stiffness, tensile strength and compressive strength of materials. In engineering use, because the pascal represents a small quantity, the megapascal is the preferred unit for these uses; the pascal is equivalent to the SI unit of energy density, J/m3. This applies not only to the thermodynamics of pressurised gases, but to the energy density of electric and gravitational fields. In measurements of sound pressure or loudness of sound, one pascal is equal to 94 decibels SPL; the quietest sound a human can hear, known as the threshold of hearing, is 20 µPa. The airtightness of buildings is measured at 50 Pa; the units of atmospheric pressure used in meteorology were the bar, close to the average air pressure on Earth, the millibar. Since the introduction of SI units, meteorologists measure pressures in hectopascals unit, equal to 100 pascals or 1 millibar.
Exceptions include Canada. In many other fields of science, the SI is preferred. Many countries use the millibars. In all other fields, the kilopascal is used instead. Atmospheric pressure which gives the usage of the hbar end the mbar Centimetre of water Meteorology Metric prefix Orders of magnitude Pascal's law Pressure measurement
Dysprosium is a chemical element with symbol Dy and atomic number 66. It is a rare earth element with a metallic silver luster. Dysprosium is never found in nature as a free element, though it is found in various minerals, such as xenotime. Occurring dysprosium is composed of seven isotopes, the most abundant of, 164Dy. Dysprosium was first identified in 1886 by Paul Émile Lecoq de Boisbaudran, but it was not isolated in pure form until the development of ion exchange techniques in the 1950s. Dysprosium has few applications where it cannot be replaced by other chemical elements, it is used for its high thermal neutron absorption cross-section in making control rods in nuclear reactors, for its high magnetic susceptibility in data storage applications, as a component of Terfenol-D. Soluble dysprosium salts are mildly toxic. Dysprosium is a rare earth element that has a bright silver luster, it is quite soft, can be machined without sparking if overheating is avoided. Dysprosium's physical characteristics can be affected by small amounts of impurities.
Dysprosium and holmium have the highest magnetic strengths of the elements at low temperatures. Dysprosium has a simple ferromagnetic ordering at temperatures below 85 K. Above 85 K, it turns into a helical antiferromagnetic state in which all of the atomic moments in a particular basal plane layer are parallel, oriented at a fixed angle to the moments of adjacent layers; this unusual antiferromagnetism transforms into a disordered state at 179 K. Dysprosium metal tarnishes in air and burns to form dysprosium oxide: 4 Dy + 3 O2 → 2 Dy2O3Dysprosium is quite electropositive and reacts with cold water to form dysprosium hydroxide: 2 Dy + 6 H2O → 2 Dy3 + 3 H2 Dysprosium metal vigorously reacts with all the halogens at above 200 °C: 2 Dy + 3 F2 → 2 DyF3 2 Dy + 3 Cl2 → 2 DyCl3 2 Dy + 3 Br2 → 2 DyBr3 2 Dy + 3 I2 → 2 DyI3 Dysprosium dissolves in dilute sulfuric acid to form solutions containing the yellow Dy ions, which exist as a 3+ complex: 2 Dy + 3 H2SO4 → 2 Dy3+ + 3 SO2−4 + 3 H2 The resulting compound, dysprosium sulfate, is noticeably paramagnetic.
Dysprosium halides, such as DyF3 and DyBr3, tend to take on a yellow color. Dysprosium oxide known as dysprosia, is a white powder, magnetic, more so than iron oxide. Dysprosium combines with various non-metals at high temperatures to form binary compounds with varying composition and oxidation states +3 and sometimes +2, such as DyN, DyP, DyH2 and DyH3. Dysprosium carbonate, Dy23, dysprosium sulfate, Dy23, result from similar reactions. Most dysprosium compounds are soluble in water, though dysprosium carbonate tetrahydrate and dysprosium oxalate decahydrate are both insoluble in water. Two of the most abundant dysprosium carbonates, tengerite- and kozoite- are known to form via a poorly ordered precursor phase with a formula of Dy23·4H2O; this amorphous precursor consists of hydrated spherical nanoparticles of 10–20 nm diameter that are exceptionally stable under dry treatment at ambient and high temperatures. Occurring dysprosium is composed of seven isotopes: 156Dy, 158Dy, 160Dy, 161Dy, 162Dy, 163Dy, 164Dy.
These are all considered stable, although 156Dy can theoretically undergo alpha decay with a half-life of over 1×1018 years. Of the occurring isotopes, 164Dy is the most abundant at 28%, followed by 162Dy at 26%; the least abundant is 156Dy at 0.06%. Twenty-nine radioisotopes have been synthesized, ranging in atomic mass from 138 to 173; the most stable of these is 154Dy, with a half-life of 3×106 years, followed by 159Dy with a half-life of 144.4 days. The least stable is 138Dy, with a half-life of 200 ms; as a general rule, isotopes that are lighter than the stable isotopes tend to decay by β+ decay, while those that are heavier tend to decay by β− decay. However, 154Dy decays by alpha decay, 152Dy and 159Dy decay by electron capture. Dysprosium has at least 11 metastable isomers, ranging in atomic mass from 140 to 165; the most stable of these is 165mDy, which has a half-life of 1.257 minutes. 149Dy has two metastable isomers. In 1878, erbium ores were found to contain the oxides of thulium. French chemist Paul Émile Lecoq de Boisbaudran, while working with holmium oxide, separated dysprosium oxide from it in Paris in 1886.
His procedure for isolating the dysprosium involved dissolving dysprosium oxide in acid adding ammonia to precipitate the hydroxide. He was only able to isolate dysprosium from its oxide after more than 30 attempts at his procedure. On succeeding, he named the element dysprosium from the Greek dysprositos, meaning "hard to get"; the element was not isolated in pure form until after the development of ion exchange techniques by Frank Spedding at Iowa State University in the early 1950s. While dysprosium is never encountered as a free element, it is found in many minerals, including xenotime, gadolinite, polycrase, blomstrandine and bastnäsite with erbium and holmium or other rare earth elements. No dysprosium-dominant mineral has yet been found. In the high-ytt
Nobelium is a synthetic chemical element with symbol No and atomic number 102. It is named in the inventor of dynamite and benefactor of science. A radioactive metal, it is the tenth transuranic element and is the penultimate member of the actinide series. Like all elements with atomic number over 100, nobelium can only be produced in particle accelerators by bombarding lighter elements with charged particles. A total of twelve nobelium isotopes are known to exist. Chemistry experiments have confirmed that nobelium behaves as a heavier homolog to ytterbium in the periodic table; the chemical properties of nobelium are not known: they are only known in aqueous solution. Before nobelium's discovery, it was predicted that it would show a stable +2 oxidation state as well as the +3 state characteristic of the other actinides: these predictions were confirmed, as the +2 state is much more stable than the +3 state in aqueous solution and it is difficult to keep nobelium in the +3 state. In the 1950s and 1960s, many claims of the discovery of nobelium were made from laboratories in Sweden, the Soviet Union, the United States.
Although the Swedish scientists soon retracted their claims, the priority of the discovery and therefore the naming of the element was disputed between Soviet and American scientists, it was not until 1997 that International Union of Pure and Applied Chemistry credited the Soviet team with the discovery, but retained nobelium, the Swedish proposal, as the name of the element due to its long-standing use in the literature. The discovery of element 102 was a complicated process and was claimed by groups from Sweden, the United States, the Soviet Union; the first complete and incontrovertible report of its detection only came in 1966 from the Joint Institute of Nuclear Research at Dubna. The first announcement of the discovery of element 102 was announced by physicists at the Nobel Institute in Sweden in 1957; the team reported that they had bombarded a curium target with carbon-13 ions for twenty-five hours in half-hour intervals. Between bombardments, ion-exchange chemistry was performed on the target.
Twelve out of the fifty bombardments contained samples emitting MeV alpha particles, which were in drops which eluted earlier than fermium and californium. The half-life reported was 10 minutes and was assigned to either 251102 or 253102, although the possibility that the alpha particles observed were from a short-lived mendelevium isotope created from the electron capture of element 102 was not excluded; the team proposed the name nobelium for the new element, approved by IUPAC, a decision which the Dubna group characterized in 1968 as being hasty. The following year, scientists at the Lawrence Berkeley National Laboratory repeated the experiment but were unable to find any 8.5 MeV events which were not background effects. In 1959, the Swedish team attempted to explain the Berkeley team's inability to detect element 102 in 1958, maintaining that they did discover it; however work has shown that no nobelium isotopes lighter than 259No with a half-life over 3 minutes exist, that the Swedish team's results are most from thorium-225, which has a half-life of 8 minutes and undergoes triple alpha decay to polonium-213, which has a decay energy of 8.53612 MeV.
This hypothesis is lent weight by the fact that thorium-225 can be produced in the reaction used and would not be separated out by the chemical methods used. Work on nobelium showed that the divalent state is more stable than the trivalent one and hence that the samples emitting the alpha particles could not have contained nobelium, as the divalent nobelium would not have eluted with the other trivalent actinides. Thus, the Swedish team retracted their claim and associated the activity to background effects; the Berkeley team, consisting of Albert Ghiorso, Glenn T. Seaborg, John R. Walton and Torbjørn Sikkeland claimed the synthesis of element 102 in 1958; the team used the new heavy-ion linear accelerator to bombard a curium target with 12C ions. They were unable to confirm the 8.5 MeV activity claimed by the Swedes but were instead able to detect decays from fermium-250 the daughter of 254102, which had an apparent half-life of ~3 s. 1963 Dubna work confirmed that 254102 could be produced in this reaction, but that its half-life was 50±10 s.
In 1967, the Berkeley team attempted to defend their work, stating that the isotope found was indeed 250Fm but the isotope that the half-life measurements related to was californium-244, granddaughter of 252102, produced from the more abundant curium-244. Energy differences were attributed to "resolution and drift problems", although these had not been reported and should have influenced other results. 1977 experiments showed. However, 1973 work showed that the 250Fm recoil could have easily been produced from the isomeric transition of 250mFm which could have been formed in the reaction at the energy used. Given this, it is probable that no nobelium was produced in this experiment. In 1959 the team continued their studies and claimed that they were able to produce an isotope that decayed predominantly by emission of an 8.3 MeV alpha particle, with a half-life of 3 s with an assoc
An electrolyte is a substance that produces an electrically conducting solution when dissolved in a polar solvent, such as water. The dissolved electrolyte separates into cations and anions, which disperse uniformly through the solvent. Electrically, such a solution is neutral. If an electric potential is applied to such a solution, the cations of the solution are drawn to the electrode that has an abundance of electrons, while the anions are drawn to the electrode that has a deficit of electrons; the movement of anions and cations in opposite directions within the solution amounts to a current. This includes most soluble salts and bases; some gases, such as hydrogen chloride, under conditions of high temperature or low pressure can function as electrolytes. Electrolyte solutions can result from the dissolution of some biological and synthetic polymers, termed "polyelectrolytes", which contain charged functional groups. A substance that dissociates into ions in solution acquires the capacity to conduct electricity.
Sodium, chloride, calcium and phosphate are examples of electrolytes. In medicine, electrolyte replacement is needed when a person has prolonged vomiting or diarrhea, as a response to strenuous athletic activity. Commercial electrolyte solutions are available for sick children and athletes. Electrolyte monitoring is important in the treatment of bulimia; the word electrolyte derives from the Greek lytós, meaning "able to be untied or loosened". Svante Arrhenius put forth, in his 1884 dissertation, his explanation of the fact that solid crystalline salts disassociate into paired charged particles when dissolved, for which he won the 1903 Nobel Prize in Chemistry. Arrhenius's explanation was that in forming a solution, the salt dissociates into charged particles, to which Michael Faraday had given the name "ions" many years earlier. Faraday's belief had been. Arrhenius proposed that in the absence of an electric current, solutions of salts contained ions, he thus proposed. Electrolyte solutions are formed when a salt is placed into a solvent such as water and the individual components dissociate due to the thermodynamic interactions between solvent and solute molecules, in a process called "solvation".
For example, when table salt, NaCl, is placed in water, the salt dissolves into its component ions, according to the dissociation reaction NaCl → Na+ + Cl−It is possible for substances to react with water, producing ions. For example, carbon dioxide gas dissolves in water to produce a solution that contains hydronium and hydrogen carbonate ions. Molten salts can be electrolytes as, for example, when sodium chloride is molten, the liquid conducts electricity. In particular, ionic liquids, which are molten salts with melting points below 100 °C, are a type of conductive non-aqueous electrolytes and thus have found more and more applications in fuel cells and batteries. An electrolyte in a solution may be described as "concentrated" if it has a high concentration of ions, or "diluted" if it has a low concentration. If a high proportion of the solute dissociates to form free ions, the electrolyte is strong; the properties of electrolytes may be exploited using electrolysis to extract constituent elements and compounds contained within the solution.
Alkaline earth metals form hydroxides that are strong electrolytes with limited solubility in water, due to the strong attraction between their constituent ions. This limits their application to situations. In physiology, the primary ions of electrolytes are sodium, calcium, chloride, hydrogen phosphate, hydrogen carbonate; the electric charge symbols of plus and minus indicate that the substance is ionic in nature and has an imbalanced distribution of electrons, the result of chemical dissociation. Sodium is the main electrolyte found in extracellular fluid and potassium is the main intracellular electrolyte. All known higher lifeforms require a subtle and complex electrolyte balance between the intracellular and extracellular environments. In particular, the maintenance of precise osmotic gradients of electrolytes is important; such gradients affect and regulate the hydration of the body as well as blood pH, are critical for nerve and muscle function. Various mechanisms exist in living species that keep the concentrations of different electrolytes under tight control.
Both muscle tissue and neurons are considered electric tissues of the body. Muscles and neurons are activated by electrolyte activity between the extracellular fluid or interstitial fluid, intracellular fluid. Electrolytes may enter or leave the cell membrane through specialized protein structures embedded in the plasma membrane called "ion channels". For example, muscle contraction is dependent upon the presence of calcium and potassium. Without sufficient levels of these key electrolytes, muscle weakness or severe muscle contractions may occur. Electrolyte balance is maintained by oral, or in emergencies, intravenous intake of electrolyte-containing substances, is regulated by hormones, in general with the kidneys flushing out excess levels. In humans, electrolyte homeostasis is regulated by hormones such as antidiuretic hormones and parathyroid hormones. Serious electrol
Neodymium is a chemical element with symbol Nd and atomic number 60. It is a soft silvery metal. Neodymium was discovered in 1885 by the Austrian chemist Carl Auer von Welsbach, it is present in significant quantities in the ore minerals bastnäsite. Neodymium is not found in metallic form or unmixed with other lanthanides, it is refined for general use. Although neodymium is classed as a rare earth, it is a common element, no rarer than cobalt, nickel, or copper, is distributed in the Earth's crust. Most of the world's commercial neodymium is mined in China. Neodymium compounds were first commercially used as glass dyes in 1927, they remain a popular additive in glasses; the color of neodymium compounds—due to the Nd3+ ion—is a reddish-purple but it changes with the type of lighting, due to the interaction of the sharp light absorption bands of neodymium with ambient light enriched with the sharp visible emission bands of mercury, trivalent europium or terbium. Some neodymium-doped glasses are used in lasers that emit infrared with wavelengths between 1047 and 1062 nanometers.
These have been used in extremely-high-power applications, such as experiments in inertial confinement fusion. Neodymium is used with various other substrate crystals, such as yttrium aluminium garnet in the Nd:YAG laser; this laser emits infrared at a wavelength of about 1064 nanometers. The Nd:YAG laser is one of the most used solid-state lasers. Another important use of neodymium is as a component in the alloys used to make high-strength neodymium magnets—powerful permanent magnets; these magnets are used in such products as microphones, professional loudspeakers, in-ear headphones, high performance hobby DC electric motors, computer hard disks, where low magnet mass or strong magnetic fields are required. Larger neodymium magnets are used in generators. Neodymium, a rare-earth metal, was present in the classical mischmetal at a concentration of about 18%. Metallic neodymium has a bright, silvery metallic luster, but as one of the more reactive lanthanide rare-earth metals, it oxidizes in ordinary air.
The oxide layer that forms peels off, exposing the metal to further oxidation. Thus, a centimeter-sized sample of neodymium oxidizes within a year. Neodymium exists in two allotropic forms, with a transformation from a double hexagonal to a body-centered cubic structure taking place at about 863 °C. Neodymium metal tarnishes in air and it burns at about 150 °C to form neodymium oxide: 4 Nd + 3 O2 → 2 Nd2O3Neodymium is a quite electropositive element, it reacts with cold water, but quite with hot water to form neodymium hydroxide: 2 Nd + 6 H2O → 2 Nd3 + 3 H2 Neodymium metal reacts vigorously with all the halogens: 2 Nd + 3 F2 → 2 NdF3 2 Nd + 3 Cl2 → 2 NdCl3 2 Nd + 3 Br2 → 2 NdBr3 2 Nd + 3 I2 → 2 NdI3 Neodymium dissolves in dilute sulfuric acid to form solutions that contain the lilac Nd ion; these exist as a 3+ complexes: 2 Nd + 3 H2SO4 → 2 Nd3+ + 3 SO2−4 + 3 H2 Neodymium compounds include halides: neodymium fluoride. Occurring neodymium is a mixture of five stable isotopes, 142Nd, 143Nd, 145Nd, 146Nd and 148Nd, with 142Nd being the most abundant, two radioisotopes, 144Nd and 150Nd.
In all, 31 radioisotopes of neodymium have been detected as of 2010, with the most stable radioisotopes being the occurring ones: 144Nd and 150Nd. All of the remaining radioactive isotopes have half-lives that are shorter than eleven days, the majority of these have half-lives that are shorter than 70 seconds. Neodymium has 13 known meta states, with the most stable one being 139mNd, 135mNd and 133m1Nd; the primary decay modes before the most abundant stable isotope, 142Nd, are electron capture and positron decay, the primary mode after is beta minus decay. The primary decay products before 142Nd are element Pr isotopes and the primary products after are element Pm isotopes. Neodymium was discovered by Baron Carl Auer von Welsbach, an Austrian chemist, in Vienna in 1885, he separated neodymium, as well as the element praseodymium, from a material known as didymium by means of fractional crystallization of the double ammonium nitrate tetrahydrates from nitric acid, while following the separation by spectroscopic analysis.
The name neodymium is derived from the Greek words neos and didymos, twin. Double nitrate crystallization was the means of commercial neodymium purification until the 1950s. Lindsay Chemical Division was the first to commercialize large-scale ion-exchange purification of neodymium. Starting in the 1950s, high purity