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to the oral microbiome, is associated with periodontal disease and has been linked with a wide variety of health outcomes, including unfavorable pregnancy outcomes. == Immune response == The immune response becomes more pronounced when bacteria are found that are not commensal. == History == Investigations into reproductive-associated microbiomes began around 1885 by Theodor Escherich. He wrote that meconium from the newborn was free of bacteria. There was a general consensus at the time and even recently that the uterus was sterile and this was referred to as the sterile womb paradigm. Other investigations used sterile diapers for meconium collection. No bacteria were able to be cultured from the samples. Other studies showed that bacteria were detected and were directly proportional to the time between birth and the passage of meconium. == Research == Investigations into the role of the uterine microbiome in the development of the infant microbiome are ongoing. In recent years, the number of articles and review publications discussing the uterine microbiome has grown. Based on a Web of Science analysis, the highest number of documents published on the topic was in 2023, with a total of 23 papers. The Daunert Lab, based at the University of Miami’s Sylvester Comprehensive Cancer Center, focuses on the role of the microbiome in endometrial cancer and the role the uterine microbiome plays in the success of an IVF cycle. Similarly, Dr. Maria Walther-Antonio’s lab at the Mayo Clinic focuses on the microbiome’s role in endometrial cancer. Notably, Dr. Walther- Antonio has confirmed that Porphyromas somerae is able to invade endometrial cells, indicating a possibility that this microbe contributes to the pathogenesis of endometrial cancer. The Carlos Simon Foundation, based in Valencia, Spain, is an women’s health research organization founded by reproductive endocrinologist Carlos Simon, MD PhD. A research team led
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{
"page_id": 50925207,
"source": null,
"title": "Uterine microbiome"
}
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by Dr. Inmaculada Moreno at the Carlos Simon Foundation studies the role of the endometrial microbiome in human reproduction. When research on the uterine microbiome was scarce, Dr. Moreno and her team analyzed the endometrial microbiota and discovered that there was a correlation between certain endometrial microbiota compositions and the outcome of implantation success or failure. Six years later, they followed up with a paper revealing that specific pathogenic bacteria and depletion of Lactobacillus spp. in the endometrium correlated with impaired fertility. == See also == Human microbiome Human Microbiome Project Human virome List of antimicrobial peptides in the female reproductive tract List of bacterial vaginosis microbiota Placental microbiome Vaginal epithelium Vaginal flora in pregnancy == References and notes ==
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{
"page_id": 50925207,
"source": null,
"title": "Uterine microbiome"
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Bánh tráng or bánh đa nem, a Vietnamese term (literally, coated bánh), sometimes called rice paper wrappers, rice crepes, rice wafers or nem wrappers, are edible Vietnamese wrappers used in Vietnamese cuisine, primarily in finger foods and appetizers such as Vietnamese nem dishes. The term rice paper wrappers can sometimes be a misnomer, as some banh trang wrappers are made from rice flour supplemented with tapioca flour or sometimes replaced completely with tapioca starch. The roasted version is bánh tráng nướng. == Description == Vietnamese banh trang are rice paper wrappers that are edible. They are made from steamed rice batter, then sun-dried. A more modern method is to use machines that can steam and dry the wrapper for a thinner and more hygienic product, suitable for the export market. == Types == Vietnamese banh trang wrappers come in various textures, shapes and types. Textures may vary from thin, soft to thick (much like a rice cracker). Banh trang wrappers come in various shapes, though circular and squared shapes are most commonly used. A plethora of local Vietnamese ingredients and spices are added to Vietnamese banh trang wrappers for the purpose of creating different flavors and textures, such as sesame seeds, chili, coconut milk, bananas, and durian, to name a few. === Bánh tráng === Southern Vietnamese term for rice wrappers, which are also commonly used overseas. These banh trang wrappers are made from a mixture of rice flour with tapioca starch, water and salt. These wrappers are thin and light in texture. They are often used for chả giò and gỏi cuốn. There are also certain rice wrappers products that are specifically for frying. === Bánh đa nướng / Bánh tráng nướng (grilled rice cracker) === Bánh đa nướng or bánh tráng nướng are roasted or grilled rice crackers.
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{
"page_id": 3477144,
"source": null,
"title": "Bánh tráng"
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Some can be thicker than the standard rice wrapper and can also include sesame seeds. It is often used in dishes like mì Quảng as a topping. Not to be confused with the street food dish from Đà Lạt (also known as Vietnamese pizza or quesadilla) with the same name. === Bánh đa nem (vỏ ram / bánh ram) === This is a northern Vietnamese term for rice paper. There is a special variety that contains only rice flour and no tapioca starch, water and salt. Brown bánh đa nem may contain cane syrup (mật mía). These wrappers are very thin and translucent. They also do not require to be soaked with water before usage. When fried, they are crispier and not as chewy like its southern counterpart. It is used as a wrapper when making nem cua bể. === Bánh tráng rế (woven banh trang) === These banh trang wrappers are made from rice flour, green beans, vegetable oil and salt. These wrappers are delicate and thin. They are lacy, net-like wrappers typically used for deep-fried cha gio rolls. === Bánh tráng mè (Sesame banh trang) === These banh trang wrappers are typically made from rice starch, then adding sesame seeds. Its texture resembles that of a rice cracker. === Bánh tráng tôm mè (Sesame-shrimp banh trang) === These banh trang wrappers are made by adding sesame seeds and dried shrimps. Its texture resembles that of a rice cracker. === Bánh tráng sữa (Milky banh trang) === Bánh tráng sữa are made by adding milk. This type of banh trang is softer, and supposed to melt on your tongue. === Bánh tráng dẻo === Bánh tráng dẻo are soft and malleable sheets. They also come in different flavors. === Bánh tráng mỏng (thin bánh tráng) === Thin sheets made
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{
"page_id": 3477144,
"source": null,
"title": "Bánh tráng"
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from only tapioca starch. They are dried and translucent and are more sticky when activated with water. === Bánh tráng chuối (Banana flavored banh trang) === These banh trang wrappers are typically made by adding bananas. Its texture resembles that of a rice cracker. === Bánh tráng dừa (Coconut flavored banh trang) === These banh trang wrappers are typically made by adding coconut milk, sugar, rice flour, sesame seeds, and water. The texture resembles that of a cracker, similar to the sesame banh trang. === Bánh tráng phơi sương === Bánh tráng phơi sương are wrappers with flexible two layered rice paper. === Bánh tráng trộn (Mixed rice paper) === Mixed rice paper has the main ingredient is rice paper cut yarn, mixed with fried dried shrimp with fat. == Culinary uses == Banh trang wrappers are typically used in Vietnamese nem dishes. These wrappers are eaten dried (khô), fried (rán), baked (nướng) or soaked (ướt). They are typically served rolled (cuộn) or baked (nướng), in salads, soups and stirred fried Vietnamese dishes. The light, translucent traditional banh trang wrappers are typically used for various Vietnamese rolls, more commonly the goi cuon (salad rolls). Though commonly used in fresh rolls, Northern Vietnamese cuisine often use these wrappers in chả giò (Northern Vietnamese: Nem rán), a crispy, fried springroll. Traditional banh trang wrappers are also used to wrap common Vietnamese dishes such as banh xeo (Vietnamese sizzling pancakes), bò 7 món (Vietnamese seven courses of beef) and cá nướng (Vietnamese grilled fish) and then dipped into a sauce. The traditional banh trang wrappers are also used to make a Vietnamese salad dish called bánh tráng trộn (stirred banh trang salad). Woven banh trang wrappers are typically deep-fried to make aesthetically appealing cha gio (Vietnamese crispy spring rolls). Sesame banh trang wrappers
|
{
"page_id": 3477144,
"source": null,
"title": "Bánh tráng"
}
|
are typically baked or soaked in water, depending on individual textural preference, then served with salads, mi Quang and various other dishes. == Outside Vietnam == Banh trang wrappers are found in countries outside Vietnam with Vietnamese diaspora (people who leave Vietnam). == References ==
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{
"page_id": 3477144,
"source": null,
"title": "Bánh tráng"
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Immunoelectrophoresis is a general name for a number of biochemical methods for separation and characterization of proteins based on electrophoresis and reaction with antibodies. All variants of immunoelectrophoresis require immunoglobulins, also known as antibodies, reacting with the proteins to be separated or characterized. The methods were developed and used extensively during the second half of the 20th century. In somewhat chronological order: Immunoelectrophoretic analysis (one-dimensional immunoelectrophoresis ad modum Grabar), crossed immunoelectrophoresis (two-dimensional quantitative immunoelectrophoresis ad modum Clarke and Freeman or ad modum Laurell), rocket-immunoelectrophoresis (one-dimensional quantitative immunoelectrophoresis ad modum Laurell), fused rocket immunoelectrophoresis ad modum Svendsen and Harboe, affinity immunoelectrophoresis ad modum Bøg-Hansen. == Methods == Immunoelectrophoresis is a general term describing many combinations of the principles of electrophoresis and reaction of antibodies, also known as immunodiffusion. Agarose as 1% gel slabs of about 1 mm thickness buffered at high pH (around 8.6) is traditionally preferred for electrophoresis and the reaction with antibodies. The agarose was chosen as the gel matrix because it has large pores allowing free passage and separation of proteins but provides an anchor for the immunoprecipitates of protein and specific antibodies. The high pH was chosen because antibodies are practically immobile at high pH. Electrophoresis equipment with a horizontal cooling plate was normally recommended for the electrophoresis. Immunoprecipitates are visible in the wet agarose gel, but are stained with protein stains like Coomassie brilliant blue in the dried gel. In contrast to SDS-gel electrophoresis, the electrophoresis in agarose allows native conditions, preserving the native structure and activities of the proteins under investigation, therefore immunoelectrophoresis allows characterization of enzyme activities and ligand binding etc. in addition to electrophoretic separation. Counterimmunoelectrophoresis is the combination of immunodiffusion with electrophoresis. In essence electrophoresis speeds up the process of moving the reactants together. The immunoelectrophoretic analysis ad modum Grabar is the
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{
"page_id": 2428570,
"source": null,
"title": "Immunoelectrophoresis"
}
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classical method of immunoelectrophoresis. Proteins are separated by electrophoresis, then antibodies are applied in a trough next to the separated proteins and immunoprecipitates are formed after a period of diffusion of the separated proteins and antibodies against each other. The introduction of the immunoelectrophoretic analysis gave a great boost to protein chemistry, some of the first results were the resolution of proteins in biological fluids and biological extracts. Among the important observations made were the great number of different proteins in serum, the existence of several immunoglobulin classes and their electrophoretic heterogeneity. Crossed immunoelectrophoresis is also called two-dimensional quantitative immunoelectrophoresis ad modum Clarke and Freeman or ad modum Laurell. In this method the proteins are first separated during the first dimension electrophoresis, then instead of the diffusion towards the antibodies, the proteins are electrophoresed into an antibody-containing gel in the second dimension. Immunoprecipitation will take place during the second dimension electrophorsis and the immunoprecipitates have a characteristic bell-shape, each precipitate representing one antigen, the position of the precipitate being dependent on the amount of protein as well as the amount of specific antibody in the gel, so relative quantification can be performed. The sensitivity and resolving power of crossed immunoelectrophoresis is than that of the classical immunoelectrophoretic analysis and there are multiple variations of the technique useful for various purposes. Crossed immunoelectrophoresis has been used for studies of proteins in biological fluids, particularly human serum, and biological extracts. Rocket immunoelectrophoresis is one-dimensional quantitative immunoelectrophoresis. The method has been used for quantitation of human serum proteins before automated methods became available. Fused rocket immunoelectrophoresis is a modification of one-dimensional quantitative immunoelectrophorsis used for detailed measurement of proteins in fractions from protein separation experiments. Affinity immunoelectrophoresis is based on changes in the electrophoretic pattern of proteins through specific interaction or complex formation
|
{
"page_id": 2428570,
"source": null,
"title": "Immunoelectrophoresis"
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with other macromolecules or ligands. Affinity immunoelectrophoresis has been used for estimation of binding constants, as for instance with lectins or for characterization of proteins with specific features like glycan content or ligand binding. Some variants of affinity immunoelectrophoresis are similar to affinity chromatography by use of immobilized ligands. Binding of ligands. The open structure of the immunoprecipitate in the agarose gel will allow additional binding of radioactively labeled antibodies and other ligands to reveal specific proteins. Application of this possibility has been used for instance for identification of allergens through reaction with immunoglobulin E (IgE) and for identification of glycoproteins with lectins. General comments. Two factors determine that immunoelectrophoretic methods are not widely used. First they are rather work intensive and require some manual expertise. Second they require rather large amounts of polyclonal antibodies. Today gel electrophoresis followed by electroblotting is the preferred method for protein characterization because its ease of operation, its high sensitivity, and its low requirement for specific antibodies. In addition proteins are separated by gel electrophoresis on the basis of their apparent molecular weight, which is not accomplished by immunoelectrophoresis, but nevertheless immunoelectrophoretic methods are still useful when non-reducing conditions are needed. == Applications == Counter-immunoelectrophoresis and its modification In comparison to other conventional methods of diagnosis e.g. for viral infection testing, counter-immunoelectrophoresis is a highly specific, simple, and speedy method that does not require sophisticated, expensive tools, input materials, or long-term capacity building. Considering the high informativeness of counter-immunoelectrophoresis, the results in practice can be dubious at times. As a result, by using a manufactured amphiphilic fluorescein-containing copolymer to increase the antigen and antibody interaction, counter-immunoelectrophoresis procedures can be improved. The use of the fluorescein copolymer-antigen mixture improved the association with plasma levels antibodies of animals immunized against hemorrhage illness and enhanced protein concentration
|
{
"page_id": 2428570,
"source": null,
"title": "Immunoelectrophoresis"
}
|
in the precipitated zone, according to the findings. The capability of the amphiphilic fluorescein copolymer to boost antigen-antibody association and see the fluorescent accumulation domain may improve the efficiency of counter-immunoelectrophoresis for infectious disease rapid diagnosis. Immunomethods The terminologies, immune-methods and immune-chemical techniques refer to a variety of immunoelectrophoresis processes whose results are identified using antibodies and immunological methodologies. As a result, immunomethods' great sensitivity is a beneficial compared to the great expense of utilizing antibodies. Many different types of agarose electrophoresis are used to see how proteins travel under diverse circumstances. Proteins are recognized after the timer has expired by incubating gels with certain antibodies, which are then stained with Comassie blue. Radial immunodiffusion The radial immunodiffusion is an immunoassay technique for determining the concentration of a particular protein in a mixture including other modules. It is made up of an agarose gel, just like the others. Furthermore, in this procedure, the materials are placed into round wells in the gel's core part and disperse through it, generating a deposition ring with a diameter relation to the number of unbound protein that has diffused. Identification of nanomaterial interaction with C3 protein complement and 2D immunoelectrophoresis 2D immunoelectrophoresis is a potential method that can be used for a range of functions involving protein flow of migrants, such as the deep examination of protein opsonization, in succession of first dimension as an activity of protein molar mass and the second dimension as a role of the isoelectric point. Despite the fact that it contains a large number of proteins, each spot on the 2D gel will symbolize a particular protein with a specific molecular mass and feature. 2D immunoelectrophoresis is also provided as a valuable implement for examining the stimulation of the signal transduction pathway, which is an essential factor in
|
{
"page_id": 2428570,
"source": null,
"title": "Immunoelectrophoresis"
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researching nanoparticles before in vivo delivery, because it will impact nanoparticle longevity, destination, and bio-distribution. This method employs two-dimensional horizontally agarose protein electrophoresis to specifically identify the association of nanoparticles with the C3 protein. Proteins can be separated in the first dimension according to their molecular mass (the shorter the protein, the far it drifts), and in the second dimension according to their abundance Some limitations of immunoelectrophoresis Though immunoelectrophoresis has a number of benefits, it also has certain drawbacks, such as when compared to other methods of electrophoresis, such as immunofixation, this method is sluggish and less precise. It can be difficult to interpret the results. Several tiny monoclonal proteins may be harder to identify. The accessibility of particular antibodies limits its utility in analytical techniques. Traditional (classical or conventional) immunoelectrophoresis has a number of drawbacks, including the fact that it is time consuming and the protocol might take up to 3 days to finish, has limited specificity and sensitivity, and the results can be difficult to read. As a result, newer immunoelectrophoresis techniques have largely supplanted the conventional immunoelectrophoresis. == References == == External links == Comprehensive text edited by Niels H. Axelsen in Scandinavian Journal of Immunology, 1975 Volume 4 Supplement Immunoelectrophoresis at the U.S. National Library of Medicine Medical Subject Headings (MeSH) https://web.archive.org/web/20070612225626/http://www.lib.mcg.edu/edu/esimmuno/ch4/immelec.htm Immuno-Electrophoresis. Immuno-Diffusion Archived 2011-12-20 at the Wayback Machine
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{
"page_id": 2428570,
"source": null,
"title": "Immunoelectrophoresis"
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Cytochrome P450, family 25, also known as CYP25, is a nematoda cytochrome P450 monooxygenase family. The first gene identified in this family is the CYP25A1 from the Caenorhabditis elegans. == References ==
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{
"page_id": 69275290,
"source": null,
"title": "CYP25 family"
}
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Probabilistic genotyping is the use of statistical methods and mathematical algorithms in DNA Profiling. It may be used instead of manual methods in difficult situations, such as when a DNA sample is very small or includes a mixture of multiple individuals' DNA. Probabilistic genotyping, unlike traditional methods, avoids the need for subjective judgment. The reliability of the method has been questioned by some defense lawyers because the source code of some probabilistic genotyping programs is proprietary. == References ==
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{
"page_id": 60165786,
"source": null,
"title": "Probabilistic genotyping"
}
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This is a list of molecules that have been detected in the interstellar medium and circumstellar envelopes, grouped by the number of component atoms. The chemical formula is listed for each detected compound, along with any ionized form that has also been observed. == Background == The molecules listed below were detected through astronomical spectroscopy. Their spectral features arise because molecules either absorb or emit a photon of light when they transition between two molecular energy levels. The energy (and thus the wavelength) of the photon matches the energy difference between the levels involved. Molecular electronic transitions occur when one of the molecule's electrons moves between molecular orbitals, producing a spectral line in the ultraviolet, optical or near-infrared parts of the electromagnetic spectrum. Alternatively, a vibrational transition transfers quanta of energy to (or from) vibrations of molecular bonds, producing signatures in the mid- or far-infrared. Gas-phase molecules also have quantised rotational levels, leading to transitions at microwave or radio wavelengths. Sometimes a transition can involve more than one of these types of energy level e.g. ro-vibrational spectroscopy changes both the rotational and vibrational energy level. Occasionally all three occur together, as in the Phillips band of C2 (diatomic carbon), in which an electronic transition produces a line in the near-infrared, which is then split into several vibronic bands by a simultaneous change in vibrational level, which in turn are split again into rotational branches. The spectrum of a particular molecule is governed by the selection rules of quantum chemistry and by its molecular symmetry. Some molecules have simple spectra which are easy to identify, whilst others (even some small molecules) have extremely complex spectra with flux spread among many different lines, making them far harder to detect. Interactions between the atomic nuclei and the electrons sometimes cause further hyperfine structure
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{
"page_id": 921245,
"source": null,
"title": "List of interstellar and circumstellar molecules"
}
|
of the spectral lines. If the molecule exists in multiple isotopologues (versions containing different atomic isotopes), the spectrum is further complicated by isotope shifts. Detection of a new interstellar or circumstellar molecule requires identifying a suitable astronomical object where it is likely to be present, then observing it with a telescope equipped with a spectrograph working at the required wavelength, spectral resolution and sensitivity. The first molecule detected in the interstellar medium was the methylidyne radical (CH•) in 1937, through its strong electronic transition at 4300 angstroms (in the optical). Advances in astronomical instrumentation have led to increasing numbers of new detections. From the 1950s onwards, radio astronomy began to dominate new detections, with sub-mm astronomy also becoming important from the 1990s. The inventory of detected molecules is highly biased towards certain types which are easier to detect. For example, radio astronomy is most sensitive to small linear molecules with a high molecular dipole. The most common molecule in the Universe, H2 (molecular hydrogen), is completely invisible to radio telescopes because it has no dipole; its electronic transitions are too energetic for optical telescopes, so detection of H2 required ultraviolet observations with a sounding rocket. Vibrational lines are often not specific to an individual molecule, allowing only the general class to be identified. For example, the vibrational lines of polycyclic aromatic hydrocarbons (PAHs) were identified in 1984, showing the class of molecules is very common in space, but it took until 2021 to identify any specific PAHs through their rotational lines. One of the richest sources for detecting interstellar molecules is Sagittarius B2 (Sgr B2), a giant molecular cloud near the centre of the Milky Way. About half of the molecules listed below were first found in Sgr B2, and many of the others have been subsequently detected there. Many
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{
"page_id": 921245,
"source": null,
"title": "List of interstellar and circumstellar molecules"
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|
of the largest molecules were first detected in another molecular cloud, TMC-1. A rich source of circumstellar molecules is CW Leonis (also known as IRC +10216), a nearby carbon star, where about 50 molecules have been identified. There is no clear boundary between interstellar and circumstellar media, so both are included in the tables below. The discipline of astrochemistry includes understanding how these molecules form and explaining their abundances. The extremely low density of the interstellar medium is not conducive to the formation of molecules, making conventional gas-phase reactions between neutral species (atoms or molecules) inefficient. Many regions also have very low temperatures (typically 10 kelvin inside a molecular cloud), further reducing the reaction rates, or high ultraviolet radiation fields, which destroy molecules through photochemistry. Explaining the observed abundances of interstellar molecules requires calculating the balance between formation and destruction rates using gas-phase ion chemistry (often driven by cosmic rays), surface chemistry on cosmic dust, radiative transfer including interstellar extinction, and sophisticated reaction networks. The use of molecular lines to determine the physical properties of astronomical objects is known as molecular astrophysics. == Molecules == The following tables list molecules that have been detected in the interstellar medium or circumstellar matter, grouped by the number of component atoms. Neutral molecules and their molecular ions are listed in separate columns; if there is no entry in the molecule column, only the ionized form has been detected. Designations (names of molecules) are those used in the scientific literature describing the detection; if none was given that field is left empty. Mass is listed in atomic mass units. Deuterated molecules, which contain at least one deuterium (2H) atom, have slightly different masses and are listed in a separate table. The total number of unique species, including distinct ionization states, is indicated in each
|
{
"page_id": 921245,
"source": null,
"title": "List of interstellar and circumstellar molecules"
}
|
section header. Most of the molecules detected so far are organic. The only detected inorganic molecule with five or more atoms is SiH4. Molecules larger than that all have at least one carbon atom, with no N−N or O−O bonds. === Diatomic (45) === === Triatomic (45) === === Four atoms (31) === === Five atoms (21) === === Six atoms (16) === === Seven atoms (16) === === Eight atoms (14) === === Nine atoms (11) === === Ten or more atoms (24) === == Deuterated molecules (22) == These molecules all contain one or more deuterium atoms, a heavier isotope of hydrogen. == Unconfirmed (16) == Evidence for the existence of the following molecules has been reported in the scientific literature, but the detections either are described as tentative by the authors, or have been challenged by other researchers. They await independent confirmation. == See also == == Notes == == References == == External links == Woon, David E. (October 1, 2010). "Interstellar and Circumstellar Molecules". Retrieved 2010-10-04. "Molecules in Space". Universität zu Köln. April 2022. Retrieved 2022-05-25. Dworkin, Jason P. (February 1, 2007). "Interstellar Molecules". NASA's Cosmic Ice Lab. Retrieved 2010-12-23. Wootten, Al (November 2005). "The 129 reported interstellar and circumstellar molecules". National Radio Astronomy Observatory. Retrieved 2007-02-13. Lovas, F. J.; Dragoset, R. A. (February 2004). "NIST Recommended Rest Frequencies for Observed Interstellar Molecular Microwave Transitions, 2002 Revision". Journal of Physical and Chemical Reference Data. 33 (1): 177. Bibcode:2004JPCRD..33..177L. doi:10.1063/1.1633275. Archived from the original on 2013-02-01. Retrieved 2007-02-13. Williams, David A.; Cecchi-Pestellini, Cesare (8 February 2023). Astrochemistry: Chemistry in Interstellar and Circumstellar Space. Royal Society of Chemistry. ISBN 978-1-83916-939-7.
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{
"page_id": 921245,
"source": null,
"title": "List of interstellar and circumstellar molecules"
}
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Oxidation and reduction in red wines can lead to a particularly undesirable brick red color in red wines (or an "orangey" color in white wines). This process is sometimes referred to as browning. In chemical terms, this is called a redox reaction because first the color of the wine deepens after fermentation through oxidation, and then the color begins to brown after bottling through reduction. Browning is not strictly considered a bad thing in every wine that is produced, as sometimes the oxidation process can contribute to the style of the wine in a positive and desirable way. Generally speaking, however, browning is a sign of the wine going stale from too much exposure to oxygen. Although wine that has gone bad is typically associated with the smell and taste of vinegar or unwanted effervescence, oxidation itself can actually lead to "nutty", "applesauce", and "burnt marshmallow" aromas. The sharp vinegar component comes from acetic acid formed through bacterial processes. A lower pH level is typically preferred in wines because it decreases the overall risk of spoilage. The lower pH reduces the effect of browning and yields better colors, particularly in red wines. == Significance of color == The color of a red wine will have many variables that influence it besides its exposure to oxygen that pertain to other viticultural aspects such as different growing climates, cultivars, and production methods. For example, wines made from hybrid grapes can range anywhere from light pink to purple due to their chemical composition. Dr. Leo McCloskey, best known for his contributions in developing the Enologix software, points out that "Chemical ecology says that a wine's flavor, color and fragrance are expressions of its ecosystem." However, when people perceive wines, the visual cues received from the color of the wine have a strong impact
|
{
"page_id": 55840419,
"source": null,
"title": "Browning in red wine"
}
|
on their opinion of the wine. In a primitive sense, visual cues will generally have a stronger impact on humans than aromatic components since humans have evolved to identify things with eyesight, rather than through scent relative to other evolved species that rely on different sensory capacities more than they do on eyesight. This idea of a particular significance of color in the perception of wine is especially true in red wines. Browning is therefore an important issue both for winemakers and for wine consumers, and is generally avoided. A well-known exception to this is sherry, which is often purposely developed with oxidation as part of the production process; however, it can be considered a white wine since it is traditionally made from Palomino grapes. == Odor effects == Through the aging of wine, its aroma compounds are affected as well. The esters in the wine that were initially created during fermentation become reduced as they break back down into alcohol and carboxyl groups so the floral, fruity smells that are associated with freshness in a wine decrease. Meanwhile, the nutty, cooked smells increase through carbohydrate degradation. == See also == Wine fault Maderisation == References ==
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{
"page_id": 55840419,
"source": null,
"title": "Browning in red wine"
}
|
The metallicity distribution function is an important concept in stellar and galactic evolution. It is a curve of what proportion of stars have a particular metallicity ([Fe/H], the relative abundance of iron and hydrogen) of a population of stars such as in a cluster or galaxy. MDFs are used to test different theories of galactic evolution. Much of the iron in a star will have come from earlier type Ia supernovae. Other [alpha] metals can be produced in core collapse supernovae. == References == == Further reading == McWilliam, Andrew. (2003). Abundance ratios and galactic chemical evolution. Annual Review of Astronomy and Astrophysics. 35. 503–556. doi:10.1146/annurev.astro.35.1.503.
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{
"page_id": 37555875,
"source": null,
"title": "Metallicity distribution function"
}
|
Oxamic acid is an organic compound with the formula NH2C(O)COOH. It is a white, water-soluble solid. It is the monoamide of oxalic acid. Oxamic acid inhibits lactate dehydrogenase A. The active site of lactate dehydrogenase (LDH) is closed off once oxamic acid attaches to the LDH-NADH complex, effectively inhibiting it. Oxamic acid also has applications in polymer chemistry. It increases the water solubility of certain polymers, including polyester, epoxide, and acrylic upon binding with them. == See also == Oxamate, the conjugate base of oxamic acid == References ==
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{
"page_id": 52432551,
"source": null,
"title": "Oxamic acid"
}
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Electrochemical energy conversion is a field of energy technology concerned with electrochemical methods of energy conversion including fuel cells and photoelectrochemical. This field of technology also includes electrical storage devices like batteries and supercapacitors. It is increasingly important in context of automotive propulsion systems. There has been the creation of more powerful, longer running batteries allowing longer run times for electric vehicles. These systems would include the energy conversion fuel cells and photoelectrochemical mentioned above. == See also == Bioelectrochemical reactor Chemotronics Electrochemical cell Electrochemical engineering Electrochemical reduction of carbon dioxide Electrofuels Electrohydrogenesis Electromethanogenesis Enzymatic biofuel cell Photoelectrochemical cell Photoelectrochemical reduction of CO2 == Notes == == External links == International Journal of Energy Research MSAL Archived 2016-03-04 at the Wayback Machine NIST scientific journal article Georgia tech
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{
"page_id": 31002280,
"source": null,
"title": "Electrochemical energy conversion"
}
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Retene, methyl isopropyl phenanthrene or 1-methyl-7-isopropyl phenanthrene, C18H18, is a polycyclic aromatic hydrocarbon present in the coal tar fraction, boiling above 360 °C. It occurs naturally in the tars obtained by the distillation of resinous woods. It crystallizes in large plates, which melt at 98.5 °C and boil at 390 °C. It is readily soluble in warm ether and in hot glacial acetic acid. Sodium and boiling amyl alcohol reduce it to a tetrahydroretene, but if it heated with phosphorus and hydriodic acid to 260 °C, a dodecahydride is formed. Chromic acid oxidizes it to retene quinone, phthalic acid and acetic acid. It forms a picrate that melts at 123-124 °C. Retene is derived by degradation of specific diterpenoids biologically produced by conifer trees. The presence of traces of retene in the air is an indicator of forest fires; it is a major product of pyrolysis of conifer trees. It is also present in effluents from wood pulp and paper mills. Retene, together with cadalene, simonellite and ip-iHMN, is a biomarker of vascular plants, which makes it useful for paleobotanic analysis of rock sediments. The ratio of retene/cadalene in sediments can reveal the ratio of the genus Pinaceae in the biosphere. == Health effects == A recent study has shown retene, which is a component of the Amazonian organic PM10, is cytotoxic to human lung cells. == References == This article incorporates text from a publication now in the public domain: Chisholm, Hugh, ed. (1911). "Retene". Encyclopædia Britannica. Vol. 23 (11th ed.). Cambridge University Press. p. 202.
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{
"page_id": 2952874,
"source": null,
"title": "Retene"
}
|
Hexanone may refer to the following ketones containing six carbon atoms: 2-Hexanone (Methyl butyl ketone, MBK) 4-Methyl-2-pentanone (Methyl isobutyl ketone, MIBK) 3-Methyl-2-pentanone (Methyl sec-butyl ketone) 3,3-Dimethyl-2-butanone (Methyl tert-butyl ketone, Pinacolone) 3-Hexanone (Ethyl propyl ketone) 2-Methyl-3-pentanone (Ethyl isopropyl ketone) == See also == Cyclohexanone
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{
"page_id": 19861162,
"source": null,
"title": "Hexanone"
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|
The Chemical History of a Candle was the title of a series of six lectures on the chemistry and physics of flames given by Michael Faraday at the Royal Institution in 1848, as part of the series of Christmas lectures for young people founded by Faraday in 1825 and still given there every year. The lectures described the different zones of combustion in the candle flame and the presence of carbon particles in the luminescent zone. Demonstrations included the production and examination of the properties of hydrogen, oxygen, nitrogen and carbon dioxide gases. An electrolysis cell is demonstrated, first in the electroplating of platinum conductors by dissolved copper, then the production of hydrogen and oxygen gases and their recombination to form water. The properties of water itself are studied, including its expansion while freezing (iron vessels are burst by this expansion), and the relative volume of steam produced when water is vaporized. Techniques for weighing gases on a balance are demonstrated. Atmospheric pressure is described, and its effects are demonstrated. Faraday emphasizes that several of the demonstrations and experiments performed in the lectures may be performed by children "at home" and makes several comments regarding proper attention to safety. The lectures were first printed as a book in 1861. In 2016, Bill Hammack published a video series of lectures supplemented by commentary and a companion book. Faraday's ideas are still used as the basis for open teaching about energy in modern primary and secondary schools == Contents of the six lectures == Lecture 1: A Candle: The Flame - Its Sources - Structure - Mobility - Brightness Lecture 2: Brightness of the Flame - Air necessary for Combustion - Production of Water Lecture 3: Products: Water from the Combustion - Nature of Water - A Compound - Hydrogen Lecture 4:
|
{
"page_id": 2690730,
"source": null,
"title": "The Chemical History of a Candle"
}
|
Hydrogen in the Candle - Burns into Water - The Other Part of Water - Oxygen Lecture 5: Oxygen present in the Air - Nature of the Atmosphere - Its Properties - Other Products from the Candle - Carbonic Acid - Its Properties Lecture 6: Carbon or Charcoal - Coal Gas Respiration and its Analogy to the Burning of a Candle - Conclusion == Reception == Intended for young beginners, for whom it is well adapted, as an introduction to the study of chemistry. According to Frank Wilczek: It is a wonderful laying-bare of surprising facts and intricate structure in a (superficially) familiar process — the burning of a candle. I think it exhibits a marvellously creative mind at work on its home ground, poking into details and following peculiarities to their root with carefully crafted experiments. According to Bill Griffith, F.R.S.C., of Imperial College London: Faraday uses the candle as a symbol to talk about the nature of combustion — how the oxygen from air is needed, how water and CO2 are produced and the hidden role of hydrogen. The text is lyrical and beautifully expressed, communicating his obvious enthusiasm, authority and sense of excitement. There were many accompanying demonstrations, often involving explosions and bright lights. Endearingly, Faraday talks about himself and the audience as ’we philosophers’ and, on one occasion, as ’we juveniles’. The book was of inspiration for the Nobel-prize winner Akira Yoshino when he was a child. Austrian philosopher Ludwig Wittgenstein admired the book and mentions it in Philosophical Investigations. Wittgenstein considered the book as an example of popular science done well, in contrast to other examples of popular science which "pander to people's curiosity to be titillated by the wonders of science without having to do any of the really hard work involved in understanding
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{
"page_id": 2690730,
"source": null,
"title": "The Chemical History of a Candle"
}
|
what science is about", such as the works of James Jeans. Faraday in contrast explains the difficulties and details of experiments. == References == == External links == Faraday, Michael (1861). W. Crookes (ed.). A Course of Six Lectures on the Chemical History of a Candle. Griffin, Bohn & Co. Full text of The Chemical History Of A Candle from Project Gutenberg Walker, Mark; Gröger, Martin; Schlüter, Kirsten; Mosler, Bernd (1 January 2008). "A Bright Spark: Open Teaching of Science Using Faraday's Lectures on Candles". Journal of Chemical Education. 85 (1): 59–62. doi:10.1021/ed085p59. Retrieved 5 November 2020. Faraday, Michael (1861). W. Crookes (ed.). A Course of Six Lectures on the Chemical History of a Candle. Griffin, Bohn & Co. ISBN 1-4255-1974-1. {{cite book}}: ISBN / Date incompatibility (help) Full text of The Chemical History Of A Candle from Internet Archive, with illustrations. Pattison, Darcy and Michael Faraday (2016). Burn: Michael Faraday's Candle. Mims House. Retrieved 5 November 2020. Picture book adaptation of Faraday's lecture. The Chemical History of a Candle public domain audiobook at LibriVox The Chemical History of a Candle playlist on YouTube
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{
"page_id": 2690730,
"source": null,
"title": "The Chemical History of a Candle"
}
|
ABALDH may refer to: Aminobutyraldehyde dehydrogenase, an enzyme 1-pyrroline dehydrogenase, an enzyme
|
{
"page_id": 21302955,
"source": null,
"title": "ABALDH"
}
|
Boron is a chemical element; it has symbol B and atomic number 5. In its crystalline form it is a brittle, dark, lustrous metalloid; in its amorphous form it is a brown powder. As the lightest element of the boron group it has three valence electrons for forming covalent bonds, resulting in many compounds such as boric acid, the mineral sodium borate, and the ultra-hard crystals of boron carbide and boron nitride. Boron is synthesized entirely by cosmic ray spallation and supernovas and not by stellar nucleosynthesis, so it is a low-abundance element in the Solar System and in the Earth's crust. It constitutes about 0.001 percent by weight of Earth's crust. It is concentrated on Earth by the water-solubility of its more common naturally occurring compounds, the borate minerals. These are mined industrially as evaporites, such as borax and kernite. The largest known deposits are in Turkey, the largest producer of boron minerals. Elemental boron is found in small amounts in meteoroids, but chemically uncombined boron is not otherwise found naturally on Earth. Several allotropes exist: amorphous boron is a brown powder; crystalline boron is silvery to black, extremely hard (9.3 on the Mohs scale), and a poor electrical conductor at room temperature (1.5 × 10−6 Ω−1 cm−1 room temperature electrical conductivity). The primary use of the element itself is as boron filaments with applications similar to carbon fibers in some high-strength materials. Boron is primarily used in chemical compounds. About half of all production consumed globally is an additive in fiberglass for insulation and structural materials. The next leading use is in polymers and ceramics in high-strength, lightweight structural and heat-resistant materials. Borosilicate glass is desired for its greater strength and thermal shock resistance than ordinary soda lime glass. As sodium perborate, it is used as a bleach.
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"page_id": 3755,
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"title": "Boron"
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A small amount is used as a dopant in semiconductors, and reagent intermediates in the synthesis of organic fine chemicals. A few boron-containing organic pharmaceuticals are used or are in study. Natural boron is composed of two stable isotopes, one of which (boron-10) has a number of uses as a neutron-capturing agent. Borates have low toxicity in mammals (similar to table salt) but are more toxic to arthropods and are occasionally used as insecticides. Boron-containing organic antibiotics are known. Although only traces are required, it is an essential plant nutrient. == History == The word boron was coined from borax, the mineral from which it was isolated, by analogy with carbon, which boron resembles chemically. Borax in its mineral form (then known as tincal) first saw use as a glaze, beginning in China circa 300 AD. Some crude borax traveled westward, and was apparently mentioned by the alchemist Jabir ibn Hayyan around 700 AD. Marco Polo brought some glazes back to Italy in the 13th century. Georgius Agricola, in around 1600, reported the use of borax as a flux in metallurgy. In 1777, boric acid was recognized in the hot springs (soffioni) near Florence, Italy, at which point it became known as sal sedativum, with ostensible medical benefits. The mineral was named sassolite, after Sasso Pisano in Italy. Sasso was the main source of European borax from 1827 to 1872, when American sources replaced it. Boron compounds were rarely used until the late 1800s when Francis Marion Smith's Pacific Coast Borax Company first popularized and produced them in volume at low cost. Boron was not recognized as an element until it was isolated by Sir Humphry Davy and by Joseph Louis Gay-Lussac and Louis Jacques Thénard. In 1808 Davy observed that electric current sent through a solution of borates produced
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{
"page_id": 3755,
"source": null,
"title": "Boron"
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a brown precipitate on one of the electrodes. In his subsequent experiments, he used potassium to reduce boric acid instead of electrolysis. He produced enough boron to confirm a new element and named it boracium. Gay-Lussac and Thénard used iron to reduce boric acid at high temperatures. By oxidizing boron with air, they showed that boric acid is its oxidation product. Jöns Jacob Berzelius identified it as an element in 1824. Pure boron was arguably first produced by the American chemist Ezekiel Weintraub in 1909. == Characteristics of the element == === Isotopes === Boron has two naturally occurring and stable isotopes, 11B (80.1%) and 10B (19.9%). The mass difference results in a wide range of δ11B values, which are defined as a fractional difference between the 11B and 10B and traditionally expressed in parts per thousand, in natural waters ranging from −16 to +59. There are 13 known isotopes of boron; the shortest-lived isotope is 7B which decays through proton emission and alpha decay with a half-life of 3.5×10−22 s. Isotopic fractionation of boron is controlled by the exchange reactions of the boron species B(OH)3 and [B(OH)4]−. Boron isotopes are also fractionated during mineral crystallization, during H2O phase changes in hydrothermal systems, and during hydrothermal alteration of rock. The latter effect results in preferential removal of the [10B(OH)4]− ion onto clays. It results in solutions enriched in 11B(OH)3 and therefore may be responsible for the large 11B enrichment in seawater relative to both oceanic crust and continental crust; this difference may act as an isotopic signature. The exotic 17B exhibits a nuclear halo, i.e. its radius is appreciably larger than that predicted by the liquid drop model. ==== NMR spectroscopy ==== Both 10B and 11B possess nuclear spin. The nuclear spin of 10B is 3 and that of 11B
|
{
"page_id": 3755,
"source": null,
"title": "Boron"
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|
is 3/2. These isotopes are, therefore, of use in nuclear magnetic resonance spectroscopy; and spectrometers specially adapted to detecting the boron-11 nuclei are available commercially. The 10B and 11B nuclei also cause splitting in the resonances of attached nuclei. === Allotropes === Boron forms four major allotropes: α-rhombohedral and β-rhombohedral (α-R and β-R), γ-orthorhombic (γ) and β-tetragonal (β-T). All four phases are stable at ambient conditions, and β-rhombohedral is the most common and stable. An α-tetragonal phase also exists (α-T), but is very difficult to produce without significant contamination. Most of the phases are based on B12 icosahedra, but the γ phase can be described as a rocksalt-type arrangement of the icosahedra and B2 atomic pairs. It can be produced by compressing other boron phases to 12–20 GPa and heating to 1500–1800 °C; it remains stable after releasing the temperature and pressure. The β-T phase is produced at similar pressures, but higher temperatures of 1800–2200 °C. The α-T and β-T phases might coexist at ambient conditions, with the β-T phase being the more stable. Compressing boron above 160 GPa produces a boron phase with an as yet unknown structure, and this phase is a superconductor at temperatures below 6–12 K. === Atomic structure === Atomic boron is the lightest element having an electron in a p-orbital in its ground state. Its first three ionization energies are higher than those for heavier group III elements, reflecting its electropositive character. == Chemistry of the element == === Preparation === Elemental boron is rare and poorly studied because the pure material is extremely difficult to prepare. Most studies of "boron" involve samples that contain small amounts of carbon. Very pure boron is produced with difficulty because of contamination by carbon or other elements that resist removal. Some early routes to elemental boron involved
|
{
"page_id": 3755,
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"title": "Boron"
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|
the reduction of boric oxide with metals such as magnesium or aluminium. However, the product was often contaminated with borides of those metals. Pure boron can be prepared by reducing volatile boron halides with hydrogen at high temperatures. Ultrapure boron for use in the semiconductor industry is produced by the decomposition of diborane at high temperatures and then further purified by the zone melting or Czochralski processes. === Reactions of the element === Crystalline boron is a hard, black material with a melting point of above 2000 °C. Crystalline boron is chemically inert and resistant to attack by boiling hydrofluoric or hydrochloric acid. When finely divided, it is attacked slowly by hot concentrated hydrogen peroxide, hot concentrated nitric acid, hot sulfuric acid or hot mixture of sulfuric and chromic acids. Since elemental boron is very rare, its chemical reactions are of little significance practically speaking. The elemental form is not typically used as a precursor to compounds. Instead, the extensive inventory of boron compounds are produced from borates. When exposed to air, under normal conditions, a protective oxide or hydroxide layer forms on the surface of boron, which prevents further corrosion. The rate of oxidation of boron depends on the crystallinity, particle size, purity and temperature. At higher temperatures boron burns to form boron trioxide: 4 B + 3 O2 → 2 B2O3 == Chemical compounds == === General trends === In some ways, boron is comparable to carbon in its capability to form stable covalently bonded molecular networks (even nominally disordered (amorphous) boron contains boron icosahedra, which are bonded randomly to each other without long-range order.). In terms of chemical behavior, boron compounds resembles silicon. Aluminium, the heavier congener of boron, does not behave analogously to boron: it is far more electropositive, it is larger, and it tends not
|
{
"page_id": 3755,
"source": null,
"title": "Boron"
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to form homoatomic Al-Al bonds. In the most familiar compounds, boron has the formal oxidation state III. These include the common oxides, sulfides, nitrides, and halides, as well as organic derivatives Boron compounds often violate the octet rule. === Halides === Boron forms the complete series of trihalides, i.e. BX3 (X = F, Cl, Br, I). The trifluoride is produced by treating borate salts with hydrogen fluoride, while the trichloride is produced by carbothermic reduction of boron oxides in the presence of chlorine gas: B2O3 + 3 C + 6 Cl2 → 2 BCl3 + 3 CO The trihalides adopt a planar trigonal structures, in contrast to the behavior of aluminium trihalides. All charge-neutral boron halides violate the octet rule, hence they typically are Lewis acidic. For example, boron trifluoride (BF3) combines eagerly with fluoride sources to give the tetrafluoroborate anion, BF4−. Boron trifluoride is used in the petrochemical industry as a catalyst. The halides react with water to form boric acid. Other boron halides include those with B-B bonding, such as B2F4 and B4Cl4. === Oxide derivatives === Boron-containing minerals exclusively exist as oxides of B(III), often associated with other elements. More than one hundred borate minerals are known. These minerals resemble silicates in some respect, although it is often found not only in a tetrahedral coordination with oxygen, but also in a trigonal planar configuration. The borates can be subdivided into two classes, anhydrous and the far more common hydrates. The hydrates contain B-OH groups and sometimes water of crystallization. A typical motif is exemplified by the tetraborate anions of the common mineral borax. The formal negative charge of the tetrahedral borate center is balanced by sodium (Na+). Some idea of the complexity of these materials is provided by the inventory of zinc borates, which are common wood
|
{
"page_id": 3755,
"source": null,
"title": "Boron"
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|
preservatives and fire retardants: 4ZnO·B2O3·H2O, ZnO·B2O3·1.12H2O, ZnO·B2O3·2H2O, 6ZnO·5B2O3·3H2O, 2ZnO·3B2O3·7H2O, 2ZnO·3B2O3·3H2O, 3ZnO·5B2O3·14H2O, and ZnO·5B2O3·4.5H2O. As illustrated by the preceding examples, borate anions tend to condense by formation of B-O-B bonds. Borosilicates, with B-O-Si, and borophosphates, with B-O-P linkages, are also well represented in both minerals and synthetic compounds. Related to the oxides are the alkoxides and boronic acids with the formula B(OR)3 and R2BOH, respectively. Boron forms a wide variety of such metal-organic compounds, some of which are used in the synthesis of pharmaceuticals. These developments, especially the Suzuki reaction, was recognized with the 2010 Nobel Prize in Chemistry to Akira Suzuki. === Hydrides === Boranes and borohydrides are neutral and anionic compounds of boron and hydrogen, respectively. Sodium borohydride is the progenitor of the boranes. Sodium borohydride is obtained by hydrogenation of trimethylborate: B(OCH3)3 + 4 Na + 2H2 → NaBH4 + 3 NaOCH3 Sodium borohydride is a white, fairly air-stable salt. Sodium borohydride converts to diborane by treatment with boron trifluoride: 3 NaBH4 + 4 BF3 → 2 (BH3)2 + 3 NaBF4 Diborane is the dimer of the elusive parent called borane, BH3. Having a formula akin to ethane's (C2H6), diborane adopts a very different structure, featuring a pair of bridging H atoms. This unusual structure, which was deduced only in the 1940s, was an early indication of the many surprises provided by boron chemistry. Pyrolysis of diborane gives boron hydride clusters, such as pentaborane(9) B5H9 and decaborane B10H14.: 164, 170, 173 A large number of anionic boron hydrides are also known, e.g. [B12H12]2−. In these cluster compounds, boron has a coordination number greater than four. The analysis of the bonding in these polyhedra clusters earned William N. Lipscomb the 1976 Nobel Prize in Chemistry for "studies on the structure of boranes illuminating problems of chemical bonding". Not only
|
{
"page_id": 3755,
"source": null,
"title": "Boron"
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are their structures unusual, many of the boranes are extremely reactive. For example, a widely used procedure for pentaborane states that it will "spontaneously inflame or explode in air". === Organoboron compounds === A large number of organoboron compounds, species with B-C bonds, are known. Many organoboron compounds are produced from hydroboration, the addition of B-H bonds to C=C and C≡C bonds. Diborane is traditionally used for such reactions, as illustrated by the preparation of trioctylborane: B2H6 + 6 H2C=CH(CH2)5CH3 → 2 B((CH2)7CH3)3 This regiochemistry, i.e. the tendency of B to attach to the terminal carbon - is explained by the polarization of the bonds in boranes, which is indicated as Bδ+-Hδ-.: 144, 166 Hydroboration opened the doors for many subsequent reactions, several of which are useful in the synthesis of complex organic compounds. The significance of these methods was recognized by the award of Nobel Prize in Chemistry to H. C. Brown in 1979. Even complicated boron hydrides, such as decaborane undergo hydroboration. Like the volatile boranes, the alkyl boranes ignite spontaneously in air. In the 1950s, several studies examined the use of boranes as energy-increasing "Zip fuel" additives for jet fuel. Triorganoboron(III) compounds are trigonal planar and exhibit weak Lewis acidity. The resulting adducts are tetrahedral. This behavior contrasts with that of triorganoaluminium compounds (see trimethylaluminium), which are tetrahedral with bridging alkyl groups. A compound with the B≡C triple bond was synthesized for the first time in 2025. === Nitrides === The boron-nitrides follow the pattern of avoiding B-B and N-N bonds: only B-N bonding is observed generally. The boron nitrides exhibit structures analogous to various allotropes of carbon, including graphite, diamond, and nanotubes. This similarity reflects the fact that B and N have eight valence electrons as does a pair of carbon atoms. In cubic boron nitride
|
{
"page_id": 3755,
"source": null,
"title": "Boron"
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(tradename Borazon), boron and nitrogen atoms are tetrahedral, just like carbon in diamond. Cubic boron nitride, among other applications, is used as an abrasive, as its hardness is comparable with that of diamond. Hexagonal boron nitride (h-BN) is the BN analogue of graphite, consisting of sheets of alternating B and N atoms. These sheets stack with boron and nitrogen in registry between the sheets. Graphite and h-BN have very different properties, although both are lubricants, as these planes slip past each other easily. However, h-BN is a relatively poor electrical and thermal conductor in the planar directions. Molecular analogues of boron nitrides are represented by borazine, (BH)3(NH)3. === Carbides === Boron carbide is a ceramic material. It is obtained by carbothermal reduction of B2O3in an electric furnace: 2 B2O3 + 7 C → B4C + 6 CO Boron carbide's structure is only approximately reflected in its formula of B4C, and it shows a clear depletion of carbon from this suggested stoichiometric ratio. This is due to its very complex structure. The substance can be seen with empirical formula B12C3 (i.e., with B12 dodecahedra being a motif), but with less carbon, as the suggested C3 units are replaced with C-B-C chains, and some smaller (B6) octahedra are present as well (see the boron carbide article for structural analysis). The repeating polymer plus semi-crystalline structure of boron carbide gives it great structural strength per weight. === Borides === Binary metal-boron compounds, the metal borides, contain only boron and a metal. They are metallic, very hard, with high melting points. TiB2, ZrB2, and HfB2 have melting points above 3000 °C. Some metal borides find specialized applications as hard materials for cutting tools. == Occurrence == Boron is rare in the universe and solar system. The amount of boron formed in the Big Bang
|
{
"page_id": 3755,
"source": null,
"title": "Boron"
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|
is negligible. Boron is not generated in the normal course of stellar nucleosynthesis and is destroyed in stellar interiors. In the high oxygen environment of the Earth's surface, boron is always found fully oxidized to borate. Boron does not appear on Earth in elemental form. Extremely small traces of elemental boron were detected in Lunar regolith. Although boron is a relatively rare element in the Earth's crust, representing only 0.001% of the crust mass, it can be highly concentrated by the action of water, in which many borates are soluble. It is found naturally combined in compounds such as borax and boric acid (sometimes found in volcanic spring waters). About a hundred borate minerals are known. == Production == Economically important sources of boron are the minerals colemanite, rasorite (kernite), ulexite and tincal. Together these constitute 90% of mined boron-containing ore. The largest global borax deposits known, many still untapped, are in Central and Western Turkey, including the provinces of Eskişehir, Kütahya and Balıkesir. Global proven boron mineral mining reserves exceed one billion metric tonnes, against a yearly production of about four million tonnes. Turkey and the United States are the largest producers of boron products. Turkey produces about half of the global yearly demand, through Eti Mine Works (Turkish: Eti Maden İşletmeleri) a Turkish state-owned mining and chemicals company focusing on boron products. It holds a government monopoly on the mining of borate minerals in Turkey, which possesses 72% of the world's known deposits. In 2012, it held a 47% share of production of global borate minerals, ahead of its main competitor, Rio Tinto Group. Almost a quarter (23%) of global boron production comes from the Rio Tinto Borax Mine (also known as the U.S. Borax Boron Mine) 35°2′34.447″N 117°40′45.412″W near Boron, California. === Market trend === The average
|
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"page_id": 3755,
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"title": "Boron"
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|
cost of crystalline elemental boron is US$5/g. Elemental boron is chiefly used in making boron fibers, where it is deposited by chemical vapor deposition on a tungsten core (see below). Boron fibers are used in lightweight composite applications, such as high strength tapes. This use is a very small fraction of total boron use. Boron is introduced into semiconductors as boron compounds, by ion implantation. Estimated global consumption of boron (almost entirely as boron compounds) was about 4 million tonnes of B2O3 in 2012. As compounds such as borax and kernite its cost was US$377/tonne in 2019. Increasing demand for boric acid has led a number of producers to invest in additional capacity. Turkey's state-owned Eti Mine Works opened a new boric acid plant with the production capacity of 100,000 tonnes per year at Emet in 2003. Rio Tinto Group increased the capacity of its boron plant from 260,000 tonnes per year in 2003 to 310,000 tonnes per year by May 2005, with plans to grow this to 366,000 tonnes per year in 2006. Chinese boron producers have been unable to meet rapidly growing demand for high quality borates. This has led to imports of sodium tetraborate (borax) growing by a hundredfold between 2000 and 2005 and boric acid imports increasing by 28% per year over the same period. The rise in global demand has been driven by high growth rates in glass fiber, fiberglass and borosilicate glassware production. A rapid increase in the manufacture of reinforcement-grade boron-containing fiberglass in Asia, has offset the development of boron-free reinforcement-grade fiberglass in Europe and the US. The recent rises in energy prices may lead to greater use of insulation-grade fiberglass, with consequent growth in the boron consumption. Roskill Consulting Group forecasts that world demand for boron will grow by 3.4% per year
|
{
"page_id": 3755,
"source": null,
"title": "Boron"
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to reach 21 million tonnes by 2010. The highest growth in demand is expected to be in Asia where demand could rise by an average 5.7% per year. == Applications == Nearly all boron ore extracted from the Earth is refined as boric acid and sodium tetraborate pentahydrate. In the United States, 70% of the boron is used for the production of glass and ceramics. The major global industrial-scale use of boron compounds (about 46% of end-use) is in production of glass fiber for boron-containing insulating and structural fiberglasses, especially in Asia. Boron is added to the glass as borax pentahydrate or boron oxide, to influence the strength or fluxing qualities of the glass fibers. Another 10% of global boron production is for borosilicate glass as used in high strength glassware. About 15% of global boron is used in boron ceramics, including super-hard materials discussed below. Agriculture consumes 11% of global boron production, and bleaches and detergents about 6%. === Boronated fiberglass === Fiberglasses, a fiber reinforced polymer sometimes contain borosilicate, borax, or boron oxide, and is added to increase the strength of the glass. The highly boronated glasses, E-glass (named for "Electrical" use) are alumino-borosilicate glass. Another common high-boron glasses, C-glass, also has a high boron oxide content, used for glass staple fibers and insulation. D-glass, a borosilicate glass, named for its low dielectric constant. Because of the ubiquitous use of fiberglass in construction and insulation, boron-containing fiberglasses consume over half the global production of boron, and are the single largest commercial boron market. === Borosilicate glass === Borosilicate glass, which is typically 12–15% B2O3, 80% SiO2, and 2% Al2O3, has a low coefficient of thermal expansion, giving it a good resistance to thermal shock. Schott AG's "Duran" and Owens-Corning's trademarked Pyrex are two major brand names for this
|
{
"page_id": 3755,
"source": null,
"title": "Boron"
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glass, used both in laboratory glassware and in consumer cookware and bakeware, chiefly for this resistance. === Elemental boron fiber === Boron fibers (boron filaments) are high-strength, lightweight materials that are used chiefly for advanced aerospace structures as a component of composite materials, as well as limited production consumer and sporting goods such as golf clubs and fishing rods. The fibers can be produced by chemical vapor deposition of boron on a tungsten filament. Boron fibers and sub-millimeter sized crystalline boron springs are produced by laser-assisted chemical vapor deposition. Translation of the focused laser beam allows production of even complex helical structures. Such structures show good mechanical properties (elastic modulus 450 GPa, fracture strain 3.7%, fracture stress 17 GPa) and can be applied as reinforcement of ceramics or in micromechanical systems. === Boron carbide ceramic === Boron carbide's ability to absorb neutrons without forming long-lived radionuclides (especially when doped with extra boron-10) makes the material attractive as an absorbent for neutron radiation arising in nuclear power plants. Nuclear applications of boron carbide include shielding, control rods and shut-down pellets. Within control rods, boron carbide is often powdered, to increase its surface area. === High-hardness and abrasive compounds === Boron carbide and cubic boron nitride powders are widely used as abrasives. Boron nitride is a material isoelectronic to carbon. Similar to carbon, it has both hexagonal (soft graphite-like h-BN) and cubic (hard, diamond-like c-BN) forms. h-BN is used as a high temperature component and lubricant. c-BN, also known under commercial name borazon, is a superior abrasive. Its hardness is only slightly smaller than, but its chemical stability is superior, to that of diamond. Heterodiamond (also called BCN) is another diamond-like boron compound. === Metallurgy === Boron is added to boron steels at the level of a few parts per million to
|
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"page_id": 3755,
"source": null,
"title": "Boron"
}
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increase hardenability. Higher percentages are added to steels used in the nuclear industry due to boron's neutron absorption ability. Boron can also increase the surface hardness of steels and alloys through boriding. Additionally metal borides are used for coating tools through chemical vapor deposition or physical vapor deposition. Implantation of boron ions into metals and alloys, through ion implantation or ion beam deposition, results in a spectacular increase in surface resistance and microhardness. Laser alloying has also been successfully used for the same purpose. These borides are an alternative to diamond coated tools, and their (treated) surfaces have similar properties to those of the bulk boride. For example, rhenium diboride can be produced at ambient pressures, but is rather expensive because of rhenium. The hardness of ReB2 exhibits considerable anisotropy because of its hexagonal layered structure. Its value is comparable to that of tungsten carbide, silicon carbide, titanium diboride or zirconium diboride. Similarly, AlMgB14 + TiB2 composites possess high hardness and wear resistance and are used in either bulk form or as coatings for components exposed to high temperatures and wear loads. === Detergent formulations and bleaching agents === Borax is used in various household laundry and cleaning products. It is also present in some tooth bleaching formulas. Sodium perborate serves as a source of active oxygen in many detergents, laundry detergents, cleaning products, and laundry bleaches. However, despite its name, "Borateem" laundry bleach no longer contains any boron compounds, using sodium percarbonate instead as a bleaching agent. === Insecticides and antifungals === Zinc borates and boric acid, popularized as fire retardants, are widely used as wood preservatives and insecticides. Boric acid is also used as a domestic insecticide. === Semiconductors === Boron is a useful dopant for such semiconductors as silicon, germanium, and silicon carbide. Having one fewer valence
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{
"page_id": 3755,
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"title": "Boron"
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electron than the host atom, it donates a hole resulting in p-type conductivity. Traditional method of introducing boron into semiconductors is via its atomic diffusion at high temperatures. This process uses either solid (B2O3), liquid (BBr3), or gaseous boron sources (B2H6 or BF3). However, after the 1970s, it was mostly replaced by ion implantation, which relies mostly on BF3 as a boron source. Boron trichloride gas is also an important chemical in semiconductor industry, however, not for doping but rather for plasma etching of metals and their oxides. Triethylborane is also injected into vapor deposition reactors as a boron source. Examples are the plasma deposition of boron-containing hard carbon films, silicon nitride–boron nitride films, and for doping of diamond film with boron. === Magnets === Boron is a component of neodymium magnets (Nd2Fe14B), which are among the strongest type of permanent magnet. These magnets are found in a variety of electromechanical and electronic devices, such as magnetic resonance imaging (MRI) medical imaging systems, in compact and relatively small motors and actuators. As examples, computer HDDs (hard disk drives), CD (compact disk) and DVD (digital versatile disk) players rely on neodymium magnet motors to deliver intense rotary power in a remarkably compact package. In mobile phones 'Neo' magnets provide the magnetic field which allows tiny speakers to deliver appreciable audio power. === Shielding and neutron absorber in nuclear reactors === Boron shielding is used as a control for nuclear reactors, taking advantage of its high cross-section for neutron capture. In pressurized water reactors a variable concentration of boronic acid in the cooling water is used as a neutron poison to compensate the variable reactivity of the fuel. When new rods are inserted the concentration of boronic acid is maximal, and is reduced during the lifetime. === Other nonmedical uses === Because
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{
"page_id": 3755,
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of its distinctive green flame, amorphous boron is used in pyrotechnic flares. Some anti-corrosion systems contain borax. Sodium borates are used as a flux for soldering silver and gold and with ammonium chloride for welding ferrous metals. They are also fire retarding additives to plastics and rubber articles. Boric acid (also known as orthoboric acid) H3BO3 is used in the production of textile fiberglass and flat panel displays and in many PVAc- and PVOH-based adhesives. Triethylborane is a substance which ignites the JP-7 fuel of the Pratt & Whitney J58 turbojet/ramjet engines powering the Lockheed SR-71 Blackbird. It was also used to ignite the F-1 Engines on the Saturn V Rocket utilized by NASA's Apollo and Skylab programs from 1967 until 1973. Today SpaceX uses it to ignite the engines on their Falcon 9 rocket. Triethylborane is suitable for this because of its pyrophoric properties, especially the fact that it burns with a very high temperature. Triethylborane is an industrial initiator in radical reactions, where it is effective even at low temperatures. Borates are used as environmentally benign wood preservatives. === Pharmaceutical and biological applications === Boron plays a role in pharmaceutical and biological applications as it is found in various antibiotics produced by bacteria, such as boromycins, aplasmomycins, borophycins, and tartrolons. These antibiotics have shown inhibitory effects on the growth of certain bacteria, fungi, and protozoa. Boron is also being studied for its potential medicinal applications, including its incorporation into biologically active molecules for therapies like boron neutron capture therapy for brain tumors. Some boron-containing biomolecules may act as signaling molecules interacting with cell surfaces, suggesting a role in cellular communication. Boric acid has antiseptic, antifungal, and antiviral properties and, for these reasons, is applied as a water clarifier in swimming pool water treatment. Mild solutions of boric acid
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{
"page_id": 3755,
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"title": "Boron"
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have been used as eye antiseptics. Bortezomib (marketed as Velcade and Cytomib). Boron appears as an active element in the organic pharmaceutical bortezomib, a new class of drug called the proteasome inhibitor, for treating myeloma and one form of lymphoma (it is currently in experimental trials against other types of lymphoma). The boron atom in bortezomib binds the catalytic site of the 26S proteasome with high affinity and specificity. A number of potential boronated pharmaceuticals using boron-10, have been prepared for use in boron neutron capture therapy (BNCT). Some boron compounds show promise in treating arthritis, though none have as yet been generally approved for the purpose. Tavaborole (marketed as Kerydin) is an Aminoacyl tRNA synthetase inhibitor which is used to treat toenail fungus. It gained FDA approval in July 2014. Dioxaborolane chemistry enables radioactive fluoride (18F) labeling of antibodies or red blood cells, which allows for positron emission tomography (PET) imaging of cancer and hemorrhages, respectively. A Human-Derived, Genetic, Positron-emitting and Fluorescent (HD-GPF) reporter system uses a human protein, PSMA and non-immunogenic, and a small molecule that is positron-emitting (boron bound 18F) and fluorescence for dual modality PET and fluorescent imaging of genome modified cells, e.g. cancer, CRISPR/Cas9, or CAR T-cells, in an entire mouse. The dual-modality small molecule targeting PSMA was tested in humans and found the location of primary and metastatic prostate cancer, fluorescence-guided removal of cancer, and detects single cancer cells in tissue margins. == Research == === MgB2 === Magnesium diboride (MgB2) is a superconductor with the transition temperature of 39 K. MgB2 wires are produced with the powder-in-tube process and applied in superconducting magnets. A project at CERN to make MgB2 cables has resulted in superconducting test cables able to carry 20,000 amperes for extremely high current distribution applications, such as the contemplated high
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{
"page_id": 3755,
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"title": "Boron"
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luminosity version of the Large Hadron Collider. === Commercial isotope enrichment === Because of its high neutron cross-section, boron-10 is often used to control fission in nuclear reactors as a neutron-capturing substance. Several industrial-scale enrichment processes have been developed; however, only the fractionated vacuum distillation of the dimethyl ether adduct of boron trifluoride (DME-BF3) and column chromatography of borates are being used. === Radiation-hardened semiconductors === Cosmic radiation will produce secondary neutrons if it hits spacecraft structures. Those neutrons will be captured in 10B, if it is present in the spacecraft's semiconductors, producing a gamma ray, an alpha particle, and a lithium ion. Those resultant decay products may then irradiate nearby semiconductor "chip" structures, causing data loss (bit flipping, or single event upset). In radiation-hardened semiconductor designs, one countermeasure is to use depleted boron, which is greatly enriched in 11B and contains almost no 10B. This is useful because 11B is largely immune to radiation damage. Depleted boron is a byproduct of the nuclear industry (see above). === Proton-boron fusion === 11B is also a candidate as a fuel for aneutronic fusion. When struck by a proton with energy of about 500 keV, it produces three alpha particles and 8.7 MeV of energy. Most other fusion reactions involving hydrogen and helium produce penetrating neutron radiation, which weakens reactor structures and induces long-term radioactivity, thereby endangering operating personnel. The alpha particles from 11B fusion can be turned directly into electric power, and all radiation stops as soon as the reactor is turned off. === Enriched boron (boron-10) === The 10B isotope is useful for capturing thermal neutrons (see neutron cross section#Typical cross sections). The nuclear industry enriches natural boron to nearly pure 10B. The less-valuable by-product, depleted boron, is nearly pure 11B. Enriched boron or 10B is used in both radiation
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{
"page_id": 3755,
"source": null,
"title": "Boron"
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|
shielding and is the primary nuclide used in neutron capture therapy of cancer. In the latter ("boron neutron capture therapy" or BNCT), a compound containing 10B is incorporated into a pharmaceutical which is selectively taken up by a malignant tumor and tissues near it. The patient is then treated with a beam of low energy neutrons at a relatively low neutron radiation dose. The neutrons, however, trigger energetic and short-range secondary alpha particle and lithium-7 heavy ion radiation that are products of the boron-neutron nuclear reaction, and this ion radiation additionally bombards the tumor, especially from inside the tumor cells. In nuclear reactors, 10B is used for reactivity control and in emergency shutdown systems. It can serve either function in the form of borosilicate control rods or as boric acid. In pressurized water reactors, 10B boric acid is added to the reactor coolant after the plant is shut down for refueling. When the plant is started up again, the boric acid is slowly filtered out over many months as fissile material is used up and the fuel becomes less reactive. === Nuclear fusion === Boron has been investigated for possible applications in nuclear fusion research. It is commonly used for conditioning the walls in fusion reactors by depositing boron coatings on plasma-facing components and walls to reduce the release of hydrogen and impurities from the surfaces. It is also being used for the dissipation of energy in the fusion plasma boundary to suppress excessive energy bursts and heat fluxes to the walls. === Neutron capture therapy === In neutron capture therapy (BNCT) for malignant brain tumors, boron is researched to be used for selectively targeting and destroying tumor cells. The goal is to deliver higher concentrations of the non-radioactive boron isotope (10B) to the tumor cells than to the surrounding
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{
"page_id": 3755,
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"title": "Boron"
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normal tissues. When these 10B-containing cells are irradiated with low-energy thermal neutrons, they undergo nuclear capture reactions, releasing high linear energy transfer (LET) particles such as α-particles and lithium-7 nuclei within a limited path length. These high-LET particles can destroy the adjacent tumor cells without causing significant harm to nearby normal cells. Boron acts as a selective agent due to its ability to absorb thermal neutrons and produce short-range physical effects primarily affecting the targeted tissue region. This binary approach allows for precise tumor cell killing while sparing healthy tissues. The effective delivery of boron involves administering boron compounds or carriers capable of accumulating selectively in tumor cells compared to surrounding tissue. BSH and BPA have been used clinically, but research continues to identify more optimal carriers. Accelerator-based neutron sources have also been developed recently as an alternative to reactor-based sources, leading to improved efficiency and enhanced clinical outcomes in BNCT. By employing the properties of boron isotopes and targeted irradiation techniques, BNCT offers a potential approach to treating malignant brain tumors by selectively killing cancer cells while minimizing the damage caused by traditional radiation therapies. BNCT has shown promising results in clinical trials for various other malignancies, including glioblastoma, head and neck cancer, cutaneous melanoma, hepatocellular carcinoma, lung cancer, and extramammary Paget's disease. The treatment involves a nuclear reaction between nonradioactive boron-10 isotope and low-energy thermal or high-energy epithermal neutrons to generate α particles and lithium nuclei that selectively destroy DNA in tumor cells. The primary challenge lies in developing efficient boron agents with higher content and specific targeting properties tailored for BNCT. Integration of tumor-targeting strategies with BNCT could potentially establish it as a practical personalized treatment option for different types of cancers. Ongoing research explores new boron compounds, optimization strategies, theranostic agents, and radiobiological advances to overcome
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{
"page_id": 3755,
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limitations and cost-effectively improve patient outcomes. == Biological role == Boron is an essential plant nutrient, required primarily for maintaining the integrity of cell walls. However, high soil concentrations of greater than 1.0 ppm lead to marginal and tip necrosis in leaves as well as poor overall growth performance. Levels as low as 0.8 ppm produce these same symptoms in plants that are particularly sensitive to boron in the soil. Nearly all plants, even those somewhat tolerant of soil boron, will show at least some symptoms of boron toxicity when soil boron content is greater than 1.8 ppm. When this content exceeds 2.0 ppm, few plants will perform well and some may not survive. Some boron-containing antibiotics exist in nature. The first one found was boromycin, isolated from streptomyces in the 1960s. Others are tartrolons, a group of antibiotics discovered in the 1990s from culture broth of the myxobacterium Sorangium cellulosum. In 2013, chemist and synthetic biologist Steve Benner suggested that the conditions on Mars three billion years ago were much more favorable to the stability of RNA and formation of oxygen-containing boron and molybdenum catalysts found in life. According to Benner's theory, primitive life, which is widely believed to have originated from RNA, first formed on Mars before migrating to Earth. === In human health === It is thought that boron plays several essential roles in animals, including humans, but the exact physiological role is poorly understood. Boron deficiency has only been clearly established in livestock; in humans, boron deficiency may affect bone mineral density, though it has been noted that additional research on the effects of bone health is necessary. Boron is not classified as an essential human nutrient because research has not established a clear biological function for it. The U.S. Food and Nutrition Board (FNB) found
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{
"page_id": 3755,
"source": null,
"title": "Boron"
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|
the existing data insufficient to derive a Recommended Dietary Allowance (RDA), Adequate Intake (AI), or Estimated Average Requirement (EAR) for boron and the U.S. Food and Drug Administration (FDA) has not established a daily value for boron for food and dietary supplement labeling purposes. While low boron status can be detrimental to health, probably increasing the risk of osteoporosis, poor immune function, and cognitive decline, high boron levels are associated with cell damage and toxicity. Still, studies suggest that boron may exert beneficial effects on reproduction and development, calcium metabolism, bone formation, brain function, insulin and energy substrate metabolism, immunity, and steroid hormone (including estrogen) and vitamin D function, among other functions. A small human trial published in 1987 reported on postmenopausal women first made boron deficient and then repleted with 3 mg/day. Boron supplementation markedly reduced urinary calcium excretion and elevated the serum concentrations of 17 beta-estradiol and testosterone. Environmental boron appears to be inversely correlated with arthritis. The exact mechanism by which boron exerts its physiological effects is not fully understood, but may involve interactions with adenosine monophosphate (ADP) and S-adenosyl methionine (SAM-e), two compounds involved in important cellular functions. Furthermore, boron appears to inhibit cyclic ADP-ribose, thereby affecting the release of calcium ions from the endoplasmic reticulum and affecting various biological processes. Some studies suggest that boron may reduce levels of inflammatory biomarkers. Congenital endothelial dystrophy type 2, a rare form of corneal dystrophy, is linked to mutations in SLC4A11 gene that encodes a transporter reportedly regulating the intracellular concentration of boron. In humans, boron is usually consumed with food that contains boron, such as fruits, leafy vegetables, and nuts. Foods that are particularly rich in boron include avocados, dried fruits such as raisins, peanuts, pecans, prune juice, grape juice, wine and chocolate powder. According to 2-day
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{
"page_id": 3755,
"source": null,
"title": "Boron"
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food records from the respondents to the Third National Health and Nutrition Examination Survey (NHANES III), adult dietary intake was recorded at 0.9 to 1.4 mg/day. === Health issues and toxicity === Elemental boron, boron oxide, boric acid, borates, and many organoboron compounds are relatively nontoxic to humans and animals (with toxicity similar to that of table salt). The LD50 (dose at which there is 50% mortality) for animals is about 6 g per kg of body weight. Substances with an LD50 above 2 g/kg are considered nontoxic. An intake of 4 g/day of boric acid was reported without incident, but more than this is considered toxic in more than a few doses. Intakes of more than 0.5 grams per day for 50 days cause minor digestive and other problems suggestive of toxicity. Boric acid is more toxic to insects than to mammals, and is routinely used as an insecticide. However, it has been used in neutron capture therapy alongside other boron compounds such as sodium borocaptate and boronophenylalanine with reported low toxicity levels. The boranes (boron hydrogen compounds) and similar gaseous compounds are quite poisonous. As usual, boron is not an element that is intrinsically poisonous, but the toxicity of these compounds depends on structure (for another example of this phenomenon, see phosphine). The boranes are also highly flammable and require special care when handling, some combinations of boranes and other compounds are highly explosive. Sodium borohydride presents a fire hazard owing to its reducing nature and the liberation of hydrogen on contact with acid. Boron halides are corrosive. Boron is necessary for plant growth, but an excess of boron is toxic to plants, and occurs particularly in acidic soil. It presents as a yellowing from the tip inwards of the oldest leaves and black spots in barley leaves,
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{
"page_id": 3755,
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but it can be confused with other stresses such as magnesium deficiency in other plants. == See also == == Notes == == References == == External links == Boron at The Periodic Table of Videos (University of Nottingham) J. B. Calvert: Boron, 2004, private website (archived version)
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{
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The molecular formula C9H8O4 (molar mass: 180.15 g/mol, exact mass: 180.0423 u) may refer to: Acetozone Aspirin 4-Hydroxyphenylpyruvic acid, a natural phenol Dihydroxycinnamic acids: Caffeic acid (3,4-Dihydroxycinnamic acid) Umbellic acid (2,4-dihydroxycinnamic acid) 2,3-Dihydroxycinnamic acid 2,5-Dihydroxycinnamic acid 3,5-Dihydroxycinnamic acid Homophthalic acid Monomethyl phthalate Uvitic acid
|
{
"page_id": 13766319,
"source": null,
"title": "C9H8O4"
}
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Meprin A (EC 3.4.24.18, endopeptidase-2, meprin-a, meprin, N-benzoyl-L-tyrosyl-p-aminobenzoic acid hydrolase, PABA-peptide hydrolase, PPH) is an enzyme that cleaves protein and peptide substrates preferentially on carboxyl side of hydrophobic residues. This metalloprotease can be associated with inflammatory responses. It can be found in the extracellular space where it can also form complex structures by joining its monomers together. Meprin A is a dimer composed of the products transcribed from the following two genes: == References ==
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{
"page_id": 30740145,
"source": null,
"title": "Meprin A"
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Bromine is a chemical element; it has symbol Br and atomic number 35. It is a volatile red-brown liquid at room temperature that evaporates readily to form a similarly coloured vapour. Its properties are intermediate between those of chlorine and iodine. Isolated independently by two chemists, Carl Jacob Löwig (in 1825) and Antoine Jérôme Balard (in 1826), its name was derived from Ancient Greek βρῶμος (bromos) 'stench', referring to its sharp and pungent smell. Elemental bromine is very reactive and thus does not occur as a free element in nature. Instead, it can be isolated from colourless soluble crystalline mineral halide salts analogous to table salt, a property it shares with the other halogens. While it is rather rare in the Earth's crust, the high solubility of the bromide ion (Br−) has caused its accumulation in the oceans. Commercially the element is easily extracted from brine evaporation ponds, mostly in the United States and Israel. The mass of bromine in the oceans is about one three-hundredth that of chlorine. At standard conditions for temperature and pressure it is a liquid; the only other element that is liquid under these conditions is mercury. At high temperatures, organobromine compounds readily dissociate to yield free bromine atoms, a process that stops free radical chemical chain reactions. This effect makes organobromine compounds useful as fire retardants, and more than half the bromine produced worldwide each year is put to this purpose. The same property causes ultraviolet sunlight to dissociate volatile organobromine compounds in the atmosphere to yield free bromine atoms, causing ozone depletion. As a result, many organobromine compounds—such as the pesticide methyl bromide—are no longer used. Bromine compounds are still used in well drilling fluids, in photographic film, and as an intermediate in the manufacture of organic chemicals. Large amounts of bromide salts
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{
"page_id": 3756,
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"title": "Bromine"
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are toxic from the action of soluble bromide ions, causing bromism. However, bromine is beneficial for human eosinophils, and is an essential trace element for collagen development in all animals. Hundreds of known organobromine compounds are generated by terrestrial and marine plants and animals, and some serve important biological roles. As a pharmaceutical, the simple bromide ion (Br−) has inhibitory effects on the central nervous system, and bromide salts were once a major medical sedative, before replacement by shorter-acting drugs. They retain niche uses as antiepileptics. == History == Bromine was discovered independently by two chemists, Carl Jacob Löwig and Antoine Balard, in 1825 and 1826, respectively. Löwig isolated bromine from a mineral water spring from his hometown Bad Kreuznach in 1825. Löwig used a solution of the mineral salt saturated with chlorine and extracted the bromine with diethyl ether. After evaporation of the ether, a brown liquid remained. With this liquid as a sample of his work he applied for a position in the laboratory of Leopold Gmelin in Heidelberg. The publication of the results was delayed and Balard published his results first. Balard found bromine chemicals in the ash of seaweed from the salt marshes of Montpellier. The seaweed was used to produce iodine, but also contained bromine. Balard distilled the bromine from a solution of seaweed ash saturated with chlorine. The properties of the resulting substance were intermediate between those of chlorine and iodine; thus he tried to prove that the substance was iodine monochloride (ICl), but after failing to do so he was sure that he had found a new element and named it muride, derived from the Latin word muria ("brine"). After the French chemists Louis Nicolas Vauquelin, Louis Jacques Thénard, and Joseph-Louis Gay-Lussac approved the experiments of the young pharmacist Balard, the results were
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{
"page_id": 3756,
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"title": "Bromine"
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presented at a lecture of the Académie des Sciences and published in Annales de Chimie et Physique. In his publication, Balard stated that he changed the name from muride to brôme on the proposal of M. Anglada. The name brôme (bromine) derives from the Greek βρῶμος (brômos, "stench"). Other sources claim that the French chemist and physicist Joseph-Louis Gay-Lussac suggested the name brôme for the characteristic smell of the vapors. Bromine was not produced in large quantities until 1858, when the discovery of salt deposits in Stassfurt enabled its production as a by-product of potash. Apart from some minor medical applications, the first commercial use was the daguerreotype. In 1840, bromine was discovered to have some advantages over the previously used iodine vapor to create the light sensitive silver halide layer in daguerreotypy. By 1864, a 25% solution of liquid bromine in .75 molar aqueous potassium bromide was widely used to treat gangrene during the American Civil War, before the publications of Joseph Lister and Pasteur. Potassium bromide and sodium bromide were used as anticonvulsants and sedatives in the late 19th and early 20th centuries, but were gradually superseded by chloral hydrate and then by the barbiturates. In the early years of the First World War, bromine compounds such as xylyl bromide were used as poison gas. == Properties == Bromine is the third halogen, being a nonmetal in group 17 of the periodic table. Its properties are thus similar to those of fluorine, chlorine, and iodine, and tend to be intermediate between those of chlorine and iodine, the two neighbouring halogens. Bromine has the electron configuration [Ar]4s23d104p5, with the seven electrons in the fourth and outermost shell acting as its valence electrons. Like all halogens, it is thus one electron short of a full octet, and is hence a
|
{
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strong oxidising agent, reacting with many elements in order to complete its outer shell. Corresponding to periodic trends, it is intermediate in electronegativity between chlorine and iodine (F: 3.98, Cl: 3.16, Br: 2.96, I: 2.66), and is less reactive than chlorine and more reactive than iodine. It is also a weaker oxidising agent than chlorine, but a stronger one than iodine. Conversely, the bromide ion is a weaker reducing agent than iodide, but a stronger one than chloride. These similarities led to chlorine, bromine, and iodine together being classified as one of the original triads of Johann Wolfgang Döbereiner, whose work foreshadowed the periodic law for chemical elements. It is intermediate in atomic radius between chlorine and iodine, and this leads to many of its atomic properties being similarly intermediate in value between chlorine and iodine, such as first ionisation energy, electron affinity, enthalpy of dissociation of the X2 molecule (X = Cl, Br, I), ionic radius, and X–X bond length. The volatility of bromine accentuates its very penetrating, choking, and unpleasant odour. All four stable halogens experience intermolecular van der Waals forces of attraction, and their strength increases together with the number of electrons among all homonuclear diatomic halogen molecules. Thus, the melting and boiling points of bromine are intermediate between those of chlorine and iodine. As a result of the increasing molecular weight of the halogens down the group, the density and heats of fusion and vaporisation of bromine are again intermediate between those of chlorine and iodine, although all their heats of vaporisation are fairly low (leading to high volatility) thanks to their diatomic molecular structure. The halogens darken in colour as the group is descended: fluorine is a very pale yellow gas, chlorine is greenish-yellow, and bromine is a reddish-brown volatile liquid that freezes at −7.2
|
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°C and boils at 58.8 °C. (Iodine is a shiny black solid.) This trend occurs because the wavelengths of visible light absorbed by the halogens increase down the group. Specifically, the colour of a halogen, such as bromine, results from the electron transition between the highest occupied antibonding πg molecular orbital and the lowest vacant antibonding σu molecular orbital. The colour fades at low temperatures so that solid bromine at −195 °C is pale yellow. Liquid bromine is infrared-transparent. Like solid chlorine and iodine, solid bromine crystallises in the orthorhombic crystal system, in a layered arrangement of Br2 molecules. The Br–Br distance is 227 pm (close to the gaseous Br–Br distance of 228 pm) and the Br···Br distance between molecules is 331 pm within a layer and 399 pm between layers (compare the van der Waals radius of bromine, 195 pm). This structure means that bromine is a very poor conductor of electricity, with a conductivity of around 5 × 10−13 Ω−1 cm−1 just below the melting point, although this is higher than the essentially undetectable conductivity of chlorine. At a pressure of 55 GPa (roughly 540,000 times atmospheric pressure) bromine undergoes an insulator-to-metal transition. At 75 GPa it changes to a face-centered orthorhombic structure. At 100 GPa it changes to a body centered orthorhombic monatomic form. === Isotopes === Bromine has two stable isotopes, 79Br and 81Br. These are its only two natural isotopes, with 79Br making up 51% of natural bromine and 81Br making up the remaining 49%. Both have nuclear spin 3/2− and thus may be used for nuclear magnetic resonance, although 81Br is more favourable. The relatively 1:1 distribution of the two isotopes in nature is helpful in identification of bromine containing compounds using mass spectroscopy. Other bromine isotopes are all radioactive, with half-lives too short
|
{
"page_id": 3756,
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"title": "Bromine"
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to occur in nature. Of these, the most important are 80Br (t1/2 = 17.7 min), 80mBr (t1/2 = 4.421 h), and 82Br (t1/2 = 35.28 h), which may be produced from the neutron activation of natural bromine. The most stable bromine radioisotope is 77Br (t1/2 = 57.04 h). The primary decay mode of isotopes lighter than 79Br is electron capture to isotopes of selenium; that of isotopes heavier than 81Br is beta decay to isotopes of krypton; and 80Br may decay by either mode to stable 80Se or 80Kr. Br isotopes from 87Br and heavier undergo beta decay with neutron emission and are of practical importance because they are fission products. == Chemistry and compounds == Bromine is intermediate in reactivity between chlorine and iodine, and is one of the most reactive elements. Bond energies to bromine tend to be lower than those to chlorine but higher than those to iodine, and bromine is a weaker oxidising agent than chlorine but a stronger one than iodine. This can be seen from the standard electrode potentials of the X2/X− couples (F, +2.866 V; Cl, +1.395 V; Br, +1.087 V; I, +0.615 V; At, approximately +0.3 V). Bromination often leads to higher oxidation states than iodination but lower or equal oxidation states to chlorination. Bromine tends to react with compounds including M–M, M–H, or M–C bonds to form M–Br bonds. === Hydrogen bromide === The simplest compound of bromine is hydrogen bromide, HBr. It is mainly used in the production of inorganic bromides and alkyl bromides, and as a catalyst for many reactions in organic chemistry. Industrially, it is mainly produced by the reaction of hydrogen gas with bromine gas at 200–400 °C with a platinum catalyst. However, reduction of bromine with red phosphorus is a more practical way to produce hydrogen
|
{
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bromide in the laboratory: 2 P + 6 H2O + 3 Br2 → 6 HBr + 2 H3PO3 H3PO3 + H2O + Br2 → 2 HBr + H3PO4 At room temperature, hydrogen bromide is a colourless gas, like all the hydrogen halides apart from hydrogen fluoride, since hydrogen cannot form strong hydrogen bonds to the large and only mildly electronegative bromine atom; however, weak hydrogen bonding is present in solid crystalline hydrogen bromide at low temperatures, similar to the hydrogen fluoride structure, before disorder begins to prevail as the temperature is raised. Aqueous hydrogen bromide is known as hydrobromic acid, which is a strong acid (pKa = −9) because the hydrogen bonds to bromine are too weak to inhibit dissociation. The HBr/H2O system also involves many hydrates HBr·nH2O for n = 1, 2, 3, 4, and 6, which are essentially salts of bromine anions and hydronium cations. Hydrobromic acid forms an azeotrope with boiling point 124.3 °C at 47.63 g HBr per 100 g solution; thus hydrobromic acid cannot be concentrated beyond this point by distillation. Unlike hydrogen fluoride, anhydrous liquid hydrogen bromide is difficult to work with as a solvent, because its boiling point is low, it has a small liquid range, its dielectric constant is low and it does not dissociate appreciably into H2Br+ and HBr−2 ions – the latter, in any case, are much less stable than the bifluoride ions (HF−2) due to the very weak hydrogen bonding between hydrogen and bromine, though its salts with very large and weakly polarising cations such as Cs+ and NR+4 (R = Me, Et, Bun) may still be isolated. Anhydrous hydrogen bromide is a poor solvent, only able to dissolve small molecular compounds such as nitrosyl chloride and phenol, or salts with very low lattice energies such as tetraalkylammonium halides.
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{
"page_id": 3756,
"source": null,
"title": "Bromine"
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=== Other binary bromides === Nearly all elements in the periodic table form binary bromides. The exceptions are decidedly in the minority and stem in each case from one of three causes: extreme inertness and reluctance to participate in chemical reactions (the noble gases, with the exception of xenon in the very unstable XeBr2); extreme nuclear instability hampering chemical investigation before decay and transmutation (many of the heaviest elements beyond bismuth); and having an electronegativity higher than bromine's (oxygen, nitrogen, fluorine, and chlorine), so that the resultant binary compounds are formally not bromides but rather oxides, nitrides, fluorides, or chlorides of bromine. (Nonetheless, nitrogen tribromide is named as a bromide as it is analogous to the other nitrogen trihalides.) Bromination of metals with Br2 tends to yield lower oxidation states than chlorination with Cl2 when a variety of oxidation states is available. Bromides can be made by reaction of an element or its oxide, hydroxide, or carbonate with hydrobromic acid, and then dehydrated by mildly high temperatures combined with either low pressure or anhydrous hydrogen bromide gas. These methods work best when the bromide product is stable to hydrolysis; otherwise, the possibilities include high-temperature oxidative bromination of the element with bromine or hydrogen bromide, high-temperature bromination of a metal oxide or other halide by bromine, a volatile metal bromide, carbon tetrabromide, or an organic bromide. For example, niobium(V) oxide reacts with carbon tetrabromide at 370 °C to form niobium(V) bromide. Another method is halogen exchange in the presence of excess "halogenating reagent", for example: FeCl3 + BBr3 (excess) → FeBr3 + BCl3 When a lower bromide is wanted, either a higher halide may be reduced using hydrogen or a metal as a reducing agent, or thermal decomposition or disproportionation may be used, as follows: 3 WBr5 + Al thermal gradient→475
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"page_id": 3756,
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°C → 240 °C 3 WBr4 + AlBr3 EuBr3 + 1/2 H2 → EuBr2 + HBr 2 TaBr4 500 °C→ TaBr3 + TaBr5 Most metal bromides with the metal in low oxidation states (+1 to +3) are ionic. Nonmetals tend to form covalent molecular bromides, as do metals in high oxidation states from +3 and above. Both ionic and covalent bromides are known for metals in oxidation state +3 (e.g. scandium bromide is mostly ionic, but aluminium bromide is not). Silver bromide is very insoluble in water and is thus often used as a qualitative test for bromine. === Bromine halides === The halogens form many binary, diamagnetic interhalogen compounds with stoichiometries XY, XY3, XY5, and XY7 (where X is heavier than Y), and bromine is no exception. Bromine forms a monofluoride and monochloride, as well as a trifluoride and pentafluoride. Some cationic and anionic derivatives are also characterised, such as BrF−2, BrCl−2, BrF+2, BrF+4, and BrF+6. Apart from these, some pseudohalides are also known, such as cyanogen bromide (BrCN), bromine thiocyanate (BrSCN), and bromine azide (BrN3). The pale-brown bromine monofluoride (BrF) is unstable at room temperature, disproportionating quickly and irreversibly into bromine, bromine trifluoride, and bromine pentafluoride. It thus cannot be obtained pure. It may be synthesised by the direct reaction of the elements, or by the comproportionation of bromine and bromine trifluoride at high temperatures. Bromine monochloride (BrCl), a red-brown gas, quite readily dissociates reversibly into bromine and chlorine at room temperature and thus also cannot be obtained pure, though it can be made by the reversible direct reaction of its elements in the gas phase or in carbon tetrachloride. Bromine monofluoride in ethanol readily leads to the monobromination of the aromatic compounds PhX (para-bromination occurs for X = Me, But, OMe, Br; meta-bromination occurs for the deactivating
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{
"page_id": 3756,
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X = –CO2Et, –CHO, –NO2); this is due to heterolytic fission of the Br–F bond, leading to rapid electrophilic bromination by Br+. At room temperature, bromine trifluoride (BrF3) is a straw-coloured liquid. It may be formed by directly fluorinating bromine at room temperature and is purified through distillation. It reacts violently with water and explodes on contact with flammable materials, but is a less powerful fluorinating reagent than chlorine trifluoride. It reacts vigorously with boron, carbon, silicon, arsenic, antimony, iodine, and sulfur to give fluorides, and will also convert most metals and many metal compounds to fluorides; as such, it is used to oxidise uranium to uranium hexafluoride in the nuclear power industry. Refractory oxides tend to be only partially fluorinated, but here the derivatives KBrF4 and BrF2SbF6 remain reactive. Bromine trifluoride is a useful nonaqueous ionising solvent, since it readily dissociates to form BrF+2 and BrF−4 and thus conducts electricity. Bromine pentafluoride (BrF5) was first synthesised in 1930. It is produced on a large scale by direct reaction of bromine with excess fluorine at temperatures higher than 150 °C, and on a small scale by the fluorination of potassium bromide at 25 °C. It also reacts violently with water and is a very strong fluorinating agent, although chlorine trifluoride is still stronger. === Polybromine compounds === Although dibromine is a strong oxidising agent with a high first ionisation energy, very strong oxidisers such as peroxydisulfuryl fluoride (S2O6F2) can oxidise it to form the cherry-red Br+2 cation. A few other bromine cations are known, namely the brown Br+3 and dark brown Br+5. The tribromide anion, Br−3, has also been characterised; it is analogous to triiodide. === Bromine oxides and oxoacids === Bromine oxides are not as well-characterised as chlorine oxides or iodine oxides, as they are all fairly unstable: it
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{
"page_id": 3756,
"source": null,
"title": "Bromine"
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was once thought that they could not exist at all. Dibromine monoxide is a dark-brown solid which, while reasonably stable at −60 °C, decomposes at its melting point of −17.5 °C; it is useful in bromination reactions and may be made from the low-temperature decomposition of bromine dioxide in a vacuum. It oxidises iodine to iodine pentoxide and benzene to 1,4-benzoquinone; in alkaline solutions, it gives the hypobromite anion. So-called "bromine dioxide", a pale yellow crystalline solid, may be better formulated as bromine perbromate, BrOBrO3. It is thermally unstable above −40 °C, violently decomposing to its elements at 0 °C. Dibromine trioxide, syn-BrOBrO2, is also known; it is the anhydride of hypobromous acid and bromic acid. It is an orange crystalline solid which decomposes above −40 °C; if heated too rapidly, it explodes around 0 °C. A few other unstable radical oxides are also known, as are some poorly characterised oxides, such as dibromine pentoxide, tribromine octoxide, and bromine trioxide. The four oxoacids, hypobromous acid (HOBr), bromous acid (HOBrO), bromic acid (HOBrO2), and perbromic acid (HOBrO3), are better studied due to their greater stability, though they are only so in aqueous solution. When bromine dissolves in aqueous solution, the following reactions occur: Hypobromous acid is unstable to disproportionation. The hypobromite ions thus formed disproportionate readily to give bromide and bromate: Bromous acids and bromites are very unstable, although the strontium and barium bromites are known. More important are the bromates, which are prepared on a small scale by oxidation of bromide by aqueous hypochlorite, and are strong oxidising agents. Unlike chlorates, which very slowly disproportionate to chloride and perchlorate, the bromate anion is stable to disproportionation in both acidic and aqueous solutions. Bromic acid is a strong acid. Bromides and bromates may comproportionate to bromine as follows: BrO−3 + 5
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{
"page_id": 3756,
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Br− + 6 H+ → 3 Br2 + 3 H2O There were many failed attempts to obtain perbromates and perbromic acid, leading to some rationalisations as to why they should not exist, until 1968 when the anion was first synthesised from the radioactive beta decay of unstable 83SeO2−4. Today, perbromates are produced by the oxidation of alkaline bromate solutions by fluorine gas. Excess bromate and fluoride are precipitated as silver bromate and calcium fluoride, and the perbromic acid solution may be purified. The perbromate ion is fairly inert at room temperature but is thermodynamically extremely oxidising, with extremely strong oxidising agents needed to produce it, such as fluorine or xenon difluoride. The Br–O bond in BrO−4 is fairly weak, which corresponds to the general reluctance of the 4p elements arsenic, selenium, and bromine to attain their group oxidation state, as they come after the scandide contraction characterised by the poor shielding afforded by the radial-nodeless 3d orbitals. === Organobromine compounds === Like the other carbon–halogen bonds, the C–Br bond is a common functional group that forms part of core organic chemistry. Formally, compounds with this functional group may be considered organic derivatives of the bromide anion. Due to the difference of electronegativity between bromine (2.96) and carbon (2.55), the carbon atom in a C–Br bond is electron-deficient and thus electrophilic. The reactivity of organobromine compounds resembles but is intermediate between the reactivity of organochlorine and organoiodine compounds. For many applications, organobromides represent a compromise of reactivity and cost. Organobromides are typically produced by additive or substitutive bromination of other organic precursors. Bromine itself can be used, but due to its toxicity and volatility, safer brominating reagents are normally used, such as N-bromosuccinimide. The principal reactions for organobromides include dehydrobromination, Grignard reactions, reductive coupling, and nucleophilic substitution. Organobromides are the most
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{
"page_id": 3756,
"source": null,
"title": "Bromine"
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common organohalides in nature, even though the concentration of bromide is only 0.3% of that for chloride in sea water, because of the easy oxidation of bromide to the equivalent of Br+, a potent electrophile. The enzyme bromoperoxidase catalyzes this reaction. The oceans are estimated to release 1–2 million tons of bromoform and 56,000 tons of bromomethane annually. An old qualitative test for the presence of the alkene functional group is that alkenes turn brown aqueous bromine solutions colourless, forming a bromohydrin with some of the dibromoalkane also produced. The reaction passes through a short-lived strongly electrophilic bromonium intermediate. This is an example of a halogen addition reaction. == Occurrence and production == Bromine is significantly less abundant in the crust than fluorine or chlorine, comprising only 2.5 parts per million of the Earth's crustal rocks, and then only as bromide salts. It is significantly more abundant in the oceans, resulting from long-term leaching. There, it makes up 65 parts per million, corresponding to a ratio of about one bromine atom for every 660 chlorine atoms. Salt lakes and brine wells may have higher bromine concentrations: for example, the Dead Sea contains 0.4% bromide ions. It is from these sources that bromine extraction is mostly economically feasible. Bromine is the tenth most abundant element in seawater. The main sources of bromine production are Israel and Jordan. The element is liberated by halogen exchange, using chlorine gas to oxidise Br− to Br2. This is then removed with a blast of steam or air, and is then condensed and purified. Today, bromine is transported in large-capacity metal drums or lead-lined tanks that can hold hundreds of kilograms or even tonnes of bromine. The bromine industry is about one-hundredth the size of the chlorine industry. Laboratory production is unnecessary because bromine is commercially
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{
"page_id": 3756,
"source": null,
"title": "Bromine"
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available and has a long shelf life. == Applications == A wide variety of organobromine compounds are used in industry. Some are prepared from bromine and others are prepared from hydrogen bromide, which is obtained by burning hydrogen in bromine. === Flame retardants === Brominated flame retardants represent a commodity of growing importance, and make up the largest commercial use of bromine. When the brominated material burns, the flame retardant produces hydrobromic acid which interferes in the radical chain reaction of the oxidation reaction of the fire. The mechanism is that the highly reactive hydrogen radicals, oxygen radicals, and hydroxyl radicals react with hydrobromic acid to form less reactive bromine radicals (i.e., free bromine atoms). Bromine atoms may also react directly with other radicals to help terminate the free radical chain-reactions that characterise combustion. To make brominated polymers and plastics, bromine-containing compounds can be incorporated into the polymer during polymerisation. One method is to include a relatively small amount of brominated monomer during the polymerisation process. For example, vinyl bromide can be used in the production of polyethylene, polyvinyl chloride or polypropylene. Specific highly brominated molecules can also be added that participate in the polymerisation process. For example, tetrabromobisphenol A can be added to polyesters or epoxy resins, where it becomes part of the polymer. Epoxies used in printed circuit boards are normally made from such flame retardant resins, indicated by the FR in the abbreviation of the products (FR-4 and FR-2). In some cases, the bromine-containing compound may be added after polymerisation. For example, decabromodiphenyl ether can be added to the final polymers. A number of gaseous or highly volatile brominated halomethane compounds are non-toxic and make superior fire suppressant agents by this same mechanism, and are particularly effective in enclosed spaces such as submarines, airplanes, and spacecraft. However,
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{
"page_id": 3756,
"source": null,
"title": "Bromine"
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they are expensive and their production and use has been greatly curtailed due to their effect as ozone-depleting agents. They are no longer used in routine fire extinguishers, but retain niche uses in aerospace and military automatic fire suppression applications. They include bromochloromethane (Halon 1011, CH2BrCl), bromochlorodifluoromethane (Halon 1211, CBrClF2), and bromotrifluoromethane (Halon 1301, CBrF3). === Other uses === Silver bromide is used, either alone or in combination with silver chloride and silver iodide, as the light sensitive constituent of photographic emulsions. Ethylene bromide was an additive in gasolines containing lead anti-engine knocking agents. It scavenges lead by forming volatile lead bromide, which is exhausted from the engine. This application accounted for 77% of the bromine use in 1966 in the US. This application has declined since the 1970s due to environmental regulations (see below). Brominated vegetable oil (BVO), a complex mixture of plant-derived triglycerides that have been reacted to contain atoms of the element bromine bonded to the molecules, is used primarily to help emulsify citrus-flavored soft drinks, preventing them from separating during distribution. Poisonous bromomethane was widely used as pesticide to fumigate soil and to fumigate housing, by the tenting method. Ethylene bromide was similarly used. These volatile organobromine compounds are all now regulated as ozone depletion agents. The Montreal Protocol on Substances that Deplete the Ozone Layer scheduled the phase out for the ozone depleting chemical by 2005, and organobromide pesticides are no longer used (in housing fumigation they have been replaced by such compounds as sulfuryl fluoride, which contain neither the chlorine or bromine organics which harm ozone). Before the Montreal protocol in 1991 (for example) an estimated 35,000 tonnes of the chemical were used to control nematodes, fungi, weeds and other soil-borne diseases. In pharmacology, inorganic bromide compounds, especially potassium bromide, were frequently used as
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{
"page_id": 3756,
"source": null,
"title": "Bromine"
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general sedatives in the 19th and early 20th century. Bromides in the form of simple salts are still used as anticonvulsants in both veterinary and human medicine, although the latter use varies from country to country. For example, the U.S. Food and Drug Administration (FDA) does not approve bromide for the treatment of any disease, and sodium bromide was removed from over-the-counter sedative products like Bromo-Seltzer, in 1975. Commercially available organobromine pharmaceuticals include the vasodilator nicergoline, the sedative brotizolam, the anticancer agent pipobroman, and the antiseptic merbromin. Otherwise, organobromine compounds are rarely pharmaceutically useful, in contrast to the situation for organofluorine compounds. Several drugs are produced as the bromide (or equivalents, hydrobromide) salts, but in such cases bromide serves as an innocuous counterion of no biological significance. Other uses of organobromine compounds include high-density drilling fluids, dyes (such as Tyrian purple and the indicator bromothymol blue), and pharmaceuticals. Bromine itself, as well as some of its compounds, are used in water treatment, and is the precursor of a variety of inorganic compounds with an enormous number of applications (e.g. silver bromide for photography). Zinc–bromine batteries are hybrid flow batteries used for stationary electrical power backup and storage; from household scale to industrial scale. Bromine is used in cooling towers (in place of chlorine) for controlling bacteria, algae, fungi, and zebra mussels. Because it has similar antiseptic qualities to chlorine, bromine can be used in the same manner as chlorine as a disinfectant or antimicrobial in applications such as swimming pools. Bromine came into this use in the United States during World War II due to a predicted shortage of chlorine. However, bromine is usually not used outside for these applications due to it being relatively more expensive than chlorine and the absence of a stabilizer to protect it from the
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{
"page_id": 3756,
"source": null,
"title": "Bromine"
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sun. For indoor pools, it can be a good option as it is effective at a wider pH range. It is also more stable in a heated pool or hot tub. == Biological role and toxicity == A 2014 study suggests that bromine (in the form of bromide ion) is a necessary cofactor in the biosynthesis of collagen IV, making the element essential to basement membrane architecture and tissue development in animals. Nevertheless, no clear deprivation symptoms or syndromes have been documented in mammals. In other biological functions, bromine may be non-essential but still beneficial when it takes the place of chlorine. For example, in the presence of hydrogen peroxide, H2O2, formed by the eosinophil, and either chloride, iodide, thiocyanate, or bromide ions, eosinophil peroxidase provides a potent mechanism by which eosinophils kill multicellular parasites (such as the nematode worms involved in filariasis) and some bacteria (such as tuberculosis bacteria). Eosinophil peroxidase is a haloperoxidase that preferentially uses bromide over chloride for this purpose, generating hypobromite (hypobromous acid), although the use of chloride is possible. α-Haloesters are generally thought of as highly reactive and consequently toxic intermediates in organic synthesis. Nevertheless, mammals, including humans, cats, and rats, appear to biosynthesize traces of an α-bromoester, 2-octyl 4-bromo-3-oxobutanoate, which is found in their cerebrospinal fluid and appears to play a yet unclarified role in inducing REM sleep. Neutrophil myeloperoxidase can use H2O2 and Br− to brominate deoxycytidine, which could result in DNA mutations. Marine organisms are the main source of organobromine compounds, and it is in these organisms that bromine is more firmly shown to be essential. More than 1600 such organobromine compounds were identified by 1999. The most abundant is methyl bromide (CH3Br), of which an estimated 56,000 tonnes is produced by marine algae each year. The essential oil of the
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{
"page_id": 3756,
"source": null,
"title": "Bromine"
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Hawaiian alga Asparagopsis taxiformis consists of 80% bromoform. Most of such organobromine compounds in the sea are made by the action of a unique algal enzyme, vanadium bromoperoxidase. The bromide anion is not very toxic: a normal daily intake is 2 to 8 milligrams. However, high levels of bromide chronically impair the membrane of neurons, which progressively impairs neuronal transmission, leading to toxicity, known as bromism. Bromide has an elimination half-life of 9 to 12 days, which can lead to excessive accumulation. Doses of 0.5 to 1 gram per day of bromide can lead to bromism. Historically, the therapeutic dose of bromide is about 3 to 5 grams of bromide, thus explaining why chronic toxicity (bromism) was once so common. While significant and sometimes serious disturbances occur to neurologic, psychiatric, dermatological, and gastrointestinal functions, death from bromism is rare. Bromism is caused by a neurotoxic effect on the brain which results in somnolence, psychosis, seizures and delirium. Elemental bromine (Br2) is toxic and causes chemical burns on human flesh. Inhaling bromine gas results in similar irritation of the respiratory tract, causing coughing, choking, shortness of breath, and death if inhaled in large enough amounts. Chronic exposure may lead to frequent bronchial infections and a general deterioration of health. As a strong oxidising agent, bromine is incompatible with most organic and inorganic compounds. Caution is required when transporting bromine; it is commonly carried in steel tanks lined with lead, supported by strong metal frames. The Occupational Safety and Health Administration (OSHA) of the United States has set a permissible exposure limit (PEL) for bromine at a time-weighted average (TWA) of 0.1 ppm. The National Institute for Occupational Safety and Health (NIOSH) has set a recommended exposure limit (REL) of TWA 0.1 ppm and a short-term limit of 0.3 ppm. The exposure
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{
"page_id": 3756,
"source": null,
"title": "Bromine"
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to bromine immediately dangerous to life and health (IDLH) is 3 ppm. Bromine is classified as an extremely hazardous substance in the United States as defined in Section 302 of the U.S. Emergency Planning and Community Right-to-Know Act (42 U.S.C. 11002), and is subject to strict reporting requirements by facilities which produce, store, or use it in significant quantities. == References == == General and cited references == Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
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"page_id": 3756,
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Barium is a chemical element; it has symbol Ba and atomic number 56. It is the fifth element in group 2 and is a soft, silvery alkaline earth metal. Because of its high chemical reactivity, barium is never found in nature as a free element. The most common minerals of barium are barite (barium sulfate, BaSO4) and witherite (barium carbonate, BaCO3). The name barium originates from the alchemical derivative "baryta", from Greek βαρὺς (barys), meaning 'heavy'. Baric is the adjectival form of barium. Barium was identified as a new element in 1772, but not reduced to a metal until 1808 with the advent of electrolysis. Barium has few industrial applications. Historically, it was used as a getter for vacuum tubes and in oxide form as the emissive coating on indirectly heated cathodes. It is a component of YBCO (high-temperature superconductors) and electroceramics, and is added to steel and cast iron to reduce the size of carbon grains within the microstructure. Barium compounds are added to fireworks to impart a green color. Barium sulfate is used as an insoluble additive to oil well drilling fluid. In a purer form it is used as X-ray radiocontrast agents for imaging the human gastrointestinal tract. Water-soluble barium compounds are poisonous and have been used as rodenticides. == Characteristics == === Physical properties === Barium is a soft, silvery-white metal, with a slight golden shade when ultrapure.: 2 The silvery-white color of barium metal rapidly vanishes upon oxidation in air yielding a dark gray layer containing the oxide. Barium has a medium specific weight and high electrical conductivity. Because barium is difficult to purify, many of its properties have not been accurately determined.: 2 At room temperature and pressure, barium metal adopts a body-centered cubic structure, with a barium–barium distance of 503 picometers, expanding with
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{
"page_id": 3757,
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heating at a rate of approximately 1.8×10−5/°C.: 2 It is a soft metal with a Mohs hardness of 1.25.: 2 Its melting temperature of 1,000 K (730 °C; 1,340 °F): 4–43 is intermediate between those of the lighter strontium (1,050 K or 780 °C or 1,430 °F): 4–86 and heavier radium (973 K or 700 °C or 1,292 °F);: 4–78 however, its boiling point of 2,170 K (1,900 °C; 3,450 °F) exceeds that of strontium (1,655 K or 1,382 °C or 2,519 °F).: 4–86 The density (3.62 g/cm3): 4–43 is again intermediate between those of strontium (2.36 g/cm3): 4–86 and radium (≈5 g/cm3).: 4–78 === Chemical reactivity === Barium is chemically similar to magnesium, calcium, and strontium, but more reactive. Its compounds are almost invariably found in the +2 oxidation state. As expected for a highly electropositive metal, barium's reaction with chalcogens is highly exothermic (release energy). Barium reacts with atmospheric oxygen in air at room temperature. For this reason, metallic barium is often stored under oil or in an inert atmosphere.: 2 Reactions with other nonmetals, such as carbon, nitrogen, phosphorus, silicon, and hydrogen, proceed upon heating.: 2–3 Reactions with water and alcohols are also exothermic and release hydrogen gas:: 3 Ba + 2 ROH → Ba(OR)2 + H2↑ (R is an alkyl group or a hydrogen atom) Barium reacts with ammonia to form the electride [Ba(NH3)6](e−)2, which near room temperature gives the amide Ba(NH2)2. The metal is readily attacked by acids. Sulfuric acid is a notable exception because passivation stops the reaction by forming the insoluble barium sulfate on the surface. Barium combines with several other metals, including aluminium, zinc, lead, and tin, forming intermetallic phases and alloys. === Compounds === Barium salts are typically white when solid and colorless when dissolved. They are denser than the strontium
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{
"page_id": 3757,
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"title": "Barium"
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or calcium analogs, except for the halides (see table; zinc is given for comparison). Barium hydroxide ("baryta") was known to alchemists, who produced it by heating barium carbonate. Unlike calcium hydroxide, it absorbs very little CO2 in aqueous solutions and is therefore insensitive to atmospheric fluctuations. This property is used in calibrating pH equipment. Barium compounds burn with a green to pale green flame, which is an efficient test to detect a barium compound. The color results from spectral lines at 455.4, 493.4, 553.6, and 611.1 nm.: 3 Organobarium compounds are a growing field of knowledge: recently discovered are dialkylbariums and alkylhalobariums.: 3 === Isotopes === Barium found in the Earth's crust is a mixture of seven primordial nuclides, barium-130, 132, and 134 through 138. Barium-130 undergoes very slow radioactive decay to xenon-130 by double beta plus decay, with a half-life of (0.5–2.7)×1021 years (about 1011 times the age of the universe). Its abundance is ≈0.1% that of natural barium. Theoretically, barium-132 can similarly undergo double beta decay to xenon-132; this decay has not been detected. The radioactivity of these isotopes is so weak that they pose no danger to life. Of the stable isotopes, barium-138 composes 71.7% of all barium; other isotopes have decreasing abundance with decreasing mass number. In total, barium has 40 known isotopes, ranging in mass between 114 and 153. The most stable artificial radioisotope is barium-133 with a half-life of approximately 10.51 years. Five other isotopes have half-lives longer than a day. Barium also has 10 meta states, of which barium-133m1 is the most stable with a half-life of about 39 hours. == History == Alchemists in the early Middle Ages knew about some barium minerals. Smooth pebble-like stones of mineral baryte were found in volcanic rock near Bologna, Italy, and so were called "Bologna
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{
"page_id": 3757,
"source": null,
"title": "Barium"
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stones". Alchemists were attracted to them because after exposure to light they would glow for years. The phosphorescent properties of baryte heated with organics were described by V. Casciorolus in 1602.: 5 Carl Scheele determined that baryte contained a new element in 1772, but could not isolate barium, only barium oxide. Johan Gottlieb Gahn also isolated barium oxide two years later in similar studies. Oxidized barium was at first called "barote" by Guyton de Morveau, a name that was changed by Antoine Lavoisier to baryte (in French) or baryta (in Latin). Also in the 18th century, English mineralogist William Withering noted a heavy mineral in the lead mines of Cumberland, now known to be witherite. Barium was first isolated by electrolysis of molten barium salts in 1808 by Sir Humphry Davy in England. Davy, by analogy with calcium, named "barium" after baryta, with the "-ium" ending signifying a metallic element. Robert Bunsen and Augustus Matthiessen obtained pure barium by electrolysis of a molten mixture of barium chloride and ammonium chloride. The production of pure oxygen in the Brin process was a large-scale application of barium peroxide in the 1880s, before it was replaced by electrolysis and fractional distillation of liquefied air in the early 1900s. In this process barium oxide reacts at 500–600 °C (932–1,112 °F) with air to form barium peroxide, which decomposes above 700 °C (1,292 °F) by releasing oxygen: 2 BaO + O2 ⇌ 2 BaO2 Barium sulfate was first applied as a radiocontrast agent in X-ray imaging of the digestive system in 1908. == Occurrence and production == The abundance of barium is 0.0425% in the Earth's crust and 13 μg/L in sea water. The primary commercial source of barium is baryte (also called barytes or heavy spar), a barium sulfate mineral.: 5 with deposits in
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{
"page_id": 3757,
"source": null,
"title": "Barium"
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many parts of the world. Another commercial source, far less important than baryte, is witherite, barium carbonate. The main deposits are located in Britain, Romania, and the former USSR.: 5 The baryte reserves are estimated between 0.7 and 2 billion tonnes. The highest production, 8.3 million tonnes, was achieved in 1981, but only 7–8% was used for barium metal or compounds.: 5 Baryte production has risen since the second half of the 1990s from 5.6 million tonnes in 1996 to 7.6 in 2005 and 7.8 in 2011. China accounts for more than 50% of this output, followed by India (14% in 2011), Morocco (8.3%), US (8.2%), Iran and Kazakhstan (2.6% each) and Turkey (2.5%). The mined ore is washed, crushed, classified, and separated from quartz. If the quartz penetrates too deeply into the ore, or the iron, zinc, or lead content is abnormally high, then froth flotation is used. The product is a 98% pure baryte (by mass); the purity should be no less than 95%, with a minimal content of iron and silicon dioxide.: 7 It is then reduced by carbon to barium sulfide:: 6 BaSO4 + 2 C → BaS + 2 CO2 The water-soluble barium sulfide is the starting point for other compounds: treating BaS with oxygen produces the sulfate, with nitric acid the nitrate, with aqueous carbon dioxide the carbonate, and so on.: 6 The nitrate can be thermally decomposed to yield the oxide.: 6 Barium metal is produced by reduction with aluminium at 1,100 °C (2,010 °F). The intermetallic compound BaAl4 is produced first:: 3 3 BaO + 14 Al → 3 BaAl4 + Al2O3 BaAl4 is an intermediate reacted with barium oxide to produce the metal. Note that not all barium is reduced.: 3 8 BaO + BaAl4 → Ba↓ + 7 BaAl2O4 The
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"page_id": 3757,
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"title": "Barium"
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remaining barium oxide reacts with the formed aluminium oxide:: 3 BaO + Al2O3 → BaAl2O4 and the overall reaction is: 3 4 BaO + 2 Al → 3 Ba↓ + BaAl2O4 Barium vapor is condensed and packed into molds in an atmosphere of argon.: 3 This method is used commercially, yielding ultrapure barium.: 3 Commonly sold barium is about 99% pure, with main impurities being strontium and calcium (up to 0.8% and 0.25%) and other contaminants contributing less than 0.1%.: 4 A similar reaction with silicon at 1,200 °C (2,190 °F) yields barium and barium metasilicate.: 3 Electrolysis is not used because barium readily dissolves in molten halides and the product is rather impure.: 3 === Gemstone === The barium mineral, benitoite (barium titanium silicate), occurs as a very rare blue fluorescent gemstone, and is the official state gem of California. === Barium in seawater === Barium exists in seawater as the Ba2+ ion with an average oceanic concentration of 109 nmol/kg. Barium also exists in the ocean as BaSO4, or barite. Barium has a nutrient-like profile with a residence time of 10,000 years. Barium shows a relatively consistent concentration in upper ocean seawater, excepting regions of high river inputs and regions with strong upwelling. There is little depletion of barium concentrations in the upper ocean for an ion with a nutrient-like profile, thus lateral mixing is important. Barium isotopic values show basin-scale balances instead of local or short-term processes. == Applications == === Metal and alloys === Barium, as a metal or when alloyed with aluminium, is used to remove unwanted gases (gettering) from vacuum tubes, such as TV picture tubes.: 4 Barium is suitable for this purpose because of its low vapor pressure and reactivity towards oxygen, nitrogen, carbon dioxide, and water; it can even partly remove noble
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"title": "Barium"
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gases by dissolving them in the crystal lattice. This application is gradually disappearing due to the rising popularity of the tubeless LCD, LED, and plasma sets.: 4 Other uses of elemental barium are minor and include an additive to silumin (aluminium–silicon alloys) that refines their structure, as well as: 4 bearing alloys; lead–tin soldering alloys – to increase the creep resistance; alloy with nickel for spark plugs; additive to steel and cast iron as an inoculant; alloys with calcium, manganese, silicon, and aluminium as high-grade steel deoxidizers. === Barium sulfate and baryte === Barium sulfate (the mineral baryte, BaSO4) is important to the petroleum industry as a drilling fluid in oil and gas wells.: 4–5 The precipitate of the compound (called "blanc fixe", from the French for "permanent white") is used in paints and varnishes; as a filler in ringing ink, plastics, and rubbers; as a paper coating pigment; and in nanoparticles, to improve physical properties of some polymers, such as epoxies.: 9 Barium sulfate has a low toxicity and relatively high density of ca. 4.5 g/cm3 (and thus opacity to X-rays). For this reason it is used as a radiocontrast agent in X-ray imaging of the digestive system ("barium meals" and "barium enemas").: 4–5 Lithopone, a pigment that contains barium sulfate and zinc sulfide, is a permanent white with good covering power that does not darken when exposed to sulfides. === Other barium compounds === Other compounds of barium find only niche applications, limited by the toxicity of Ba2+ ions (see § Toxicity), which is not a problem for the insoluble BaSO4. Barium oxide coating on the electrodes of fluorescent lamps facilitates the release of electrons. By its great atomic density, barium carbonate increases the refractive index and luster of glass: 4–5 and reduces leaks of X-rays from CRT
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{
"page_id": 3757,
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"title": "Barium"
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screens.: 12–13 Barium, typically as barium nitrate imparts a yellow or "apple" green color to fireworks when no chlorine donors are present.; emerald greens are generated using chlorine donors (barium chlorate acting as the donor and oxidizer in many formulas) to produce barium chloride in-situ. Barium peroxide is a catalyst in the aluminothermic reaction (thermite) for welding rail tracks. It is also a green flare in tracer ammunition and a bleaching agent. Barium titanate is a promising electroceramic. Barium fluoride is used for optics in infrared applications because of its wide transparency range of 0.15–12 micrometers. YBCO was the first high-temperature superconductor cooled by liquid nitrogen, with a transition temperature of 93 K (−180.2 °C; −292.3 °F) greater than the boiling point of nitrogen (77 K or −196.2 °C or −321.1 °F). Ferrite, a type of sintered ceramic composed of iron oxide (Fe2O3) and barium oxide (BaO), is both electrically nonconductive and ferrimagnetic, and can be temporarily or permanently magnetized. === Palaeoceanography === The lateral mixing of barium is caused by water mass mixing and ocean circulation. Global ocean circulation reveals a strong correlation between dissolved barium and silicic acid. The large-scale ocean circulation combined with remineralization of barium show a similar correlation between dissolved barium and ocean alkalinity. Dissolved barium's correlation with silicic acid can be seen both vertically and spatially. Particulate barium shows a strong correlation with particulate organic carbon or POC. Barium is becoming more popular as a base for palaeoceanographic proxies. With both dissolved and particulate barium's links with silicic acid and POC, it can be used to determine historical variations in the biological pump, carbon cycle, and global climate. The barium particulate barite (BaSO4), as one of many proxies, can be used to provide a host of historical information on processes in different oceanic settings
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"page_id": 3757,
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"title": "Barium"
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(water column, sediments, and hydrothermal sites). In each setting there are differences in isotopic and elemental composition of the barite particulate. Barite in the water column, known as marine or pelagic barite, reveals information on seawater chemistry variation over time. Barite in sediments, known as diagenetic or cold seeps barite, gives information about sedimentary redox processes. Barite formed via hydrothermal activity at hydrothermal vents, known as hydrothermal barite, reveals alterations in the condition of the earth's crust around those vents. == Toxicity == Soluble barium compounds have LD50 near 10 mg/kg (oral rats). Symptoms include "convulsions... paralysis of the peripheral nerve system ... severe inflammation of the gastrointestinal tract".: 18 The insoluble sulfate is nontoxic and is not classified as a dangerous goods in transport regulations.: 9 Little is known about the long term effects of barium exposure. The US EPA considers it unlikely that barium is carcinogenic when consumed orally. Inhaled dust containing insoluble barium compounds can accumulate in the lungs, causing a benign condition called baritosis. Barium carbonate has been used as a rodenticide. Though considered obsolete, it may still be in use in some countries. == See also == Han purple and Han blue – synthetic barium copper silicate pigments developed and used in ancient and imperial China == References == == External links == Barium at The Periodic Table of Videos (University of Nottingham) Elementymology & Elements Multidict 3-D Holographic Display Using Strontium Barium Niobate
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"title": "Barium"
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Berkelium is a synthetic chemical element; it has symbol Bk and atomic number 97. It is a member of the actinide and transuranium element series. It is named after the city of Berkeley, California, the location of the Lawrence Berkeley National Laboratory (then the University of California Radiation Laboratory) where it was discovered in December 1949. Berkelium was the fifth transuranium element discovered after neptunium, plutonium, curium and americium. The major isotope of berkelium, 249Bk, is synthesized in minute quantities in dedicated high-flux nuclear reactors, mainly at the Oak Ridge National Laboratory in Tennessee, United States, and at the Research Institute of Atomic Reactors in Dimitrovgrad, Russia. The longest-lived and second-most important isotope, 247Bk, can be synthesized via irradiation of 244Cm with high-energy alpha particles. Just over one gram of berkelium has been produced in the United States since 1967. There is no practical application of berkelium outside scientific research which is mostly directed at the synthesis of heavier transuranium elements and superheavy elements. A 22-milligram batch of berkelium-249 was prepared during a 250-day irradiation period and then purified for a further 90 days at Oak Ridge in 2009. This sample was used to synthesize the new element tennessine for the first time in 2009 at the Joint Institute for Nuclear Research, Russia, after it was bombarded with calcium-48 ions for 150 days. This was the culmination of the Russia–US collaboration on the synthesis of the heaviest elements on the periodic table. Berkelium is a soft, silvery-white, radioactive metal. The berkelium-249 isotope emits low-energy electrons and thus is relatively safe to handle. It decays with a half-life of 330 days to californium-249, which is a strong emitter of ionizing alpha particles. This gradual transformation is an important consideration when studying the properties of elemental berkelium and its chemical compounds, since
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the formation of californium brings not only chemical contamination, but also free-radical effects and self-heating from the emitted alpha particles. == Characteristics == === Physical === Berkelium is a soft, silvery-white, radioactive actinide metal. In the periodic table, it is located to the right of the actinide curium, to the left of the actinide californium and below the lanthanide terbium with which it shares many similarities in physical and chemical properties. Its density of 14.78 g/cm3 lies between those of curium (13.52 g/cm3) and californium (15.1 g/cm3), as does its melting point of 986 °C, below that of curium (1340 °C) but higher than that of californium (900 °C). Berkelium is relatively soft and has one of the lowest bulk moduli among the actinides, at about 20 GPa (2×1010 Pa). Berkelium(III) ions shows two sharp fluorescence peaks at 652 nanometers (red light) and 742 nanometers (deep red – near-infrared) due to internal transitions at the f-electron shell. The relative intensity of these peaks depends on the excitation power and temperature of the sample. This emission can be observed, for example, after dispersing berkelium ions in a silicate glass, by melting the glass in presence of berkelium oxide or halide. Between 70 K and room temperature, berkelium behaves as a Curie–Weiss paramagnetic material with an effective magnetic moment of 9.69 Bohr magnetons (μB) and a Curie temperature of 101 K. This magnetic moment is almost equal to the theoretical value of 9.72 μB calculated within the simple atomic L-S coupling model. Upon cooling to about 34 K, berkelium undergoes a transition to an antiferromagnetic state. The enthalpy of dissolution in hydrochloric acid at standard conditions is −600 kJ/mol, from which the standard enthalpy of formation (ΔfH°) of aqueous Bk3+ ions is obtained as −601 kJ/mol. The standard electrode potential Bk3+/Bk is
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"page_id": 3758,
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−2.01 V. The ionization potential of a neutral berkelium atom is 6.23 eV. === Allotropes === At ambient conditions, berkelium assumes its most stable α form which has a hexagonal symmetry, space group P63/mmc, lattice parameters of 341 pm and 1107 pm. The crystal has a double-hexagonal close packing structure with the layer sequence ABAC and so is isotypic (having a similar structure) with α-lanthanum and α-forms of actinides beyond curium. This crystal structure changes with pressure and temperature. When compressed at room temperature to 7 GPa, α-berkelium transforms to the β modification, which has a face-centered cubic (fcc) symmetry and space group Fm3m. This transition occurs without change in volume, but the enthalpy increases by 3.66 kJ/mol. Upon further compression to 25 GPa, berkelium transforms to an orthorhombic γ-berkelium structure similar to that of α-uranium. This transition is accompanied by a 12% volume decrease and delocalization of the electrons at the 5f electron shell. No further phase transitions are observed up to 57 GPa. Upon heating, α-berkelium transforms into another phase with an fcc lattice (but slightly different from β-berkelium), space group Fm3m and the lattice constant of 500 pm; this fcc structure is equivalent to the closest packing with the sequence ABC. This phase is metastable and will gradually revert to the original α-berkelium phase at room temperature. The temperature of the phase transition is believed to be quite close to the melting point. === Chemical === Like all actinides, berkelium dissolves in various aqueous inorganic acids, liberating gaseous hydrogen and converting into the berkelium(III) state. This trivalent oxidation state (+3) is the most stable, especially in aqueous solutions, but tetravalent (+4), pentavalent (+5), and possibly divalent (+2) berkelium compounds are also known. The existence of divalent berkelium salts is uncertain and has only been reported in mixed
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"page_id": 3758,
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"title": "Berkelium"
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lanthanum(III) chloride-strontium chloride melts. A similar behavior is observed for the lanthanide analogue of berkelium, terbium. Aqueous solutions of Bk3+ ions are green in most acids. The color of Bk4+ ions is yellow in hydrochloric acid and orange-yellow in sulfuric acid. Berkelium does not react rapidly with oxygen at room temperature, possibly due to the formation of a protective oxide layer surface. However, it reacts with molten metals, hydrogen, halogens, chalcogens and pnictogens to form various binary compounds. In 2025 an organometallic compound containing berkelium was synthesized from 0.3 mg of berkelium and named berkelocene. === Isotopes === Nineteen isotopes and six nuclear isomers (excited states of an isotope) of berkelium have been characterized, with mass numbers ranging from 233 to 253 (except 235 and 237). All of them are radioactive. The longest half-lives are observed for 247Bk (1,380 years), 248Bk (over 300 years), and 249Bk (330 days); the half-lives of the other isotopes range from microseconds to several days. The isotope which is the easiest to synthesize is berkelium-249. This emits mostly soft β-particles which are inconvenient for detection. Its alpha radiation is rather weak (1.45×10−3%) with respect to the β-radiation, but is sometimes used to detect this isotope. The second important berkelium isotope, berkelium-247, is an alpha-emitter, as are most actinide isotopes. === Occurrence === All berkelium isotopes have a half-life far too short to be primordial. Therefore, any primordial berkelium − that is, berkelium present on the Earth during its formation − has decayed by now. On Earth, berkelium is mostly concentrated in certain areas, which were used for the atmospheric nuclear weapons tests between 1945 and 1980, as well as at the sites of nuclear incidents, such as the Chernobyl disaster, Three Mile Island accident and 1968 Thule Air Base B-52 crash. Analysis of the debris
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"page_id": 3758,
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at the testing site of the first United States' first thermonuclear weapon, Ivy Mike, (1 November 1952, Enewetak Atoll), revealed high concentrations of various actinides, including berkelium. For reasons of military secrecy, this result was not published until 1956. Nuclear reactors produce mostly, among the berkelium isotopes, berkelium-249. During the storage and before the fuel disposal, most of it beta decays to californium-249. The latter has a half-life of 351 years, which is relatively long compared to the half-lives of other isotopes produced in the reactor, and is therefore undesirable in the disposal products. The transuranium elements from americium to fermium, including berkelium, occurred naturally in the natural nuclear fission reactor at Oklo, but no longer do so. Berkelium is also one of the elements that have theoretically been detected in Przybylski's Star. == History == Although very small amounts of berkelium were possibly produced in previous nuclear experiments, it was first intentionally synthesized, isolated and identified in December 1949 by Glenn T. Seaborg, Albert Ghiorso, Stanley Gerald Thompson, and Kenneth Street Jr. They used the 60-inch cyclotron at the University of California, Berkeley. Similar to the nearly simultaneous discovery of americium (element 95) and curium (element 96) in 1944, the new elements berkelium and californium (element 98) were both produced in 1949–1950. The name choice for element 97 followed the previous tradition of the Californian group to draw an analogy between the newly discovered actinide and the lanthanide element positioned above it in the periodic table. Previously, americium was named after a continent as its analogue europium, and curium honored scientists Marie and Pierre Curie as the lanthanide above it, gadolinium, was named after the explorer of the rare-earth elements Johan Gadolin. Thus, the discovery report by the Berkeley group reads: "It is suggested that element 97 be given
|
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"page_id": 3758,
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"title": "Berkelium"
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the name berkelium (symbol Bk) after the city of Berkeley in a manner similar to that used in naming its chemical homologue terbium (atomic number 65) whose name was derived from the town of Ytterby, Sweden, where the rare earth minerals were first found." This tradition ended with berkelium, though, as the naming of the next discovered actinide, californium, was not related to its lanthanide analogue dysprosium, but after the discovery place. The most difficult steps in synthesising berkelium were its separation from the final products and the production of sufficient quantities of americium for the target material. First, americium (241Am) nitrate solution was coated on a platinum foil, the solution was evaporated and the residue converted by annealing to americium dioxide (AmO2). This target was irradiated with 35 MeV alpha particles for 6 hours in the 60-inch cyclotron at the Lawrence Radiation Laboratory, University of California, Berkeley. The (α,2n) reaction induced by the irradiation yielded the 243Bk isotope and two free neutrons: 24195Am + 42He → 24397Bk + 2 10n After the irradiation, the coating was dissolved with nitric acid and then precipitated as the hydroxide using concentrated aqueous ammonia solution. The product was centrifugated and re-dissolved in nitric acid. To separate berkelium from the unreacted americium, this solution was added to a mixture of aqueous ammonia and ammonium sulfate and heated in the presence of atmospheric oxygen to convert all the dissolved americium into the oxidation state +6. Unoxidized residual americium was precipitated by the addition of hydrofluoric acid as americium(III) fluoride (AmF3). This step yielded a mixture of the accompanying product curium and the expected element 97 in form of trifluorides. The mixture was converted to the corresponding hydroxides by treating it with potassium hydroxide, and after centrifugation, was dissolved in perchloric acid. Further separation was carried
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{
"page_id": 3758,
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out in the presence of a citric acid/ammonium buffer solution in a weakly acidic medium (pH ≈ 3.5), using ion exchange at elevated temperature. The chromatographic separation behavior was unknown for element 97 at the time but was anticipated by analogy with terbium. The first results were disappointing because no alpha-particle emission signature could be detected from the elution product. With further analysis, searching for characteristic X-rays and conversion electron signals, a berkelium isotope was eventually detected. Its mass number was uncertain between 243 and 244 in the initial report, but was later established as 243. == Synthesis and extraction == === Preparation of isotopes === Berkelium is produced by bombarding lighter actinides uranium (238U) or plutonium (239Pu) with neutrons in a nuclear reactor. In a more common case of uranium fuel, plutonium is produced first by neutron capture (the so-called (n,γ) reaction or neutron fusion) followed by beta-decay: U 92 238 → ( n , γ ) U 92 239 → 23.5 min β − Np 93 239 → 2.3565 d β − Pu 94 239 {\displaystyle {\ce {^{238}_{92}U ->[{\ce {(n,\gamma)}}] ^{239}_{92}U ->[\beta^-][23.5 \ {\ce {min}}] ^{239}_{93}Np ->[\beta^-][2.3565 \ {\ce {d}}] ^{239}_{94}Pu}}} (The times are half-lives.) Plutonium-239 is further irradiated by a source that has a high neutron flux, several times higher than a conventional nuclear reactor, such as the 85-megawatt High Flux Isotope Reactor (HFIR) at the Oak Ridge National Laboratory in Tennessee, US. The higher flux promotes fusion reactions involving not one but several neutrons, converting 239Pu to 244Cm and then to 249Cm: Pu 94 239 → 4 ( n , γ ) Pu 94 243 → 4.956 h β − Am 95 243 → ( n , γ ) Am 95 244 → 10.1 h β − Cm 96 244 Cm 96 244 → 5 (
|
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n , γ ) Cm 96 249 {\displaystyle {\begin{aligned}{\ce {^{239}_{94}Pu ->[{\ce {4(n,\gamma)}}] ^{243}_{94}Pu ->[\beta^-][4.956 \ {\ce {h}}] ^{243}_{95}Am ->[{\ce {(n,\gamma)}}] ^{244}_{95}Am ->[\beta^-][10.1 \ {\ce {h}}]}}&{\ce {^{244}_{96}Cm}}\\&{\ce {^{244}_{96}Cm ->[{\ce {5(n,\gamma)}}] ^{249}_{96}Cm}}\end{aligned}}} Curium-249 has a short half-life of 64 minutes, and thus its further conversion to 250Cm has a low probability. Instead, it transforms by beta-decay into 249Bk: Cm 96 249 → 64.15 min β − 97 249 Bk → 330 d β − 98 249 Cf {\displaystyle {\ce {^{249}_{96}Cm->[{\beta ^{-}}][64.15\ {\ce {min}}]_{97}^{249}Bk->[\beta ^{-}][330\ {\ce {d}}]_{98}^{249}Cf}}} The thus-produced 249Bk has a long half-life of 330 days and thus can capture another neutron. However, the product, 250Bk, again has a relatively short half-life of 3.212 hours and thus does not yield any heavier berkelium isotopes. It instead decays to the californium isotope 250Cf: Bk 97 249 → ( n , γ ) Bk 97 250 → 3.212 h β − Cf 98 250 {\displaystyle {\ce {^{249}_{97}Bk ->[{\ce {(n,\gamma)}}] ^{250}_{97}Bk ->[\beta^-][3.212 \ {\ce {h}}] ^{250}_{98}Cf}}} Although 247Bk is the most stable isotope of berkelium, its production in nuclear reactors is very difficult because its potential progenitor 247Cm has never been observed to undergo beta decay. Thus, 249Bk is the most accessible isotope of berkelium, which still is available only in small quantities (only 0.66 grams have been produced in the US over the period 1967–1983) at a high price of the order 185 USD per microgram. It is the only berkelium isotope available in bulk quantities, and thus the only berkelium isotope whose properties can be extensively studied. The isotope 248Bk was first obtained in 1956 by bombarding a mixture of curium isotopes with 25 MeV α-particles. Although its direct detection was hindered by strong signal interference with 245Bk, the existence of a new isotope was proven by the growth of the decay product
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248Cf which had been previously characterized. The half-life of 248Bk was estimated as 23±5 hours, though later 1965 work gave a half-life in excess of 300 years (which may be due to an isomeric state). Berkelium-247 was produced during the same year by irradiating 244Cm with alpha-particles: { Cm 96 244 → ( α , n ) Cf 98 247 → 3.11 h ϵ Bk 97 247 Cm 96 244 → ( α , p ) Bk 97 247 {\displaystyle {\begin{cases}{\ce {^{244}_{96}Cm ->[{\ce {(\alpha,n)}}] ^{247}_{98}Cf ->[\epsilon][3.11 \ {\ce {h}}] ^{247}_{97}Bk}}\\{\ce {^{244}_{96}Cm ->[{\ce {(\alpha,p)}}] ^{247}_{97}Bk}}\end{cases}}} Berkelium-242 was synthesized in 1979 by bombarding 235U with 11B, 238U with 10B, 232Th with 14N or 232Th with 15N. It converts by electron capture to 242Cm with a half-life of 7.0±1.3 minutes. A search for an initially suspected isotope 241Bk was then unsuccessful; 241Bk has since been synthesized. { U 92 235 + B 5 11 ⟶ Bk 97 242 + 4 0 1 n Th 90 232 + N 7 14 ⟶ Bk 97 242 + 4 0 1 n U 92 238 + B 5 10 ⟶ Bk 97 242 + 6 0 1 n Th 90 232 + N 7 15 ⟶ Bk 97 242 + 5 0 1 n {\displaystyle {\begin{cases}{\ce {^{235}_{92}U + ^{11}_{5}B -> ^{242}_{97}Bk + 4^{1}_{0}n}}&{\ce {^{232}_{90}Th + ^{14}_{7}N -> ^{242}_{97}Bk + 4^{1}_{0}n}}\\{\ce {^{238}_{92}U + ^{10}_{5}B -> ^{242}_{97}Bk + 6^{1}_{0}n}}&{\ce {^{232}_{90}Th + ^{15}_{7}N -> ^{242}_{97}Bk + 5^{1}_{0}n}}\end{cases}}} === Separation === The fact that berkelium readily assumes oxidation state +4 in solids, and is relatively stable in this state in liquids, greatly assists separation of berkelium from many other actinides. These are produced in relatively large amounts during the nuclear synthesis and often favor the +3 state. This fact was not yet known in the initial experiments, which used a
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more complex separation procedure. Various inorganic oxidation agents can be applied to the berkelium(III) solution to convert it to the +4 state, such as bromates (BrO−3), bismuthates (BiO−3), chromates (CrO2−4 and Cr2O2−7), silver(I) thiolate (Ag2S2O8), lead(IV) oxide (PbO2), ozone (O3), or photochemical oxidation procedures. More recently, it has been discovered that some organic and bio-inspired molecules, such as the chelator 3,4,3-LI(1,2-HOPO), can also oxidize Bk(III) and stabilize Bk(IV) under mild conditions. Berkelium(IV) is then extracted with ion exchange, extraction chromatography or liquid-liquid extraction using HDEHP (bis-(2-ethylhexyl) phosphoric acid), amines, tributyl phosphate or various other reagents. These procedures separate berkelium from most trivalent actinides and lanthanides, except for the lanthanide cerium (lanthanides are absent in the irradiation target but are created in various nuclear fission decay chains). A more detailed procedure adopted at the Oak Ridge National Laboratory was as follows: the initial mixture of actinides is processed with ion exchange using lithium chloride reagent, then precipitated as hydroxides, filtered and dissolved in nitric acid. It is then treated with high-pressure elution from cation exchange resins, and the berkelium phase is oxidized and extracted using one of the procedures described above. Reduction of the thus-obtained berkelium(IV) to the +3 oxidation state yields a solution, which is nearly free from other actinides (but contains cerium). Berkelium and cerium are then separated with another round of ion-exchange treatment. === Bulk metal preparation === In order to characterize chemical and physical properties of solid berkelium and its compounds, a program was initiated in 1952 at the Material Testing Reactor, Arco, Idaho, US. It resulted in preparation of an eight-gram plutonium-239 target and in the first production of macroscopic quantities (0.6 micrograms) of berkelium by Burris B. Cunningham and Stanley Gerald Thompson in 1958, after a continuous reactor irradiation of this target for six years.
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This irradiation method was and still is the only way of producing weighable amounts of the element, and most solid-state studies of berkelium have been conducted on microgram or submicrogram-sized samples. The world's major irradiation sources are the 85-megawatt High Flux Isotope Reactor at the Oak Ridge National Laboratory in Tennessee, USA, and the SM-2 loop reactor at the Research Institute of Atomic Reactors (NIIAR) in Dimitrovgrad, Russia, which are both dedicated to the production of transcurium elements (atomic number greater than 96). These facilities have similar power and flux levels, and are expected to have comparable production capacities for transcurium elements, although the quantities produced at NIIAR are not publicly reported. In a "typical processing campaign" at Oak Ridge, tens of grams of curium are irradiated to produce decigram quantities of californium, milligram quantities of berkelium-249 and einsteinium, and picogram quantities of fermium. In total, just over one gram of berkelium-249 has been produced at Oak Ridge since 1967. The first berkelium metal sample weighing 1.7 micrograms was prepared in 1971 by the reduction of berkelium(III) fluoride with lithium vapor at 1000 °C; the fluoride was suspended on a tungsten wire above a tantalum crucible containing molten lithium. Later, metal samples weighing up to 0.5 milligrams were obtained with this method. BkF3 + 3 Li → Bk + 3 LiF Similar results are obtained with berkelium(IV) fluoride. Berkelium metal can also be produced by the reduction of berkelium(IV) oxide with thorium or lanthanum. == Compounds == === Oxides === Two oxides of berkelium are known, with the berkelium oxidation state of +3 (Bk2O3) and +4 (BkO2). Berkelium(IV) oxide is a brown solid, while berkelium(III) oxide is a yellow-green solid with a melting point of 1920 °C and is formed from BkO2 by reduction with molecular hydrogen: 2 BkO2 +
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H2 → Bk2O3 + H2O Upon heating to 1200 °C, the oxide Bk2O3 undergoes a phase change; it undergoes another phase change at 1750 °C. Such three-phase behavior is typical for the actinide sesquioxides. Berkelium(II) oxide, BkO, has been reported as a brittle gray solid but its exact chemical composition remains uncertain. === Halides === In halides, berkelium assumes the oxidation states +3 and +4. The +3 state is the most stable, especially in solutions, while the tetravalent halides BkF4 and Cs2BkCl6 are only known in the solid phase. The coordination of berkelium atom in its trivalent fluoride and chloride is tricapped trigonal prismatic, with the coordination number of 9. In trivalent bromide, it is bicapped trigonal prismatic (coordination 8) or octahedral (coordination 6), and in the iodide it is octahedral. Berkelium(IV) fluoride (BkF4) is a yellow-green ionic solid and is isotypic with uranium tetrafluoride or zirconium tetrafluoride. Berkelium(III) fluoride (BkF3) is also a yellow-green solid, but it has two crystalline structures. The most stable phase at low temperatures is isotypic with yttrium(III) fluoride, while upon heating to between 350 and 600 °C, it transforms to the structure found in lanthanum trifluoride. Visible amounts of berkelium(III) chloride (BkCl3) were first isolated and characterized in 1962, and weighed only 3 billionths of a gram. It can be prepared by introducing hydrogen chloride vapors into an evacuated quartz tube containing berkelium oxide at a temperature about 500 °C. This green solid has a melting point of 600 °C, and is isotypic with uranium(III) chloride. Upon heating to nearly melting point, BkCl3 converts into an orthorhombic phase. Two forms of berkelium(III) bromide are known: one with berkelium having coordination 6, and one with coordination 8. The latter is less stable and transforms to the former phase upon heating to about 350 °C. An
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"page_id": 3758,
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"title": "Berkelium"
}
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important property of radioactive solids has been studied on these two crystal forms: the structure of fresh and aged 249BkBr3 samples was probed by X-ray diffraction over a period longer than 3 years, so that various fractions of berkelium-249 had beta decayed to californium-249. No change in structure was observed upon the 249BkBr3—249CfBr3 transformation. However, other differences were noted for 249BkBr3 and 249CfBr3. For example, the latter could be reduced with hydrogen to 249CfBr2, but the former could not – this result was reproduced on individual 249BkBr3 and 249CfBr3 samples, as well on the samples containing both bromides. The intergrowth of californium in berkelium occurs at a rate of 0.22% per day and is an obstacle to studying berkelium properties. Beside a chemical contamination, 249Cf, being an alpha emitter, brings undesirable self-damage of the crystal lattice and the resulting self-heating. The chemical effect however can be avoided by performing measurements as a function of time and extrapolating the obtained results. === Other inorganic compounds === The pnictides of berkelium-249 of the type BkX are known for the elements nitrogen, phosphorus, arsenic and antimony. They crystallize in the rock-salt structure and are prepared by the reaction of either berkelium(III) hydride (BkH3) or metallic berkelium with these elements at elevated temperature (about 600 °C) under high vacuum. Berkelium(III) sulfide, Bk2S3, is prepared by either treating berkelium oxide with a mixture of hydrogen sulfide and carbon disulfide vapors at 1130 °C, or by directly reacting metallic berkelium with elemental sulfur. These procedures yield brownish-black crystals. Berkelium(III) and berkelium(IV) hydroxides are both stable in 1 molar solutions of sodium hydroxide. Berkelium(III) phosphate (BkPO4) has been prepared as a solid, which shows strong fluorescence under excitation with a green light. Berkelium hydrides are produced by reacting metal with hydrogen gas at temperatures about 250 °C.
|
{
"page_id": 3758,
"source": null,
"title": "Berkelium"
}
|
They are non-stoichiometric with the nominal formula BkH2+x (0 < x < 1). Several other salts of berkelium are known, including an oxysulfide (Bk2O2S), and hydrated nitrate (Bk(NO3)3·4H2O), chloride (BkCl3·6H2O), sulfate (Bk2(SO4)3·12H2O) and oxalate (Bk2(C2O4)3·4H2O). Thermal decomposition at about 600 °C in an argon atmosphere (to avoid oxidation to BkO2) of Bk2(SO4)3·12H2O yields the crystals of berkelium(III) oxysulfate (Bk2O2SO4). This compound is thermally stable to at least 1000 °C in inert atmosphere. === Organoberkelium compounds === Berkelium forms a trigonal (η5–C5H5)3Bk metallocene complex with three cyclopentadienyl rings, which can be synthesized by reacting berkelium(III) chloride with the molten beryllocene (Be(C5H5)2) at about 70 °C. It has an amber color and a density of 2.47 g/cm3. The complex is stable to heating to at least 250 °C, and sublimates without melting at about 350 °C. The high radioactivity of berkelium gradually destroys the compound (within a period of weeks). One cyclopentadienyl ring in (η5–C5H5)3Bk can be substituted by chlorine to yield [Bk(C5H5)2Cl]2. The optical absorption spectra of this compound are very similar to those of (η5–C5H5)3Bk. Berkelium also forms Berkelocene, an actinocene complex, with substituted cyclooctatetraenides. == Applications == There is currently no use for any isotope of berkelium outside basic scientific research. Berkelium-249 is a common target nuclide to prepare still heavier transuranium elements and superheavy elements, such as lawrencium, rutherfordium and bohrium. It is also useful as a source of the isotope californium-249, which is used for studies on the chemistry of californium in preference to the more radioactive californium-252 that is produced in neutron bombardment facilities such as the HFIR. A 22 milligram batch of berkelium-249 was prepared in a 250-day irradiation and then purified for 90 days at Oak Ridge in 2009. This target yielded the first 6 atoms of tennessine at the Joint Institute for Nuclear
|
{
"page_id": 3758,
"source": null,
"title": "Berkelium"
}
|
Research (JINR), Dubna, Russia, after bombarding it with calcium ions in the U400 cyclotron for 150 days. This synthesis was a culmination of the Russia-US collaboration between JINR and Lawrence Livermore National Laboratory on the synthesis of elements 113 to 118 which was initiated in 1989. == Nuclear fuel cycle == The nuclear fission properties of berkelium are different from those of the neighboring actinides curium and californium, and they suggest berkelium to perform poorly as a fuel in a nuclear reactor. Specifically, berkelium-249 has a moderately large neutron capture cross section of 710 barns for thermal neutrons, 1200 barns resonance integral, but very low fission cross section for thermal neutrons. In a thermal reactor, much of it will therefore be converted to berkelium-250 which quickly decays to californium-250. In principle, berkelium-249 can sustain a nuclear chain reaction in a fast breeder reactor. Its critical mass is relatively high at 192 kg, which can be reduced with a water or steel reflector but would still exceed the world production of this isotope. Berkelium-247 can maintain a chain reaction both in a thermal-neutron and in a fast-neutron reactor, however, its production is rather complex and thus the availability is much lower than its critical mass, which is about 75.7 kg for a bare sphere, 41.2 kg with a water reflector and 35.2 kg with a steel reflector (30 cm thickness). == Health issues == Little is known about the effects of berkelium on human body, and analogies with other elements may not be drawn because of different radiation products (electrons for berkelium and alpha particles, neutrons, or both for most other actinides). The low energy of electrons emitted from berkelium-249 (less than 126 keV) hinders its detection, due to signal interference with other decay processes, but also makes this isotope relatively
|
{
"page_id": 3758,
"source": null,
"title": "Berkelium"
}
|
harmless to humans as compared to other actinides. However, berkelium-249 transforms with a half-life of only 330 days to the strong alpha-emitter californium-249, which is rather dangerous and has to be handled in a glovebox in a dedicated laboratory. Most available berkelium toxicity data originate from research on animals. Upon ingestion by rats, only about 0.01% of berkelium ends in the blood stream. From there, about 65% goes to the bones, where it remains for about 50 years, 25% to the lungs (biological half-life about 20 years), 0.035% to the testicles or 0.01% to the ovaries where berkelium stays indefinitely. The balance of about 10% is excreted. In all these organs berkelium might promote cancer, and in the skeleton, its radiation can damage red blood cells. The maximum permissible amount of berkelium-249 in the human skeleton is 0.4 nanograms. == References == == Bibliography == Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. ISBN 978-0-08-037941-8. Holleman, Arnold F.; Wiberg, Nils (2007). Textbook of Inorganic Chemistry (102nd ed.). Berlin: de Gruyter. ISBN 978-3-11-017770-1. Peterson, J. R.; Hobart, D. E. (1984). "The Chemistry of Berkelium". In Emeléus, Harry Julius (ed.). Advances in inorganic chemistry and radiochemistry. Vol. 28. Academic Press. pp. 29–64. doi:10.1016/S0898-8838(08)60204-4. ISBN 978-0-12-023628-2. == External links == Berkelium at The Periodic Table of Videos (University of Nottingham)
|
{
"page_id": 3758,
"source": null,
"title": "Berkelium"
}
|
Caterina Vitale (1566–1619) was the first female pharmacist and chemist in Malta, and the first female pharmacist of the Knights Hospitaller. Caterina Vitale was originally from Greece. She married Ettore Vitale, pharmacist of the Knights Hospitaller, when she was a teenager. Upon his death in 1590, she inherited his pharmacy and the task of providing pharmacies to the Sacra Infermeria. She was described as a successful businessperson, became very rich, and is known as a benefactor of the Carmelites. Being in an uncommon position for a woman, she was a controversial person and the subjects of legends, libelous slander and rumours, and was accused of being an enterprising prostitute, litigator and sadistic torturer of slaves. She died in 1619 at Syracuse and her body was brought to Valletta and buried at the Carmelite Church. To the left and right as you enter the church are her tombstone, and the tombstone of Caterina Scappi, the founder of the first hospital for women in Malta. Upon her death she bequeathed part of her fortune and her property Selmun Palace to the Monte della Redenzione degli Schiavi, a charity founded during the reign of Grand Master Alof de Wignacourt in 1607 to finance the redemption of the Maltese who had fallen into slavery. She also left legacies to the Order of Malta, to her niece, to the Carmelites, to the Greek Church, but left nothing to her daughter. == See also == Timeline of women in science == Further reading == The nuns who lived off the proceeds of prostitution == References ==
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{
"page_id": 52104883,
"source": null,
"title": "Caterina Vitale"
}
|
The molecular formula C36H44N4O8 (molar mass: 660.75 g/mol, exact mass: 660.315914 u) may refer to: Coproporphyrinogen I Coproporphyrinogen III
|
{
"page_id": 24121012,
"source": null,
"title": "C36H44N4O8"
}
|
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