Ianbruceite

Ianbruceite is a rare hydrated zinc arsenate with the formula [Zn2(OH)(H2O)(AsO4)](H2O)2; material from the Driggith mine has traces of cobalt. It was first discovered at Tsumeb, approved by the International Mineralogical Association as a new mineral species in 2011, reference IMA2011-49, and named for Ian Bruce, who founded "Crystal Classics" in the early 1990s, and was heavily involved in attempts to reopen the famous Tsumeb mine for specimen mining.
In 2013 new occurrences of ianbruceite were reported from the neighbouring Driggith and Potts Gill mines on High Pike in the Caldbeck Fells, Cumbria, England. Here the mineral is probably a post-mining product. Caldbeck Fells and Tsumeb are the only reported localities for ianbruceite to date (May 2013).
Structure.
There are four formula units per unit cell (Z = 4). Unit cell parameters are the lengths of the sides of the unit cell, a, b and c, and the angle β between the a and c crystal axes. For ianbruceite a = 11.793 Å, b = 9.1138 Å, c = 6.8265 Å, and β = 103.859°. It belongs in the monoclinic, prismatic crystal Class 2/m, with space group P21/c.
Ianbruceite consists of layers with composition [Zn2(OH)(H2O)(AsO4)] alternating with an inter-layer space containing loosely bound water molecules. The arsenic is tetrahedrally coordinated by four oxygen anions.
Appearance.
Ianbruceite occurs as aggregates of tiny, radiating, lath-like crystals with a distinctive diamond-shaped outline, up to 0.2 mm across. Material from Tsumeb is blue, ranging from a very pale color to deep blue. Material from the Caldbeck Fells is white to pale pink. The crystals have a vitreous luster and a white streak.
Optical properties.
The refractive indices of ianbruceite are similar to that of ordinary window glass, with nα = 1.601, nβ = 1.660 and nγ = 1.662. The refractive index varies with the wavelength (color) of light, so the positions of the optic axes in biaxial crystals, and the angle 2V between them, will change when the color of the incident light is changed. This effect may be expressed in the form r > v, indicating that the angle 2V is greater for red than for violet light, or vice versa. For ianbruceite 2V is greater for violet light than for red light, v > r. The angle 2V has been measured as 18°; it can also be calculated from the values of the refractive indices α, β and γ, giving a value of 20°. The mineral does not fluoresce in ultraviolet light.
Physical properties.
Cleavage is perfect parallel to (100). Crystals of ianbruceite are flexible and deform plastically. The mineral is very soft, with Mohs hardness only 1, the same as talc. Measurements of the specific gravity have not been reported, but from the formula and the cell dimensions it has been calculated as 3.197. It is insoluble in dilute acid, which distinguishes it from hydrozincite and aragonite, which are also supergene minerals that form white crusts. The loosely bound water between the layers is easily lost, so the mineral may effloresce. It also readily absorbs alcohol.
Occurrence.
The type locality is the Tsumeb Mine, Otjikoto Region, Namibia. Type material is conserved in the Department of Natural History, Royal Ontario Museum, Toronto, Ontario, Canada, catalogue number M53150.
At localities in the Caldbeck Fells ianbruceite occurs as a late-stage supergene mineral in fractures in partly oxidised sulfide-rich carbonate or quartz-carbonate matrix. The mines of Caldbeck Fells are famous for their supergene minerals.
Driggith and Potts Gill Mines worked low temperature lead-zinc-copper veins on the eastern and northern slopes of High Pike. Arsenopyrite is abundant there, and its oxidation in veins that also contain primary lead, zinc and copper sulfides has produced a range of supergene arsenates.
At the Driggith Mine, Caldbeck Fells, Cumbria, on dumps from the 30 fathom level, radiating aggregates of lath-like crystals of ianbruceite have been found in fractures in quartz-dolomite matrix with sphalerite, chalcopyrite and cobalt-bearing köttigite as rounded pink aggregates, or more rarely pale pink to colorless monoclinic blades, as well as irregular black patches of a cobalt-bearing manganese oxide, adamite and an unidentified copper silicate.
At the nearby Potts Gill Mine, Caldbeck Fells, Cumbria, on dumps from the Endeavor level, it has been found as minute, pearly, lath-like crystals in fractures in arsenopyrite-rich dolomite. One block of material contained abundant marcasite, galena, sphalerite and minor fine-grained arsenopyrite, in dolomitic matrix.
In Namibia ianbruceite has been found in the zinc pocket at the 44 Level, Tsumeb Mine, as sky blue to very pale blue platy crystals associated with leiteite, köttigite, legrandite and adamite. It occurs as thin platy crystals up to 80 μm long and a few μm thick, which form flattened aggregates up to 0.10 mm across, and ellipsoidal aggregates up to 0.5 mm across, associated with coarse white leiteite, dark blue köttigite, minor legrandite and adamite.


Untribium


Arsenic pentachloride

Arsenic pentachloride is a chemical compound of arsenic and chlorine. This compound was first prepared in 1976 through the UV irradiation of arsenic trichloride, AsCl3, in liquid chlorine at −105 °C. AsCl5 decomposes at around −50 °C. The structure of the solid was finally determined in 2001. AsCl5 is similar to phosphorus pentachloride, PCl5 in having a trigonal bipyramidal structure where the equatorial bonds are shorter than the axial bonds (As-Cleq = 210.6 pm, 211.9 pm; As-Clax= 220.7 pm).
The pentachlorides of the elements above and below arsenic in group 15, phosphorus pentachloride and antimony pentachloride are much more stable and the instability of AsCl5 appears anomalous. The cause is believed to be due to incomplete shielding of the nucleus in the 4p elements following the first transition series (i.e. gallium, germanium, arsenic, selenium, bromine, and krypton) which leads to stabilisation of their 4s electrons making them less available for bonding. This effect has been termed the d-block contraction and is similar to the f-block contraction normally termed the lanthanide contraction.


Potassium formate

Potassium formate, HCO2K, HCOOK, or CHKO2, is the potassium salt of formic acid. This strongly hygroscopic white solid is an intermediate in the formate potash process for the production of potassium. Potassium formate has also been studied as a potential environmentally friendly deicing salt for use on roads. It has also been suggested for use in a less corrosive liquid desiccant. A 52% solution of potassium formate has a freezing point of .


Deuterated benzene

Deuterated benzene (C6D6) is an isotopologue of benzene (C6H6) in which the hydrogen atom ("H") is replaced with deuterium (heavy hydrogen) isotope ("D"). Deuterated benzene is a common solvent used in NMR spectroscopy.


Octadecyltrimethoxysilane

Octadecyltrimethoxysilane (OTMS) is an organosilicon compound. This colorless liquid is used for preparing hydrophobic coatings and self-assembled monolayers. It is sensitive toward water, irreversibly degrading to a siloxane polymer. It places a C18H39SiO3 "cap" on oxide surfaces. The formation of OTMS monolayers is used for converting hydrophilic surfaces to hydrophobic surfaces, e.g. for use in certain areas of nanotechnology and analytical chemistry.


Organofluorine chemistry

Organofluorine chemistry describes the chemistry of the organofluorines, organic compounds that contain the carbon–fluorine bond. Organofluorine compounds find diverse applications ranging from oil and water repellents to pharmaceuticals, refrigerants, and reagents in catalysis. In addition to these applications, some organofluorine compounds are pollutants because of their contributions to ozone depletion, global warming, bioaccumulation, and toxicity. The area of organofluorine chemistry often requires special techniques associated with the handling of fluorinating agents.
The carbon–fluorine bond.
Fluorine has several distinctive differences from all other substituents encountered in organic molecules. As a result, the physical and chemical properties of organofluorines can be distinctive in comparison to other organohalogens.
In comparison to aryl chlorides and bromides, aryl fluorides form Grignard reagents only reluctantly. On the other hand, aryl fluorides, e.g. fluoroanilines and fluorophenols, often undergo nucleophilic substitution efficiently.
Types of organofluorine compounds.
Fluorocarbons.
Formally, fluorocarbons only contain carbon and fluorine. Sometimes they are called perfluorocarbons. They can be gases, liquids, waxes, or solids, depending upon their molecular weight. The simplest fluorocarbon is the gas tetrafluoromethane (CF4). Liquids include perfluorooctane and perfluorodecalin. While fluorocarbons with single bonds are stable, unsaturated fluorocarbons are more reactive, especially those with triple bonds. Fluorocarbons are more chemically and thermally stable than hydrocarbons, reflecting the relative inertness of the C-F bond. They are also relatively lipophobic. Because of the reduced intermolecular van der Waals interactions, fluorocarbon-based compounds are sometimes used as lubricants or are highly volatile. Fluorocarbon liquids have medical applications as oxygen carriers.
The structure of organofluorine compounds can be distinctive. As shown below, perfluorinated aliphatic compounds tend to segregate from hydrocarbons. This "like dissolves like effect" is related to the usefulness of fluorous phases and the use of PFOA in processing of fluoropolymers. In contrast to the aliphatic derivatives, perfluoroaromatic derivatives tend to form mixed phases with nonfluorinated aromatic compounds, resulting from donor-acceptor interactions between the pi-systems.
Fluoropolymers.
Polymeric organofluorine compounds are numerous and commercially significant. They range from fully fluorinated species, e.g. PTFE to partially fluorinated, e.g. polyvinylidene fluoride ([CH2CF2]n) and polychlorotrifluoroethylene ([CFClCF2]n). The fluoropolymer polytetrafluoroethylene (PTFE/Teflon) is a solid.
Hydrofluorocarbons.
Hydrofluorocarbons (HFCs), organic compounds that contain fluorine and hydrogen atoms, are the most common type of organofluorine compounds. They are commonly used in air conditioning and as refrigerants in place of the older chlorofluorocarbons such as R-12 and hydrochlorofluorocarbons such as R-21. They do not harm the ozone layer as much as the compounds they replace; however, they do contribute to global warming. Their atmospheric concentrations and contribution to anthropogenic greenhouse gas emissions are rapidly increasing, causing international concern about their radiative forcing.
Fluorocarbons with few C-F bonds behave similarly to the parent hydrocarbons, but their reactivity can be altered significantly. For example, both uracil and 5-fluorouracil are colourless, high-melting crystalline solids, but the latter is a potent anti-cancer drug. The use of the C-F bond in pharmaceuticals is predicated on this altered reactivity. Several drugs and agrochemicals contain only one fluorine center or one trifluoromethyl group.
Unlike other greenhouse gases in the Paris Agreement, hydrofluorocarbons have other international negotiations.
In September 2016, the so-called New York Declaration urged a global reduction in the use of HFCs. On 15 October 2016, due to these chemicals contribution to climate change, negotiators from 197 nations meeting at the summit of the United Nations Environment Programme in Kigali, Rwanda reached a legally-binding accord to phase out hydrofluorocarbons (HFCs) in an amendment to the Montreal Protocol.
Fluorocarbenes.
As indicated throughout this article, fluorine-substituents lead to reactivity that differs strongly from classical organic chemistry. The premier example is difluorocarbene, CF2, which is a singlet whereas carbene (CH2) has a triplet ground state. This difference is significant because difluorocarbene is a precursor to tetrafluoroethylene.
Perfluorinated compounds.
Perfluorinated compounds are fluorocarbon derivatives, as they are closely structurally related to fluorocarbons. However, they also possess new atoms such as nitrogen, iodine, or ionic groups, such as perfluorinated carboxylic acids.
Methods for preparation of C–F bonds.
Organofluorine compounds are prepared by numerous routes, depending on the degree and regiochemistry of fluorination sought and the nature of the precursors. The direct fluorination of hydrocarbons with F2, often diluted with N2, is useful for highly fluorinated compounds:
Such reactions however are often unselective and require care because hydrocarbons can uncontrollably "burn" in , analogous to the combustion of hydrocarbon in . For this reason, alternative fluorination methodologies have been developed. Generally, such methods are classified into two classes.
Electrophilic fluorination.
Electrophilic fluorination rely on sources of "F+". Often such reagents feature N-F bonds, for example F-TEDA-BF4. Asymmetric fluorination, whereby only one of two possible enantiomeric products are generated from a prochiral substrate, rely on electrophilic fluorination reagents. Illustrative of this approach is the preparation of a precursor to anti-inflammatory agents:
Electrosynthetic methods.
A specialized but important method of electrophilic fluorination involves electrosynthesis. The method is mainly used to perfluorinate, i.e. replace all C–H bonds by C–F bonds. The hydrocarbon is dissolved or suspended in liquid HF, and the mixture is electrolyzed at 5–6 V using Ni anodes. The method was first demonstrated with the preparation of perfluoropyridine () from pyridine (). Several variations of this technique have been described, including the use of molten potassium bifluoride or organic solvents.
Nucleophilic fluorination.
The major alternative to electrophilic fluorination is, naturally, nucleophilic fluorination using reagents that are sources of "F−," for Nucleophilic displacement typically of chloride and bromide. Metathesis reactions employing alkali metal fluorides are the simplest.
Alkyl monofluorides can be obtained from alcohols and Olah reagent (pyridinium fluoride) or another fluoridating agents.
The decomposition of aryldiazonium tetrafluoroborates in the Sandmeyer or Schiemann reactions exploit fluoroborates as F− sources.
Although hydrogen fluoride may appear to be an unlikely nucleophile, it is the most common source of fluoride in the synthesis of organofluorine compounds. Such reactions are often catalysed by metal fluorides such as chromium trifluoride. 1,1,1,2-Tetrafluoroethane, a replacement for CFC's, is prepared industrially using this approach:
Notice that this transformation entails two reaction types, metathesis (replacement of Cl− by F−) and hydrofluorination of an alkene.
Deoxofluorination.
Deoxofluorination agents effect the replacement hydroxyl and carbonyl groups with one and two fluorides, respectively. One such reagent, useful for fluoride for oxide exchange in carbonyl compounds, is sulfur tetrafluoride:
Alternates to SF4 include the diethylaminosulfur trifluoride (DAST, NEt2SF3) and bis(2-methoxyethyl)aminosulfur trifluoride (deoxo-fluor). These organic reagents are easier to handle and more selective:
From fluorinated building blocks.
Many organofluorine compounds are generated from reagents that deliver perfluoroalkyl and perfluoroaryl groups. (Trifluoromethyl)trimethylsilane, CF3Si(CH3)3, is used as a source of the trifluoromethyl group, for example. Among the available fluorinated building blocks are CF3X (X = Br, I), C6F5Br, and C3F7I. These species form Grignard reagents that then can be treated with a variety of electrophiles. The development of fluorous technologies (see below, under solvents) is leading to the development of reagents for the introduction of "fluorous tails".
A special but significant application of the fluorinated building block approach is the synthesis of tetrafluoroethylene, which is produced on a large-scale industrially via the intermediacy of difluorocarbene. The process begins with the thermal (600-800 °C) dehydrochlorination of chlorodifluoromethane:
Sodium fluorodichloroacetate (CAS# 2837-90-3) is used to generate chlorofluorocarbene, for cyclopropanations.
18F-Delivery methods.
The usefulness of fluorine-containing radiopharmaceuticals in 18F-positron emission tomography has motivated the development of new methods for forming C–F bonds. Because of the short half-life of 18F, these syntheses must be highly efficient, rapid, and easy. Illustrative of the methods is the preparation of fluoride-modified glucose by displacement of a triflate by a labeled fluoride nucleophile:
Biological role.
Biologically synthesized organofluorines have been found in microorganisms and plants, but not animals. The most common example is fluoroacetate, which occurs as a plant defence against herbivores in at least 40 plants in Australia, Brazil and Africa. Other biologically synthesized organofluorines include ω-fluoro fatty acids, fluoroacetone, and 2-fluorocitrate which are all believed to be biosynthesized in biochemical pathways from the intermediate fluoroacetaldehyde. Adenosyl-fluoride synthase is an enzyme capable of biologically synthesizing the carbon–fluorine bond. Man made carbon–fluorine bonds are commonly found in pharmaceuticals and agrichemicals because it adds stability to the carbon framework; also, the relatively small size of fluorine is convenient as fluorine acts as an approximate bioisostere of the hydrogen. Introducing the carbon–fluorine bond to organic compounds is the major challenge for medicinal chemists using organofluorine chemistry, as the carbon–fluorine bond increases the probability of having a successful drug by about a factor of ten. An estimated 20% of pharmaceuticals, and 30–40% of agrichemicals are organofluorines, including several of the top drugs. Examples include 5-fluorouracil, fluoxetine (Prozac), paroxetine (Paxil), ciprofloxacin (Cipro), mefloquine, and fluconazole.
Applications.
Organofluorine chemistry impacts many areas of everyday life and technology. The C-F bond is found in pharmaceuticals, agrichemicals, fluoropolymers, refrigerants, surfactants, anesthetics, oil-repellents, catalysis, and water-repellents, among others.
Pharmaceuticals and agrochemicals.
The carbon-fluorine bond is commonly found in pharmaceuticals and agrochemicals because it is generally metabolically stable and fluorine acts as a bioisostere of the hydrogen atom. An estimated one fifth of pharmaceuticals contain fluorine, including several of the top drugs. Examples include 5-fluorouracil, flunitrazepam (Rohypnol), fluoxetine (Prozac), paroxetine (Paxil), ciprofloxacin (Cipro), mefloquine, and fluconazole. Fluorine-substituted ethers are volatile anesthetics, including the commercial products methoxyflurane, enflurane, isoflurane, sevoflurane and desflurane. Fluorocarbon anesthetics reduce the hazard of flammability with diethyl ether and cyclopropane. Perfluorinated alkanes are used as blood substitutes.
Inhaler propellant.
Fluorocarbons are also used as a propellant for metered-dose inhalers used to administer some asthma medications. The current generation of propellant consists of hydrofluoroalkanes (HFA), which has replaced CFC-propellant-based inhalers. CFC inhalers were banned as part of the Montreal Protocol because of environmental concerns with the ozone layer. HFA propellant inhalers like FloVent and ProAir ( Salbutamol ) have no generic versions available as of October 2014.
Fluorosurfactants.
Fluorosurfactants, which have a polyfluorinated "tail" and a hydrophilic "head", serve as surfactants because they concentrate at the liquid-air interface due to their lipophobicity. Fluorosurfactants have low surface energies and dramatically lower surface tension. The fluorosurfactants perfluorooctanesulfonic acid (PFOS) and perfluorooctanoic acid (PFOA) are two of the most studied because of their ubiquity, toxicity, and long residence times in humans and wildlife.
Solvents.
Fluorinated compounds often display distinct solubility properties. Dichlorodifluoromethane and chlorodifluoromethane were widely used refrigerants. CFCs have potent ozone depletion potential due to the homolytic cleavage of the carbon-chlorine bonds; their use is largely prohibited by the Montreal Protocol. Hydrofluorocarbons (HFCs), such as tetrafluoroethane, serve as CFC replacements because they do not catalyze ozone depletion. Oxygen exhibits a high solubility in perfluorocarbon compounds, reflecting again on their lipophilicity. Perfluorodecalin has been demonstrated as a blood substitutes, transporting oxygen to the lungs.
The solvent 1,1,1,2-tetrafluoroethane has been used for extraction of natural products such as taxol, evening primrose oil, and vanillin. 2,2,2-trifluoroethanol is an oxidation-resistant polar solvent.
Organofluorine reagents.
The development of organofluorine chemistry has contributed many reagents of value beyond organofluorine chemistry. Triflic acid (CF3SO3H) and trifluoroacetic acid (CF3CO2H) are useful throughout organic synthesis. Their strong acidity is attributed to the electronegativity of the trifluoromethyl group that stabilizes the negative charge. The triflate-group (the conjugate base of the triflic acid) is a good leaving group in substitution reactions.
Fluorous phases.
Of topical interest in the area of "Green Chemistry," highly fluorinated substituents, e.g. perfluorohexyl (C6F13) confer distinctive solubility properties to molecules, which facilitates purification of products in organic synthesis. This area, described as "fluorous chemistry," exploits the concept of like-dissolves-like in the sense that fluorine-rich compounds dissolve preferentially in fluorine-rich solvents. Because of the relative inertness of the C-F bond, such fluorous phases are compatible with even harsh reagents. This theme has spawned techniques of "fluorous tagging" and "fluorous protection". Illustrative of fluorous technology is the use of fluoroalkyl-substituted tin hydrides for reductions, the products being easily separated from the spent tin reagent by extraction using fluorinated solvents.
Hydrophobic fluorinated ionic liquids, such as organic salts of bistriflimide or hexafluorophosphate, can form phases that are insoluble in both water and organic solvents, producing multiphasic liquids.
Organofluorine ligands in transition metal chemistry.
Organofluorine ligands have long been featured in organometallic and coordination chemistry. One advantage to F-containing ligands is the convenience of 19F NMR spectroscopy for monitoring reactions. The organofluorine compounds can serve as a "sigma-donor ligand," as illustrated by the titanium(III) derivative [(C5Me5)2Ti(FC6H5)]BPh4. Most often, however, fluorocarbon substituents are used to enhance the Lewis acidity of metal centers. A premier example is "Eufod," a coordination complex of europium(III) that features a perfluoroheptyl modified acetylacetonate ligand. This and related species are useful in organic synthesis and as "shift reagents" in NMR spectroscopy.
In an area where coordination chemistry and materials science overlap, the fluorination of organic ligands is used to tune the properties of component molecules. For example, the degree and regiochemistry of fluorination of metalated 2-phenylpyridine ligands in platinum(II) complexes significantly modifies the emission properties of the complexes.
The coordination chemistry of organofluorine ligands also embraces fluorous technologies. For example, triphenylphosphine has been modified by attachment of perfluoroalkyl substituents that confer solubility in perfluorohexane as well as supercritical carbon dioxide. As a specific example, [(C8F17C3H6-4-C6H4)3P.
C-F bond activation.
An active area of organometallic chemistry encompasses the scission of C-F bonds by transition metal-based reagents. Both stoichiometric and catalytic reactions have been developed and are of interest from the perspectives of organic synthesis and remediation of xenochemicals. C-F bond activation has been classified as follows "(i) oxidative addition of fluorocarbon, (ii) M–C bond formation with HF elimination, (iii) M–C bond formation with fluorosilane elimination, (iv) hydrodefluorination of fluorocarbon with M–F bond formation, (v) nucleophilic attack on fluorocarbon, and (vi) defluorination of fluorocarbon". An illustrative metal-mediated C-F activation reaction is the defluorination of fluorohexane by a zirconium dihydride, an analogue of Schwartz's reagent:
Fluorocarbon anions in Ziegler-Natta catalysis.
Fluorine-containing compounds are often featured in noncoordinating or weakly coordinating anions. Both tetrakis(pentafluorophenyl)borate, B(C6F5)4−, and the related tetrakis[3,5-bis(trifluoromethyl)phenyl]borate, are useful in Ziegler-Natta catalysis and related alkene polymerization methodologies. The fluorinated substituents render the anions weakly basic and enhance the solubility in weakly basic solvents, which are compatible with strong Lewis acids.
Materials science.
Organofluorine compounds enjoy many niche applications in materials science. With a low coefficient of friction, fluid fluoropolymers are used as specialty lubricants. Fluorocarbon-based greases are used in demanding applications. Representative products include Fomblin and Krytox, manufactured by Solvay Solexis and DuPont, respectively. Certain firearm lubricants such as "Tetra Gun" contain fluorocarbons. Capitalizing on their nonflammability, fluorocarbons are used in fire fighting foam. Organofluorine compounds are components of liquid crystal displays. The polymeric analogue of triflic acid, nafion is a solid acid that is used as the membrane in most low temperature fuel cells. The bifunctional monomer 4,4'-difluorobenzophenone is a precursor to PEEK-class polymers.
Biosynthesis of organofluorine compounds.
In contrast to the many naturally-occurring organic compounds containing the heavier halides, chloride, bromide, and iodide, only a handful of biologically synthesized carbon-fluorine bonds are known. The most common natural organofluorine species is fluoroacetate, a toxin found in a few species of plants. Others include fluorooleic acid, fluoroacetone, nucleocidin (4'-fuoro-5'-O-sulfamoyladenosine), fluorothreonine, and 2-fluorocitrate. Several of these species are probably biosynthesized from fluoroacetaldehyde. The enzyme fluorinase catalyzed the synthesis of 5'-fluoro-5-deoxyadenosine (see scheme to right).
History.
Organofluorine chemistry began in the 1800s with the development of organic chemistry. The first organofluorine compounds were prepared using antimony trifluoride as the F− source. The nonflammability and nontoxicity of the chlorofluorocarbons CCl3F and CCl2F2 attracted industrial attention in the 1920s. on April 6, 1938, Roy J. Plunkett a young research chemist who worked at DuPont's Jackson Laboratory in Deepwater, New Jersey, accidentally discovered polytetrafluoroethylene (PTFE). Subsequent major developments, especially in the US, benefited from expertise gained in the production of uranium hexafluoride.
Starting in the late 1940s, a series of electrophilic fluorinating methodologies were introduced, beginning with CoF3. Electrochemical fluorination ("electrofluorination") was announced, which Joseph H. Simons had developed in the 1930s to generate highly stable perfluorinated materials compatible with uranium hexafluoride. These new methodologies allowed the synthesis of C-F bonds without using elemental fluorine and without relying on metathetical methods.
In 1957, the anticancer activity of 5-fluorouracil was described. This report provided one of the first examples of rational design of drugs. This discovery sparked a surge of interest in fluorinated pharmaceuticals and agrichemicals. The discovery of the noble gas compounds, e.g. XeF4, provided a host of new reagents starting in the early 1960s. In the 1970s, fluorodeoxyglucose was established as a useful reagent in 18F positron emission tomography. In Nobel Prize-winning work, CFC's were shown to contribute to the depletion of atmospheric ozone. This discovery alerted the world to the negative consequences of organofluorine compounds and motivated the development of new routes to organofluorine compounds. In 2002, the first C-F bond-forming enzyme, fluorinase, was reported.
Environmental and health concerns.
Only a few organofluorine compounds are acutely bioactive and highly toxic, such as fluoroacetate and perfluoroisobutene.
Some organofluorine compounds pose significant risks and dangers to health and the environment. CFCs and HCFCs (hydrochlorofluorocarbon) deplete the ozone layer and are potent greenhouse gases. HFCs are potent greenhouse gases and are facing calls for stricter international regulation and phase out schedules as a fast-acting greenhouse emission abatement measure, as are perfluorocarbons (PFCs), and sulphur hexafluoride (SF6).
Because of the compound's effect on climate, the G-20 major economies agreed in 2013 to support initiatives to phase out use of HCFCs. They affirmed the roles of the Montreal Protocol and the United Nations Framework Convention on Climate Change in global HCFC accounting and reduction. The U.S. and China at the same time announced a bilateral agreement to similar effect.
Persistence and bioaccumulation.
Because of the strength of the carbon–fluorine bond, many synthetic fluorocarbons and fluorcarbon-based compounds are persistent in the environment. Fluorosurfactants, such as PFOS and PFOA, are persistent global contaminants. Fluorocarbon based CFCs and tetrafluoromethane have been reported in igneous and metamorphic rock. PFOS is a persistent organic pollutant and may be harming the health of wildlife; the potential health effects of PFOA to humans are under investigation by the C8 Science Panel.


Isotopes of platinum

Naturally occurring platinum (78Pt) consists of five stable isotopes (192Pt, 194Pt, 195Pt, 196Pt, 198Pt) and one very long-lived (half-life 6.50×1011 years) radioisotope (190Pt). There are also 34 known synthetic radioisotopes, the longest-lived of which is 193Pt with a half-life of 50 years. All other isotopes have half-lives under a year, most under a day.


Bararite

Bararite is a natural form of ammonium fluorosilicate (also known as hexafluorosilicate or fluosilicate). It has chemical formula (NH4)2SiF6 and trigonal crystal structure. This mineral was once classified as part of cryptohalite. Bararite is named after the place where it was first described, Barari, India. It is found at the fumaroles of volcanoes (Vesuvius, Italy), over burning coal seams (Barari, India), and in burning piles of anthracite (Pennsylvania, U.S.). It is a sublimation product that forms with cryptohalite, sal ammoniac, and native sulfur.
History.
A. Scacchi first discovered cryptohalite in 1873. It appeared in a volcanic sublimate from the Vesuvian eruption of 1850. In 1926, W.A.K. Christie reported his own chemical study. A microscope was used to pick out enough material for analysis. Distilling with sodium hydroxide (NaOH) produced ammonia (NH3). The anions of hexafluorosilicic acid (H2SiF6) precipitated as potassium fluorosilicate (K2SiF6). Barium sulfate (BaSO4) was thrown into the filtrate, and then calcium fluoride (CaF2). Christie found 20.43% (NH4)+ and 78.87% (SiF6)2−.
Bararite is named after Barari, a locality in India. This was where the species was first completely described. Earlier, bararite was recognized as part of mixtures with cryptohalite. However, it did not receive its own name until 1951. The East Indian Coal Company provided the sample that Christie used to evaluate bararite.
Bararite has not received a quantitative chemical analysis in its natural form. Christie received far too little for more than "qualitative" analysis through microchemistry. He utilized F. Emich's methods with capillary tube centrifuges.
Structure.
Bararite is the beta, trigonal (scalenohedral) form of ammonium hexafluorosilicate. Its symmetry is 2/m. The space group is Pm1. The "a"-axes in the unit cell are 5.784 ± 0.005 Å (angstroms), and the "c"-axis is 4.796 ± 0.006 Å. The unit lattice is primitive. (Note: Data for the space group come from synthetic crystals.) Cryptohalite has the cubic (isometric) crystal structure and corresponds to the alpha form. Both minerals have the chemical formula (NH4)2SiF6. The halides of form "AmBX"6 fall into two groups: hieratite and malladrite. The hieratite group is isometric whereas the malladrite is hexagonal.
The (SiF6)2− is octahedral—one fluorine atom at each vertex. In bararite, the (NH4)+’s are trigonally coordinated. They all appear at sites of C"3v" (3m) symmetry. The (NH4)+ has 12 fluorine neighbors, which form four triangles. Three of these triangles are isosceles. These triangles themselves form a triangle—around the 3-fold axis containing the nitrogen atom. One triangle is equilateral. Its symmetry axis is the same axis that goes through the nitrogen atom. (For structural diagrams, see link to unit cell and downloadable articles in “References.”)
The silicon atoms of cryptohalite, α-(NH4)2SiF6 (alpha), have cubic close(st) packing (CCP). A third form (gamma, γ) of (NH4)2SiF6 uses hexagonal close(st) packing (HCP). Bararite, β-(NH4)2SiF6, utilizes hexagonal primitive (HP) packing. Layers with distorted octahedral gaps separate those with the anions. The (NH4)+ ions appear a little below and above the (SiF6)2−. In all three phases, 12 fluorine atoms neighbor the (NH4)+. Distances range from about 3.0 to 3.2 Å. The (NH4)+ has no free rotation. It only librates (oscillates)—at least when vibrationally excited.
As a salt, bararite is an ionic compound. The ions, of course, have ionic bonding. The atoms of polyatomic ions are held together covalently. The orientation of (NH4)+ is sustained by four trifurcated (three-branch) hydrogen bonds. These bonds point toward the triangles containing the 12 fluorine neighbors. Three H bonds are equivalent. The fourth bond, pointing toward the equilateral triangle, has a shorter distance.
The intermolecular distances between fluorine atoms are smaller in bararite (3.19 and 3.37 Å) than cryptohalite. In cryptohalite, each anion is coordinated to 12 others. Bararite has (2+6)-fold coordination. The two Si-Si distances between layers (4.796 ± 0.006 Å) do not equal the six within a layer (5.784 ± 0.005 Å). Bararite is more compressible along the "c"-axis than the "a"-axis.
Bararite has no known solution or exsolution, but it is always mixed with other substances (cryptohalite, sal ammoniac, and sulfur). Due to thermal motion, atomic behavior of ammonium salts can be very hard to evaluate. The anions, however, are ordered and have no unusual motion from heat.
A third form of (NH4)2SiF6 was discovered in 2001 and identified with the 6mm symmetry (hexagonal). In all three arrangements, the (SiF6)2− octahedra come in layers. In the cubic form (cryptohalite), these layers are perpendicular to [111]. In the trigonal (bararite) and hexagonal (gamma, γ) forms, the layers are perpendicular to the c-axis. (Note: Trigonal crystals are part of the hexagonal group. But not all hexagonal crystals are trigonal.)
Although bararite was claimed to be metastable at room temperature, it does not appear one polymorph has ever turned into another. Still, bararite is fragile enough that grinding it for spectroscopy will produce a little cryptohalite. Even so, ammonium fluorosilicate assumes a trigonal form at pressures of 0.2 to 0.3 giga-pascals (GPa). The reaction is irreversible. If this phase is not bararite, it is at least very closely related.
The hydrogen bonding in (NH4)2SiF6 allows this salt to change phases in ways that normal salts cannot. Interactions between cations and anions are especially important in how ammonium salts change phase.
Physical properties.
Bararite forms tabular crystals. They are flattened, sometimes elongated, on {0001} (perpendicular to "c"). Christie reported tiny, transparent crystals of bararite that looked like paddlewheels and darts. Each had four barbs at 90°. The crystals reached up to 1 mm long, the barbs up to 0.2 mm wide. They were interpenetration twins, the twin axis perpendicular to the "c"-axis. Visually, cryptohalite crystals are almost impossible to discern from sal ammoniac (NH4Cl). Inclusions of bararite in cryptohalite can be seen only with plane-polarized light.
Bararite has perfect cleavage on the {0001} plane. The hardness is probably . The anions (as already shown) are bonded much more strongly within layers than between layers. Also, ionic bonds are not the strongest bonds, and halides cannot normally scratch glass plates.
Bararite has a measured density of 2.152 g/mL (synthetic)—but a calculated density of 2.144 g/mL. It tastes salty, and it dissolves in water. Its luster is vitreous (like glass). Bararite is white to colorless. These properties are similar to halite (NaCl)—which gave the halide group its name.
Whereas cryptohalite belongs to the isotropic optical class, bararite is uniaxial negative. At 1.391 ± 0.003, the refractive index through "c" is smaller than through "a" (1.406 ± 0.001). The "c"-axis in bararite is shorter than the "a"-axes (see “Structure”). Furthermore, only this path lets light hit nothing but the same ion in the same orientation (all the layers have the same structure and orientation).
Bararite has about a 6% greater density than cryptohalite. As discussed before, its structure is more packed. This substance can be produced easily from aqueous solution, but only below 5 °C (41 °F) will pure bararite form. Above 13 °C (55 °F), almost pure cryptohalite emerges. Bararite sublimes without leaving residue.
Geologic occurrence.
In nature, bararite appears with cryptohalite, sal ammoniac, and native sulfur. It is found over a burning coal seam in Barari, India, and as a sublimation product in Vesuvius, Italy, at fumaroles (opening in or near a volcano where hot sulfurous gases come out). It also is found in the United States, in Pennsylvania. It appears in burning piles of anthracite (highest grade of coal)—again as a sublimation product.
Christie found translucent arborescent (treelike) crystals, with vitreous luster. He found white, opaque lumps that were a mixture of (NH4)2SiF6 with SiO2. They were irregularly shaped but usually had a mammillary surface (several convex surfaces smoothly rounded). These hold primarily cryptohalite but also some bararite. In Pennsylvania, bararite normally comes as tiny inclusions in cryptohalite crystals. It appears that first, bararite forms through direct sublimation. Afterward, it quickly changes to cryptohalite.
In Barari, burning-coal gases go through a dike (igneous intrusion) of mica and peridotite. The sulfur dioxide must attack apatite in the dike, which produces hydrofluoric acid that attacks the abundant silicates. Silicon fluoride is formed. Ammonia also comes from burning coal. From there, ammonium fluorosilicate can form. A slight excess of ammonia could lead to the white lumps of silica and cryptohalite. Bararite and cryptohalite in their pure forms, for the most part, grow out of these nodules. Recrystallization from the rain is probably responsible.
Fluorosilicate minerals are thermodynamically unstable in soil. Still, intense heat promotes the formation of (NH4)2SiF6 to some degree—as seen in some experiments by Rehim. But this compound will break up at 320 to 335 °C. Both burning coal and volcanoes are important sources of SO2 and SiF4.
Chemical properties and uses.
Fluorosilicic acid and its salts are poisonous. Ammonium fluorosilicate, however, is very rare in nature and apparently much easier to synthesize.


Carbonium ion

In chemistry, a carbonium ion is any cation that has a pentavalent carbon atom. The name carbonium may also be used for the simplest member of the class, properly called methanium (), where the five valences are filled with hydrogen atoms.
The next simplest carbonium ions after methanium have two carbon atoms. Ethynium, or protonated acetylene , and ethenium are usually classified in other families. The ethanium ion has been studied as an extremely rarefied gas by infrared spectroscopy. The isomers of octonium (protonated octane, ) have been studied. The carbonium ion has a planar geometry.
In older literature, the name "carbonium ion" was used for what is today called carbenium. The current definitions were proposed by the chemist George Andrew Olah in 1972 and are now widely accepted.
A stable carbonium ion is the complex pentakis(triphenylphosphinegold(I))methanium , produced by Schmidbauer and others.
Preparation.
Carbonium ions can be obtained by treating alkanes with very strong acids. Industrially, they are formed in the refining of petroleum during primary thermal cracking.


Tuperssuatsiaite

Tuperssuatsiaite is a rare clay mineral found in Greenland, Namibia and Brazil. It is a hydrated phyllosilicate (sheet silicate) of sodium and iron.
Discovery.
Tuperssuatsiaite was first found by Karup-Moller and Petersen in Greenland, in 1984, and given the International Mineralogical Association designation IMA1984-002. It was later named after the type locality, Tuperssuatsiat Bay, Ilimaussaq, Greenland. In 1992 Karup-Moller and Petersen, together with von Knorring and Leonardsen, found more specimens from a second find in the Aris Quarry in Namibia, that allowed a better definition of the properties and composition of the mineral. Later still, in 2005, more material was found by a group of researchers from the University of São Paulo at the Bortolan Quarry, Pocos de Caldas, Brazil, and although the crystals were small, accurate determinations were made of their physical and optical properties, which differed slightly from those of the specimens from Greenland and Namibia.
Mineral group.
Tuperssuatsiaite is a member of the palygorskite-sepiolite group, palygorskite subgroup.
Subgroup members (formulae according to the IMA):
Iron occurs both in the ferric state Fe3+ and the ferrous state Fe2+. The formula for tuperssuatsiaite contains only ferric iron, and the Greenland material is consistent with this. Analysis of the Namibian material, however, shows that part of the iron is in the ferrous state. Manganese is also present as a substitute for iron, and a zinc-rich material has been reported from Greenland.
Structure.
The mineral belongs to the monoclinic crystal class 2/m, meaning that it has a twofold axis of rotational symmetry perpendicular to a mirror plane. At one time it was thought that the material from Namibia might belong to the monoclinic class 2, without the mirror plane, but a more recent study gives it as 2/m, the same as the material from Greenland.
The space group is B2/m, meaning that in the unit cell there is one structural unit at each vertex, and one in the centre of each B face. The palygorskite-sepiolite minerals are clay minerals with a layered structure. In tuperssuatsiaite ribbons of SiO4 tetrahedra, similar to those in the amphibole structure, are aligned parallel to the c crystal axis, and they link to form layers parallel to the plane containing the a and b axes. The spacing between the layers, in the c direction, is about 5 Å, which is typical for minerals with an amphibole-type structure, due to the repeat distance along the chains of tetrahedra. Channels occur that could be occupied by H2O as in palygorskite.
Unit cell.
There are two formula units per unit cell (Z = 2), and the cell dimensions vary slightly for specimens from different locations; for all specimens, to the nearest Å, a = 14 Å, b = 18 Å and c = 5 Å, and the angle β = 103° to 105°.
For the three main localities the reported values are:
Appearance.
Tuperssuatsiaite occurs as fan-shaped aggregates up to several centimeters across, as rosettes and as fibers elongated parallel to the c axis. It is red-brown in reflected light, and colorless to light yellowish brown in transmitted light, with a brownish yellow streak. Crystals are transparent with a bright vitreous luster, but aggregates may be dull and translucent.
Optical properties.
The mineral is biaxial (+), with refractive indices Nx ~ 1.54, Ny ~ 1.56 and Nz ~ 1.58 to 1.60.
For the three main localities the reported values are:
It is mildly pleochroic, with X colorless, Y colorless to pale brown or green and Z generally reddish brown. No fluorescence has been observed.
Physical properties.
Cleavage is good on a plane containing the b and c crystal axes, parallel to the layers within the structure, and twinning is common.
Fracture is uneven to conchoidal (shell-like) and the mineral is brittle; it is quite light, with specific gravity 2.465, which is similar to that of quartz.
Type locality.
The type locality is Tuperssuatsiat Bay, Tunugdliarfik Firth (Eriksfjord), Ilimaussaq complex, Narsaq, Kitaa (West Greenland) Province, Greenland, and type material is conserved at the University of Copenhagen, Copenhagen, Denmark, and at the National Museum of Natural History, Washington DC, US, reference number 162402.


Copper(II) carbonate

Copper(II) carbonate or cupric carbonate is a chemical compound with formula . At ambient temperatures, it is an ionic solid (a salt) consisting of copper(II) cations and carbonate anions .
This compound is rarely encountered because it is difficult to prepare and readily reacts with water moisture from the air. The terms "copper carbonate", "copper(II) carbonate", and "cupric carbonate" almost always refer (even in chemistry texts) to a basic copper carbonate (or copper(II) carbonate hydroxide), such as ()2 (which occurs naturally as the mineral malachite) or ()2()2 (azurite). For this reason, the qualifier neutral may be used instead of "basic" to refer specifically to .
Preparation.
Reactions that may be expected to yield , such as mixing solutions of copper(II) sulfate and sodium carbonate in ambient conditions, yield instead a basic carbonate and , due to the great affinity of the ion for the hydroxide anion .
Thermal decomposition of the basic carbonate at atmospheric pressure yields copper(II) oxide rather than the carbonate.
In 1960, C. W. F. T. Pistorius claimed synthesis by heating basic copper carbonate at 180 °C in an atmosphere of carbon dioxide (450 atm) and water (50 atm) for 36 hours. The bulk of the products was well-crystallized malachite ()2, but a small yield of a rhombohedral substance was also obtained, claimed to be . However, this synthesis was apparently not reproduced.
Reliable synthesis of true copper(II) carbonate was reported for the first time in 1973 by Hartmut Ehrhardt and others. The compound was obtained as a gray powder, by heating basic copper carbonate in an atmosphere of carbon dioxide (produced by the decomposition of silver oxalate ) at 500 °C and 2 GPa (20,000 atm). The compound was determined to have a monoclinic structure.
Chemical and physical properties.
The stability of dry depends critically on the partial pressure of carbon dioxide (pCO2). It is stable for months in dry air, but decomposes slowly into and if pCO2 is less than 0.11 atm.
In the presence of water or moist air at 25 °C, is stable only for pCO2 above 4.57 atmospheres and pH between about 4 and 8. Below that partial pressure, it reacts with water to form a basic carbonate (azurite, ()2()2).
In highly basic solutions, the complex anion ()22− is formed instead.
The solubility product of the true copper(II) carbonate was measured by Reiterer and others as pKso = 11.45 ± 0.10 at 25 °C.
Structure.
In the crystal structure of CuCO3, copper adopts a distorted square pyramidal coordination environment with coordination number 5. Each carbonate ion bonds to 5 copper centres.


Manganese(II) oxide

Manganese(II) oxide is an inorganic compound with chemical formula MnO. It forms green crystals. The compound is produced on a large scale as a component of fertilizers and food additives.
Structure, stoichiometry, reactivity.
Like many monoxides, MnO adopts the rock salt structure, where cations and anions are both octahedrally coordinated. Also like many oxides, manganese(II) oxide is often nonstoichiometric: its composition can vary from MnO to MnO1.045.
Below 118 K MnO is antiferromagnetic. MnO has the distinction of being one of the first compounds to have its magnetic structure determined by neutron diffraction, the report appearing in 1951. This study showed that the Mn2+ ions form a face centered cubic magnetic sub-lattice where there are ferromagnetically coupled sheets that are anti-parallel with adjacent sheets.
Manganese(II) oxide undergoes the chemical reactions typical of an ionic oxide. Upon treatment with acids, it converts to the corresponding manganese(II) salt and water. Oxidation of manganese(II) oxide gives manganese(III) oxide.
Preparation and occurrence.
MnO occurs in nature as the rare mineral manganosite. Commercially it is prepared by reduction of MnO2 with hydrogen, carbon monoxide or methane, e.g.:
MnO can also be prepared by decarbonation of the carbonate:
This calcining process is conducted anaerobically, lest Mn2O3 form.
An alternative route, mostly for demonstration purposes, is the oxalate method, which also applicable to the synthesis of ferrous oxide and stannous oxide. Upon heating in an oxygen-free atmosphere (usually CO2), manganese(II) oxalate decomposes into MnO:
Applications.
Together with manganese sulfate, MnO is a component of fertilizers and food additives. Many thousands of tons are consumed annually for this purpose. Other uses include: a catalyst in the manufacture of allyl alcohol, ceramics, paints, colored glass, bleaching tallow and textile printing.


Yttrium(III) oxide

Yttrium oxide, also known as yttria, is Y2O3. It is an air-stable, white solid substance.
The thermal conductivity of yttrium oxide is 27 W/(m·K).
Uses.
Phosphors.
Yttria is widely used to make Eu:YVO4 and Eu:Y2O3 phosphors that give the red color in color TV picture tubes.
Yttria lasers.
Y2O3 is a prospective solid-state laser material. In particular, lasers with ytterbium as dopant allow the efficient operation both in continuous operation
and in pulsed regimes.
At high concentration of excitations (of order of 1%) and poor cooling, the quenching of emission at laser frequency and avalanche broadband emission takes place. (Yttria-based lasers are not to be confused with YAG lasers using yttrium aluminum garnet, a widely used crystal host for rare earth laser dopants).
Gas Lighting.
The original use of the mineral yttria and the purpose of its extraction from mineral sources was as part of the process of making gas mantles and other products for turning the flames of artificially-produced gases (initially hydrogen, later coal gas, paraffin, or other products) into human-visible light. This use is almost obsolete - thorium and cerium oxides are larger components of such products these days.
Dental ceramics.
Yttrium oxide is used to stabilize the Zirconia in late-generation porcelain-free metal-free dental ceramics. This is a very hard ceramic used as a strong base material in some full ceramic restorations. The zirconia used in dentistry is zirconium oxide which has been stabilized with the addition of yttrium oxide. The full name of zirconia used in dentistry is "yttria-stabilized zirconia" or YSZ.
Microwave filters.
Yttrium oxide is also used to make yttrium iron garnets, which are very effective microwave filters.
Superconductors.
Y2O3 is used to make the high temperature superconductor YBa2Cu3O7, known as "1-2-3" to indicate the ratio of the metal constituents:
This synthesis is typically conducted at 800 °C.
Inorganic synthesis.
Yttrium oxide is an important starting point for inorganic compounds. For organometallic chemistry it is converted to YCl3 in a reaction with concentrated hydrochloric acid and ammonium chloride.
Natural occurrence.
Yttriaite-(Y), approved as a new mineral species in 2010, is the natural form of yttria. It is exceedingly rare, occurring as inclusions in native tungsten particles in a placer deposit of the Bol’shaja Pol’ja () river, Prepolar Ural, Siberia. As a chemical component of other minerals, the oxide yttria was first isolated in 1789 by Johan Gadolin, from rare-earth minerals in a mine at the Swedish town of Ytterby, near Stockholm.


Disulfur dichloride

Disulfur dichloride is the inorganic compound of sulfur and chlorine with the formula S2Cl2.
Some alternative names for this compound are "sulfur monochloride" (the name implied by its empirical formula, SCl), "disulphur dichloride" (British English Spelling) and "sulphur monochloride" (British English Spelling). S2Cl2 has the structure implied by the formula Cl−S−S−Cl, wherein the angle between the Cla−S−S and S−S−Clb planes is 90°. This structure is referred to as "gauche", and is akin to that for H2O2. A rare isomer of S2Cl2 is S=SCl2; this isomer forms transiently when S2Cl2 is exposed to UV-radiation (see thiosulfoxides).
Synthesis, basic properties, reactions.
Pure disulfur dichloride is a yellow liquid that "smokes" in moist air due to reaction with water. An idealized (but complicated) equation is:
It is produced by partial chlorination of elemental sulfur. The reaction proceeds at usable rates at room temperature. In the laboratory, chlorine gas is led into a flask containing elemental sulfur. As disulfur dichloride is formed, the contents become a golden yellow liquid:
Excess chlorine produces sulfur dichloride, which causes the liquid to become less yellow and more orange-red:
The reaction is reversible, and upon standing, SCl2 releases chlorine to revert to the disulfur dichloride. Disulfur dichloride has the ability to dissolve large quantities of sulfur, which reflects in part the formation of polysulfanes:
Disulfur dichloride can be purified by distillation from excess elemental sulfur.
S2Cl2 also arises from the chlorination of CS2 as in the synthesis of thiophosgene.
Reactions.
S2Cl2 hydrolyzes to sulfur dioxide and elemental sulfur. When treated with hydrogen sulfide, polysulfanes are formed as indicated in the following idealized formula:
It reacts with ammonia to give heptasulfur imide (S7NH) and related S−N rings S8−x(NH)x (x = 2, 3).
Applications.
S2Cl2 has been used to introduce C−S bonds. In the presence of aluminium chloride (AlCl3), S2Cl2 reacts with benzene to give diphenyl sulfide:
Anilines (1) react with S2Cl2 in the presence of NaOH to give 1,2,3-benzodithiazolium salts (2) (Herz reaction) which can be transformed into "ortho"-aminothiophenolates (3), these species are precursors to thioindigo dyes.
It is also used to prepare sulfur mustard, commonly known as "mustard gas", by reaction with ethylene at 60 °C (the Levinstein process):
Other uses of S2Cl2 include the manufacture of sulfur dyes, insecticides, and synthetic rubbers. It is also used in cold vulcanization of rubbers, as polymerization catalyst for vegetable oils and for hardening soft woods.
Safety and regulation.
This substance is listed in Schedule 3 Part B – Precursor Chemicals of the Chemical Weapons Convention (CWC). Facilities that produce and/or process and/or consume Scheduled chemicals may be subject to control, reporting mechanisms and inspection by the OPCW (Organisation for the Prohibition of Chemical Weapons).


Unbiquadium

Unbiquadium, also known as element 124 or eka-uranium, is the hypothetical chemical element with atomic number 124 and placeholder symbol Ubq. "Unbiquadium" and "Ubq" are the temporary IUPAC name and symbol, respectively, until the element is discovered, confirmed, and a permanent name is decided upon. In the periodic table, unbiquadium is expected to be a g-block superactinide and the sixth element in the 8th period. Unbiquadium has attracted attention, as it may lie within the island of stability, leading to longer half-lives, especially for 308Ubq which is predicted to have a magic number of neutrons (184).
Despite several searches, unbiquadium has not been synthesized, nor have any naturally occurring isotopes been found to exist. It is believed that the synthesis of unbiquadium will be far more challenging than that of lighter undiscovered elements, and nuclear instability may pose further difficulties in identifying unbiquadium, unless the island of stability has a stronger stabilizing effect than predicted in this region.
As a member of the superactinide series, unbiquadium is expected to bear some resemblance to its possible lighter congener uranium. The valence electrons of unbiquadium are expected to participate in chemical reactions fairly easily, though relativistic effects may significantly influence some of its properties; for example, the electron configuration has been calculated to differ considerably from the one predicted by the Aufbau principle.
History.
Synthesis attempts.
Because complete nuclear shells (or, equivalently, a magic number of protons or neutrons) may confer additional stability on the nuclei of superheavy elements, moving closer to the center of the island of stability, it was thought that the synthesis of element 124 or nearby elements would populate longer-lived nuclei within the island. Scientists at GANIL (Grand Accélérateur National d'Ions Lourds) attempted to measure the direct and delayed fission of compound nuclei of elements with "Z" = 114, 120, and 124 in order to probe shell effects in this region and to pinpoint the next spherical proton shell. In 2006, with full results published in 2008, the team provided results from a reaction involving the bombardment of a natural germanium target with uranium ions:
The team reported that they had been able to identify compound nuclei fissioning with half-lives > 10−18 s. This result suggests a strong stabilizing effect at "Z" = 124 and points to the next proton shell at "Z" > 120, not at "Z" = 114 as previously thought. A compound nucleus is a loose combination of nucleons that have not arranged themselves into nuclear shells yet. It has no internal structure and is held together only by the collision forces between the target and projectile nuclei. It is estimated that it requires around 10−14 s for the nucleons to arrange themselves into nuclear shells, at which point the compound nucleus becomes a nuclide, and this number is used by IUPAC as the minimum half-life a claimed isotope must have to potentially be recognised as being discovered. Thus, the GANIL experiments do not count as a discovery of element 124.
The fission of the compound nucleus 312124 was also studied in 2006 at the tandem ALPI heavy-ion accelerator at the Laboratori Nazionali di Legnaro (Legnaro National Laboratories) in Italy:
Similarly to previous experiments conducted at the JINR (Joint Institute for Nuclear Research), fission fragments clustered around doubly magic nuclei such as 132Sn ("Z" = 50, "N" = 82), revealing a tendency for superheavy nuclei to expel such doubly magic nuclei in fission. The average number of neutrons per fission from the 312124 compound nucleus (relative to lighter systems) was also found to increase, confirming that the trend of heavier nuclei emitting more neutrons during fission continues into the superheavy mass region.
Possible natural occurrence.
A study in 1976 by a group of American researchers from several universities proposed that primordial superheavy elements, mainly livermorium, unbiquadium, unbihexium, and unbiseptium, could be a cause of unexplained radiation damage (particularly radiohalos) in minerals. Unbiquadium was then suggested to exist in nature with its possible congener uranium in detectable quantities, at a relative abundance of 10−11. Such unbiquadium nuclei were thought to undergo alpha decay with very long half-lives down to flerovium, which would then exist in natural lead at a similar concentration (10−11) and undergo spontaneous fission. This prompted many researchers to search for them in nature from 1976 to 1983. A group led by Tom Cahill, a professor at the University of California at Davis, claimed in 1976 that they had detected alpha particles and X-rays with the right energies to cause the damage observed, supporting the presence of these elements. Others claimed that none had been detected, and questioned the proposed characteristics of primordial superheavy nuclei. In particular, they cited that the magic number "N" = 228 necessary for enhanced stability would create a neutron-excessive nucleus in unbiquadium that would not be beta-stable. This activity was also proposed to be caused by nuclear transmutations in natural cerium, raising further ambiguity upon this claimed observation of superheavy elements.
The possible extent of primordial superheavy elements on Earth today is uncertain. Even if they are confirmed to have caused the radiation damage long ago, they might now have decayed to mere traces, or even be completely gone. It is also uncertain if such superheavy nuclei may be produced naturally at all, as spontaneous fission is expected to terminate the r-process responsible for heavy element formation between mass number 270 and 290, well before elements such as unbiquadium may be formed.
Naming.
Using the 1979 IUPAC recommendations, the element should be temporarily called "unbiquadium" (symbol "Ubq") until it is discovered, the discovery is confirmed, and a permanent name chosen. Although widely used in the chemical community on all levels, from chemistry classrooms to advanced textbooks, the recommendations are mostly ignored among scientists who work theoretically or experimentally on superheavy elements, who call it "element 124", with the symbol "E124", "(124)", or "124". Some researchers have also referred to unbiquadium as "eka-uranium", a name derived from the system Dmitri Mendeleev used to predict unknown elements, though such an extrapolation might not work for g-block elements with no known congeners and "eka-uranium" would instead refer to element 144 or 146 when the term is meant to denote the element directly below uranium.
Difficulties in synthesis.
Every element from mendelevium onward was produced in fusion-evaporation reactions, culminating in the discovery of the heaviest known element oganesson in 2002 and more recently tennessine in 2010. These reactions approached the limit of current technology; for example, the synthesis of tennessine required 22 milligrams of 249Bk and an intense 48Ca beam for six months. The intensity of beams in superheavy element research cannot exceed 1012 projectiles per second without damaging the target and detector, and producing larger quantities of increasingly rare and unstable actinide targets is impractical.
Consequently, future experiments must be done at facilities such as the under-construction superheavy element factory (SHE-factory) at the Joint Institute for Nuclear Research (JINR) or RIKEN, which will allow experiments to run for longer stretches of time with increased detection capabilities and enable otherwise inaccessible reactions. Even so, it is expected to be a great challenge to continue past elements 120 or 121 given short predicted half-lives and low predicted cross sections.
The production of new superheavy elements will require projectiles heavier than 48Ca, which was successfully used in the discovery of elements 114-118, though this necessitates more symmetric reactions which are less favorable. Hence, it is likely that the reactions between 58Fe and a 249Cf or newly available 251Cf target are most promising. Studies on the fission of various superheavy compound nuclei have found that the dynamics of 48Ca- and 58Fe-induced reactions are similar, suggesting that 58Fe projectiles may be viable in producing superheavy nuclei up to "Z" = 124 or possibly 125. It is also possible that a reaction with 251Cf will produce the compound nucleus 309Ubq* with 185 neutrons, immediately above the "N" = 184 shell closure. For this reason, the compound nucleus is predicted to have relatively high survival probability and low neutron separation energy, leading to the 1n–3n channels and isotopes 306–308Ubq with a relatively high cross section. These dynamics are highly speculative, as the cross section may be far lower should trends in the production of elements 112–118 continue or the fission barriers be lower than expected, regardless of shell effects, leading to decreased stability against spontaneous fission (which is of growing importance). Nonetheless, the prospect of reaching the "N" = 184 shell on the proton-rich side of the chart of nuclides by increasing proton number has long been considered; already in 1970, Soviet nuclear physicist Georgy Flyorov suggested bombarding a plutonium target with zinc projectiles to produce isotopes of element 124 at the "N" = 184 shell.
Predicted properties.
Nuclear stability and isotopes.
Unbiquadium is of interest to researchers because of its possible location near the center of an island of stability, a theoretical region comprising longer-lived superheavy nuclei. Such an island of stability was first proposed by University of California professor Glenn Seaborg, specifically predicting a region of stability centered at element 126 (unbihexium) and encompassing nearby elements, including unbiquadium, with half-lives possibly as long as 109 years. In known elements, the stability of nuclei decreases greatly with the increase in atomic number after uranium, the heaviest primordial element, so that all observed isotopes with an atomic number above 101 decay radioactively with a half-life under a day, with an exception of dubnium-268 with half-life of 28 hours (1 day and 4 hours). Nevertheless, there is a slight increase in nuclear stability in nuclides around atomic numbers 110–114, which suggests the presence of an island of stability. This is attributed to the possible closure of nuclear shells in the superheavy mass region, with stabilizing effects that may lead to half-lives on the order of years or longer for some as-yet undiscovered isotopes of these elements. While still unproven, the existence of superheavy elements as heavy as oganesson provides evidence of such stabilizing effects, as elements with an atomic number greater than approximately 104 are extremely unstable in models neglecting magic numbers.
In this region of the periodic table, "N" = 184 and "N" = 228 have been proposed as closed neutron shells, and various atomic numbers have been proposed as closed proton shells, including "Z" = 124. The island of stability is characterized by longer half-lives of nuclei located near these magic numbers, though the extent of stabilizing effects is uncertain due to predictions of weakening of the proton shell closures and possible loss of double magicity. More recent research predicts the island of stability to instead be centered at beta-stable copernicium isotopes 291Cn and 293Cn, which would place unbiquadium well above the island and result in short half-lives regardless of shell effects. A 2016 study on the decay properties of unbiquadium isotopes 284–339Ubq predicts that 284–304Ubq lie outside the proton drip line and thus may be proton emitters, 305–323Ubq may undergo alpha decay, with some chains terminating as far as flerovium, and heavier isotopes will decay by spontaneous fission. These results, as well as those from a quantum-tunneling model, predict no half-lives over a millisecond for isotopes lighter than 319Ubq, as well as especially short half-lives for 309–314Ubq in the sub-microsecond range due to destabilizing effects immediately above the shell at "N" = 184. This renders the identification of many unbiquadium isotopes nearly impossible with current technology, as detectors cannot distinguish rapid successive signals from alpha decays in a time period shorter than microseconds.
Increasingly short spontaneous fission half-lives of superheavy nuclei and the possible domination of fission over alpha decay will also probably determine the stability of unbiquadium isotopes. While some fission half-lives constituting a "sea of instability" may be on the order of 10−18 s as a consequence of very low fission barriers, especially in even–even nuclei due to pairing effects, stabilizing effects at "N" = 184 and "N" = 228 may allow the existence of relatively long-lived isotopes. For "N" = 184, fission half-lives may increase, though alpha half-lives are still expected to be on the order of microseconds or less, despite the shell closure at 308Ubq. It is also possible that the island of stability may shift to the "N" = 198 region, where total half-lives may be on the order of seconds, in contrast to neighboring isotopes that would undergo fission in less than a microsecond. In the neutron-rich region around "N" = 228, alpha half-lives are also predicted to increase with increasing neutron number, meaning that the stability of such nuclei would primarily depend on the location of the beta-stability line and resistance to fission. One early calculation by P. Moller, a physicist at Los Alamos National Laboratory, estimates the total half-life of 352Ubq (with "N" = 228) to be around 67 seconds, and possibly the longest in the "N" = 228 region.
Chemical.
Unbiquadium is the fourth member of the superactinide series and should be similar to uranium: both elements have six valence electrons over a noble gas core. In the superactinide series, the Aufbau principle is expected to break down due to relativistic effects, and an overlap of the 5g, 6f, 7d, and 8p orbitals is expected. The ground state electron configuration of unbiquadium is thus predicted to be [Og] 6f3 8s2 8p1 or 6f2 8s2 8p2, in contrast to [Og] 5g4 8s2 derived from Aufbau. This predicted overlap of orbitals and uncertainty in order of filling, especially for f and g orbitals, renders predictions of chemical and atomic properties of these elements very difficult.
One predicted oxidation state of unbiquadium is +6, which would exist in the halides UbqX6 (X = a halogen), analogous to the known +6 oxidation state in uranium. Like the other early superactinides, the binding energies of unbiquadium's valence electrons are predicted to be small enough that all six should easily participate in chemical reactions. The predicted electron configuration of the Ubq5+ ion is [Og] 6f1.


List of countries by manganese production

Manganese is a chemical element that is designated by the symbol Mn and has an atomic number of 25. It is found as the free element in nature (often in combination with iron), and in many minerals. The free element is a metal with important industrial metal alloy uses. Manganese ions are variously colored, and are used industrially as pigments and as oxidation chemicals. Manganese (II) ions function as cofactors for a number of enzymes; the element is thus a required trace mineral for all known living organisms.
List of Mining Countries.
This is a list of countries by manganese ore mining in 2015, based on United States Geological Survey accessed in 2016.


Triphenylmethyl hexafluorophosphate

Triphenylmethyl hexafluorophosphate is an organic salt with the formula or , consisting of the triphenylmethyl cation and the hexafluorophosphate anion] . The cation is also called triphenylcarbenium, trityl cation, or tritylium.
Triphenylmethyl hexafluorophosphate is a brown powder that hydrolyzes readily to triphenylmethanol. It is used as a catalyst and reagent in organic syntheses.
Preparation.
Triphenylmethyl hexafluorophosphate can be prepared by combining silver hexafluorophosphate with triphenylmethyl chloride:
A second method involves protonolysis of generating triphenylmethanol:
Structure and reactions.
Triphenylmethyl hexafluorophosphate readily hydrolyzes, in a reaction that is the reverse of one of its syntheses:
Triphenylmethyl hexafluorophosphate has been used for abstracting hydride () from organic compounds. Treatment of metal-alkene and diene complexes one can generate allyl and pentadienyl complexes, respectively.
Triphenylmethyl perchlorate is a common substitute for triphenylmethyl hexafluorophosphate. However, the perchlorate is not used as widely, because, like other organic perchlorates, it is potentially explosive.


Zhanghengite

Zhanghengite is a mineral consisting of 80% copper and zinc, 10% iron with the balance made up of chromium and aluminium. Its color is golden yellow. It was discovered in 1986 during the analysis of the Bo Xian meteorite and is named after Zhang Heng, an ancient Chinese astronomer.


Titanium nitrate

Titanium nitrate is the inorganic compound with formula Ti(NO3)4. It is a colorless, diamagnetic solid that sublimes readily. It is an unusual example of a volatile binary transition metal nitrate. Ill defined species called titanium nitrate are produced upon dissolution of titanium or its oxides in nitric acid.
Preparation.
Similarly to its original method, Ti(NO3)4 is prepared by the nitration of titanium tetrachloride using dinitrogen pentoxide:
A hydrated titanium nitrate is produced upon dissolution of titanium compounds in nitric acid.
Structure.
The complex has D2d symmetry, with four bidentate nitrate ligands. The N-O distances are 1·29 Å and 1·185 Å (noncoordinated).
Physical properties.
In the infrared spectrum, it absorbs strongly at 1635 cm−1, assigned to a N-O vibrational mode.
It is soluble in nonpolar solvents silicon tetrachloride and carbon tetrachloride.
Reactions.
Titanium nitrate is hygroscopic, converting to ill-defined hydrates. The anhydrous material is highly reactive, even toward hydrocarbons. Titanium nitrate also reacts with n-dodecane, p-dichlorobenzene, anisole, biphenyl,
It decomposes thermally to titanium dioxide.


Mesothorium


Zinc gluconate

Zinc gluconate is the zinc salt of gluconic acid. It is an ionic compound consisting of two anions of gluconate for each zinc(II) cation. Zinc gluconate is a popular form for the delivery of zinc as a dietary supplement providing 14.35% elemental zinc by weight.
Gluconic acid is found naturally, and is industrially made by the fermentation of glucose, typically by "Aspergillus niger", but also by other fungi, e.g. "Penicillium", or by bacteria, e.g. "Acetobacter", "Pseudomonas" and "Gluconobacter". In its pure form, it is a white to off-white powder. It can also be made by electrolytic oxidation, although this is a more expensive process. The advantages are a lower microbiological profile, and a more complete reaction, yielding a product with a longer shelf life.
Zinc gluconate and the common cold.
Zinc gluconate has been used in lozenges for treating the common cold. However, controlled trials with lozenges which include zinc acetate have found it has the greatest effect on the duration of colds. Zinc has also been administered nasally for treating the common cold, but has been reported to cause anosmia in some cases.
Safety concerns.
Instances of anosmia (loss of smell) have been reported with intranasal use of some products containing zinc gluconate. In September 2003, Zicam faced lawsuits from users who claimed that the product, a nasal gel containing zinc gluconate and several inactive ingredients, negatively affected their sense of smell and sometimes taste. Some plaintiffs alleged experiencing a strong and very painful burning sensation when they used the product. Matrixx Initiatives, Inc., the maker of Zicam, responded that only a small number of people had experienced problems and that anosmia can be caused by the common cold itself. In January 2006, 340 lawsuits were settled for $12 million.
The U.S. Food and Drug Administration (FDA) considers zinc gluconate to be generally recognized as safe (GRAS) when used in accordance with good manufacturing practice, although this does not constitute a finding by the FDA that the substance is a useful dietary supplement. On 16 June 2009 the FDA "warned consumers to stop using and discard three zinc-containing Zicam intranasal products. The products may cause a loss of sense of smell. ... FDA is concerned that the loss of sense of smell may be permanent." Matrixx responded that the FDA's allegations were "unfounded and misleading", citing a lack of evidence from controlled tests that Zicam causes anosmia. In its warning, the FDA stated, "This warning does not involve oral zinc tablets and lozenges taken by mouth. Dietary zinc is also not subject to this warning."
Veterinary use.
A zinc gluconate-based product, also containing arginine, is used as a veterinary chemical castration drug. For dogs, the product is injected directly into the testicles. It has been sold under various brand names, including Neutersol and Esterilsol.


Tungsten(III) oxide

Tungsten(III) oxide (W2O3) is a compound of tungsten and oxygen. It has been reported (2006) as being grown as a thin film by atomic layer deposition at temperatures between 140 and 240 °C using W2(N(CH3)2)6 as a precursor. It is not referred to in major textbooks.
Some older literature refers to the compound W2O3 but as the atomic weight of tungsten was believed at the time to be 92, i.e., approximately half the modern accepted value of 183.84, the compound actually being referred to was WO3.
Usage.
Tungsten(III) oxide is used in various types of infrared absorbing coatings and foils.


Sodium hyponitrite

Sodium hyponitrite is a solid ionic compound with formula or ()2[ON=NO]2−.
There are "cis" and "trans" forms of the hyponitrite ion . The "trans" form is more common, but the "cis" form can be obtained too, and it is more reactive than the "trans" form.
"Trans" isomer.
The "trans" isomer is colorless and soluble in water and insoluble in ethanol and ether.
Preparation.
Sodium hyponitrite ("trans") is conventionally prepared by reduction of sodium nitrite with sodium amalgam.
Sodium hyponitrite ("trans") was prepared in 1927 by A. W. Scott by reacting alkyl nitrites, hydroxylammonium chloride, and sodium ethoxide
An earlier method, published by D. Mendenhall in 1974, reacted gaseous nitric oxide () with sodium metal in 1,2-dimethoxyethane, toluene, and benzophenone. The salt was then extracted with water. The method was later modified to use pyridine.
Other methods included oxidation of a concentrated solution of hydroxylamine with sodium nitrite in an alkaline medium; or electrolysis of sodium nitrite.
Hydrates.
A variety of hydrates ()"x" of the "trans" isomer have been reported, with "x" including 2, 3.5, 4, 5, 6, 7, 8, and 9; but there is some dispute.
The hydration water seems to be just trapped in the crystal lattice rather than coordinated to the ions. The anhydrous substance can be obtained by drying the hydrates over phosphorus pentoxide and then heating them to 120 °C.
Reactions.
Sodium hyponitrite ("trans") in solution is decomposed by carbon dioxide from air to form sodium carbonate.
Liquid N2O4 oxidises sodium hyponitrite ("trans") to give sodium peroxohyponitrite [ON=NOO]2−).
"Cis" isomer.
The "cis" isomer of sodium hyponitrite is a white crystalline solid, insoluble in aprotic solvents, and (unlike the "trans" isomer) decomposed by water and other protic solvents.
Preparation.
The "cis" isomer of can be prepared by passing nitric oxide () through a solution of sodium metal in liquid ammonia at −50 °C.
The "cis" isomer was also obtained in 1996 by C. Feldmann and M. Jansen by heating sodium oxide with 77 kPa of nitrous oxide (laughing gas) in a sealed tube at 360 °C for 2 hours. The two reagents combined to yield the "cis" hyponitrite quantitatively as white microcrystals.
Properties and reactions.
The anhydrous "cis" salt is stable up to 325 °C, when it disproportionates to nitrogen and sodium orthonitrite:
It is generally more reactive than the "trans" isomer.


Barium orthotitanate

Barium orthotitanate is the inorganic compound with the chemical formula Ba2TiO4. It is a colourless solid that is of interest because of its relationship to barium titanate, a useful electroceramic.
Structure.
The solid has two known phases: a low-temperature (β) phase with "P"21/"n" symmetry and a high-temperature (α′) phase with "P"21"nb" symmetry. The structure of Ba2TiO4 is unusual among the titanates because its titanium atoms sit in a four-oxygen tetrahedron rather than a six-oxygen octahedron.
Production.
It forms as white crystals from a melt of BaCl2, BaCO3 and TiO2 or from just sintering BaCO3 and TiO2. Another method of preparation is heating pellets of Ba(OH)2 and TiO2. Additionally, there are polymer precursor, sol-gel and reverse micellar routes to Ba2TiO4 synthesis. Ba2TiO4 has also been successfully grown as a thin film with chemical vapor deposition.
Properties.
Barium orthotitanate can remove up to 99.9% of from a high-temperature gas stream by the reaction:


Tetraiodonickelate

Tetraiodonickelate is a complex ion of nickel with four iodide atoms [NiI4]2− arranged in a tetrahedron. [NiI4]2− is red in solution. This colour is due to absorption around 530 nm and below 450 nm. Maximum light transmission is around 620 nm, which is red. A broad weak absorption in the near infrared is at 740 nm. The magnetic moment is anomalously low.
A mixture of lithium iodide and nickel iodide in water or methanol can partition NiI42− ions into a cyclohexane-amine mixture. The solution formed is blood red.
History.
Already in 1909 Cambi had noticed that a mixture of nickel iodide and sodium iodide dissolved in acetone has a red colour. This red colour was due to the presence of tetraiodonickelate.
Salts.
Bis-triphenylmethylarsonium tetraiodoronickelate [(C6H5)3CH3As]2NiI4 is red in colour. It can be made from triphenylmethylarsonium iodide and nickel iodide in hot ethanol. The red flakes that precipitate must be filtered before the alcohol cools, else the compound decomposes.
Bis(tetraethylammonium) tetraiodonickelate has a molecular weight of 826.8135 and a CAS number of 13927-28-1.
1,2,6-trimethyl-pyrazinium-tetraiodonickelate has CAS 88227-96-7.


Unquadquadium


Metallocarbohedryne

A metallocarbohedryne (or met-car for short) is any one of a family of chemical compounds with the generic molecular formula , where M is a transition metal such as titanium, vanadium, zirconium, niobium, hafnium, molybdenum, chromium, or iron.
These compounds have similar properties and a similar molecular structure, with the eight metal atoms at the corners of a somewhat distorted cube, and the twelve carbon atoms, in pairs, placed diagonally across the faces of the cube. The structure can also be described as two intersecting tetrahedra of metal atoms, with the carbon atoms placed in pairs along the edges of one tetrahedron. They have been extensively studied in the gas phase, and sometimes dispersed in solid materials, but so far have not been produced in bulk or in solution. Nevertheless, they have attracted interest because of their stability and symmetry, a relatively low ionization potential, delayed ionization, and possibly interesting magnetic properties. Some authors suggest that they may eventually find applications in electronics and catalysis.
The name is also used for the corresponding cations and anions .
The first papers used the name metallo-carbohedrene (with or without the hyphen) for this type of compound.
History.
The earliest known member of this family is the cation , discovered by Guo, kerns, and Castleman in 1992 while researching the dehydrogenation of various hydrocarbons (including methane, acetylene, ethylene, benzene, and propylene) with titanium atoms, in the gas phase. Although fullerenes like were already known, that may have been the first cage-like molecule with metal atoms replacing carbon at some corners of the mesh. They observed that the cluster would bind eight ammonia molecules, indicating that the eight titanium atoms were exposed. They also observed the analogous cations with vanadium, zirconium, or hafnium substituted for titanium, the corresponding neutral molecules, and the anion .
Synthesis.
Metallocarbohedrynes can be readily generated by vaporizing the desired metal with a laser, in an atmosphere containing the suitable hydrocarbon. The technique can produce mixed clusters, such as .
They have been also detected, at a concentration of 1% or less, in the soot generated by an electric arc between two Ti-C electrodes.
Structure.
The structure of these clusters has been extensively investigated since their discovery. At first, the 20 atoms of were conjectured to be arranged as the vertices of a dodecahedron, with the titanium atoms at the corners of a cube, and two carbon atom pairs, on opposite faces, aligned with each set of four parallel edges of the cube. This structure was conjectured to be analogous to that of the hypothetical dodecahedral fullerene . However, this claim was soon disputed by Linus Pauling who proposed an alternative arrangement—with the titanium atoms still at the corners of a cube, but with the carbon atoms pushed inwards so as to be nearly coplanar with the faces of that cube.
Theoretical studies.
The first ab initio theoretical investigations of the structure of (by Li and others, Methfessel and others, in 1993) indicated a slightly distorted version of the dodecahedron proposed by Guo and others, with C-C distances 139 pm and Ti-C distances 199 pm. In this model, the eight titanium atoms were still equivalent and located at the corners of a cube, with C-C pairs parallel to edges, so that the molecule would have the symmetry group formula_1. Nevertheless, they found the atoms are almost equidistant from the center, (260 pm for C, 262 pm for Ti). The electronic structure however was quite unlike that of graphite and .
Several other models were proposed. Ceulemans and Fowler proposed a ring of 12 carbon atoms capped by two tetrahedra. Khan proposed a cage of 12 carbons at the vertices of a cuboctahedron, surrounded by an elongated cage of metal atoms.
Eventually a consensus was reached on a structure proposed by Dance and others, in which the metal atoms are divided in two groups of four ("outer" or "o-", and "inner" or "i-"), at the vertices of two intersecting concentric regular tetrahedra, with different radii and opposite orientations; and the six carbon pairs are aligned with the edges of the larger tetrahedron. This structure can be seen as a deformation of the original proposal, by pulling four vertices of the cube slightly outwards, and rotating the carbon pairs by 45 degrees. Its symmetry group is formula_2 instead of formula_1, and it was predicted to have considerably lower energy (by 300 kcal/mol). Indeed, the formation of with the Dance structure was predicted to be energetically favored (exothermic) relative to metallic titanium and graphite.
Acceptance of this structure was delayed because the yields of the various clusters in Guo's process suggested that the eight metal atom sites were equivalent. In particular, the cluster did not seem to be exceptionally stable. However, the energy difference between placing the four zirconium atoms in the inner positions, rather than the outer ones, was eventually computed to be only 0.5 kcal/mol.
In 2003, Hou and others predicted a slight displacement of two of the carbon pairs, that reduced the symmetry group to formula_4 A similar conclusion had been reached by Chen and others. However, later studies by Lou and Nordlander concluded that the formula_2 form had lower energy (by about 70 kcal/mol) However, the zinc cluster was predicted to have the symmetrical dodecahedral (formula_1) structure suggested by Guo for the titanium cluster.
Electronically, is believed to have a metallic character, with 80 delocalized valence electrons. Its static polarizability was computed to be of the same order of magnitude as that of the fullerene .
Spectroscopy and ionization.
Pilgrim and Duncan observed in 1993 that can be dissociated by visible light is a fragment of +
In 1998, Sakurai and Castleman measured ionization potentials of via near threshold photoionization spectroscopy. In particular, they got 4.40 eV of for and 3.95 eV for . The former value was said to be more consistent with the formula_2 structure than the formula_1 one.
The infrared spectrum of neutral and of cations was studied by van Heijnsbergen and others, starting 1999. They measured clusters in the gas phase, accumulated as cations in an ion trap. They saw evidence that the loss of one electron from to does not change the structure significantly.
In 2004, Martínez and others computed from theoretical models the optical absorption spectrum of and . They predicted a broad spectrum for both, with high absorption starting at about 8 eV and centered around 12–14 eV.
Reactions.
The chemistry of and it analogs was studied in the gas phase, already by Castleman's and others. After creation, the ionized clusters were separated from other species by mass spectrometry, and injected into a drift tube containing the gaseous reactant, diluted in helium.
With theoretical computations, Huo and others predicted that the clusters and could bind 4 carbonyls, at outer metal atoms.
Potential applications.
While the clusters have yet to be produced in bulk, they have been investigated theoretically for possible use as catalysts.
Desulfurization of oil.
Specifically, in 2004 Liu and others have simulated the decomposition of thiophene by three hydrogen molecules to 2-butene and hydrogen disulfide , catalyzed by a neutral . This reaction is an important step in the removal of sulfur from oil. They predicted that the first molecule would spontaneously dissociate in contact with the pairs, and each H atom would then migrate to the adjacent outer titanium atom ("o-Ti"). The thiophene would then react exothermally with each H atom in turn, yielding a butadiene attached to an o-Ti and the sulfur atom attached at the nearby inner titanium ("i-Ti") atom. A second molecule would then dissociate at the o-Ti site and turn butadiene into 2-butene. A third would dissociate at an o-Ti site, and the two atoms would migrate to the i-Ti atom bearing the sulfur atom, and convert it into .


Okavango Blue

The Okavango Blue is a 20.46 carat blue diamond of exceptional clarity.
The Okavango Blue originated from the Orapa diamond mine in Botswana. The original stone was found in May 2018, weighing 41.11 carat, and was cut in an oval shape. The diamond is classified as a Type IIb "fancy blue" diamond, and was graded by the Gemological Institute of America (GIA) as "Very, Very Slightly Included", or VVS2. The stone was unveiled in April 2019, and is the largest ever blue diamond found in the country. The stone is anticipated to be sold in November 2019.


Stannocene

Stannocene is an organometallic compound with the formula .
It is a metallocene that can be produced efficiently from cyclopentadienyl sodium and tin(II) chloride. Unlike in ferrocene the two cyclopentadienyl rings are not parallel.