Methanetetracarboxylate

In chemistry, methanetetracarboxylate is a tetravalent anion with formula or C(COO−)4. It has four carboxylate groups attached to a central carbon atom; so it has the same carbon backbone as neopentane. It is an oxocarbon anion, that is, consists only of carbon and oxygen.
The term is also used for any salt with that anion; or for any ester with the C(COO)4 moiety.
The salts and esters are relatively uncommon, and their uses appear to be limited to chemical research. The sodium salt Na4C(COO)4 can be obtained by oxidation of pentaerythritol C(CH2OH)4 with oxygen in sodium hydroxide solution at pH 10 and about 60 °C, in the presence of palladium as a catalyst. The tetraethyl ester C(COO-C2H5)4 is traded as a specialty chemical and has been used in organic synthesis.
The anion can be seen as the result of removing four protons from methanetetracarboxylic acid, a hypothetical organic compound with formula C5H4O8 or C(COOH)4. However this acid has not been synthesised (as of 2009), and is believed to be unstable.


Borinic acid

Borinic acid, also known as boronous acid, is an oxyacid of boron with formula . Borinate is the associated anion of borinic acid with formula ; however, being a Lewis acid, the form in basic solution is .
Borinic acid can be formed as the first step in the hydrolysis of diborane:
Borinic acid itself is unstable and only lasts for a few seconds during the hydrolysis reaction. However, by using microwave spectroscopy various properties can be determined. The B-O distance is 1.352 Å, O-H distance 0.96 Å, B-H length is probably 1.2 Å. The angle between bonds at the oxygen atom BOH = 112° and the angles at boron are cis-HBO 121°, and trans-HBO = 117°. The dipole moment is 1.506 Debye.
Borinic acid can form esters such as methoxyborane. This too is unstable only lasting about ten seconds. It can be formed by heating diborane and methanol gas together.
By substituting organic components instead of hydrogen, more generic borinic acids (containing RR'BOH) or borinic esters (RR'BOR") can be formed. Esters will tend to be stable in acidic conditions, but in alkaline conditions the boron atom can gain a negative charge and attach two hydroxyl groups, or two ester bonds. RR'B−(OH)2 or RR'B−(OR")2. The anionic borinate ion can very easily form esters with diols such as ethylene glycol or sugars.
Naming.
The IUPAC name "borinic acid" is a unique name for the acid. The anhydrides are named diboroxanes, H2BOBH2, as the base compound and H being able to be substituted, e.g. tetraethyldiboroxane, as the anhydride for diethylborinic acid. Organic naming standard in the "Bluebook" allows skeletal replacement naming where the name is shorter, 3-borapentan-3-ol versus diethyl borinic acid. The grouping -BH-O-BH2 is called diboroxanyl. Substituting sulfur for oxygen gives borinothioic acid (H2BSH). (Dimethylboranyl)oxy is used for the group (CH3)2B-O− and methyl(hydroxy)boranyl for the grouping CH3B(OH)-.
Formation.
There are several ways to produce substituted borinic acids.
Firstly borinic acids can be made from oxidizing trialkyl borane starting materials [R3B] with exposure to moist air, or treatment with iodine, which makes a dialkyliodoborane [R2BI]. Hydrolysis then results in the boronic acid (R2BOH). Trialkylborates [(RO)3B] or trialkoxyboroxine [(ROBO)3] can be reduced to borinic acid by us of a Grignard reagent. Grignard reagents can also reduce a boronic ester [RB(OR')2] to a borinic ester.
("Z")-Enolates give syn aldol product when reacted with an aldehyde, whereas ("E")-enolates give and anti aldol products.
Dialkyl boron chloride (R2BCl) with tertiary amine react with ketones to form an enol borinate.
A trialkoxyborane can react with lithium containing organic molecules to eliminate lithium and one or two alkoxy groups to make boronic and borinic esters.
Purification of the mixtures that result from the reactions is required, as often boronic esters will also be produced and mixed in with the borinic esters. The method of Letsinger is dissolve the mixture in ether and precipitate the borinic ester by forming a complex with ammonia. Treatment with ethanolamine ends up making an aminoetylborinate.
2-APB.
2-Aminoethoxydiphenyl borate (2-APB) inhibits transient receptor potential channels. This kind of inhibition, particularly inhibition of TRPM7, is being studied to find treatments for prostate cancer. 2-APB can work as a catalyst to add an alkyl group from an alkyl halide to a polyol or carbohydrate that contains a cis-vicinal diol to a precise position. It does this by first combining with the two hydroxy groups to make a ring containing OCCOB−. It can also calatlyse acid chloride or chloroformate reaction a specific region of the diol.
Diphenylborinic acid.
Diphenylborinic acid was discovered in 1894 by Michaelis who produced it by hydrolyzing the chloride. Letsinger determined its properties in 1955.
Diphenylborinic acid has an extra high affinity for catechols compared with carbohydrates.
Diphenylborinic acid can catalyse the condensation of pyruvic acids with aldehydes to yield substituted isotetronic acid.
Diphenylborinic acid is an inhibitor of several enzymes such as α-chymotrypsin, subtilisin BPN', and trypsin.
Borinate radical.
Borinate radicals (RR'BO·) can be formed from peroxyborinate decomposition.
Other compounds.
Other compounds include methoxy(dimethyl)borane, methoxy(methyl)boron, methoxy(methylidene)borane (with a C=B double bond).
[C5H5BR]− uses a B− to be equivalent to carbon in an aromatic benzene like ring. This too is called borinate. The 1-methyl and 1-phenyl borinates can form some of the few organo-thallium(I) compounds.
HB(C6F5)2 + phosphino alcohol → tBu2P+HCH2C(CH3)2OB−H(C6F5)2 → H2 + tBu2PCH2C(CH3)2OB(C6F5)2 and same for tBu2PCH2C(CF3)2OB(C6F5)2
di-Tris(tert-butoxy)siloxy borinic acid HOB[OSi(O(t)Bu)3]2 can be made from tributoxyborate and tributoxysiloxane. It can form a very complex crystal with Cp2Zr(Me)[OB[OSi(O(t)Bu)3]2]2.
Diborinic acids have two RBOH groups linked together by an organic connection such as diphenyl or phenyl.
Applications.
1,1,1,3,3,3-Hexafluoroisopropylbis(pentafluorophenyl)borinate can greatly increase solubility of LiF by complexing the F− anion. This has potential to improve lithium batteries.
Borinic esters are being researched as bacterial growth inhibitors due to their ability to disable some bacterial enzymes such as menaquinone methyltransferase and CcrM. This may result in development of treatments for topical application on skin.


Trichloro(dichlorophenyl)silane

Trichloro(dichlorophenyl)silane is a compound with formula Si(C6H3Cl2)Cl3.


Neodymium(III) sulfate

Neodymium(III) sulfate is a salt of the rare earth metal neodymium when mixed with sulfuric acid. Its octahydrate forms violet-red crystals. It is moderately soluble in water and has a monoclinic crystal structure. It can be used to treat water but because of its rarity, it is not used that way.
Preparation.
Neodymium sulfate is produced by dissolving neodymium metal in sulfuric acid.
Uses.
It is used in glass for extremely powerful lasers. It can be used to treat water but is not used because of its rarity.


Boron (disambiguation)

Boron is a chemical element with symbol B and atomic number 5.
Boron may also refer to:


Calcium morphenate

Calcium morphenate is a calcium salt of morphine which is produced by using calcium bases to raise the pH of an aqueous solution of opium alkaloids to around 9.0. This was a method used in pharmaceutical manufacturing to separate morphine from other alkaloids and inert materials from the opium solution. Variations on this route are still used in Afghanistan. When poppy straw concentrate or opium latex is dissolved in hot water and the calcium base is added, calcium morphenate is formed. Treatment with a weak acid such as ammonium chloride then causes morphine freebase to precipitate, leaving codeine and other alkaloids of the plant in solution.


Erbium(III) oxide

Erbium(III) oxide, is synthesized from the lanthanide metal erbium. It was partially isolated by Carl Gustaf Mosander in 1843, and first obtained in pure form in 1905 by Georges Urbain and Charles James. It has a pink color with a cubic crystal structure. Under certain conditions erbium oxide can also have a hexagonal form.
Reactions.
Erbium burns readily to form erbium (III) oxide:
The formation of erbium oxide is accomplished via the reaction 4 Er + 3 O2 → 2 Er2O3. Erbium oxide is insoluble in water and soluble in mineral acids. Er2O3 readily absorbs moisture and carbon dioxide from the atmosphere. It can react with acids to form the corresponding erbium(III) salts.
For example, with hydrochloric acid, the oxide follows the reaction Er2O3 + 6 HCl → 2 ErCl3 + 3 H2O to form erbium chloride.
Properties.
One interesting property of erbium oxides is their ability to up convert photons. Photon upconversion takes place when infrared or visible radiation, low energy light, is converted to ultraviolet or violet radiation higher energy light via multiple transfer or absorption of energy. Erbium oxide nanoparticles also possess photoluminescence properties. Erbium oxide nanoparticles can be formed by applying ultrasound (20 kHz, 29 W·cm−2) in the presence of multiwall carbon nanotubes. The erbium oxide nanoparticles that have been successfully made by employing ultrasound are erbium carboxioxide, hexagonal and spherical geometry erbium oxide. Each ultrasonically formed erbium oxide is photoluminescence in the visible region of the electromagnetic spectrum under excitation of 379 nm in water. Hexagonal erbium oxide photoluminescence is long lived and allows higher energy transitions (4S3/2 - 4I15/2). Spherical erbium oxide does not experience 4S3/2 - 4I15/2 energy transitions.
Uses.
The applications of Er2O3 are varied due to their electrical, optical and photoluminescence properties. Nanoscale materials doped with Er3+ are of much interest because they have special particle-size-dependent optical and electrical properties. Erbium oxide doped nanoparticle materials can be dispersed in glass or plastic for display purposes, such as display monitors. The spectroscopy of Er3+ electronic transitions in host crystals lattices of nanoparticles combined with ultrasonically formed geometries in aqueous solution of carbon nanotubes is of great interest for synthesis of photoluminescence nanoparticles in 'green' chemistry. Erbium oxide is among the most important rare earth metals used in biomedicine. The photoluminescence property of erbium oxide nanoparticles on carbon nanotubes makes them useful in biomedical applications. For example, erbium oxide nanoparticles can be surface modified for distribution into aqueous and non-aqueous media for bioimaging. Erbium oxides are also used as gate dielectrics in semi conductor devices since it has a high dielectric constant (10–14) and a large band gap. Erbium is sometimes used as a coloring for glasses and erbium oxide can also be used as a burnable neutron poison for nuclear fuel.


Untriquadium


Plutonium(III) bromide

Plutonium(III) bromide is an inorganic salt of bromine and plutonium with the formula PuBr3. This radioactive green solid has few uses, however its crystal structure is often used as a structural archetype in crystallography.
Crystal structure.
The PuBr3 crystal structure was first published in 1948 by William Houlder Zachariasen. The compound forms orthorhombic crystals, a type of square antiprism, within which the Pu atoms adopt an 8 coordinate bicapped trigonal prismatic arraignment. It's Pearson symbol is oS16 with the corresponding space group No. 63 (in International Union of Crystallography classification) or Cmcm (in Hermann–Mauguin notation). The majority of trivalent chloride and bromide salts of lanthanide and actinides crystallise in the PuBr3 structure.


Magnesium taurate

Magnesium taurate, also known as magnesium ditaurate, is the magnesium salt of taurine, and a mineral supplement.
It contains 8.9% elemental magnesium by mass. Accordingly, 100 mg of magnesium is contained in 1121 mg of magnesium taurate.
Research.
Magnesium taurate has been studied in rats for delaying the onset and progression of cataract.
Safety.
Due to the expected dissociation of magnesium taurate in the body before absorption, safety data on magnesium and taurine can be used to evaluate the safety of magnesium taurate.
Taurine has an observed safe level of supplemental intake in normal healthy adults at up to 3 g/day. Using the same level as an approximation for taurate yields a limit of 3.3 g/day for magnesium taurate, or alternatively 300 mg/day for elemental magnesium as taurate.


Isotopes of silver

Naturally occurring silver (47Ag) is composed of the two stable isotopes 107Ag and 109Ag in almost equal proportions, with 107Ag being slightly more abundant (51.839% natural abundance). Twenty-eight radioisotopes have been characterized with the most stable being 105Ag with a half-life of 41.29 days, 111Ag with a half-life of 7.45 days, and 112Ag with a half-life of 3.13 hours.
All of the remaining radioactive isotopes have half-lives that are less than an hour, and the majority of these have half-lives that are less than 3 minutes. This element has numerous meta states, with the most stable being 108mAg ("t"* 418 years), 110mAg ("t"* 249.79 days) and 106mAg ("t"* 8.28 days).
Isotopes of silver range in atomic weight from 92.950 u (93Ag) to 129.950 u (130Ag). The primary decay mode before the most abundant stable isotope, 107Ag, is electron capture and the primary mode after is beta decay. The primary decay products before 107Ag are palladium (element 46) isotopes and the primary products after are cadmium (element 48) isotopes.
The palladium isotope 107Pd decays by beta emission to 107Ag with a half-life of 6.5 million years. Iron meteorites are the only objects with a high enough palladium/silver ratio to yield measurable variations in 107Ag abundance. Radiogenic 107Ag was first discovered in the Santa Clara meteorite in 1978.
The discoverers suggest that the coalescence and differentiation of iron-cored small planets may have occurred 10 million years after a nucleosynthetic event. 107Pd versus 107Ag correlations observed in bodies, which have clearly been melted since the accretion of the solar system, must reflect the presence of live short-lived nuclides in the early solar system.


Hubeite

The mineral hubeite, , is a sorosilicate of the group. Structurally it also belongs to the Akatoreite group. It was found and named after the province of Hubei, China. It is common to iron ores in a mine of that region. It occurs mainly as aggregates of fan like crystals. It is dark to pale brown, has orange-brown streak and is vitreous. Hubeite has a hardness of 5.5 in the Mohs scale, one good cleavage and conchoidal fracture. It is triclinic with a space group of P1*. The structure of hubeite is very uncommon, and in fact there is only one other mineral that fits the group, which is ruizite.
Background.
Hubeite was discovered by Hawthorne et al. (2002) at the Daye mines in the Hubei province of China. It is classified as a sorosilicate, based on its formula (Hawthorn et al., 2004). Other related minerals would be inesite, (Hawthorne et al., 2004), ruizite (Hawthorne et al., 2002) and Akatoreite (Burns et al., 1993).
Composition.
To analyze the composition, an electron microprobe was used in the wavelength-dispersion mode (Hawthorn et al., 2002). The quantity of (OH) and (H2O) was acquired by solid solution and refinement, based on previous work by Hawthorne et al., 1990. To assure the presence of (OH) and (H2O) groups, an infrared spectrum was also recorded (Hawthorn et al., 2002).
Physical and optical properties.
Hubeite is most common as aggregates of intergrown crystals (Fig.1) that are usually less than 5 mm across and that have individual crystals with well-developed faces that are as long as 1 mm (Hawthorne et al., 2002). The color ranges from pale to dark brown, depending on the crystal size (Fig.2). Other properties consist of a pale orange-brown streak, vitreous luster, non-fluorescence, and one good cleavage parallel to the c-axis. It is also brittle with conchoidal fracture, has a hardness of 5.5 in the Mohs scale and a specific gravity of 3.02 (Hawthorn et al., 2002). As for optical features, it is important to note that hubeite is strongly pleochroic, biaxial with an indeterminate optic sign and has a birefringence of 0.023 (γ-α) (Hawthorne et al., 2002).
Structure.
The crystals used for structure study were acquired at the Daye Mine (Hawthorne et al., 2004). To get a first general idea of the mineral structure it went through X-ray intensity data analysis and then, for a more detailed study, an electron microprobe was used (Hawthorne et al., 2004). Hubeite is triclinic (P1*). Basically, there are two Ca sites in the structure of hubeite, with site one being and octahedron and the second site is coordinated by 6 oxygen atoms at the same distance and one extra oxygen atom further out and arranged in an augmented octahedron (Hawthorne et al., 2004). There are also 4 sites for Si in tetrahedral arrangement, and the fourth site bonds to an OH group forming an acid-silicate group (SiO3(OH)) (Hawthorne et al., 2004). There are 2 oxygen sites that connect 2 Si atoms, thus creating a sorosilicate (Hawthorne et al., 2002). [Si4O13] corresponds to a four membered chain fragment of tetrahedra according to Hawthorne et al. (2004). The only other sorosilicate mineral that has that same four membered configuration is ruizite (Moore et al., 1985). The main difference of the two minerals is the valence of Mn and the existence of Fe3+ for Hubeite (Hawthorne et al., 2002). Ruizite is of the [Si4O13] sorosilicate group (Hawthorne, 1984) and when it was discovered, it did not much any other Ca-Mn silicate already known (Willams et al., 1977), and now with the discovery of hubeite it is easier to understand the [Si4O13] sorosilicate group. The other two sites left in the hubeite structure are filled with Fe with CN=6 and Mn with CN=6, being one of the bonds to OH in the Mn case. The structure of hubeite is heteropolyhedra, with alternating layers of tetrahedra and different polyhedra parallel to (001) (Hawthorne et al., 2004). The tetrahedral layers are formed by [Si4O13] sharing corners, and the other alternating layer is formed by the [6], [7] and [8] Ca, Mn2+ and Fe3+ polyhedral sharing edges (Hawthorne et al., 2004). This last feature is what relates hubeite to the akatoreite group. Akatoreite, like hubeite, is triclinic with space group P1*(Burns et al., 1993). Akatoreite’ structure is layered as well with alternating sheets of octahedra and tetrahedra, parallel to (101) (Burns et al., 1993). The octahedra groups, as well as one Mn tetrahedra group, are sharing edges and linked by the corner sharing tetrahedral. The same happens in ruizite, except that they are linked by the [Si4O13] group. The inesite structure also relates very well to the hubeite structure. It is also based on layers of edge sharing polyhedra alternating with corner sharing tetrahedra (Hawthrone et al., 2004). The main difference is that inesite is a cyclosilicate, and in fact, by omitting 2 of the 6 tetrahedra that form the tetrahedra ring, and if the other 8 membered ring is broken and hydoxylated, the new arrangement becomes a hubeite (Hawthorne et al., 2004). This just confirms the association of hubeite and inesite in the Daye mines (Hawthorn et al., 2004).
Geological occurrence.
Hubeite is mainly associated to a skarn assemblage with pink inesite, colorless apophyllite, quartz, pyrite and colorless-white calcite (Hawthorne et al., 2004). They all occur together at the Daye Mine. Usually hubeite appears in two different situations. It may occur as isolated aggregates of crystals perched on white quartz, or it may occur covering both sides of thick specimens, that are usually pink inesite and apophyllite (Hawthorne et al., 2002). Figures 3 and 4 illustrate both situations.
The localities where ruizite is found, associated with apophyllite, inesite and pyrite as well, and there is no hubeite, which leads the conclusion that hubeite needs oxidized environments and sufficient concentration of iron to occur.
The Daye mine is an iron ore deposit (Dingyu et al., 1982). This specific area is characterized by deposits of late Paleozoic carbonate rocks in contact with plutons aging between the middle Jurassic to middle Cretaceous (Dingyu et al., 1982). According to Dingyu et al. (1982), the iron rich magma injections are the main cause for the formation of the ore deposits of the region. These polymetallic deposits form a belt that crosses China in the west–east direction (Ottens, 2007). Curiously, the mine where hubeite was first found is in fact a wollastonite source for minerals collectors.
Locations.
Hawthorne et al. (2002) discovered hubeite in the Daye mine, in the Hubei province of China. This mine became famous after that discovery and despite this specific breakthrough, the mine is most popular for its crystals of inesite and wollastonite (Ottens, 2007). It opened up in 1966 for copper exploration, but after the lack of profits, it became a major source for wollastonite (Ottens, 2007). Luckily, along the area of the Daye mine, there are other skarn-type iron and copper deposits that are big contributors to the overall reserves of copper and iron in China (Ottens, 2007). Daye County is also rich in non-metallic mineral deposits, but it is the metallic ores that make it special and it is an important city for the manufacture of bronze (Ottens, 2007).
The Hubei province has gold and silver production as the main revenue source (Ottens, 2007). This province is also one of the birthplaces of the Chinese bronze-age culture, represented in the artwork of the Yangtze River culture (Ottens, 2007). The copper extraction began in this area is related to the Yin Dynasty and the iron extraction began in the Qing Dynasty, making these mines a “symbol” in the Chinese culture (Ottens, 2007).


Pentafluoride

Pentafluoride may refer to:


Recycled diamond

A recycled diamond is a diamond, which had a prior use and has re-entered the diamond supply chain.
Diamonds in recent times have become good candidates for recycling, due to a number of reasons including their hardness and the gemstone being a finite and valuable resource. Diamonds are the hardest known naturally occurring substance and rate a 10 on the Mohs scale of mineral hardness.
It has been speculated that the value of the market could be as high as $1 trillion US dollars.
History.
Traditional history.
Traditionally, diamonds are considered a valuable commodity and large quantities of rough diamonds have been mined in the last few hundred years via pit mining or alluvial mining. According to Reuters, the lack of quality rough diamonds has been one of the main reasons for the rise of the recycled diamond industry in recent times.
Mined production for 2014 was estimated at around 135 million carats of rough diamonds with a value of $17.8 billion. Mine production over the last 150 years is estimated at over half a trillion carats, with a value close to half a trillion dollars. It has been forecast that the mining of diamonds will slow in the coming years, as mines are nearing the end of their productive lives and few new mines are being established.
Market expansion.
From 2000 onwards, businesses such as pawnbrokers and jewelers began to experience large increases in scrap gold buying from the public. A boom in this industry saw the gold price jump from $250 to $1,900. This led to an increase in the number of pawnbrokers receiving jewelry trade ins.
These exchanges led to a huge surplus of diamonds, during the recycling of the gold. Many refiners merely burned up the small "melee" diamonds, while others used an acid removal process known as aqua regia, (unless the gold buyers had physically removed the diamonds prior to melting the gold), to harmlessly remove the stones.
According to Tacy, a leading provider of research and news on the diamond industry, this urban diamond mine is said to be worth up to US$1 trillion and those diamonds could eventually find their way into the recycled market in the future. Since then, the market of recycled diamonds has continued to expand.
De Beers CEO stated in 2014 that one of the biggest issues to the diamond industry was the diamond trade-ins market or recycling. He stated when consumers receive low prices for their recycled gems, it can lead to a low opinion of the product. Following that statement, De Beers were responsible for launching the International Institute of Diamond Valuation.
Recycling of diamonds.
It is impossible to tell the difference between newly mined and polished diamonds and recycled diamonds. As a result, the price should be similar. A recycled diamond might be in excellent condition, though quite possibly have a dated cut or look. As such, a lot of recycled diamonds go through a re-cutting or re-polishing process to make them “as new”. If they require a re-cut, they will be less valuable than a similar, newly polished diamond.
Within the diamond industry, recycled diamonds of smaller sizes, typically one-fifth of a carat or smaller, are traded in parcels of mixed diamonds, known as a 'melee’ and are typically bought by jewelry manufacturers, diamond manufacturers and dealers. A diamond laboratory, such as GIA, AGS, EGL USA or IGI, certifies larger diamonds, and these diamonds are sold individually, directly to the public or to jewelry retailers and diamond dealers.
According to an independent diamond analyst, the global diamond market saw a huge change in the demand for recycled diamonds. It was suggested that recycled diamonds were at close to 10% of global supply in 2014, compared to a decade earlier when they were well below 5%. Tacy estimated that the recycled diamond market was valued at $1.2 billion in 2012, with growth planned over the coming years.also
Non-recyclable diamonds.
Some diamonds, no matter their size, cannot be recycled. Certain industries that commonly use diamonds for grinding or even the re-polishing of other diamonds mean that they wear over time and effectively become dust.


Organouranium chemistry

Organouranium chemistry is the science exploring the properties, structure and reactivity of organouranium compounds, which are organometallic compounds containing a carbon to uranium chemical bond. The field is of some importance to the nuclear industry and of theoretical interest in organometallic chemistry.
History.
The development of organouranium compounds started in World War II when the Manhattan Project required volatile uranium compounds for 235U/238U isotope separation. For example, Henry Gilman attempted to synthesize compounds like tetramethyluranium, and others worked on uranium metal carbonyls, but none of the efforts met success due to organouranium instability. After the discovery of ferrocene in 1951, Todd Reynolds and Geoffrey Wilkinson in 1956 synthesized the uranium metallocene Cp3UCl from sodium cyclopentadienide and uranium tetrachloride as a stable but extremely air-sensitive compound. In it, the U-Cl bond is an ionic bond, while the bonds with the three cyclopentadienyl ligands are covalent of the type found in sandwich compounds with involvement of the uranium 5f atomic orbitals.
Ernst Otto Fischer in 1962 discovered tetracyclopentadienyluranium Cp4U by reaction of KCp with UCl4 (6% yield) as a compound stable in air as a solid but not in solution. A zero molecular dipole moment and IR spectroscopy revealed that it was also a sandwich compound with uranium in a tetrahedral molecular geometry. In 1970, Fischer added Cp3U to the list of known organouranium compounds by reduction of Cp4U with elemental uranium.
In 1968, the group of Andrew Streitwieser prepared the stable but pyrophoric compound uranocene (COT)2U, which has an atom of uranium sandwiched between two cyclooctatetraenide anions (D8h molecular symmetry). The uranium f orbitals interact substantially with the aromatic rings just as the d-orbitals in ferrocene interact with the Cp ligands. Uranocene differs from ferrocene because its HOMO and LUMO are centered on the metal and not on the rings and all reactions thus involve the metal often resulting in ligand - metal cleavage.
Uranocenes show ease of reduction of U(IV) compounds to U(III) compounds; otherwise they are fairly unreactive. A close relative that does have sufficient reactivity, obtained by reaction of uranocene with uranium borohydride is the half-sandwich compound (COT)U(BH4)2 discovered in 1983 by the group of M.J. Ephritikhine. Compounds of this type react in many different ways, for instance alkylation at uranium with organolithium reagents or conversion to hybrid sandwich compounds.
Other organouranium compounds are inverted uranocenes with a COT ligand in between two uranium atoms or uranium sandwich compounds with pentalenide ligands instead of COT ligands.


Gold(III) nitrate

Gold(III) nitrate is a crystalline gold compound. This salt can be used as a source of gold in lower (acidic) pH conditions. When mixed with hydrocarbons, it can form a highly flammable mixture. It decomposes at 72 °C to gold, nitrogen dioxide, and oxygen.
Preparation.
Gold nitrate is prepared by reacting chloroauric acid and silver nitrate.


Organoiridium compound

Organoiridium chemistry is the chemistry of organometallic compounds containing an iridium-carbon chemical bond. Organoiridium compounds are relevant to many important processes including olefin hydrogenation and the industrial synthesis of acetic acid. They are also of great academic interest because of the diversity of the reactions and their relevance to the synthesis of fine chemicals.
Classification based on principal oxidation states.
Organoiridium compounds share many characteristics with those of rhodium, but less so with cobalt. Iridium can exist in oxidation states of -III to +V, but iridium(I) and iridium(III) are the more common. iridium(I) compounds (d8 configuration) usually occur with square planar or trigonal bipyramidal geometries, whereas iridium(III) compounds (d6 configuration) typically have an octahedral geometry.
Iridium(0).
Iridium(0) complexes are binary carbonyls, the principal member being tetrairidium dodecacarbonyl, Ir4(CO)12. Unlike the related Rh4(CO)12, all CO ligands are terminal in Ir4(CO)12, analogous to the difference between Fe3(CO)12 and Ru3(CO)12.
Iridium(I).
A well known example is Vaska's complex, bis(triphenylphosphine)iridium carbonyl chloride. Although iridium(I) complexes are often useful homogeneous catalysts, Vaska' complex is not. Rather it is iconic in the diversity of its reactions. Other common complexes include Ir2Cl2(cyclooctadiene)2, chlorobis(cyclooctene)iridium dimer, The analogue of Wilkinson's catalyst, IrCl(PPh3)3), undergoes orthometalation:
This difference between RhCl(PPh3)3 and IrCl(PPh3)3 reflects the generally greater tendency of iridium to undergo oxidative addition. A similar trend is exhibited by RhCl(CO)(PPh3)2 and IrCl(CO)(PPh3)2, only the latter oxidatively adds O2 and H2. The olefin complexes chlorobis(cyclooctene)iridium dimer and cyclooctadiene iridium chloride dimer are often used as sources of "IrCl", exploiting the lability of the alkene ligands or their susceptibility to removal by hydrogenation. Crabtree's catalyst ([Ir(P(C6H11)3)(pyridine)(cyclooctadiene)]PF6) is a versatile homogeneous catalyst for hydrogenation of alkenes.
(η5-Cp)Ir(CO)2 oxidatively adds C-H bonds upon photolytic dissociation of one CO ligand.
Iridium(II).
As is the case for rhodium(II), iridium(II) is rarely encountered. One example is iridocene, IrCp2. As with rhodocene, iridocene dimerises at room temperature.
Iridium(III).
Iridium is usually supplied commercially in the Ir(III) and Ir(IV) oxidation states. Important starting reagents being hydrated iridium trichloride and ammonium hexachloroiridate. These salts are reduced upon treatment with CO, hydrogen, and alkenes. Illustrative is the carbonylation of the trichloride:
IrCl3(H2O)x + 3 CO → [Ir(CO)2Cl2]− + CO2 + 2 H+ + Cl− + (x-1) H2O
Many organoiridium(III) compounds are generated from pentamethylcyclopentadienyl iridium dichloride dimer. Many of derivatives feature kinetically inert cyclometalated ligands. Related half-sandwich complexes were central in the development of C-H activation.
Iridium(V).
Oxidation states greater than III are more common for iridium than rhodium. They typically feature strong-field ligands. One often cited example is oxotrimesityliridium(V).
Uses.
The dominant application of organoiridium complexes is as catalyst in the Cativa process for carbonylation of methanol to produce acetic acid.
Optical devices and photoredox.
Iridium complexes such as cyclometallated derived from 2-phenylpyridines are used as phosphorescent organic light-emitting diodes. Related complexes are photoredox catalysts.
Potential applications.
Iridium complexes are highly active for hydrogenation both directly and via transfer hydrogenation. The asymmetric versions of these reactions are widely studied.
Many half-sandwich complexes have been investigated as possible anti-cancer drugs. Related complexes are electrocatalysts for the conversion of carbon dioxide to formate. In academic laboratories, iridium complexes are widely studied because its complexes promote C-H activation, but such reactions are not employed in any commercial process.


Kosnarite

Kosnarite is a alkali zirconium phosphate mineral (KZr2(PO4)3) named after an expert of pegmatites Richar A. Kosnar. Kosnarite contains potassium, oxygen, phosphorus, and zirconium with sodium, rubidium, hafnium, manganese and fluorine (Na, Rb, Hf, Mn, and F) being common impurities found in kosnarite. It was discovered in nature for the first time in 1991 by Vandall T. King. Samples that were found in granitic pegmatites from the Mount Mica Quarry, Paris, Oxford County, Maine, USA were sent to Eugene E. Foord for study. This became the first recorded case of naturally occurring kosnarite.
Occurrence.
The first naturally occurring kosnarite was discovered in northern Maine in the United States. Another deposit was later found in Black Mountain, Oxford County, Maine. Both of these deposits were found in zoned granitic pegmatites associated with several minerals such as quartz, lepidolite, and beryl. Another deposit was found in Wycheproof, Northern Victoria, Australia, and this sample of kosnarite was also found in granitic pegmatite. Pegmatite is the term for a form of igneous rock with relatively large interlocking crystals, and there are three popular theories on how pegmatites named metamorphic, magmatic, and metasomatic. The pegmatites that the kosnarite are found in are believed to be formed by a mixture of magmatic and metamorphic as the kosnarite forms in the later stages of paragenesis by the alterations of hydrothermal fluids.
Analysis.
Due to the rarity of kosnarite, forms of analysis that involve the powdering of the sample such as X-ray diffraction (XRD) can not be used, so other methods have to be used. To find the data that would have been provided by the XRD, a Gandolfi camera with an 114.6mm diameter was used to find the d-spacing and intensity. Kosnarite's density was found using the sink-float method while using an acetone mixture and methylene iodide. In order to find the chemical composition of kosnarite without damaging any sample, an ARL-SEMQ electron microprobe using an Opus microprobe automation system was used. To help find the chemical composition, CITZAF correction procedures were added to the study. Additionally, emission spectrographic analysis was completed by using a Jarrel-Ash 3.2-m spectrograph using a laser energy source.
Properties of Kosnarite.
The results of the testing found that kosnarite had large amounts of zirconium, phosphorus, and potassium. In the samples, there were also traces of calcium, zinc, and manganese, but these elements make up less than one percent, so they were classified as impurities. Physically, kosnarite from the pegmatites found in Maine occurred as rhombohedral crystals with a hexagonal unit cell and were pseudocubic with a maximum size of around 0.9mm. This is important as zirconium is in the form of a six coordinated octahedron and potassium is structured in a shape called a trigonal antiprism that is also six coordinated. Kosnarite's structure was then determined by using [100] Patterson projection, and interatomic vector projection was also used to help determine the crystal structure. Special positions in the mineral were then found for potassium, zirconium, and phosphate while the two oxygen atoms had general positions. Multiple refinement cycles using a unit weighing system, and least-squares refinements were used to minimize the deviation of the shifting of the atoms. Later tests showed that zirconium polyhedra groups were joined together by groups of phosphate these are both connected to each other by oxygen atoms. However, the oxygens are always bridges between the zirconium and phosphate and one oxygen is never shared between two the same groups and half of the oxygens are also shared with K groups. The pattern of two zirconium polyhedron, one potassium polyhedron, two zirconium polyhedron, and the oxygen and phosphate groups filling in the gaps creates kosnarite's unique crystal structure.
Depending on the impurities present in the sample, the color of kosnarite can range from pale blue to blue-green depending on the amount of iron, manganese, or other impurities, and kosnarite can sometimes appear to be nearly colorless. Other physical properties of kosnarite include its vitreous lust, non-fluorescence, a hardness of 4.5 on the mohs scale of mineral hardness, conchoidal fracturing, and perfect cleavage in the {102} direction. Structurally, kosnarite is part of the hexagonal crystal family meaning that the crystals have three or six-fold symmetry and has a space group of "R"c. The unit cell of kosnarite was calculated to be a = 8.687 Å, c = 23.877 Å; V= 1,560.45 Å³, and no twinning growth has been observed. It was discovered that kosnarite is uniaxial (+) with its axes being Nw = 1.656(2), Nc = 1.682(2), and being nonpleochronic. Due to the small number of available samples, some tests have not been carried out, such as infrared spectra.
Related Minerals.
Kosnarite is part of the alkali zirconium phosphates in which there are only two other known members of this group: gainesite, and a Cs analogue of gainesite.


Trimethylstibine

Trimethylstibine is an organoantimony compound with the formula Sb(CH3)3. It is a colorless pyrophoric and toxic liquid. It is synthesized by treatment of antimony trichloride and methyl Grignard reagent. It is produced by anaerobic bacteria in antimony-rich soils. In contrast to trimethylphosphine, trimethylstibine is a weaker Lewis base. It is used in the production of some III-V semiconductors.


Amorphous uranium(VI) oxide

Amorphous uranium(VI) oxide ("am"-U2O7) is an orange diuranyl compound, most commonly obtained from the thermal decomposition of uranyl peroxide tetrahydrate at temperatures between 150 and 500 °C. It exists at room temperature as a powder. "Am"-U2O7 does not comprise a regular, long-range atomic structure, as demonstrated by its characteristic diffuse scattering pattern obtained by X-ray diffraction. As a result, the molecular structure of this material is little understood, although experimental and computational attempts to elucidate a local atomic environment have yielded some success.
Production.
"Am"-U2O7 is produced by the thermal decomposition of uranyl peroxide tetrahydrate at temperatures between 150 and 500 °C, in either an air or nitrogen atmosphere. The resultant powder is tan orange in colour. Further heating results in the formation of alpha uranium trioxide (α-UO3).
Structure and Reactivity.
Due to the amorphous nature of a"m"-U2O7, the long-range atomic structure of this compound has not been determined. However, recent computational investigations, chiefly accomplished using density functional theory (DFT), have helped to predict a local structure. Resembling a regular uranate compound, two uranyl ([UO2]2+) groups are bridged by a μ2-O atom, where both uranium atoms are bonded to a O-O peroxo unit. In this case, a tetrameric ring would be the most stable conformation of the compound. The presence of a peroxide bond in species obtained in this temperature range is unusual; uranyl peroxide has previously been considered to be the only peroxide bearing uranium compound. Developments on this structure propose a two-site metastudtite and UO3-like bonding environment, including the bond types already mentioned. Few other suggestions for the local atomic structure of a"m"-U2O7 have been made, although a crystalline form of U2O7, calculated as a two-site 6 and 8-coordinate structure, has been reported. In the same study, it was again found that the U2O7 species contained peroxide bonding. "Am"-U2O7 is known to undergo hydrolysis in the presence of water, to produce a crystalline metaschoepite powder. In addition to a change in crystallinity, this reaction involves a change in colour from orange to bright yellow.


3-Bromocytisine

3-Bromocytisine is a derivative of the toxic alkaloid cytisine that acts as a highly potent agonist at neural nicotinic acetylcholine receptors, binding primarily to the α4β2 and α7 subtypes. 3-Bromocytisine is a full agonist at the α7 subtype while it is only a partial agonist at α4β2, but has an extremely strong binding affinity at α4β2 with 200-fold selectivity for α4β2 over α7. In animal studies 3-bromocytisine stimulates the release of dopamine and noradrenaline and increases locomotor activity.


(Terpyridine)ruthenium trichloride

(Terpyridine)ruthenium trichloride is the coordination complex with the formula RuCl3(terpy), where terpy is terpyridine. It is a brown paramagnetic solid that is a precursor to other complexes of ruthenium, mainly by substitution of the chloride ligands. The complex has octahedral geometry. It is prepared by heating ruthenium trichloride with a DMF solution of terpyridine.


Silsesquioxane

A silsesquioxane is an organosilicon compound with the chemical formula [RSiO3/2]n (R = H, alkyl, aryl or alkoxyl). Silsesquioxanes are colorless solids that adopt cage-like or polymeric structures with Si-O-Si linkages and tetrahedral Si vertices. Silsesquioxanes are members of polyoctahedral silsesquioxanes ("POSS"), which have attracted attention as preceramic polymer precursors to ceramic materials and nanocomposites. Diverse substituents (R) can be attached to the Si centers. The molecules are unusual because they feature an inorganic silicate core and an organic exterior. The silica core confers rigidity and thermal stability.
Structure.
Silsesquioxanes are known in molecular form with 6, 8, 10, and 12 Si vertices, as well as polymers. The cages are sometimes labeled T6 T8, T10, and T12, respectively (T = tetrahedral vertex). The T8 cages, the most widely studied members, have the formula [RSiO3/2]8, or equivalently R8Si8O12. In all cases each Si center is bonded to three oxo groups, which in turn connect to other Si centers. The fourth group on Si is usually an alkyl, halide, hydride, alkoxide, etc. In the cubic clusters with "Oh" symmetry the Si-O-Si angles are in the range 145–152°, being bowed out, allowing the Si centers to better adopt tetrahedral geometry. The O-Si-O angle are in the range: 107–112°, Si-O bond: 1.55–1.65 Å.
Synthesis.
Silsesquioxanes are usually synthesized by hydrolysis of organotrichlorosilanes. An idealized synthesis is:
The formation of HCl negatively impacts the relative rates of hydrolysis and condensation of intermediate silanols. Consequently, silsesquioxanes can be obtained directly by condensation of the corresponding silanetriols which occurs at neutral pH and works even for sterically very bulky substituents.
Depending on the R substituent, the exterior of cage can be further modified. When R = H, the Si-H group can undergo hydrosilylation or oxidation to the silanol. Bridged polysilsesquioxanes are most readily prepared from clusters that contain two or more trifunctional silyl groups attached to non-hydrolysable silicon-carbon bonds, with typical sol-gel processing. Vinyl-substituted silsesquioxanes can be linked by the alkene metathesis.
Reactivity.
Cage-Rearrangement.
Reorganization of the siloxane
cage-like core (T8 → T10) can be performed, including
isolation of intermediates, and cage rearrangement achieved by using Bronsted
superacid, trifluoromethanesulfonic acid (CF3SO3H). In this case, reaction of hexahedral silsesquioxane and CF3SO3H in DMSO conducted in 1 : 12 molar ratio gives heptahedral silsesquioxane. In the first step CF3SO3H acid attacks siloxane Si-O-Si bonds and the formation of Si-O-SO2CF3 bond parallel with cage opening process is observed and compound B is obtained (Figure below). Such an inversion is observed at silicon atom during nucleophilic displacement reaction that is usually noticed when leaving groups are replaced by soft nucleophiles. Uponfurther acid attack, both T6(OH)4 C and siloxane dimer D are formed. Because this reaction takes place in an aqueous conditions, compound E of general formula T8(OH)4 as a consequence of hydrolysis reaction was obtained. E is prone to reaction with D
and due to this, the abstraction of CF3SO anion occurs and the closure frame with the spontaneous cage-rearrangement to heptahedral T10 structure F is observed. Although, heptahedral F is less favorable energetically (MM2 data), in this case its creation is forces by the formation of a new Si4O4 moiety from much more less stable substrates D and E.
Polymeric silsesquioxanes.
Polymeric silsesquioxanes have been reported, first by Brown. High molecular weight tractable polymeric phenyl silsesquioxane featured a ladder-type structure. Brown's findings were the basis for many future investigations. Brown's synthesis proceeded in three-steps:
(1) the hydrolysis of phenyltrichlorosilane,
(2) equilibration of this hydrolyzate with potassium hydroxide at a low concentration and temperature to give the prepolymer, and
(3) equilibration of the prepolymer at a high concentration and temperature to give the final polymer.
Other notable silsesquioxane polymers include the soluble polymethylsilsesquioxane with high molecular weights described by Japan Synthetic Rubber. This polymer which, unlike its phenyl derivative, gels easily during the course of its synthesis, has found applications in cosmetics, resins, and lithography.
Hydridosilsesquioxanes.
A well known hydrogen silsesquioxane is [HSiO3/2]8. Early syntheses involved treatment of trichlorosilane with concentrated sulfuric acid, and fuming sulfuric acid, affording T10-T16 oligomers. The T8 cluster was also synthesized by the reaction of trimethylsilane with a mixture of acetic acid, cyclohexane, and hydrochloric acid. The Si-H groups are amenable to hydrosilylation.
Potential applications.
Electronic materials.
Films of organosilsesquioxane, e.g., poly(methylsilsesquioxane), have been examined for semiconducting devices. Poly(hydridosilsesquioxane), which has a linked-cage structure, was sold under the name Fox Flowable Oxide.
Methylsilsesquioxanes have been examined for spin-on-glass (SOG) dielectrics. Bridged silsesquioxanes have been used for quantum confined nano-size semiconductors. Silsesquioxane resins have also been used for these applications because they have high dielectric strengths, low dielectric constants, high volume resistivities, and low dissipation factors, making them very suitable for electronics applications. These resins have heat and fire resistant properties, which can be used to make fiber-reinforced composites for electrical laminates.
Polyhedral oligomeric silsesquioxanes have been examined as a means to give improved mechanical properties and stability, with an organic matrix for good optical and electrical properties. The mechanisms of degradation in these devices is not well understood, but it is believed that material defect understanding is important for understanding the optical and electronic properties.
Hydridosilsesquioxanes can be converted to silica coatings for potential application in integrated circuits.
LEDs.
For potential applications to light emitting diodes, cubic silsesquioxanes. have been functionalized. One of the first precursors used in light emitting application was octadimethylsiloxysilsesquioxane, which can be prepared in yields of >90% by treating tetraethoxysilane or rice hull ash with tetramethylammonium hydroxide followed by dimethylchlorosilane. The general method of hydrolyzing organotrichlorosilanes is still effective here. When brominated or aminated, these structures can be coupled with epoxies, aldehydes, and bromoaromatics, which enable attachment of these silsesquioxanes to π-conjugated polymers. These methods can use copolymerization techniques, Grignard reagents, and different coupling strategies. There has also been research on the ability of conjugated dendrimer silsesquioxanes to behave as light emitting materials. Though, highly branched substituents tend to have π-π interactions, which hinder high luminescent quantum yield.
Sensors.
For chemosensor applications, silsesquioxane cages conjugated with fluorescent molecules can be directly used to detect fluoride ions under a cage-encapsulation showing a change of color under naked eyes and other anions.
Antimicrobial silsesquioxanes.
Silsesquioxanes have been functionalized with biocidal quaternary ammonium (QASs) groups to produce antimicrobial coatings. QASs are disinfectants, antiseptics, and antifoulants that kill bacteria, fungi, and algae. The relatively small size of the silsesquioxane molecule, 2-5 nm, allows a QAS functionalized molecule to have a charge density similar to dendrimers and thus the antimicrobial efficacy is prominent. Dimethyl-n-octylamine was quaternized by octa(3-chloropropylsilsesquioxane), (T-ClPr)8. The resulting material exhibited antimicrobial efficacy for the prevention of growth of both Gram-positive and Gram-negative bacteria.
Array of QAS functionalized polyhedral oligomeric silsesquioxanes (Q-POSS) have been reported. These researchers varied the alkyl chain length from –C12H25 to –C18H37 and varied the counter ion between chloride, bromide, and iodine. The first reaction was the hydrosilylation between allydimethlamine and octasilane polyhedral oligomeric silsesquioxane via Karstedt's catalyst to make a tertiaryamino-functinoalized silsesquioxane. The second step was the quaternization of the tertiaryamino groups with an alkyl halide. The alkyl halides used were 1-iodooctadecane, 1-bromohexadecane, and 1-chloroctadecane.
The silsesquioxane core in these hybrid materials provides an increased glass transition temperature, improved mechanical properties, higher use temperature, and lower flammability. These desirable properties combined with the ability to readily functionalize a silsesquioxane with multiple antimicrobial groups allows for robust biocides with higher charge densities while maintaining a compact molecular structure. The organic functionalities provide high compatibility with polymers allowing for easy incorporation into many mediums. Of particular interest are silicone paints and coatings used in hospitals. Typical biocidal ammonium functionalized polymers are incompatible, but silsesquioxanes closely mimic the silicone structure. A silicone-based paint combined with QAS-functionalized silsesquioxanes could be used to paint medical and sanitary devices, biomedical devices, exam equipment, medical storage rooms, hospital rooms, clinics, doctor offices, etc. to prevent the formation and spread of bacteria. For example, the Q-POSS developed was combined with polydimethylsiloxane and catalysis to form a crosslinked network. The researchers found that coatings based on bromide and chloride had the best antimicrobial efficacy.
Partially condensed silsesquioxanes: Si7 species.
A well studied example of a partially condensed silsesquioxanes is the trisilanol Cy7Si7O9(OH)3, prepared by the slow (months) hydrolysis of trichlorocyclohexylsilane (C6H11SiCl3). The same cage can be prepared by acid-mediated cleavage of fully condensed silsesquioxane. This process results in silanediols that can further be used to create new metallasilsesquioxanes. These partially condensed silsesquioxanes are intermediates en route to the fully condensed cages.
In general, such silsesquioxane trisilanols form discrete dimers in the solid held together by cooperatively enhanced cyclic hydrogen bonded networks. These dimers are retained in solution and a dynamic equilibrium has been elucidated using NMR.
Other partially condensed silsesquioxanes.
Other partially condensed species adopt ladder structures wherein in which two long chains composed of RSiO3/2 units are connected at regular intervals by Si-O-Si bonds. Amorphous structures include RSiO3/2 unit connections without any organized structure formation.
Metal complexes of partially condensed silsesquioxanes.
The incompletely condensed silsesquioxanes bind numerous metals, including Na+, Li+, and Be2+ as well as transition metals. Cubic metal-silsesquioxane derivatives of the core stoichiometry MSi7O12 can be prepared by treating the incomplete cage with a metal halide in the presence of a base such as triethylamine. Another route of synthesis involves first deprotonating the trisilanol group using LiN(SiMe3)2. Aspinall et al. later succeeded in doing the same using three equivalents of n-BuLi in hexanes and further results indicate that alkali metal derivatives of deprotonated silsesquioxanes could also be prepared using alkali metalbis(trimethylsilyl) amides.
Catalytic properties.
Although lacking commercial applications, metallasilsesquioxanes have been investigated as catalysts. The coordination environment provided by Cy7Si7O9(OH)3 has been proposed to approximate β-tridymite and β-cristobalite. Some of these complexes are active as catalysts for alkene metathesis, polymerization, epoxidation and Diels-Alder reactions of enones, as well as other Lewis acid-catalyzed reactions like Oppenauer oxidation and Meerwein-Pondorf-Verley reductions.
. A number of metallasilsesquioxanes have been reported that can polymerize ethene, akin to the Phillips catalyst. The catalyst can be easily activated with trimethylaluminum and typically proceeds with high turnover number. Vanadium complexes as well as Ziegler–Natta type catalysts also catalyze the polymerization of ethylene. The coordination of metals to the silsesquioxane framework gives electrophilic centers that are approximately as electron-withdrawing as a trifluoromethyl group, leading to increased catalytic activity.


Lead(II) oxalate

Lead(II) oxalate is an organic compound with the formula PbC2O4. It is naturally found as a heavy white solid.
Preparation.
This compound is commercially available. It may be prepared by the metathesis reaction between lead(II) nitrate and sodium oxalate:
Solubility.
Lead(II) oxalate is sparingly soluble in water. Its solubility is increased in presence of excess oxalate anions, due to the formation of the Pb(C2O4)22− complex ion.


Weissite

Weissite is a telluride mineral, a copper telluride. Its chemical formula is . Weissite has hexagonal crystal structure. Its specific gravity is 6 and its Mohs hardness is 3. Occurrence is in Gunnison County, Colorado, Arizona and New Mexico in the United States. It is also reported from Kalgoorlie, Western Australia and Dalarna and Värmland, Sweden.
Weissite occurs in hydrothermal deposits associated with pyrite, native tellurium, sylvanite, petzite, rickardite, native sulfur, native gold, calaverite and krennerite.
It was first described in 1927 for an occurrence in the Good Hope Mine in the Vulcan District of Gunnison County, Colorado. It was named for mine owner Louis Weiss.


Polysulfane

A polysulfane is a chemical compound of formula H2S"x", where x > 1, although sometimes disulfane, H2S2 is excluded. Polysulfanes consist of unbranched chains of sulfur atoms terminated with hydrogen. Compounds containing 2–8 concatenated sulfur atoms have been isolated, longer chain compounds have been detected, but only in solution. H2S2 is colourless, higher members are yellow with the colour increasing with the sulfur content. Even a trace of alkali will cause chemical decomposition, and containers need to be treated with acid to remove any traces of alkali.In the chemical literature the term polysulfanes is sometimes used for compounds containing −(Sn)−, e.g. organic polysulfanes R−(Sn)−R.
Since sulfur is a chalcogen, polysulfanes can be classed as hydrogen chalcogenides.
Chemistry and properties.
Polysulfanes are thermodynamically unstable with respect to decomposition (disproportionation) to H2S and sulfur:
However, the production of polysulfide ions from S2− and S is thermodynamically favourable:
Attempting to make a polysulfane by acidifying an alkali metal polysulfide salt (e.g. sodium disulfide) simply produces hydrogen sulfide, H2S, and sulfur. Polysulfanes can be made from polysulfides by pouring polysulfide into cooled concentrated hydrochloric acid to produce a mixture of metastable polysulfanes as a yellow oil, from which individual compounds may be separated by fractional distillation. Other more selective syntheses are:
The reaction of polysulfanes with sulfur dichloride or disulfur dichloride produces long-chain dichloropolysulfanes:
The reaction with a sulfite salt (a base) quantitatively decomposes the polysulfane to produce thiosulfate and hydrogen sulfide:


Samarium monochalcogenides

Samarium monochalcogenides are chemical compounds with the composition SmX, where Sm stands for the lanthanide element samarium and X denotes any one of three chalcogen elements, sulfur, selenium or tellurium, resulting in the compounds SmS, SmSe or SmTe. In these compounds, samarium formally exhibits oxidation state +2, whereas it usually assumes the +3 state, resulting in chalcogenides with the chemical formula Sm2X3.
Synthesis.
Single crystals or polycrystals of samarium monochalcogenides can be obtained by reacting the metal with sulfur, selenium or tellurium vapors at high temperature. Thin films can be obtained by magnetron sputtering or electron beam physical vapor deposition, that is bombardment of samarium metal target with electrons in and appropriate gas atmosphere (e.g. hydrogen disulfide for SmS).
Properties.
Samarium monochalcogenides are black semiconducting solids with rock-salt cubic crystal structure. Application of moderate hydrostatic pressure converts them into metals. Whereas the transition is continuous and occurs at about 45 and 60 kbar in SmSe and SmTe, respectively, it is abrupt in SmS and requires only 6.5 kbar. Similar effect is observed in monochalcogenides of another lanthanide, thulium. This results in spectacular change in color from black to golden yellow when scratching or mechanically polishing SmS. The transition does not change the crystal structure, but there is a sharp decrease (about 15%) in the crystal volume. A hysteresis is observed, that is when the pressure is released, SmS returns to semiconducting state at much lower pressure of about 0.5 kbar.
Not only color and electrical conductivity, but also other properties change in samarium monochalcogenides with increasing pressure. Their metallic behavior results from the decreasing band gap, which amounts at zero pressure to 0.15, 0.45 and 0.65 eV in SmS, SmSe and SmTe, respectively.
At the transition pressure (6.5 kbar in SmS) the gap is still finite and the low resistivity originates from thermally activated generation of carriers across a narrow band gap. The gap collapses at about 20 kbar when SmS becomes a true metal. At this pressure, the material also changes from paramagnetic to a magnetic state.
The semiconductor-metal transition in samarium monochalcogenides requires application of pressure or presence of intrinsic stress, for example in thin films, and the reverse changes occur upon release of this stress. Such release can be triggered by various means, such as heating to about 200 °C or irradiation with a pulsed, high-intensity laser beam.
Potential applications.
The change in electrical resistivity in samarium monochalcogenides can be used in a pressure sensor or in a memory device triggered between a low-resistance and high-resistance state by external pressure, and such devices are being developed commercially. Samarium monosulfide also generates electric voltage upon moderate heating to about 150 °C that can be applied in thermoelectric power converters.


Lead(II) sulfate

Lead(II) sulfate (PbSO4) is a white solid, which appears white in microcrystalline form. It is also known as "fast white", "milk white", "sulfuric acid lead salt" or "anglesite".
It is often seen in the plates/electrodes of car batteries, as it is formed when the battery is discharged (when the battery is recharged, then the lead sulfate is transformed back to metallic lead and sulfuric acid on the negative terminal or lead dioxide and sulfuric acid on the positive terminal). Lead sulfate is poorly soluble in water.
Manufacturing.
Lead(II) sulfate is prepared by treating lead oxide, hydroxide or carbonate with warm sulfuric acid or by treating a soluble lead salt with sulfuric acid.
Alternatively, it can be made by the interaction of solutions of lead nitrate and sodium sulfate.
Toxicology.
Lead sulfate is toxic by inhalation, ingestion and skin contact. It is a cumulative poison, and repeated exposure may lead to anemia, kidney damage, eyesight damage or damage to the central nervous system (especially in children). It is also corrosive - contact with the eyes can lead to severe irritation or burns. Typical threshold limit value is 0.15 mg/m3.
Mineral.
The naturally occurring mineral anglesite, PbSO4, occurs as an oxidation product of primary lead sulfide ore, galena.
Basic and hydrogen lead sulfates.
A number of lead basic sulfates are known: PbSO4·PbO; PbSO4·2PbO; PbSO4·3PbO; PbSO4·4PbO. They are used in manufacturing of active paste for lead–acid batteries. A related mineral is leadhillite, 2PbCO3·PbSO4·Pb(OH)2.
At high concentration of sulfuric acid (>80%), lead hydrogensulfate, Pb(HSO4)2, forms.
Chemical properties.
Lead(II) sulfate can be dissolved in concentrated HNO3, HCl, H2SO4 producing acidic salts or complex compounds, and in concentrated alkali giving soluble tetrahydroxidoplumbate(II) [Pb(OH)4]2− complexes.
Lead(II) sulfate decomposes when heated above 1000 °C:


Oxosilanol

Oxosilanol is a chemical compound. It is the silicon-analogue to formic acid, with silicon replacing carbon.


Hydrocalcite


Brown carbon

In chemistry, brown carbon (Cbrown/BrC) is brown smoke released by the combustion of organic matter. It coexists with black carbon when released in the atmosphere.
Black carbon is primarily released by high-temperature combustion and brown carbon is emitted mainly by biomass combustion. These two are the two most important light absorbing substances in the atmosphere. The climate and radiative transfer are highly impacted by the absorptive properties of these substances.
Overview.
Discovery.
Light absorbing aerosols have become an interest of study because of its effects on atmospheric warming.
While the function of black carbon created by biomass burning in relation to the warming of air has as of now been known, the lesser-known part of brown carbon has recently been discovered.
Brown carbon aerosols.
The way light reflects off brown carbon causes the material to appear brown or yellow. In respect to soot carbon (as an aerosol called black carbon) these light reflecting particles are collectively called brown carbon, highlighting their optical properties.
It was seen that particles from combustion or from residential coal usage can contain substantial amounts of brown carbon. This particulate matter appears light brown, and not black as would be expected for pure soot particles.
Effects.
Influence on Earth's atmosphere.
Aerosols are one of the most important contributors to climate change in the atmosphere. Many of the particles suspended in the atmosphere originate from a multitude of sources. Some of these sources are natural and some of them are man-made. Most aerosols reflect sunlight, and some also absorb it. A lot of these nanoparticles cause severe health effects in addition to climate effects. There are many human activities that produce these particles.
Black carbon particles (a component of soot) originating from combustion processes have been known for some time to absorb sunlight and warm the atmosphere, and pollution controls have been put into place to reduce their emissions and their effects.
Brown carbon has attracted interest as a possible cause of climate change. This class of organic carbon, known for its brownish color, absorbs strongly in the ultraviolet wavelengths and less significantly going into the visible wavelengths. Types of brown carbon include tar materials from smoldering fires or coal combustion, breakdown products from biomass burning, a mixture of organic compounds emitted from soil, and volatile organic compounds given off by vegetation.
Black and brown carbon.
Brown carbon in relation to black carbon.
Simulation models suggest that brown carbon contributes about 19% of the total atmospheric absorption by aerosols, while 72% is attributed to black carbon and 9% is due to the coating effect of sulfate and organic aerosols on black carbon. It seems that brown carbon can play an important role in photochemistry and the hydrologic cycle, especially over regions dominated by biomass combustion. Therefore, brown carbon needs to be considered in global climate change simulations for a more accurate understanding.
Mixture of black and brown.
Up to about 70% of light absorption is by black carbon. When Brown carbon is present independently it has nearly 15% potential to warm the atmosphere by absorbing light.


Fulleride

Fullerides are chemical compounds containing fullerene anions. Common fullerides are derivatives of the most common fullerenes, i.e. C60 and C70. The scope of the area is large because multiple charges are possible, i.e., [C60]"n"− ("n" = 1, 2...6), and all fullerenes can be converted to fullerides. The suffix "-ide" implies their negatively charged nature.
Fullerides can be isolated as derivatives with a wide range of cations. Most heavily studied derivatives are those with alkali metals, but fullerides have been prepared with organic cations. Fullerides are typically dark colored solids that generally dissolve in polar organic solvents.
Structure and bonding.
According to electronic structure calculations, the LUMO of C60 is a triply degenerate orbital of t1u symmetry. Using the technique cyclic voltammetry, C60 can be shown to undergo six reversible reductions starting at −1 V referenced to the Fc+/Fc couple. Reduction causes only subtle changes in the structure and many derivatives exhibit disorder, which obscures these effects. Many fullerides are subject to Jahn–Teller distortion. In certain cases, e.g. [PPN]2C60, the structures are highly ordered and slight (10 pm) elongation of some C−C bonds is observed.
Preparation.
Fullerides have been prepared in various ways:
The fulleride salt ([K(crypt-222)]+)2[C60]2− salt is synthesized by treating C60 with metallic potassium in the presence of [2.2.2]cryptand.
Alkali metal derivatives.
Particular attention has been paid to alkali metal (Na+, K+, Rb+, Cs+) derivatives of C603− because these compounds exhibit physical properties resulting from intercluster interactions such as metallic behavior. In contrast, in C60, the individual molecules interact only weakly, i.e. with essentially nonoverlapping bands. These alkali metal derivatives are sometimes viewed as arising by intercalation of the metal into C60 lattice. Alternatively, these materials are viewed as n-doped fullerenes.
Alkali metal salts of this trianion are superconducting. In M3C60 (M = Na, K, Rb), the M+ ions occupy the interstitial holes in a lattice composed of ccp lattice composed of nearly spherical C60 anions. In Cs3C60, the cages are arranged in a bcc lattice.
In 1991, it was revealed that potassium-doped C60 becomes superconducting at . This was the highest transition temperature for a molecular superconductor. Since then, superconductivity has been reported in fullerene doped with various other alkali metals. It has been shown that the superconducting transition temperature in alkaline-metal-doped fullerene increases with the unit-cell volume V. As Cs+ is the largest alkali ion, caesium-doped fullerene is an important material in this family. Superconductivity at has been reported in bulk Cs3C60, but only under applied pressure. The highest superconducting transition temperature of at ambient pressure is reported for Cs2RbC60.
The increase of transition temperature with the unit-cell volume had been believed to be evidence for the BCS mechanism of C60 solid superconductivity, because inter C60 separation can be related to an increase in the density of states on the Fermi level, "N"("ε"F). Therefore, efforts have been made to increase the interfullerene separation, in particular, intercalating neutral molecules into the A3C60 lattice to increase the interfullerene spacing while the valence of C60 is kept unchanged. However, this ammoniation technique has revealed a new aspect of fullerene intercalation compounds: the Mott transition and the correlation between the orientation/orbital order of C60 molecules and the magnetic structure.
Fourfold-reduced materials, i.e., those with the stoichiometry A4C60, are insulating, even though the t1u band is only partially filled. This apparent anomaly may be explained by the Jahn–Teller effect, where spontaneous deformations of high-symmetry molecules induce the splitting of degenerate levels to gain the electronic energy. The Jahn–Teller type electron-phonon interaction is strong enough in C60 solids to destroy the band picture for particular valence states.
A narrow band or strongly correlated electronic system and degenerated ground states are relevant to explaining superconductivity in fulleride solids. When the interelectron repulsion "U" is greater than the bandwidth, an insulating localized electron ground state is produced in the simple Mott–Hubbard model. This explains the absence of superconductivity at ambient pressure in caesium-doped C60 solids. Electron-correlation-driven localization of the t1u electrons exceeds the critical value, leading to the Mott insulator. The application of high pressure decreases the interfullerene spacing, therefore caesium-doped C60 solids turn to metallic and superconducting.
A fully developed theory of C60 solids superconductivity is lacking, but it has been widely accepted that strong electronic correlations and the Jahn–Teller electron–phonon coupling produce local electron pairings that show a high transition temperature close to the insulator–metal transition.


Nitrogen monofluoride

Nitrogen monofluoride (fluoroimidogen) is a metastable species that has been observed in laser studies. It is isoelectronic with O2. Like boron monofluoride, it has unusual higher than single-bonded fluorine. It is unstable with respect to its formal dimer, dinitrogen difluoride, as well as to its elements, nitrogen and fluorine. Nitrogen monofluoride can be formed by the decomposition of fluorine azide into N2F2 and N2. It is also produced when various radical species (H, O, N, CH3) react with nitrogen difluoride (NF2) to abstract one fluorine atom. Many of the reactions give the product in an excited state that has a characteristic chemiluminescence, and have thus been investigated for development as a chemical laser. The reactions are highly efficient and the product has a long lifetime. The reaction with molecular hydrogen (H2) involves a chain propagation via regeneration of the atomic hydrogen radical that can perpetuate for many cycles. An alternative azide process involves the reaction of atomic fluorine with hydrazoic acid to give an azide radical, which then reacts with another atomic fluorine to produce NF with N2 as a byproduct. This route avoids the necessity of using atomic hydrogen, a chemical that could otherwise cause the decomposition of NF.


Carbon planet

A carbon planet is a theoretical type of planet that contains more carbon than oxygen. Carbon is the fourth most abundant element in the universe by mass after hydrogen, helium, and oxygen.
Marc Kuchner and Sara Seager coined the term "carbon planet" in 2005 and investigated such planets following the suggestion of Katharina Lodders that Jupiter formed from a carbon-rich core.
Prior investigations of planets with high carbon-to-oxygen ratios include Fegley & Cameron 1987. Carbon planets could form if protoplanetary discs are carbon-rich and oxygen-poor. They would develop differently from Earth, Mars, and Venus, which are composed mostly of silicon–oxygen compounds. Different planetary systems have different carbon-to-oxygen ratios, with the Solar System's terrestrial planets closer to being "oxygen planets" with C/O molar ratio of 0.55. In 2020, survey of the 249 nearby solar analog stars found 12% of stars have C/O ratios above 0.65, making them candidates for the carbon-rich planetary systems. The exoplanet 55 Cancri e, orbiting a host star with C/O molar ratio of 0.78, is a possible example of a carbon planet.
Definition.
Such a planet would probably have an iron- or steel-rich core like the known terrestrial planets. Surrounding that would be molten silicon carbide and titanium carbide. Above that, a layer of carbon in the form of graphite, possibly with a kilometers-thick substratum of diamond if there is sufficient pressure. During volcanic eruptions, it is possible that diamonds from the interior could come up to the surface, resulting in mountains of diamonds and silicon carbides. The surface would contain frozen or liquid hydrocarbons (e.g. tar and methane) and carbon monoxide. A weather cycle is theoretically possible on carbon planets with an atmosphere, provided that the average surface temperature is below 77 °C.
However, carbon planets will probably be devoid of water, which cannot form because any oxygen delivered by comets or asteroids will react with the carbon on the surface. The atmosphere on a relatively cool carbon planet would consist primarily of carbon dioxide or carbon monoxide with a significant amount of carbon smog.
Composition.
Carbon planets are predicted to be of similar diameter to silicate and water planets of the same mass, potentially making them difficult to distinguish. The equivalents of geologic features on Earth may also be present, but with different compositions. For instance, the rivers might consist of oils. If the temperature is low enough (below 350 K), then gasses may be able to photochemically synthesize into long-chain hydrocarbons, which could rain down onto the surface.
In 2011, NASA cancelled a mission, called TPF, which was to be an observatory much bigger than the Hubble Space Telescope that would have been able to detect such planets. The spectra of carbon planets would lack water, but show the presence of carbonaceous substances, such as carbon monoxide.
Possible candidates.
The pulsar PSR 1257+12 may have carbon planets that formed from the disruption of a carbon-producing star. Carbon planets might also be located near the Galactic Center or globular clusters orbiting the galaxy, where stars have a higher carbon-to-oxygen ratio than the Sun. When old stars die, they spew out large quantities of carbon. As time passes and more and more generations of stars end, the concentration of carbon, and carbon planets, will increase.
In October 2012, it was announced that 55 Cancri e showed evidence for being a carbon planet. It has eight times the mass of Earth, and twice the radius. Research indicates that the planet is "covered in graphite and diamond rather than water and granite". It orbits the star 55 Cancri once every 18 hours.
Other carbon-rich objects.
In August 2011, Matthew Bailes and his team of experts from Swinburne University of Technology in Australia reported that the millisecond pulsar PSR J1719-1438 may have a binary companion star that has been crushed into a much smaller planet made largely of solid diamond. They deduced that a small companion planet must be orbiting the pulsar and causing a detectable gravitational pull. Further examination revealed that although the planet is relatively small (60,000 km diameter, or five times bigger than the Earth) its mass is slightly more than that of Jupiter. The high density of the planet gave the team a clue to its likely makeup of carbon and oxygen—and suggested the crystalline form of the elements. However, this "planet" is theorized to be the remains of an evaporated white dwarf companion, being only the remnant inner core. According to some definitions of planet, this would not qualify because it formed as a star.
At a distance of 267 pc (approximately 870 light-years), PSR J2222−0137 is a nearby intermediate-mass binary pulsar whose low-mass neutron star's companion is a white dwarf (PSR J2222−0137 B) that, with a temperature less than 3,000 K and a mass of 1.05 formula_1 0.06 formula_2, is likely crystallized, leading to this Earth-sized white dwarf being described as a "diamond-star".
Brown dwarfs.
Planets around brown dwarfs are likely to be carbon planets depleted of water.


Pentavalent uranyl complexes

UV is the +5 oxidation state of uranium which is found in the form of [UO2]1+. This species is known as pentavalent uranyl cation and has a low stability due to the disproportionation into tetravalent and hexavalent uranium species.
The oxidation states of uranium compounds vary from +2 to +6 (UII - UVI). Out of all these oxidation states +6 is the most common one and it is found in the form of [UO2]2+ (uranyl dication). This hexavalent form of uranium species is known to be the most thermodynamically stable species and it is soluble in water. Even though the uranium chemistry is well explored pentavalent uranyl chemistry is under developed due to the low stability of these compounds. There are two naturally occurring pentavalent uranium-containing minerals reported.
Synthesis.
Pentavalent uranyl complexes are obtained by the oxidation of UIV or other low-valent uranium compounds and these reactions are done under oxygen-free conditions in organic solvents. The oxidation of UIII and UIV precursors under non-aqueous conditions or controlled hydrolysis conditions gives large polyoxo uranium clusters containing UV. Hydrothermal synthesis is another method used to obtain UV silicates where the reactions are carried out under high temperatures like 600 °C.
Another way to make pentavalent uranyl complexes is the reduction of UVI containing [UO2]2+. In this method hexavalent uranyl complexes are reduced to UV complexes in aqueous solutions using electrochemical or photochemical reductions in the presence of a stabilizing carbonate or organic ligand.
An example for the use of electrochemical reduction method to obtain pentavalent uranyl complexes is shown below:
Hexavalent uranyl complexes with different β-diketonate ligands (R= Me, Ph, CF3) are used for this reduction of uranyl dication ([UO2]2+). The introduction of electron withdrawing substituents at the R position of the β-diketonate ligand favors the reduction of hexavalent uranyl complexes. These electrochemical reductions are carried out with
CoCp2⋅ or CoCp2 in toluene and the reduction potential is measured using cyclic voltammetry.
Coordination Chemistry.
Given below are some of the pentavalent uranyl compounds reported recently.
The electrochemical reduction of [UO2]2+ in an anaerobic solution is used to obtain pentavalent uranyl complexes (Figure 2. Compound 1-3). The coordination numbers of complexes 1-4 are 7, 7, 8 and 6. The rate of disproportionation of UV which results UIV and UVI is reported to be slow in an aqueous medium with pH 2-4 hence UV is obtained via electrochemical reduction. In these compounds tetradentate Schiff-base phenoxide (salophen) and β-diketonate ligands are bound along the equatorial plane, which stabilizes the UV species obtained by the single electron reduction of UVI species. Anhydrous uranyl triflate compounds (Figure 2, compound 4) contains discrete cation-anion pairs and it is synthesized by the crystallization of [UO2(OPPh3)4][OTf]2. UV containing cation in this uranyl triflate compound shows an octahedral geometry.
Hazards.
All uranium compounds are radioactive and toxic. Uranium compounds cause blood disorders, chronic kidney disease, acute tubular necrosis, chronic fatigue and congenital disorders. The target organs are kidneys, liver, lungs and brain. Uranyl salts are neurotoxins and accumulates in bones, liver, kidneys, and reproductive tissues. UVI in the form of [UO2]2+ being soluble in water acts as a ground water contaminant.


Formate

Formate (IUPAC name: methanoate) is the anion derived from formic acid. Its formula is represented in various equivalent ways: HCOO− or CHOO− or HCO2−. It is the product of deprotonation of formic acid. It is the simplest carboxylate anion. A formate (compound) is a salt or ester of formic acid.
Biochemistry.
Formate is reversibly oxidized by the enzyme formate dehydrogenase from "Desulfovibrio gigas":
Formate esters.
Formate esters have the formula ROC(O)H (alternative way of writing formula RO2CH). Many form spontaneously when alcohols dissolve in formic acid.
The most important formate ester is methyl formate, which is produced as an intermediate en route to formic acid. Methanol and carbon monoxide react in the presence of a strong base, such as sodium methoxide:
Hydrolysis of methyl formate gives formic acid and regenerates methanol:
Formic acid is used for many applications in industry.
Formate esters often are fragrant or have distinctive odors. Compared to the more common ethyl esters, formate esters are less commonly used commercially because they are less stable. Ethyl formate is found in some confectionaries.
Formate salts.
Formate salts have the formula M(O2CH)(H2O)"x". Such salts are prone to decarboxylation. For example, hydrated nickel formate decarboxylates at about 200 °C to give finely powdered nickel metal:
Such fine powders are useful as hydrogenation catalysts.


Wurtzite

Wurtzite is a zinc and iron sulfide mineral ((Zn,Fe)S), a less frequently encountered structural polymorph form of sphalerite. The iron content is variable up to eight percent. It is trimorphous with matraite and sphalerite.
It occurs in hydrothermal deposits associated with sphalerite, pyrite, chalcopyrite, barite and marcasite. It also occurs in low-temperature clay-ironstone concretions.
It was first described in 1861 for an occurrence in the San José Mine, Oruro City, Cercado Province, Oruro Department, Bolivia, and named for French chemist Charles-Adolphe Wurtz. It has widespread distribution. In Europe it is reported from Příbram, Czech Republic; Hesse, Germany; and Liskeard, Cornwall, England. In the US it is reported from Litchfield County, Connecticut; Butte, Silver Bow County, Montana; at Frisco, Beaver County, Utah; and from the Joplin district, Jasper County, Missouri.
Wurtzite structure.
The wurtzite group includes cadmoselite (CdSe), greenockite (CdS), mátraite (ZnS), and rambergite (MnS), in addition to wurtzite.
Its crystal structure is called the wurtzite crystal structure, to which it lends its name. This structure is a member of the hexagonal crystal system and consists of tetrahedrally coordinated zinc and sulfur atoms that are stacked in an ABABABABAB pattern.
The unit cell parameters of wurtzite are (-2H polytype):


Metaphosphate

A metaphosphate ion is an oxyanion that has the empirical formula . The structure of a metaphosphate ion can be described as being made up of PO4 structural units in which each unit shares two corners with another unit. This can come about in two ways.
Metaphosphates can be considered as salts of the corresponding metaphosphoric acids (H"n"P"n"O3"n") although none of these acids has been isolated. The metaphosphoric acids can be formulated as H2O·P2O5. In comparison, phosphoric acid, H3PO4 can be formulated as 3H2O·P2O5 and pyrophosphoric acid, H4P2O7, as 2H2O·P2O5.
Metaphosphates can be used as an alternative of white phosphorus in organic syntheses.