Serrabrancaite

Serrabrancaite is a mineral with the chemical formula MnPO4•H2O and which is named for the locality where it was found, the Alto Serra Branca Pegmatite. The Alto Serra Branca mine has been in operation since the 1940s. It is located in Paraiba, Brazil near a village named Pedra Lavrada. Tantalite is the main mineral mined here. Specimens of serrabrancaite are kept in the Mineralogical Collections of both the Bergakademie Freiberg, Germany and the Martin-Luther Universität Halle, Institut für Geologische Wissenschaften.
Occurrence.
The Alto Serra Branca Mine occurs where a granitic pegmatite intruded upon a biotite schist. There are many minerals found within this pegmatite. The main constituents are quartz, muscovite, microcline, and albite and amblygonite are common in smaller quantities. The minerals apatite, beryl, uraninite, and elbaite along with the tantalite group and other secondary uranium minerals are found throughout the mine. Near the center of the open portion of the mine, there are growths of two phosphates: triphylite and triplite. Serrabrancaite and other minerals (purpurite, rockbridgeite, tavorite, barbosalite, hureaulite, eosphorite, and phosphosiderite) occur as alteration products of these two phosphates. Serrabrancaite is altered triplite that is closely associated with vernadite. There were also associations with carlhintzeite, colquiriite, pachnolite, ralstonite, and fluellite.
Physical and optical properties.
Serrabrancaite has monoclinic crystals that range in size up to 0.3 mm. The dark brown to dark greenish-black crystals have an adamantine luster. It will streak olive green, and, on the mohs hardness scale, is a 3.5. In thin fragments, they appear translucent. It is a brittle mineral that has uneven fracture, and it does not demonstrate cleavage. Using the float-sink method with sodiumpolytungstate, it was found to have a density of 3.17(1) g/cm3. The calculated value, 3.16 g/cm3, is very similar to the value found using the float-sink method. Using wavelengths of 589 nm and immersion liquids, the refractive indices were measured at ɑ= 1.75(1) β= 1.79(1) and ɤ< 1.79; however, ɤ could not be measured with great accuracy. Its pleochroism ranges from dark brown to olive greenish-brown.
Mineral chemistry and thermal analysis.
In order to analyze the chemical composition of serrabrancaite, both energy-dispersive X-ray spectroscopy (EDS) and wavelength-dispersive X-ray spectroscopy (WDS) were used. The EDS results showed only manganese (Mn), phosphate (P), and oxygen (O). A few crystals displayed trace quantities of iron (Fe). For the WDS analysis, a crystal free of Fe was chosen. A CAMECA SX 100 run at a voltage of 15 kV using a beam current at 10 nA collected the data. Rhodonite was used as the standard for Mn while apatite was the standard used for P. The WDS test was run at four locations on the crystal. After using the PAP corrections determined by Pouchou and Pichoir in 1988 to reduce the data, it was determined that the weight percent of serrabrancaite is roughly 46.85 percent Mn2O3, 42.72 percent P2O5, and 9.80 percent H2O. These results account for 99.37 percent of the total composition of the tested crystal.
A weight loss of 14.6 percent was observed after running a thermal analysis of serrebancaite at 520 °C along with a color change from its original dark greenish-brown to a light yellowish-white. The product of the thermal analysis according to X-ray diffraction (XRD) results is Mn2P2O7. The chemical equation for this reaction is 2 MnPO4·H2O → Mn2P2O7 + 2 H2O + 1/2O2. The empirical formula determined through microprobe and thermal analysis is Mn0.98P1.00O3.98·0.90 H2O, which is very close to the ideal composition of MnPO4·H2O.
Crystal structure.
A synthetic material was used as the standard for the X-ray diffraction (XRD) analysis. The synthetic standard had a unit cell with a= 6.912(1)  Å; b= 7.470(1)  Å; and c= 7.357  Å;, β= 112.29(3)°, and Z= 4. The standard is in the space group C2/c. The XRD analysis was run with a Siemens D5000 diffractometer using CuKɑ radiation and a secondary monochromator to collect the data. The data from the analysis showed that serrabrancaite has unit cell dimensions of a= 6.914(2)  Å; b= 7.468(2)  Å; and c= 7.364(2)  Å;, β= 112.29°, and V= 351.8(1)  Å3. This puts serrabrancaite in the monoclinic crystal system and the point group C2/c after a structural analysis. Serrabrancaite's structure is built of chains of (MnO6)-octahedra that share their corners and run in the [101] direction. These chains are then held together by (PO4)-tetrahedra.
Relation to other minerals.
Serrabrancaite does not have an isostructural relationship to any natural phosphate, arsenate, or vanadate minerals. The diffraction pattern of serrabrancaite does appear to be similar to those of synthetic MnPO4•H2O, synthetic MnAsO4·H2O, and the minerals in the kieserite group (Mg, Mn, Fe, Ni, Cu, Zn)SO4·H2O.
Origin.
The Alto Serra Branca mine is characterized by two phosphates (triphylite and triplite) and their secondary minerals. Triphylite and is secondaries are iron-rich and manganese-poor, while triplite and its secondaries are manganese-rich and iron-poor. Since Serrabrancaite is an alteration of triplite, it is very low in iron. The common associate of serrabrancaite, vernadite, is also very low in iron when found in this pegmatite.


N-Butylsodium

"n"-Butylsodium CH3CH2CH2CH2Na is an organometallic compound with the idealized formula NaC4H9. Like other simple organosodium compounds, it is polymeric and highly basic. In contrast to "n"-butyllithium, "n"-butylsodium is only of specialized academic interest.
Preparation.
"n"-Butylsodium is prepared from "n"-butyllithium and "t"-butoxysodium.
Properties.
In "n"-Butylsodium, the Na-C bond has ionic character, with a negative charge on the end carbon atom. "n"-Butylsodium is insoluble in saturated hydrocarbons. It reacts with unsaturated hydrocarbons. Soluble adducts are produced with Lewis bases such as tetramethylethylenediamine or tetrahydrofuran.
Reactions.
"n"-Butylsodium reacts with alkylbenzene or allylbenzene compounds to give new organosodium compounds. With toluene the major product is benzylsodium.
"n"-Butylsodium reacts with 1-bromonaphthalene to make 1-sodiumnapthalene and 1-bromobutane, but there are few such metathesis reactions.


Khinite

Khinite is a rare tellurate mineral with the formula Pb2+Cu2+3TeO6(OH)2 It crystallizes in the orthorhombic system and has a bottle-green colour. It is often found as dipyramidal, curved or corroded crystals no more than 0.15 mm in size. The tetragonal dimorph of khinite is called parakhinite.
Occurrence and name.
Both khinite and parakhinite were first identified in 1978 in the Old Guard Mine (Royal Guard Mine), Tombstone District, Cochise County, Arizona, US They were named after Ba-Saw Khin, a Burmese-American mineralogist. They are often found together with tenorite, quetzalcoatlite, quartz, gold, dugganite, chrysocolla, chlorargyrite, bromargyrite, xocomecatlite, and tlapallite. Khinite and parakhinite are found in multiple mines across Mexico and the USA.
Parakhinite.
Parakhinite crystallizes in the tetragonal system. Khinite and parakhinite are also called khinite-4O and khinite-3T, respectively. Khinite and parakhinite are identical in colour and many other properties, like reactivity. They do differ in optical properties: Khinite is biaxial (+), while parakhinite is uniaxial (-). They also have different unit cells.


Lithium carbide

Lithium carbide, , often known as dilithium acetylide, is a chemical compound of lithium and carbon, an acetylide. It is an intermediate compound produced during radiocarbon dating procedures. is one of an extensive range of lithium-carbon compounds which include the lithium-rich , , , , , , and the graphite intercalation compounds , , and . is the most thermodynamically-stable lithium-rich compound and the only one that can be obtained directly from the elements. It was first produced by Moissan, in 1896 who reacted coal with lithium carbonate.
The other lithium-rich compounds are produced by reacting lithium vapor with chlorinated hydrocarbons, e.g. CCl4. Lithium carbide is sometimes confused with the drug lithium carbonate, , because of the similarity of its name.
Preparation and chemistry.
In the laboratory samples may be prepared by treating acetylene with a solution of lithium in ammonia, on −40°C, with creation of addition compound of Li2C2 • C2H2 • 2NH3 that decomposes in stream of hydrogen at room temperature giving white powder of Li2C2.
Samples prepared in this manner generally are poorly crystalline. Crystalline samples may be prepared by a reaction between molten lithium and graphite at over 1000 °C. Li2C2 can also be prepared by reacting CO2 with molten lithium.
Other method for production of Li2C2 is heating of metallic lithium in atmosphere of ethylene.
Lithium carbide hydrolyzes readily to form acetylene:
Lithium hydride reacts with graphite at 400°C forming lithium carbide.
Also Li2C2 can be formed when organometallic compound n-Butyllithium reacts with ethyne in THF or Et2O used as a solvent, reaction is rapid and highly exothermic.
Lithium carbide reacts with acetylene in liquid ammonia rapidly to give a clear solution of lithium acetylide.
LiC≡CLi + HC≡CH → 2 LiC≡CH
Preparation of the reagent in this way sometimes improves the yield in an ethynylation over that obtained with reagent prepared from lithium and acetylene.
Structure.
 is a Zintl phase compound and exists as a salt, . Its reactivity, combined with the difficulty in growing suitable single crystals, has made the determination of its crystal structure difficult. It adopts a distorted "anti"-fluorite crystal structure, similar to that of rubidium peroxide () and caesium peroxide (). Each Li atom is surrounded by six carbon atoms from 4 different acetylides, with two acetylides co-ordinating side -on and the other two end-on. The observed C-C distance of 120 pm indicates the presence of a C≡C triple bond.
At high temperatures transforms reversibly to a cubic anti-fluorite structure.
Use in radiocarbon dating.
There are a number of procedures employed, some that burn the sample producing CO2 that is then reacted with lithium, and others where the carbon containing sample is reacted directly with lithium metal. The outcome is the same: Li2C2 is produced, which can then be used to create species easy to use in mass spectroscopy, like acetylene and benzene. Note that lithium nitride may be formed and this produces ammonia when hydrolyzed, which contaminates the acetylene gas.


Hydrostannylation

In chemistry, hydrostannylation is the insertion of unsaturated substrates into an Sn-H bond. The reaction occurs under free-radical conditions, but the stereochemistry and regiochemistry are often complex. The reaction gained synthetic importance with the discovery that palladium complexes catalyze the reaction. The reaction is analogous to hydrosilylation and is a subset of hydroelementation. Hydrostannylation is a versatile route to organotin compounds, many of which are versatile synthetic intermediates, e.g. in Stille coupling.
Substrates and scope.
The typical Pd-based catalyst is tetrakis(triphenylphosphine)palladium(0). The typical unsaturated substrates are alkynes. The typical stannane is tributyltin hydride. The reaction mechanism is assumed to operate via oxidative addition of the stannane to give a stannyl palladium hydride.
Often an excess of the hydride reagent is required since some is consumed through competing dehydrocoupling reactions:


Strontium nitride

Strontium nitride, Sr3N2, is produced by burning strontium metal in air (resulting in a mixture with strontium oxide) or in nitrogen. Like other metal nitrides, it reacts with water to give strontium hydroxide and ammonia:


Tincalconite

Tincalconite is a hydrous sodium borate mineral closely related to borax, and is a secondary mineral that forms as a dehydration product of borax. Its formula is Na2B4O7·5H2O or Na2[B4O5(OH)4]·3H2O.
Tincalconite typically occurs as a fine grained white powder. It crystallizes in the hexagonal crystal system and has been found as primary euhedral di-rhombohedral, pseudo-octahedral crystals. It is also found pseudomorphically replacing borax crystals. It has a specific gravity of 1.88 and a Mohs hardness of 2. Refractive index values are nω=1.460 and nε=1.470.
While most tincalconite is created by man through exposing borax to dry air, there are natural occurrences of tincalconite, as in Searles Lake, California where it was first described in 1878. In addition to several California and Nevada locations it is reported from Argentina, Italy, Turkey and Ukraine.
The name comes from "tincal", Sanskrit for borax, and Greek, "konis", meaning powder, for its composition and typical powdery nature.
References.
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Magnesium silicide

Magnesium silicide, Mg2Si, is an inorganic compound consisting of magnesium and silicon. As-grown Mg2Si usually forms black crystals; they are semiconductors with n-type conductivity and have potential applications in thermoelectric generators.
Crystal structure.
Mg2Si crystallizes in the antifluorite structure. In the face-centered cubic lattice Si centers occupy the corners and face-centered positions of the unit cell and Mg centers occupy eight tetrahedral sites in the interior of the unit cell. The coordination numbers of Si and Mg are eight and four, respectively.
Synthesis.
It can be produced by heating silicon dioxide, SiO2, found in sand, with excess magnesium. The process first forms silicon metal and magnesium oxide, and, if an excess of SiO2 is used, then elemental silicon is formed:
If an excess of Mg is present, Mg2Si is formed from the reaction of the remaining magnesium with the silicon:
These reactions proceed exothermically, even explosively.
Reactions.
Magnesium silicide can be viewed as consisting of Si4− ions. As such it is reactive toward acids. Thus, when magnesium silicide is treated with hydrochloric acid, silane (SiH4) and magnesium chloride are produced:
Sulfuric acid can be used as well. These protonolysis reactions are typical of a Group 2 alkaline earth metal and Group 1 alkali metal silicides. The early development of silicon hydrides relied on this reaction.
Uses.
Magnesium silicide is used to create aluminium alloys of the 6000 series, containing up to approximately 1.5% Mg2Si. An alloy of this group can be age-hardened to form Guinier-Preston zones and a very fine precipitate, both resulting in increased strength of the alloy.
Magnesium silicide is a narrow-gap semiconductor. Its as-grown crystal exhibit n-type conductivity, but it can changed to p-type by doping with Ag, Ga, Sn and possibly Li (at high doping level). The major potential electronic application of Mg2Si is in thermoelectric generators.


Graphene quantum dot

Graphene quantum dots (GQDs) are graphene nanoparticles with a size less than 100 nm. Due to their exceptional properties such as low toxicity, stable photoluminescence, chemical stability and pronounced quantum confinement effect, GQDs are considered as a novel material for biological, opto-electronics, energy and environmental applications.
Properties.
Graphene quantum dots (GQDs) consist of one or a few layers of graphene and are smaller than 100 nm in size. They are chemically and physically stable, have a large surface to mass ratio and can be dispersed in water easily due to functional groups at the edges. The fluorescence emission of GQDs can extend across a broad spectral range, including the UV, visible, and IR. The origin of GQD fluorescence emission is a subject of debate, as it has been related to quantum confinement effects, defect states and functional groups that might depend on the pH, when GQDs are dispersed in water. Their electronic structure depends sensitively on the crystallographic orientation of their edges, for example zigzag-edge GQDs with 7-8 nm diameter show a metallic behavior. In general, their energy gap decreases, when the number of graphene layers or the number of carbon atoms per graphene layer is increased.
Health and Safety.
The toxicity of graphene-family nanoparticles is a matter of ongoing research. The toxicity (both in vivo and cytotoxicity) of GQDs are related to a variety of factors including particle size, methods of synthesis, chemical doping and so on. Many authors claim, that GQDs are biocompatible and cause only low toxicity as they are just composed of organic materials, which should lead to an advantage over semiconductor quantum dots. Several in vitro studies, based on cell cultures, show only marginal effects of GQDs on the viability of human cells. An in-depth look at the gene expression changes caused by GQDs with a size of 3 nm revealed that only one, namely the selenoprotein W, 1 out of 20 800 gene expressions was affected significantly in primary human hematopoietic stem cells. On the contrary, other in vitro studies observe a distinct decrease of cell viability and the induction of autophagy after exposure of the cells to GQDs and one in vivo study in zebrafish larvae observed the alteration of 2116 gene expressions. These inconsistent findings may be attributed to the diversity of the used GQDs, as the related toxicity is dependent on particle size, surface functional groups, oxygen content, surface charges and impurities. Currently, the literature is insufficient to draw conclusions about the potential hazards of GQDs.
Preparation.
Presently, a range of techniques have been developed to prepare GQDs. These methods are normally classified into two groups top down and bottom up. Top down approaches applied different techniques to cut bulk graphitic materials into GQDs including graphite, graphene, carbon nanotubes, coal, carbon black and carbon fibres. These techniques mainly include electron beam lithography, chemical synthesis, electrochemical preparation, graphene oxide (GO) reduction, C60 catalytic transformation, the microwave assisted hydrothermal method (MAH), the Soft-Template method, the hydrothermal method, and the ultrasonic exfoliation method. Top down methods usually need intense purification as strong mixed acids are used in these methods. On the other hand, bottom up methods assemble GQDs from small organic molecules such as citric acid and glucose. These GQDs have better biocompatibility.
Application.
Graphene quantum dots are studied as an advanced multifunctional material due to their unique optical, electronic, spin, and photoelectric properties induced by the quantum confinement effect and edge effect. They have possible applications in bioimaging, cancer therapeutics, temperature sensing, drug delivery, LEDs lighter converters, photodetectors, OPV solar cells, and photoluminescent material, biosensors fabrication.


Bis(cyclooctadiene)nickel(0)

Bis(cyclooctadiene)nickel(0) is the organonickel compound with the formula Ni(C8H12)2, also written Ni(cod)2. It is a diamagnetic coordination complex featuring tetrahedral nickel(0) bound to the alkene groups in two 1,5-cyclooctadiene ligands. This highly air-sensitive yellow solid is a common source of Ni(0) in chemical synthesis.
Preparation and properties.
The complex is prepared by reduction of anhydrous nickel(II) acetylacetonate in the presence of the diolefin:
Ni(cod)2 is moderately soluble in several organic solvents. One or both 1,5-cyclooctadiene ligands are readily displaced by phosphines, phosphites, bipyridine, and isocyanides. If exposed to air, the solid oxidizes to nickel(II) oxide. As a result, this compound is generally handled in a glovebox.


Nickel oxyacid salts

The Nickel oxyacid salts are a class of chemical compounds of nickel with an oxyacid. The compounds include a number of minerals and industrially important nickel compounds.
Nickel(II) sulfate can crystallise with six water molecules yielding Retgersite or with seven making Morenosite which is isomorphic to Epsom salts. These contain the hexaquanickel(II) ion.
There is also an anhydrous form, a dihydrate and a tetrahydrate, the last two crystallised from sulfuric acid. The hexahydrate has two forms, a blue tetragonal form, and a green monoclinic form, with a transition temperature around 53 °C. The heptahydrate crystallises from water below 31.5 above this blue hexhydrate forms, and above 53.3 the green form. Heating nickel sulfate dehydrates it, and then 700° it loses sulfur trioxide, sulfur dioxide and oxygen.
Nickel sulfite can be formed by bubbling sulfur dioxide through nickel carbonate suspended in water. A solution is formed that slowly loses sulfur dioxide, and which crystallises nickel sulfite hexahydrate. Crystals are frequently in the shape of stars, caused by the two opposite triangular enantiomorphs growing base to base. nickel sulfite hexahydrate is highly piezoelectric. Optically it is uniaxial negative with refractive indexes ω=1.552 ε=1.509. When heated it dehydrates and then ends up making nickel oxide and nickel sulfate.
Nickel thiosulfate NiS2O3 has the same structure as the magnesium salt. It has alternating layers of octahedral shaped nickel2+ hexahydrate, and tetrahedral shaped S2O32− perpendicular to the β direction. When heated to 90 °C it decomposes to form NiS. NiS2O3 can be made from BaS2O3 and NiSO4. Nickel sulfamate can be used for nickel or mixed nickel-tungsten plating. It can be formed by the action of sulfamic acid on nickel carbonate.
Nickel selenite NiSeO3 has many different hydrates, anhydrous NiSeO3, NiSeO3, NiSeO3 (which is also a mineral called ahlfeldite), and NiSeO3.
Nickel nitrate commonly crystallises with six water molecules, but can also be anhydrous, or with two, four or nine waters.
triphenylphosphine oxide nickel nitrate [(C6H6)3PO]2Ni(NO3)2 is non ionic, with nitrato as a ligand. It can be made from nickel perchlorate. It is yellow and melts at 266 °C.
Nickel carbonate NiCO3, hellyerite, crystallising with six water molecules, precipitates when an alkali bicarbonate is added to a Ni aqueous solution. Basic nickel carbonate, zaratite, with the formula Ni4CO3(OH)6(H2O)4, is produced when alkali carbonates are added to a nickel solution. Nickel phosphate, Ni3(PO4)2 is also insoluble. A number of other phosphates have been made, including nanoporous substances resembling zeolites named with "Versailles Santa Barbara" or VSB. The nanoporous nickel phosphates can accommodate sufficiently small molecules and selectively catalyze reactions on them. A nickel arsenate, Ni3(AsO4)2·8H2O occurs as the mineral annabergite.
Nickel perchlorate, Ni(ClO4)2, nickel chlorate, Ni(ClO3)2, nickel chromate (NiCrO4), nickel chromite (NiCr2O4), nickel(II) titanate, nickel bromate Ni(BrO3)2, nickel iodate (Ni(IO3)2), nickel stannate (NiSnO3) are some other oxy-salts.
The uranates include NiU2O6, NiUO4 α and β forms (orthorhombic a=6.415 Å; b=6.435 Å; c=6.835 Å), and NiU3O10.


Lithium iron phosphate

Lithium iron phosphate (LFP) is an inorganic compound with the formula . It is a gray, red-grey, brown or black solid that is insoluble in water. The material has attracted attention as a component of lithium iron phosphate batteries, a type of Li-ion battery. This battery chemistry is targeted for use in power tools, electric vehicles, and solar energy installations. It is also used in OLPC XO education laptops.
Most lithium batteries (Li-ion) used in 3C (computer, communication, consumer electronics) products use cathodes made of other lithium compounds, such as lithium cobalt oxide (), lithium manganese oxide (), and lithium nickel oxide (). The anodes are generally made of graphite.
Lithium iron phosphate exists naturally in the form of the mineral triphylite, but this material has insufficient purity for use in batteries.
With general chemical formula of , compounds in the family adopt the olivine structure. M includes not only Fe but also Co, Mn and Ti. As the first commercial was C/, the whole group of is informally called “lithium iron phosphate” or “”. However, more than one olivine-type phase may be used as a battery's cathode material. Olivine compounds such as , , and have the same crystal structures as and may replace in a cathode. All may be referred to as “LFP”.
Manganese, phosphate, iron, and lithium also form an olivine structure. This structure is a useful contributor to the cathode of lithium rechargeable batteries. This is due to the olivine structure created when lithium is combined with manganese, iron, and phosphate (as described above). The olivine structures of lithium rechargeable batteries are significant, for they are affordable, stable, and can be safely stored as energy.
History and production.
Arumugam Manthiram and John B. Goodenough first identified the polyanion class of cathode materials for lithium ion batteries. was then identified as a cathode material belonging to the polyanion class for use in batteries in 1996 by Padhi et al. Reversible extraction of lithium from and insertion of lithium into was demonstrated. Neutron diffraction confirmed that LFP was able to ensure the security of large input/output current of lithium batteries.
The material can be produced by heating a variety of iron and lithium salts with phosphates or phosphoric acid. Many related routes have been described including those that use hydrothermal synthesis.
Physical and chemical properties.
In , lithium has a +1 charge, iron +2 charge balancing the -3 charge for phosphate. Upon removal of Li, the material converts to the ferric form FePO4.
The iron atom and 6 oxygen atoms form an octahedral coordination sphere, described as , with the Fe ion at the center. The phosphate groups, , are tetrahedral. The three-dimensional framework is formed by the octahedra sharing O corners. Lithium ions reside within the octahedral channels in a zigzag manner. In crystallography, this structure is thought to belong to the Pmnb space group of the orthorhombic crystal system. The lattice constants are: a = 6.008 Å, b = 10.334 Å, and c = 4.693 Å. The volume of the unit cell is 291.4 Å3.
In contrast to two traditional cathode materials - and , lithium ions of migrate in the lattice's one-dimensional free volume. During charge/discharge, the lithium ions are extracted concomitant with oxidation of Fe:
Extraction of lithium from produces with a similar structure. adopts a Pmnb space group with a unit cell volume of 272.4 3, only slightly smaller than that of its lithiated precursor. Extraction of lithium ions reduces the lattice volume, as is the case with lithium oxides. 's corner-shared octahedra are separated by the oxygen atoms of the 3- tetrahedra and cannot form a continuous network, reducing conductivity.
A nearly close-packed hexagonal array of oxides centers provides relatively little free volume for Li+ ions to migrate within. For this reason, the ionic conductivity of Li+ is relatively low at ambient temperate. The details of the lithiation of FePO4 and the delithiation of LiFePO4 have been examined. Two phases of the lithiated material are implicated.
Applications.
LFP batteries have an operating voltage of 3.3 V, charge density of 170mAh/g, high power density, long cycle life and stability at high temperatures.
LFP's major commercial advantages are that it poses few safety concerns such as overheating and explosion, as well as long cycle lifetimes, high power density and has a wider operating temperature range. Power plants and automobiles use LFP.
BAE has announced that their HybriDrive Orion 7 hybrid bus uses about 180 kW LFP battery cells. AES has developed multi-trillion watt battery systems that are capable of subsidiary services of the power network, including spare capacity and frequency adjustment. In China, BAK and Tianjin Lishen are active in the area.
Comparison.
Although LFP has 25% less capacity than other lithium batteries due to its operational voltage (3.2 volts vs 3.7 for Co based cathode chemistries), it has 70% more than nickel-hydrogen batteries.
The major differences between LFP batteries and other lithium ion battery types is that LFP batteries contain no cobalt (removing ethical questions about material sourcing) and have a flat discharge curve.
LFP batteries have drawbacks, including higher costs due to lower energy density. The energy density is significantly lower than (although higher than the nickel-metal hydride battery).
Lithium cobalt oxide based battery chemistries are more prone to thermal runaway if overcharged and Cobalt is both expensive and not widely available. Other chemistries such as NMC Ni Mn Co have supplanted LiCo chemistry cells in most applications. The original ratio of Ni to Mn to Co was 3:3:3 today, cells are being made with ratios of 8:1:1 whereby the Co content has been drastically reduced.
Intellectual property.
The root patents of LFP compounds are held by four organizations. University of Texas-Austin for the discovery of the material. Hydro-Québec, Université de Montréal and the French National Center for Scientific Research (CNRS) for the carbon coating that enhance its conductivity and actually makes LFP suitable for industrial developments. These patents underlie mature mass production technologies. The largest production capacity is up to 250 tons per month. The key feature of from A123 is the nano-LFP, which modifies its physical properties and adds noble metals in the anode, as well as the use of special graphite as the cathode.
The main feature of from Phostech is increased capacitance and conductivity by an appropriate carbon coating. The special feature of • zM from Aleees a high capacitance and low impedance obtained by the stable control of the ferrites and crystal growth. This improved control is realized by applying strong mechanical stirring forces to the precursors in high oversaturation states, which induces crystallization of the metal oxides and LFP.
In patent lawsuits in the US in 2005 and 2006, the University of Texas at Austin|University of Texas-Austin and Hydro-Québec claimed that as the cathode infringed their patents, and . The patent claims involved a unique crystal structure and a chemical formula of the battery cathode material.
On April 7, 2006, A123 filed an action seeking a declaration of non-infringement and invalidity UT's patents. A123 separately filed two ex parte Reexamination Proceedings before the United States Patent and Trademark Office (USPTO), in which they sought to invalidate the patents based upon prior art.
In a parallel court proceeding, UT sued Valence Technology, Inc. ("Valence") - a company that commercializes LFP products that alleged infringement.
The USPTO issued a Reexamination Certificate for the '382 patent on April 15, 2008 and for the '640 patent on May 12, 2009, by which the claims of these patents were amended. This allowed the current patent infringement suits filed by Hydro-Quebec against Valence and A123 to proceed. After a markman hearing, on April 27, 2011 the Western District Court of Texas held that the claims of the reexamined patents had a narrower scope than as originally granted.
On Dec 9th, 2008, the European Patent Office revoked Dr. Goodenough’s patent numbered 0904607. This decision basically reduced the patent risk of using LFP in European automobile applications. The decision is believed to be based on the lack of novelty.
The first major large settlement was the lawsuit between NTT and the University of Texas-Austin (UT). In October 2008, NTT announced that they would settle the case in the Japan Supreme Civil Court for $30 million. As part of the agreement UT agreed that NTT did not steal the information and that NTT would share its LFP patents with UT. NTT’s patent is also for an olivine LFP, with the general chemical formula of (A is for alkali metal and M for the combination of Co and Fe), now used by BYD Company. Although chemically the materials are nearly the same, from the viewpoint of patents, of NTT is different from the materials covered by UT. has higher capacity than . At the heart of the case was that NTT engineer Okada Shigeto, who had worked in the UT labs developing the material, was accused of stealing UT’s intellectual property.
As of 2020, an organisation named LifePO+C claims to own the key IP and offers licenses. It is a consortium between Johnson Matthey, the CNRS, University of Montreal, and Hydro Quebec.
Research.
Power density.
LFP has two shortcomings: low conductivity and low lithium diffusion constant, both of which limit the charge/discharge rate. Adding conducting particles in delithiated raises its electron conductivity. For example, adding conducting particles with good diffusion capability like graphite and carbon to powders significantly improves conductivity between particles, increases the efficiency of and raises its reversible capacity up to 95% of the theoretical values. shows good cycling performance even under charge/discharge current as large as 5C.
Stability.
Coating LFP with inorganic oxides can make LFP’s structure more stable and increase conductivity. Traditional with oxide coating shows improved cycling performance. This coating also inhibits dissolution of Co and slows the decay of capacity. Similarly, with an inorganic coating such as ZnO and , has a better cycling lifetime, larger capacity and better characteristics under rapid discharge. The addition of a conductive carbon increases efficiency. Mitsui Zosen and Aleees reported that addition of conducting metal particles such as copper and silver increased efficiency. with 1 wt% of metal additives has a reversible capacity up to 140mAh/g and better efficiency under high discharge current.
Metal substitution.
Substituting other metals for the iron or lithium in can also raise efficiency. Substituting zinc for iron increases crystallinity of because zinc and iron have similar ion radii. Cyclic voltammetry confirms that , after metal substitution, has higher reversibility of lithium ion insertion and extraction. During lithium extraction, Fe (II) is oxidized to Fe (III) and the lattice volume shrinks. The shrinking volume changes lithium’s returning paths.
Synthesis processes.
Mass production with stable and high quality still faces many challenges.
Similar to lithium oxides, may be synthesized by a variety of methods, including: solid-phase synthesis, emulsion drying, sol-gel process, solution coprecipitation, vapor-phase deposition, electrochemical synthesis, electron beam irradiation, microwave process, hydrothermal synthesis, ultrasonic pyrolysis and spray pyrolysis.
In the emulsion drying process, the emulsifier is first mixed with kerosene. Next, the solutions of lithium salts and iron salts are added to this mixture. This process produces nanocarbon particles. Hydrothermal synthesis produces with good crystallinity. Conductive carbon is obtained by adding polyethylene glycol to the solution followed by thermal processing. Vapor phase deposition produces a thin film . In flame spray pyrolysis FePO4 is mixed with Lithium carbonate and glucose and charged with electrolytes. The mixture is then injected inside a flame and filtered to collect the synthesized .


Peroxynitrous acid

Peroxynitrous acid (HNO3) is a reactive nitrogen species (RNS). It is the conjugate acid of peroxynitrite (ONOO−). It has a p"K"a of approximately 6.8. It is formed "in vivo" from the diffusion-controlled reaction of nitrogen monoxide (ON•) and superoxide (). It is an isomer of nitric acid and isomerises with a rate constant of "k" = 1.2 s−1, a process whereby up to 5% of hydroxyl and nitrogen dioxide radicals may be formed. It oxidises and nitrates aromatic compounds in low yield. The mechanism may involve a complex between the aromatic compound and ONOOH, and a transition from the "cis"- to the "trans"-configuration of ONOOH. Peroxynitrous acid is also important in atmospheric chemistry.


LR-87

The LR87 was an American liquid-propellant rocket engine, which was used on the first stages of Titan intercontinental ballistic missiles and launch vehicles. Composed of twin motors with separate combustion chambers and turbopump machinery, it is considered a single unit. The LR87 first flew in 1959.
The LR87 was developed in the late 1950s by Aerojet. It was the first production rocket engine capable (in its various models) of burning the three most common liquid rocket propellant combinations: liquid oxygen/RP-1, nitrogen tetroxide/Aerozine 50 (a 50:50 mixture by mass of hydrazine and UDMH), and liquid oxygen/liquid hydrogen. The engine operated on an open gas-generator cycle and utilized a regenerative cooled nozzle and combustion chamber. Later versions had additional ablative-cooled flanges. The LR87 served as a template for the LR-91, which was used in the second stage of the Titan missile.
It was a fixed-thrust engine, which could not be throttled or restarted in flight. The LR87 delivered approximately 1,900 kilonewtons (430,000 pounds) of thrust. Early LR87 engines used on the Titan I burned RP-1 and liquid oxygen. Because liquid oxygen is cryogenic, it could not be stored in the missile for long periods of time, and had to be loaded before the missile could be launched. For the Titan II, the engine was converted to use Aerozine 50 and nitrogen tetroxide, which are hypergolic and storable at room temperature. This allowed Titan II missiles to be kept fully fueled and ready to launch on short notice.
Variants.
LR87-3.
Used on the Titan I, the LR87-3 burned liquid oxygen and RP-1. Following the retirement of the Titan missile program, these engines saw no further use. The LR87-3 was also tested with LOX/H2 and NTO/Aerozine 50 making it one of very few engines to have been run on three different propellant combinations.
LR87-5.
Modified to burn nitrogen tetroxide and Aerozine 50 for the Titan II. The engine was generally lighter and simpler than its predecessor, partly due to the use of hypergolic propellants, which did not need an independent ignition system.
LR87-7.
Modified versions of LR87-5 adapted to the needs of the Gemini program. The performance was similar to the previous version, only reducing the chamber pressure and nozzle thrust to meet human-rating requirements. This version was only used on the Titan II GLV.
LR87-9.
Used on the Titan IIIA, IIIB, and IIIC
LR87-11/LR-87-11A.
Used on Titan 24B, 34B, IIIBS, IIID, 34D, 34D7, IIIE. The LR-87-11A was used on the Titan IV A/B.
LR87 LH2.
Modified to burn liquid oxygen and liquid hydrogen. The development coincided with other variants of the late 1950s. Compared to the -3, it had a number of changes associated with the use of lighter and colder liquid hydrogen. The fuel injector turbo was changed along with the fuel turbopump among other things. A total of 52 static tests were performed without serious issue. Aerojet took part in the selection process for a new engine for the second stage of the Saturn IB and Saturn V. Though LR87 LH2 was the best in 10 out of 11 criteria, NASA selected Rocketdyne's J-2. Lessons learned were used during development of the Aerojet M-1


Niobium chloride

Niobium chloride may refer to:


Yttrium(III) nitrate

Yttrium(III) nitrate is an inorganic compound with the formula Y(NO3)3. The hexahydrate is the most common form commercially available.
Preparation.
Yttrium(III) nitrate can be prepared by dissolving corresponding metal oxide in 6 mol/L nitric acid:
Properties.
Yttrium(III) nitrate hexahydrate loses crystallized water at relatively low temperature. Upon further heating, basic salt YONO3 is formed. At 600 C, the thermal decomposition is complete. Y2O3 is the final product.
Y(NO3)3·3TBP is formed when tributyl phophate is used as the extracting solvent.
Uses.
Yttrium(III) nitrate mainly uses as a source of Y3+ cation. It is a precursor of some yttrium-containing materials, such as Y4Al2O9, YBa2Cu3O6.5+x and yttrium-based metal-organic frameworks.
 It can also be used as a catalyst in organic synthesis.


Vanadinite

Vanadinite is a mineral belonging to the apatite group of phosphates, with the chemical formula Pb5(VO4)3Cl. It is one of the main industrial ores of the metal vanadium and a minor source of lead. A dense, brittle mineral, it is usually found in the form of red hexagonal crystals. It is an uncommon mineral, formed by the oxidation of lead ore deposits such as galena. First discovered in 1801 in Mexico, vanadinite deposits have since been unearthed in South America, Europe, Africa, and North America.
Origins.
Vanadinite is an uncommon mineral, only occurring as the result of chemical alterations to a pre-existing material. It is therefore known as a secondary mineral. It is found in arid climates and forms by oxidation of primary lead minerals. Vanadinite is especially found in association with the lead sulfide, galena. Other associated minerals include wulfenite, limonite, and barite.
It was originally discovered in Mexico by the Spanish mineralogist Andrés Manuel del Río in 1801. He called the mineral "brown lead" and asserted that it contained a new element, which he first named pancromium and later, erythronium. However, he was later led to believe that this was not a new element but merely an impure form of chromium. In 1830, Nils Gabriel Sefström discovered a new element, which he named vanadium. It was later revealed that this was identical to the metal discovered earlier by Andrés Manuel del Río. Del Río's "brown lead" was also rediscovered, in 1838 in Zimapan, Hidalgo, Mexico, and was named vanadinite because of its high vanadium content. Other names that have since been given to vanadinite are johnstonite and lead vanadate.
Occurrence.
Vanadinite occurs as a secondary mineral in the oxidized zone of lead-bearing deposits, the vanadium is leached from wall-rock silicates. Associated minerals include mimetite, pyromorphite, descloizite, mottramite, wulfenite, cerussite, anglesite, calcite, barite, and various iron oxide minerals.
Deposits of vanadinite are found worldwide including Austria, Spain, Scotland, the Ural Mountains, South Africa, Namibia, Morocco, Argentina, Mexico, and 4 states of the United States: Arizona, Colorado, New Mexico, and South Dakota.
Vanadinite deposits are found in over 400 mines across the world. Notable vanadinite mines include those at Mibladen and Touisset in Morocco; Tsumeb, Namibia; Cordoba, Argentina; and Sierra County, New Mexico, and Gila County, Arizona, in the United States.
Structure.
Vanadinite is a lead chlorovanadate with the chemical formula Pb5(VO4)3Cl. It is composed (by weight) of 73.15% lead, 10.79% vanadium, 13.56% oxygen, and 2.50% chlorine. Each structural unit of vanadinite contains a chlorine ion surrounded by six divalent lead ions at the corners of a regular octahedron, with one of the lead ions provided by an adjoining vanadinite molecule. The distance between each lead and chlorine ion is 317 picometres. The shortest distance between each lead ion is 4.48 Å. The octahedron shares two of its opposite faces with that of neighbouring vanadinite units, forming a continuous chain of octahedrons. Each vanadium atom is surrounded by four oxygen atoms at the corners of an irregular tetrahedron. The distance between each oxygen and vanadium atom is either 1.72 or 1.76 Å. Three oxygen tetrahedrons adjoin each of the lead octahedrons along the chain.
Crystals of vanadinite conform to a hexagonal system of symmetry. This internal structure is often reflected in the hexagonal external shape of the crystals. The crystals are usually in the form of short hexagonal prisms, but can also be found as hexagonal pyramids, rounded masses or crusts. A unit cell of vanadinite, the smallest divisible unit that possesses the same symmetry and properties, is in the form of a hexagonal prism. The unit cell of vanadinite is composed of two of its molecules and has the dimensions "a" = 10.331 Å and "c" = 7.343 Å), where "a" is the length of each side of the hexagon and "c" is the height of the prism. The volume of each unit cell of vanadinite, given by the formula V = "a"2"c" sin(60°), is 678.72 Å3.
Characteristics.
Vanadinite is in the apatite group of phosphates, and forms a chemical series with the minerals pyromorphite (Pb5(PO4)3Cl) and mimetite (Pb5(AsO4)3Cl), with both of which it may form solid solutions. Whereas most chemical series involve the substitution of metallic ions, this series substitutes its anion groups; phosphate (PO4), arsenate (AsO4) and vanadate (VO4). Common impurities of vanadinite include phosphorus, arsenic and calcium, where these may act as an isomorphic substitute for vanadium. Vanadinite when containing a high amount of the arsenic impurity is known as endlichite.
Vanadinite is usually bright-red or orange-red in colour, although sometimes brown, red-brown, grey, yellow, or colourless. Its distinctive colour makes it popular among mineral collectors. Its streak can be either pale yellow or brownish-yellow. Vanadinite may be transparent, translucent or opaque, and its lustre can range from resinous to adamantine. Vanadinite is anisotropic, meaning that some of its properties differ when measured along different axes. When measured perpendicular and parallel to its axis of anisotropy, its refractive indices are 2.350 and 2.416 respectively. This gives it a birefringence of 0.066.
Vanadinite is very brittle, producing small, conchoidal fragments when fractured. Its hardness is 3–4 on the Mohs scale, about the same as a copper coin. Vanadinite is particularly heavy for a translucent mineral. It has a molar mass of 1416.27 g/mole and its specific gravity can range between 6.6 and 7.2 because of impurities.
Uses.
Along with carnotite and roscoelite, vanadinite is one of the main industrial ores of the element vanadium, which can be extracted by roasting and smelting. Vanadinite is also occasionally used as a source of lead. A common process for extracting the vanadium begins with the heating of vanadinite with salt (NaCl) or sodium carbonate (Na2CO3) at about 850 °C to produce sodium vanadate (NaVO3). This is dissolved in water and then treated with ammonium chloride to give an orange-coloured precipitate of ammonium metavanadate. This is then melted to form a crude form of vanadium pentoxide (V2O5). Reduction of vanadium pentoxide with calcium gives pure vanadium.


Lithium stearate

Lithium stearate is a chemical compound with the formula LiO2C(CH2)16CH3. It is formally classified as a soap (a salt of a fatty acid). Lithium stearate is a white soft solid, prepared by the reaction of lithium hydroxide and stearic acid.
Lithium stearate and lithium 12-hydroxystearate are lithium soaps, and are components of lithium grease.


Cobalt(II) oxide

Cobalt(II) oxide is an inorganic compound that has been described as an olive-green or gray solid. It is used extensively in the ceramics industry as an additive to create blue colored glazes and enamels as well as in the chemical industry for producing cobalt(II) salts. A related material is cobalt(II,III) oxide, a black solid with the formula Co3O4.
Structure and properties.
CoO crystals adopt the periclase (rock salt) structure with a lattice constant of 4.2615 Å.
It is antiferromagnetic below 16 °C.
Preparation.
Cobalt(II) oxide is prepared by oxidation of cobalt powder with air or by thermal decomposition of cobalt(II) nitrate or the carbonate.
Cobalt(II,III) oxide decomposes to cobalt(II) oxide at 950 °C:
It may also be prepared by precipitating the hydroxide, followed by thermal dehydration:
Reactions.
As can be expected, cobalt(II) oxide reacts with mineral acids to form the corresponding cobalt salts:
Applications.
Cobalt(II) oxide has for centuries used as a coloring agent on kiln fired pottery. The additive provides a deep shade of blue named cobalt blue. The band gap (CoO) is around 2.4 eV.
It also is used in cobalt blue glass.


Heazlewoodite

Heazlewoodite, Ni3S2, is a rare sulfur-poor nickel sulfide mineral found in serpentinitized dunite. It occurs as disseminations and masses of opaque, metallic light bronze to brassy yellow grains which crystallize in the trigonal crystal system. It has a hardness of 4, a specific gravity of 5.82. Heazlewoodite was first described in 1896 from Heazlewood, Tasmania, Australia.
Paragenesis.
Heazlewoodite is formed within terrestrial rocks by metamorphism of peridotite and dunite via a process of nucleation. Heazlewoodite is the least sulfur saturated of nickel sulfide minerals and is only formed via orphic exsolution of sulfur from the lattice of metamorphic olivine.
Heazlewoodite is thought to form from sulfur and nickel which exist in pristine olivine in trace amounts, and which are driven out of the olivine during metamorphic processes. Magmatic olivine generally has up to ~4000 ppm Ni and up to 2500 ppm S within the crystal lattice, as contaminants and substituting for other transition metals with similar ionic radii (Fe2+ and Mg2+).
During metamorphism, sulfur and nickel within the olivine lattice are reconstituted into metamorphic sulfide minerals, chiefly millerite, during serpentinization and talc carbonate alteration.
When metamorphic olivine is produced, the propensity for this mineral to resorb sulfur, and for the sulfur to be removed via the concomitant loss of volatiles from the serpentinite, tends to lower sulfur fugacity.
In this environment, nickel sulfide mineralogy converts to the lowest-sulfur state available, which is heazlewoodite.
Occurrence.
Heazlewoodite is known from few ultramafic intrusions within terrestrial rocks. The Honeymoon Well ultramafic intrusive, Western Australia is known to contain heazlewoodite-millerite sulfide assemblages within serpentinized olivine adcumulate dunite, formed from the metamorphic process.
The mineral is also reported, again in association with millerite, from the ultramafic rocks of New Caledonia.
This mineral has been found in meteorites including irons and CV carbonaceous chondrites.


Tetravinyltin

Tetravinyltin (also known as "'tetravinylstannane") is an organotin compound with a chemical formula of C8H12Sn.
Uses.
Upon heating, a mixture of tetravinyltin and tin tetrachloride undergo disproportionation to form divinyltin dichloride, vinyltin trichloride, and trivinyltin chloride in high yields. A study about this can be found in the Journal of American Chemical Society.
Tetravinyltin cannot be used for therapeutic or diagnostic purposes and must only be used for research. It can also be used for thin film deposition.
Hazards.
According to the European Chemicals Agency, tetravinyltin is flammable in liquid and gas form. It is also toxic when in contact with skin, inhaled, and swallowed. Therefore, personal protective equipment must be used in handling and proper caution applied during use.


Benzyl fluoride

Benzyl fluoride is an organic compound consisting of a benzene ring substituted with a fluoromethyl group.


Disorders of calcium metabolism

Disorders of calcium metabolism occur when the body has too little or too much calcium. The serum level of calcium is closely regulated within a fairly limited range in the human body. In a healthy physiology, extracellular calcium levels are maintained within a tight range through the actions of parathyroid hormone, vitamin D and the calcium sensing receptor. Disorders in calcium metabolism can lead to hypocalcemia, decreased plasma levels of calcium or hypercalcemia, elevated plasma calcium levels.
Hypocalcemia.
Hypocalcemia is common and can occur unnoticed with no symptoms or, in severe cases, can have dramatic symptoms and be life-threatening. Hypocalcemia can be parathyroid related or vitamin D related. Parathyroid related hypocalcemia includes post-surgical hypoparathyroidism, inherited hypoparathyroidism, pseudohypoparathyroidism, and pseudo-pseudohypoparathyroidism. Post-surgical hypoparathyroidism is the most common form, and can be temporary (due to suppression of tissue after removal of a malfunctioning gland) or permanent, if all parathyroid tissue has been removed. Inherited hypoparathyroidism is rare and is due to a mutation in the calcium sensing receptor. Pseudohypoparathyroidism is maternally inherited and is categorized by hypocalcemia and hyperphosphatemia. Finally, pseudo-pseudohypoparathyroidism is paternally inherited. Patients display normal parathyroid hormone action in the kidney, but exhibit altered parathyroid hormone action in the bone.
Vitamin D related hypocalcemia may be associated with a lack of vitamin D in the diet, a lack of sufficient UV exposure, or disturbances in renal function. Low vitamin D in the body can lead to a lack of calcium absorption and secondary hyperparathyroidism (hypocalcemia and raised parathyroid hormone). Symptoms of hypocalcemia include numbness in fingers and toes, muscle cramps, irritability, impaired mental capacity and muscle twitching.
Hypercalcemia.
Hypercalcemia is suspected to occur in approximately 1 in 500 adults in the general adult population. Like hypocalcemia, hypercalcemia can be non-severe and present with no symptoms, or it may be severe, with life-threatening symptoms. Hypercalcemia is most commonly caused by hyperparathyroidism and by malignancy, and less commonly by vitamin D intoxication, familial hypocalciuric hypercalcemia and by sarcoidosis. Hyperparathyroidism occurs most commonly in postmenopausal women. Hyperparathyroidism can be caused by a tumor, or adenoma, in the parathyroid gland or by increased levels of parathyroid hormone due to hypocalcemia. Approximately 10% of cancer sufferers experience hypercalcemia due to malignancy. Hypercalcemia occurs most commonly in breast cancer, lymphoma, prostate cancer, thyroid cancer, lung cancer, myeloma, and colon cancer. It may be caused by secretion of parathyroid hormone-related peptide by the tumor (which has the same action as parathyroid hormone), or may be a result of direct invasion of the bone, causing calcium release.
Symptoms of hypercalcemia include anorexia, nausea, vomiting, constipation, abdominal pain, lethargy, depression, confusion, polyuria, polydipsia and generalized aches and pains.
Plasma Calcium.
The amount of biologically active calcium varies with the level of serum albumin, a protein to which calcium is bound, and therefore levels of "ionized calcium" are better measures than a "total calcium"; however, one can correct a "total calcium" if the albumin level is known.


Osmocene

Osmocene is an organoosmium compound found as a white solid. It is a metallocene with the formula Os(C5H5)2.
Synthesis.
Osmocene is commercially available. It may be prepared by the reaction of osmium tetroxide with hydrobromic acid followed by zinc and cyclopentadiene.
It was first synthesized by Ernst Otto Fischer and Heinrich Grumbert via the reaction of osmium(IV) chloride with excess sodium cyclopentadienide in dimethoxyethane, where osmium(II) chloride is presumed to be an intermediate formed "in situ". Alternatively, cyclopentadienyl magnesium bromide could be reacted with osmium(IV) chloride, though this has worse yields.
Properties.
Osmocene is a white solid. The molecular structure features an osmium ion sandwiched between two cyclopentadienyl rings. It is isomorphous to the lighter homologue ruthenocene, both crystallizing in an eclipsed conformation. This is in contrast to ferrocene, which crystallizes with its rings staggered.
Compared to ferrocene and ruthenocene, osmocene is less reactive towards electrophilic aromatic substitution but has the greatest tendency towards adduct formation with Lewis acids.
The osmocenium cation [Os(C5H5)2]+ dimerizes, forming a binuclear complex with an Os-Os bond. In contrast, the decamethylosmocenium cation [Os(C5(CH3)5)2]+ is stable as the monomer.
Uses.
In 2009, Horst Kunkely and Arnd Vogler reported the possibility of photocatalytic water splitting with osmocene as a catalyst.


Potassium polyacrylate

Potassium polyacrylate is a potassium salt of polyacrylic acid with the chemical formula [−CH2−CH(CO2K)−]n. As a type of superabsorbent polymer, it can absorb hundreds times its original weight in purified water.
Different from sodium polyacrylate, potassium polyacrylate can be used as water retaining agent in agriculture and won't cause soil salinization.
Uses.
Water retaining agent.
As a type of superabsorbent polymer for plants (Ag-SAP), potassium polyacrylate can increase moisture availability to plants.
It mixes with soil to increase the soil’s capacity for holding water (with water gel form which stay in soil for months) and making it available to plants. This improved soil readily releases moisture, along with water-soluble nutrients, to plant roots on demand. The specific retention of potassium polyacrylate is weaker than roots of most plants.
It can be used in seed coating, grow seedlings, planting crops, topdressing for crops, planting/transplanting trees, flower transport etc.
Mechanism.
Potassium polyacrylate works similar to a sponge under the soil surface. It is composed of a set of polymeric chains, which are linked together chemically to become a water-insoluble, net-like matrix that gently attracts and holds hydrogen molecules. The immense size and weight of its molecular structure allows each potassium polyacrylate granule to absorb over 500 times its original weight in purified water.
It does not ‘bind’ water tightly. The potassium polyacrylate granules release just the right amount of water in response to a plant’s root suction. There is no waterlogging or other ill effects caused by ‘free’ water filling air cavities in the soil.
Potassium polyacrylate maximizes plant growth by reducing plant stress. It also absorbs and releases soil nutrients, water-soluble fertilizer and chemicals in the same manner as water, creating a healthy microenvironment in the plant root zone. The result is faster germination, quicker emergence of seeds, consistent growth and higher, better-quality yields of edibles with less water and fewer inputs.
As a soil amendment, potassium polyacrylate can improve the soil water holding capacity.


Unpentquadium


Carbocation

A carbocation () is an ion with a positively charged carbon atom. Among the simplest examples are the methenium , methanium and vinyl cations. Occasionally, carbocations that bear more than one positively charged carbon atom are also encountered (e.g., ethylene dication ).
Until the early 1970s, all carbocations were called "carbonium ions". In the present-day definition given by the IUPAC, a carbocation is any even-electron cation with significant partial positive charge on a carbon atom. They are further classified in two main categories according to the coordination number of the charged carbon: three in the carbenium ions and five in the carbonium ions. This nomenclature was proposed by G. A. Olah. Carbonium ions, as originally defined by Olah, are characterized by a three-center two-electron delocalized bonding scheme and are essentially synonymous with so-called 'non-classical carbocations', which are carbocations that contain bridging C–C or C–H σ-bonds. However, others have more narrowly defined the term 'carbonium ion' as formally protonated or alkylated alkanes (i.e., CR5+, where R is hydrogen or alkyl), to the exclusion of non-classical carbocations like the 2-norbornyl cation.
Definitions.
According to the IUPAC, a "carbocation" is any cation containing an even number of electrons in which a significant portion of the positive charge resides on a carbon atom. Prior to the observation of five-coordinate carbocations by Olah and coworkers, "carbocation" and "carbonium ion" were used interchangeably. Olah proposed a redefinition of "carbonium ion" as a carbocation featuring any type of three-center two-electron bonding, while a "carbenium ion" was newly coined to refer to a carbocation containing only two-center two-electron bonds with a three-coordinate positive carbon. Subsequently, others have used the term "carbonium ion" more narrowly to refer to species that are derived (at least formally) from electrophilic attack of H+ or R+ on an alkane, in analogy to other main group onium species, while a carbocation that contains any type of three-centered bonding is referred to as a "non-classical carbocation". In this usage, 2-norbornyl cation is not a carbonium ion, because it is formally derived from protonation of an alkene (norbornene) rather than an alkane, although it is a non-classical carbocation due to its bridged structure. The IUPAC acknowledges the three divergent definitions of carbonium ion and urges care in the usage of this term. For the remainder of this article, the term "carbonium ion" will be used in this latter restricted sense, while "non-classical carbocation" will be used to refer to any carbocation with C–C and/or C–H σ-bonds delocalized by bridging.
Since the late 1990s, most textbooks have stopped using the term carbonium ion for the classical three-coordinate carbocation. However, some university-level textbooks continue to use the term carbocation as if it were synonymous with carbenium ion, or discuss carbocations with only a fleeting reference to the older terminology of carbonium ions or carbenium and carbonium ions. One textbook retains the older name of carbonium ion for carbenium ion to this day, and uses the phrase "hypervalent carbonium ion" for .
A carbocation with a two-coordinate positive carbon derived from formal removal of a hydride ion (H−) from an alkene is known as a "vinyl cation". In the absence of geometric constraints, most substituted vinyl cations carry the formal positive charge on an sp-hydridized carbon atom of linear geometry. A two-coordinate approximately sp2-hybridized cation resulting from the formal removal of a hydride ion from an arene is termed an "aryl cation". These carbocations are relatively unstable (aryl cations especially so) and are infrequently encountered. Hence, they are frequently omitted from introductory and intermediate level textbooks. The IUPAC definition stipulates that carbocations are even-electron species; hence, radical cations like CH4•+ that are frequently encountered in mass spectrometry are not considered to be carbocations.
History.
The history of carbocations dates back to 1891 when G. Merling reported that he added bromine to tropylidene (cycloheptatriene) and then heated the product to obtain a crystalline, water-soluble material, . He did not suggest a structure for it; however, Doering and Knox convincingly showed that it was tropylium (cycloheptatrienylium) bromide. This ion is predicted to be aromatic by Hückel's rule.
In 1902, Norris and Kehrman independently discovered that colorless triphenylmethanol gives deep-yellow solutions in concentrated sulfuric acid. Triphenylmethyl chloride similarly formed orange complexes with aluminium and tin chlorides. In 1902, Adolf von Baeyer recognized the salt-like character of the compounds formed. Trityl carbocation (shown below) as a stable carbocationic system has been used as homogeneous organocatalyst in organic synthesis.
He dubbed the relationship between color and salt formation halochromy, of which malachite green is a prime example.
Carbocations are reactive intermediates in many organic reactions. This idea, first proposed by Julius Stieglitz in 1899, was further developed by Hans Meerwein in his 1922 study of the Wagner–Meerwein rearrangement. Carbocations were also found to be involved in the SN1 reaction, the E1 reaction, and in rearrangement reactions such as the Whitmore 1,2 shift. The chemical establishment was reluctant to accept the notion of a carbocation and for a long time the Journal of the American Chemical Society refused articles that mentioned them.
The first NMR spectrum of a stable carbocation in solution was published by Doering et al. in 1958. It was the heptamethylbenzenium ion, made by treating hexamethylbenzene with methyl chloride and aluminium chloride. The stable 7-norbornadienyl cation was prepared by Story et al. in 1960 by reacting norbornadienyl chloride with silver tetrafluoroborate in sulfur dioxide at −80 °C. The NMR spectrum established that it was non-classically bridged (the first stable non-classical ion observed).
In 1962, Olah directly observed the "tert"-butyl carbocation by nuclear magnetic resonance as a stable species on dissolving "tert"-butyl fluoride in magic acid. The NMR of the norbornyl cation was first reported by Schleyer et al. and it was shown to undergo proton-scrambling over a barrier by Saunders et al.
Structure and properties.
Carbonium ions can be thought of as protonated alkanes. Although alkanes are usually considered inert, under superacid conditions (e.g., HF/SbF5), the C-H sigma bond can act as a donor to H+"." This results in a species that contains a 3c-2e bond between a carbon and two hydrogen atoms, a type of bonding common in boron chemistry, though relatively uncommon for carbon. As an alternative view point, the 3c-2e bond of carbonium ions could be considered as a molecule of H2 coordinated to a carbenium ion ("see below"). Indeed, carbonium ions frequently decompose by loss of molecular hydrogen to form the corresponding carbenium ion. Structurally, the methanium ion CH5+ is computed to have a minimum energy structure of "C"s symmetry. However, the various possible structures of the ion are close in energy and separated by shallow barriers. Hence, the structure of the ion is often described as fluxional. Although there appear to be five bonds to carbon in carbonium ions, they are not hypervalent, as the electron count around the central carbon is only eight, on account of the 3c-2e bond.
In contrast, at least in a formal sense, carbenium ions are derived from the protonation (addition of H+) or alkylation (addition of R+) of a carbene or alkene. Thus, in at least one of their resonance depictions, they possess a carbon atom bearing a formal positive charge that is surrounded by a sextet of electrons (six valence electrons) instead of the usual octet required to fill the valence shell of carbon (octet rule). Therefore, carbenium ions (and carbocations in general) are often reactive, seeking to fill the octet of valence electrons as well as regain a neutral charge. In accord with VSEPR and Bent's rule, unless geometrically constrained to be pyramidal (e.g., 1-adamantyl cation), 3-coordinate carbenium ions are usually trigonal planar, with a pure p character empty orbital as its lowest unoccupied molecular orbital (LUMO) and CH/CC bonds formed from C(sp2) orbitals. A prototypical example is the methyl cation, . For the same reasons, carbocations that are 2-coordinate (vinyl cations) are generally linear in geometry, with CH/CC bonds formed from C(sp) orbitals.
Alkyl-substituted carbocations follow the order 3° > 2° > 1° > methyl in stability, as can be inferred by the hydride ion affinity values (231, 246, 273, and 312 kcal/mol for (CH3)3C+, (CH3)2CH+, C2H5+, and CH3+). The effect of alkyl substitution is a strong one: tertiary cations are stable and many are directly observable in superacid media, but secondary cations are usually transient and only the isopropyl, "s"-butyl, and cyclopentyl cations have been observed in solution. There is seldom any experimental support for primary carbocations in the solution phase, even as transient intermediates (the ethyl cation has been proposed for reactions in 99.9% sulfuric acid and in FSO2OH-SbF5), and methyl cation has only been unambiguously identified in the gas phase. In most, if not all cases, the ground state of alleged primary carbocations consist of bridged structures in which positive charge is shared by two or more carbon atoms and are better described as side-protonated alkenes, edge-protonated cyclopropanes, or corner-protonated cyclopropanes rather than true primary cations. Even the simple ethyl cation, C2H5+, has been demonstrated experimentally and computationally to be bridged and can be thought of as a symmetrically protonated ethylene molecule. The same is true for higher homologues like 1-propyl and 1-butyl cations. Neopentyl derivatives are thought to ionize with concomitant migration of a methyl group (anchimeric assistance); thus, in most if not all cases, a discrete neopentyl cation is not believed to be involved.
The stabilization by alkyl groups is explained by hyperconjugation. The donation of electron density from a β C-H or C-C bond into the unoccupied p orbital of the carbocation (a σCH/CC → p interaction) allows the positive charge to be delocalized.
Based on hydride ion affinity, the parent vinyl cation is less stable than even a primary sp2-hybridized carbocation, while an α alkyl-substituted vinyl cation has a stability that is comparable to the latter. Hence, vinyl cations are relatively uncommon intermediates. They can be generated by the ionization of a vinyl electrophile, provided the leaving group is sufficiently good (e.g., TfO–, IPh, or N2). They have been implicated as intermediates in some vinyl substitution reactions (designated as SN1(vinyl)) and as intermediates in the electrophilic addition reactions of arylalkynes. With the exception of the parent vinyl cation, which is believed to be a bridged species, and geometrically constrained cyclic vinyl cations, most vinyl cations take on sp hybridization and are linear.
Aryl cations are more unstable than vinyl cations, due to the ring-enforced distortion to a nonlinear geometry and approximately sp2-character of the unoccupied orbital. Only N2 in aryldiazonium salts is a good enough leaving group for the chemical generation of aryl cations.
Alkynyl cations are extremely unstable, much less stable than even CH3+ (hydride ion affinity 386 kcal/mol vs. 312 kcal/mol for CH3+) and cannot be generated by purely chemical means. They can, however, be generated radiochemically via the beta decay of tritium: (RC≡CT → [RC≡C3He]+ + e– + e → RC≡C+ + 3He + e– + e).
In terms of reactivity, carbocations are susceptible to attack by nucleophiles, like water, alcohols, carboxylates, azide, and halide ions, to form the addition product. Strongly basic nucleophiles, especially hindered ones, favor elimination over addition. Because even weak nucleophiles will react with carbocations, most can only be directly observed or isolated in non-nucleophilic media like superacids.
Carbocations typically undergo rearrangement reactions from less stable structures to equally stable or more stable ones by migration of an alkyl group or hydrogen to the cationic center to form a new carbocationic center. This often occurs with rate constants in excess of 1010 s−1 at ambient temperature and still takes place rapidly (compared to the NMR timescale) at temperatures as low as –120 °C ("see Wagner-Meerwein shift"). In especially favorable cases like the 2-norbornyl cation, hydrogen shifts may still take place at rates fast enough to interfere with X-ray crystallography at 86 K (–187 °C). Typically, carbocations will rearrange to give a tertiary isomer. For instance, all isomers of C6H11+ rapidly rearrange to give the 1-methyl-1-cyclopentyl cation. This fact often complicates synthetic pathways. For example, when 3-pentanol is heated with aqueous HCl, the initially formed 3-pentyl carbocation rearranges to a statistical mixture of the 3-pentyl and 2-pentyl. These cations react with chloride ion to produce about 3-chloropentane and 2-chloropentane. The Friedel-Crafts alkylation suffers from this limitation; for this reason, the acylation (followed by Wolff-Kishner or Clemmensen reduction to give the alkylated product) is more frequently applied.
A carbocation may be stabilized by resonance by a carbon-carbon double bond or by the lone pair of a heteroatom adjacent to the ionized carbon. In order for a carbocation to be resonance-stabilized, the molecular orbital of the donating group must have the proper symmetry, orientation, and energy level to interact with the empty 2p orbital of the carbocation. Such cations as "allyl" cation CH2=CH–CH2+ and "benzyl" cation C6H5–CH2+ are more stable than most other carbocations due to donation of electron density from π systems to the cationic center. Furthermore, carbocations present in aromatic molecules are especially stabilized, largely due to the delocalized π electrons characteristic of aromatic rings. Molecules that can form allyl or benzyl carbocations are especially reactive. These carbocations where the C+ is adjacent to another carbon atom that has a double or triple bond have extra stability because of the overlap of the empty p orbital of the carbocation with the p orbitals of the π bond. This overlap of the orbitals allows the positive charge to be dispersed and electron density from the π system to be shared with the electron-deficient center, resulting in stabilization. The doubly- and triply-benzylic carbocations, diphenylcarbenium and triphenylcarbenium (trityl) cation, are particularly stable. For the same reasons, the partial p character of strained C–C bonds in cyclopropyl groups also allows for donation of electron density and stabilizes the "cyclopropylmethyl" (cyclopropylcarbinyl) cation.
The stability order of carbocations, from most stable to least stable as reflected by hydride ion affinity (HIA) values, are as follows (HIA values in kcal/mol in parentheses):
As noted in the history section, the tropylium cation (C7H7+) was one of the first carbocations to be discovered, due to its aromatic stability. Tropylium's carbocation is so stabilized that the molecule can be isolated and sold as a salt. On the other hand, the anti-aromatic cyclopentadienyl cation (C5H5+) is destabilized by some 40 kcal/mol. The cyclopropenium cation (C3H3+), although somewhat destabilized by angle strain, is still clearly stabilized by aromaticity when compared to its open-chain analog, allyl cation. These varying cation stabilities, depending on the number of π electrons in the ring system, can furthermore be crucial factors in reaction kinetics. The formation of an aromatic carbocation is much faster than the formation of an anti-aromatic or open-chain carbocation. Given the role of carbocations in many reaction schemes, such as SN1 for example, choosing the conjugation of starting materials can be a powerful method for conferring kinetic favorability or disfavorability, as the rate constant for any given step is dependent on the step's activation energy according to the Arrhenius equation.
The effect of hyperconjugation is strongly stabilizing for carbocations: hyperconjugation with alkyl substituents is often as stabilizing or even more so than conjugation with a π system. Although conjugation to unsaturated groups results in significant stabilization by the mesomeric effect (resonance), the benefit is partially offset by the presence of a more electronegative sp2 or sp carbon next to the carbocationic center. Thus, as reflected by hydride ion affinities, a secondary carbocation is more stabilized than the allyl cation, while a tertiary carbocation is more stabilized than the benzyl cation — results that may seem counterintuitive on first glance.
Oxocarbenium and iminium ions have important secondary canonical forms (resonance structures) in which carbon bears a positive charge. As such, they are carbocations according to the IUPAC definition although some chemists do not regard them to be "true" carbocations, as their most important resonance contributors carry the formal positive charge on an oxygen or nitrogen atom, respectively.
Non-classical ions.
Some carbocations such as the 2-norbornyl cation exhibit more or less symmetrical three-center two-electron bonding. Such structures, referred to as non-classical carbocations, involve the delocalization of the bonds involved in the σ-framework of the molecule, resulting in C–C and C–H bonds of fractional bond order. This delocalization results in additional stabilization of the cation. For instance, depicted as a classical carbenium ion, 2-norbornyl cation appears to be a secondary carbocation. However, it is more stable than a typical "secondary" carbocation, being roughly as stable as a tertiary carbocation like "t"-butyl cation, according to hydride ion affinity.
The existence of non-classical carbocations was once the subject of great controversy. On opposing sides were Brown, who believed that the what appeared to be a non-classical carbocation represents the average of two rapidly equilibrating classical species (or possibly two structures exhibiting some degree of bridging or leaning but is nevertheless not symmetric) and that the true non-classical structure is a transition state between the two potential energy minima, and Winstein, who believed that a non-classical structure that possessed a plane of symmetry was the sole potential energy minimum and that the classical structures merely two contributing resonance forms of this non-classical species. George Olah's discovery of superacidic media to allow carbocations to be directly observed, together with a very sensitive NMR technique developed by Martin Saunders to distinguish between the two scenarios, played important roles in resolving this controversy. At least for the 2-norbornyl cation itself, the controversy has been settled overwhelmingly in Winstein's favor, with no sign of the putative interconverting classical species, even at temperatures as low as 6 K, and a 2013 crystal structure showing a distinctly non-classical structure. A variety of carbocations (e.g., ethyl cation, "see above") are now believed to adopt non-classical structures. However, in many cases, the energy difference between the two possible "classical" structures and the "non-classical" one is very small, and it may be difficult to distinguish between the two possibilities experimentally.
Specific carbocations.
A non-classical structure for C4H7+ is supported by substantial experimental evidence from solvolysis experiments and NMR studies conducted in non-nucleophilic media. One or both of two structures, the cyclopropylcarbinyl cation and the bicyclobutonium cation, were invoked to account for the observed reactivity in various experiments, while the NMR data point to a highly fluxional system that undergoes rapid rearrangement to give an averaged spectrum consisting of only two 13C NMR signals, even at temperatures as low as –132 °C. Computationally, it was confirmed that the energetic landscape of the C4H7+ system is very flat, and that the two isomers postulated based on experimental data are very close in energy, the bicyclobutonium structure being computed to be just 0.4 kcal/mol more stable than the cyclopropylcarbinyl structure. In the solution phase (SbF5-SO2ClF-SO2F2, with SbF6– as the counterion), the bicyclobutonium structure predominates over the cyclopropylcarbinyl structure in a 84:16 ratio at –61 °C.
Three other possible structures, two classical structures (the homoallyl cation and cyclobutyl cation) and a more highly delocalized non-classical structure (the tricyclobutonium ion), are now known to be less stable isomers (or merely a transition state rather than an energy minimum in the case of the cyclobutyl cation).
Substituted cyclopropylcarbinyl cations have also been studied by NMR:
In the NMR spectrum of a dimethyl derivative, two nonequivalent signals are found for the two methyl groups, indicating that the molecular conformation of this cation is not perpendicular (as in A), which possesses a mirror plane, but is bisected (as in B) with the empty p-orbital parallel to the cyclopropyl ring system:
In terms of bent bond theory, this preference is explained by assuming favorable orbital overlap between the filled cyclopropane bent bonds and the empty p-orbital.


Jarosite

Jarosite is a basic hydrous sulfate of potassium and ferric iron (Fe-III) with a chemical formula of KFe3(SO4)2(OH)6. This sulfate mineral is formed in ore deposits by the oxidation of iron sulfides. Jarosite is often produced as a byproduct during the purification and refining of zinc and is also commonly associated with acid mine drainage and acid sulfate soil environments.
Physical properties.
Jarosite has a trigonal crystal structure and is brittle, with basal cleavage, a hardness of 2.5-3.5, and a specific gravity of 3.15-3.26. It is translucent to opaque with a vitreous to dull luster, and is colored dark yellow to yellowish-brown. It can sometimes be confused with limonite or goethite with which it commonly occurs in the gossan (oxidized cap over an ore body). Jarosite is an iron analogue of the potassium aluminium sulfate, alunite.
Solid solution series.
The alunite supergroup includes the alunite, jarosite, beudantite, crandallite and florencite subgroups. The alunite supergroup minerals are isostructural with each other and substitution between them occurs, resulting in several solid solution series. The alunite supergroup has the general formula AB3(TO4)2(OH)6. In the alunite subgroup B is Al, and in the jarosite subgroup B is Fe3+. The beudantite subgroup has the general formula AB3(XO4)(SO4)(OH)6, the crandallite subgroup AB3(TO4)2(OH)5•H2O and the florencite subgroup AB3(TO4)2(OH)5 or 6.
In the jarosite-alunite series Al may substitute for Fe and a complete solid solution series between jarosite and alunite, KAl3(SO4)2(OH)6, probably exists, but intermediate members are rare. The material from Kopec, Czech Republic, has about equal Fe and Al, but the amount of Al in jarosite is usually small.
When jarosite forms from pyrite oxidation in sedimentary clays, the main sources of K+ are illite, a non-swelling clay, or K-feldspar. In other geological settings mica's alteration can also be a source of potassium.
In the jarosite-natrojarosite series Na substitutes for K to at least Na/K = 1:2.4 but the pure sodium end member NaFe3(SO4)2(OH)6 is not known in nature. Minerals with Na > K are known as natrojarosite. End member formation (jarosite and natrojarosite) is favoured by a low temperature environment, less than 100 °C, and is illustrated by the oscillatory zoning of jarosite and natrojarosite found in samples from the Apex Mine, Arizona, and Gold Hill, Utah. This indicates that there is a wide miscibility gap between the two end members, and it is doubtful whether a complete series exists between jarosite and natrojarosite.
In hydroniumjarosite the hydronium ion H3O+ can also substitute for K+, with increased hydronium ion content causing a marked decrease in the lattice parameter "c", although there is little change in "a". Hydroniumjarosite will only form from alkali-deficient solutions, as alkali-rich jarosite forms preferentially.
Divalent cations may also substitute for the monovalent cation K+ in the A site. Charge balance may be achieved in three ways.
History.
Jarosite was first described in 1852 by August Breithaupt in the Barranco del Jaroso in the Sierra Almagrera (near Los Lobos, Cuevas del Almanzora, Almería, Spain). The name jarosite is also directly derived from Jara, the Spanish name of a yellow flower that belongs to the genus "Cistus" and grows in this sierra. The mineral and the flower have the same color.
Mysterious spheres of clay 1.5 to 5 inches in diameter and covered with jarosite have been found beneath the Temple of the Feathered Serpent an ancient six level stepped pyramid 30 miles from Mexico City.
Mars exploration.
Ferric sulfate and jarosite have been detected by three martian rovers: "Spirit", "Opportunity" and "Curiosity". These substances are indicative of strongly oxidizing conditions prevailing at the surface of Mars. In May 2009, the "Spirit" rover became stuck when it drove over a patch of soft ferric sulfate that had been hidden under a veneer of normal-looking soil.
Because iron sulfate has very little cohesion, the rover's wheels could not gain sufficient traction to pull the body of the rover out of the iron sulfate patch. Multiple techniques were attempted to extricate the rover, but the wheels eventually sank so deeply into the iron sulfate that the body of the rover came to rest on the Martian surface, preventing the wheels from exerting any force on the material below them. As the JPL team failed to recover the mobility of "Spirit", it signified the end of the journey for the rover.
Antarctica deep borehole.
On Earth, jarosite is mainly associated with the ultimate stage of pyrite oxidation in clay environment and can also be found in mine tailings waste where acidic conditions prevail. Against all expectations, jarosite has also been fortuitously discovered in minute quantities in the form of small dust particules in ice cores recovered from a deep borehole in Antarctica. This surprising discovery was made by geologists who were searching for specific minerals capable to indicate ice age cycles within the layers of a 1620 meters-long ice core. Geologists speculate that jarosite dust could also have accumulated within ice on glaciers on Mars. However, this hypothesis is a matter of controverses, because on Mars jarosite deposits can be very thick (up to 10 meters); but Mars is also a very dusty planet and in the absence of plate tectonic on Mars glacial dust deposits might have accumulated during long periods of time.
Use in materials science.
Jarosite is also a more generic term denoting an extensive family of compounds of the form AM3(OH)6(SO4)2, where A+ = Na, K, Rb, NH4, H3O, Ag, Tl and M3+ = Fe, Cr, V. In condensed matter physics and materials science they are renowned for containing layers with kagome lattice structure, relating to geometrically frustrated magnets.


Mucobromic acid

Mucobromic acid is an organic compound that consists of a dibrominated alkene with aldehyde and carboxylic acid functional groups. It easily tautomerizes to a furanone hemiacetal form. This compound, and the analogous mucochloric acid (CAS #87-56-9), form the group of known mucohalic acids. The bromide appears to behave similarly to the more heavily studied chloride.
Synthesis and structure.
Mucobromic acid can be synthesized by bromination of furfural via an oxidation/decarboxylation process:
Mucobromic acid exists as a mixture acyclic and cyclic isomers. The compound can be reduced using sodium borohydride to give the lactone.
Hydrolysis under basic conditions of either the chloro or bromo compound involves substitution of the halide adjacent to the acid. The resulting mucoxyhalic acids exist as a mixture of keto and enol forms. The reaction occurs via a conjugate addition/elimination of the alkene–aldehyde part of the structure.
Hazards.
Mucohalic acids have received attention since they are products of the halogenation of biomass. They are genotoxins and potential carcinogens. They have the ability to alkylate certain DNA bases, specifically guanosine, adenosine, and cytosine.


Ethanium

In chemistry, ethanium or protonated ethane is a highly reactive positive ion with formula . It can be described as a molecule of ethane () with one extra proton (hydrogen nucleus), that gives it a +1 electric charge.
Ethanium is one of the simplest carbonium ions (after methanium ). It was first detected as a rarefied gas in 1960 by S. Wexler and N. Jesse. It easily dissociates into ethenium and molecular hydrogen .
Production.
Ethanium was first detected by infrared spectroscopy among the ions produced by electrical discharges in rarefied methane or ethane gas.
Ethanium can also be produced by irradiating methane containing traces of ethane with an electron beam at low pressure (about 2 mmHg). The electron beam first creates methanium and methenium ions. The former rapidly transfer their proton to ethane:
The latter reaction is also observed when , or ions are injected into ethane at somewhat lower pressure.
Stability and reactions.
At about 1 mmHg and 30 °C, ethanium dissociates very slowly to ethenium and , across an energy barrier of about 10 kcal/mol; the decomposition is considerably faster at 92 °C. The decomposition has been claimed to be nearly athermal but with 8 kcal/mol of free energy due to increase of entropy.
Structure.
Like its "unsaturated" relatives ethenium and ethynium , the ethanium ion was conjectured to have (at least momentarily) a proton bound simultaneously to the two carbon atoms, and the electrical charge evenly spread between them, as in other non-classical ions. The alternative "classical" structure would have the charge and the extra hydrogen bound to only one of the two atoms, i.e. a methylated methanium ion.
Earlier calculations had predicted that the energies of the two forms should be 4 to 12 kcal/mol lower than the dissociated state + , and they should be separated by a slightly positive energy barrier. Gas-phase infrared spectroscopy by Yeh and others (1989) has shown that both forms are stable. The bridged structure has the lowest energy, 4 to 8 kcal/mol lower than the classical one.
Refined calculations by Obata and Hirao (1993) predict that the most stable form has three orthogonal planes of symmetry (C2"v") with the two subgroups in the eclipsed configuration (unlike ethane, whose ground state has the staggered configuration). Four "bottom" H atoms lie on a plane opposite to the bridging H atom and the other two "top" H atoms. The approximate computed distances are C–C 0.211 nm, C–H 0.124 nm (bridging), 0.107 nm (bottom) and 0.108 nm (top); the C–H–C angle at the bridge is about 116 degrees, the H–C–H angles are 116 degrees (bottom-bottom) and 114 degrees (bottom-top). However, there are other configurations with near-minimum energy, including one where the two subgroups are slightly staggered (with C"s" symmetry), another where one of the carbons of a ion is loosely bound to an molecule 0.250 nm away.


S-Nitroso-N-acetylpenicillamine

"S"-Nitroso-"N"-acetylpenicillamine (SNAP) is the chemical compound with the formula ONSC(CH3)2CH(NHAc)CO2H.
SNAP is an "S"-nitrosothiol and is used as a model for the general class of "S"-nitrosothiols which have received much attention in biochemistry because nitric oxide and some organic nitroso derivatives serve as signaling molecules in living systems, especially related to vasodilation.
SNAP is derived from the amino acid penicillamine. "S"-Nitrosoglutathione is a related agent.


Dimethylmagnesium

Dimethylmagnesium is an organomagnesium compound. It is a white pyrophoric solid.
Preparation.
Like other dialkylmagnesium compounds, it is prepared by adding at least one equivalent of dioxane to a solution of methylmagnesium halide:
Addition of 0.5 molar equivalents of 1,4-dioxane causes precipitation of MgX2 as an insoluble dioxane-adduct, driving the Schlenk equilibrium to (CH3)2Mg.
It can also be prepared by treating calcium, magnesium and methyl iodide in diethyl ether or by combining dimethyl mercury and magnesium.
Properties.
The structure of this compound has been determined by X-ray crystallography. The material is a polymer with the same connectivity as silicon disulfide, featuring tetrahedral magnesium centres, each surrounded by bridging methyl groups. The Mg-C distances are 224 pm. Dimethylberylium adopts the same structure.