Mechanical properties of carbon nanotubes

The mechanical properties of carbon nanotubes reveal them as one of the strongest materials in nature. Carbon nanotubes (CNTs) are long hollow cylinders of graphene. Although graphene sheets have 2D symmetry, carbon nanotubes by geometry have different properties in axial and radial directions. It has been shown that CNTs are very strong in the axial direction. Young's modulus on the order of and tensile strength of were obtained.
Strength.
Carbon nanotubes are the strongest and stiffest materials yet discovered in terms of tensile strength and elastic modulus respectively. This strength results from the covalent sp2 bonds formed between the individual carbon atoms. In 2000, a multi-walled carbon nanotube was tested to have a tensile strength of 63 gigapascals (9,100,000 psi). (For illustration, this translates into the ability to endure tension of a weight equivalent to 6,422 kilograms-force (62,980 N; 14,160 lbf) on a cable with cross-section of 1 square millimetre (0.0016 sq in).) Further studies, such as one conducted in 2008, revealed that individual CNT shells have strengths of up to ≈100 gigapascals (15,000,000 psi), which is in agreement with quantum/atomistic models. Since carbon nanotubes have a low density for a solid of 1.3 to 1.4 g/cm3, its specific strength of up to 48,000 kN·m·kg−1 is the best of known materials, compared to high-carbon steel's 154 kN·m·kg−1.
Under excessive tensile strain, the tubes will undergo plastic deformation, which means the deformation is permanent. This deformation begins at strains of approximately 5% and can increase the maximum strain the tubes undergo before fracture by releasing strain energy.
Although the strength of individual CNT shells is extremely high, weak shear interactions between adjacent shells and tubes lead to significant reduction in the effective strength of multi-walled carbon nanotubes and carbon nanotube bundles down to only a few GPa. This limitation has been recently addressed by applying high-energy electron irradiation, which crosslinks inner shells and tubes, and effectively increases the strength of these materials to ≈60 GPa for multi-walled carbon nanotubes and ≈17 GPa for double-walled carbon nanotube bundles.
CNTs are not nearly as strong under compression. Because of their hollow structure and high aspect ratio, they tend to undergo buckling when placed under compressive, torsional, or bending stress.
EExperimental observation; TTheoretical prediction
Radial elasticity.
On the other hand, there was evidence that in the radial direction they are rather soft. The first transmission electron microscope observation of radial elasticity suggested that even the van der Waals forces can deform two adjacent nanotubes. Later, nanoindentations with atomic force microscope were performed by several groups to quantitatively measure radial elasticity of multiwalled carbon nanotubes and tapping/contact mode atomic force microscopy was also performed on single-walled carbon nanotubes. Young's modulus of on the order of several GPa showed that CNTs are in fact very soft in the radial direction.
Radial direction elasticity of CNTs is important especially for carbon nanotube composites where the embedded tubes are subjected to large deformation in the transverse direction under the applied load on the composite structure.
One of the main problems in characterizing the radial elasticity of CNTs is the knowledge about the internal radius of the CNT; carbon nanotubes with identical outer diameter may have different internal diameter (or the number of walls). In 2008, a method using an atomic force microscope was introduced to determine the exact number of layers and hence the internal diameter of the CNT. In this way, mechanical characterization is more accurate.
Hardness.
Standards single-walled carbon nanotubes can withstand a pressure up to 25 GPa without [plastic/permanent] deformation. They then undergo a transformation to superhard phase nanotubes. Maximum pressures measured using current experimental techniques are around 55 GPa. However, these new superhard phase nanotubes collapse at an even higher, albeit unknown, pressure.
The bulk modulus of superhard phase nanotubes is 462 to 546 GPa, even higher than that of diamond (420 GPa for single diamond crystal).
Wettability.
The surface wettability of CNT is of importance for its applications in various settings. Although the intrinsic contact angle of graphite is around 90°, the contact angles of most as-synthesized CNT arrays are over 160°, exhibiting a superhydrophobic property. By applying a voltage as low as 1.3V, the extreme water repellant surface can be switched to a superhydrophilic one.
Kinetic properties.
Multi-walled nanotubes are multiple concentric nanotubes precisely nested within one another. These exhibit a striking telescoping property whereby an inner nanotube core may slide, almost without friction, within its outer nanotube shell, thus creating an atomically perfect linear or rotational bearing. This is one of the first true examples of molecular nanotechnology, the precise positioning of atoms to create useful machines. Already, this property has been utilized to create the world's smallest rotational motor. Future applications such as a gigahertz mechanical oscillator are also envisioned.
Defects.
As with any material, the existence of a crystallographic defect affects the material properties. Defects can occur in the form of atomic vacancies. High levels of such defects can lower the tensile strength by up to 85%. An important example is the Stone Wales defect, otherwise known as the 5-7-7-5 defect because it creates a pentagon and heptagon pair by rearrangement of the bonds. Because of the very small structure of CNTs, the tensile strength of the tube is dependent on its weakest segment in a similar manner to a chain, where the strength of the weakest link becomes the maximum strength of the chain.
Plastic deformation.
A typical 3D material undergoes plastic deformation, meaning that the deformation is permanent, by the movement of 1D dislocations through the material. During this process, these dislocations can interact with each other and multiply. Because CNTs are themselves 1D materials, the well-known generation and multiplication mechanisms (such as a Frank-Read source) for 1D dislocations do not apply.
Instead, CNTs undergo plastic deformation through the formation and movement of defects, primarily topological defects such as the Stone Wales defect or 5-7-7-5 defect. The 5-7-7-5 defect can also be thought of as a pair of 5-7 defects, in which each defect is adjacent to one 5-membered and two 7-membered rings. This defect structure is metastable, so it takes an energy of several eV to nucleate, or form. In addition, the defect moves by the separate migration of the 5-7 defect pairs. This motion is also associated with an energy barrier. The exact energy depends on the configuration and chirality of the particular CNT. The activation energy for the formation of these defects in a CNT of diameter formula_1 and chiral angle formula_2 can be estimated as formula_3eV, where formula_4 is the external strain. This activation energy barrier partially explains the low ductility of CNTs (~6-15%) at room temperature. However, it can be overcome at high temperatures or with the application of suitable strains. For example, the defect is nucleated at positions experiencing high tensile stress in armchair-type CNTs, and at positions experiencing high compressive stress in zigag-type CNTs.
Applied stresses can overcome the energy barrier needed to move 5-7 defect pairs. Another way of understanding this is that when strained, a CNT releases strain by forming these defects spontaneously. For example, in (5,5) tubes, a critical tensile strain of ~5% results in defect generation. The defect structure reduces strain because the heptagon geometry is able to stretch more than the original hexagonal rings, while the C-C bond remains about the same length. Bending the tubes beyond a critical curvature has the same effect. This behavior can be approximated by a simple, semi-quantitative analysis. Applying a stress formula_5 over a tube of length formula_6 and diameter formula_6 does work approximately equal to
formula_8 on the tube, where formula_9 is the Burgers vectors for the defect, formula_10 is the bending curvature, and formula_11 relates the Young's modulus of the CNT to that of graphene. The energy increase resulting from the defect creation and the separation of the 5-7 pairs is approximately given by
formula_12. Here, formula_13 is the dislocation core energy and formula_14 gives the interaction energy between the defect pairs. Defect motion occurs when the work done by an applied stress overcomes it, such that the required bending curvature is inversely proportional to the diameter of the CNT:
formula_15. Similarly, thermal vibrations can provide the energy required for defect nucleation and motion. In fact, a combination of stress and high temperature is required to induce observable plastic deformation in CNTs. This has been achieved in the literature via the application of a current, which causes resistive heating in the material. For CNTs subjected to temperatures above 1500K, elongations up to 280% have been reported. This kind of behavior is called superplasticity. At these high temperatures, kinks may form and move by climb as well as glide. Climb of kinks is evidenced by the fact that they do not always move along the close-packed planes in CNTs, but rather along the length of a tube. When kinks do glide along close-packed planes in CNTs, they follow a helical path. It is proposed that elevated temperatures allow for the diffusion of vacancies, so that defects climb through a process similar to that observed in 3D crysalline materials.


Stannatrane

A stannatrane (IUPAC: 1-aza-5-stannabicyclo[3.3.3]undecane) is a tin-based atrane belonging to the larger class of organostannanes. Though the term stannatrane is often used to refer to the more commonly employed carbastannatrane, azastannatranes have also been synthesized (prefix refers to the identity of the atom bound directly to tin center). Stannatrane reagents offer highly selective methods for the incorporation of "R" substituents in complex molecules for late-stage diversification. These reagents differ from their tetraalkyl organostannane analogues in that there is no participation of dummy ligands in the transmetalation step, offering selective alkyl transfer in Stille Coupling reactions. These transmetalating agents are known to be air- and moisture-stable, as well as generally less toxic than their tetraalkyl counterparts.
History and structural properties.
The first carbastannatrane was reported in 1984 by Jurkschat and Tzschach. By reaction of an amino-triGrignard reagent with tin(IV) chloride to yield the stannatrane chloride, which was treated with methyl lithium to yield the corresponding methyl stannatrane. Based on a very small methyl "J"(119Sn–13C) coupling constant of 171 Hz, it was determined that the tin center of methyl stannatrane was indeed pentacoordinate, indicating nitrogen coordination.
The crystal and molecular structure was elucidated by X-ray crystallography. The X-ray diffraction study confirmed the tricyclic ring structure and gave insight toward the geometry of the complex. With a tin-nitrogen distance of 2.624 Å, the formal bond order was calculated to be about 0.46. The presence of the tin-nitrogen interaction, albeit weaker than anticipated, led to a few key discoveries: (1) the distortion from ideal trigonal bipyramidal toward monocapped tetrahedron geometry; (2) the lengthening of the apical tin-methyl bond by ~ 0.1 Å (largest known value for any existing tetraorganotin compounds); (3) the observation of unusual hybridization at the apical tin-methyl bond.
Syntheses of alkyl stannatranes.
A modified synthesis of atrane tricycle utilizes Schwartz's Reagent, triallylamine, and tin(IV) chloride in a one-pot method.
Though the latter step is still commonly used for functionalization of stannatrane chloride to simple alkyl derivatives "via" transmetalation, Biscoe and coworkers have developed a lithiation method that provides access to a variety of enantioenriched alkyl substituents from optically active mesylates (2). After treatment of stannatrane chloride with lithium napthalide, a lithium carbastannatrane was quenched with the corresponding enantiopure mesylate to yield the desired enantioenriched alkyl carbastannatrane in moderate yield with high enantiomeric excess.
Applications in cross-coupling.
The earliest reported use of carbastannatranes in palladium-catalyzed Stille coupling reactions in 1992 compared the efficiency of methyl stannatrane with tetramethyltin in the presence of aryl bromides and alkenyl iodides. Tetramethyltin only resulted in less than five percent conversion, whereas methyl stannatrane resulted in 67% yield under the same conditions. This difference was attributed to the nitrogen lone pair lengthening the tin-methyl bond, increasing its lability toward transmetalation. A method was developed for Stille couplings of aziridinyl stannatranes with aryl electrophiles.
Palladium also catalyzes Stille coupling of secondary alkyl carbastannatranes and aryl electrophiles. This report serves as the first example of employing chiral alkyl carbastannatrane reagents in enantioselective synthesis. Related methodology enable selective acyl substitution using enantioenriched stannatranes as an alternative to classical enolate chemistry.
A stannatrane-mediated Stille coupling was utilized for the synthesis of an anti-methicillin-resistant carbapenem to incorporate an entire side-chain in a single step.


Phenylmagnesium bromide

Phenylmagnesium bromide, with the simplified formula , is a magnesium-containing organometallic compound. It is commercially available as a solution in diethyl ether or tetrahydrofuran (THF). Phenylmagnesium bromide is a Grignard reagent. It is often used as a synthetic equivalent for the phenyl "Ph−" synthon.
Preparation.
Phenylmagnesium bromide is commercially available as solutions of diethyl ether or THF. Laboratory preparation involves treating bromobenzene with magnesium metal, usually in the form of turnings. A small amount of iodine may be used to activate the magnesium to initiate the reaction.
Coordinating solvents such as ether or THF, are required to solvate (complex) the magnesium(II) center. The solvent must be aprotic since alcohols and water contain an acidic proton and thus react with phenylmagnesium bromide to give benzene. Carbonyl-containing solvents, such as acetone and ethyl acetate, are also incompatible with the reagent.
Structure.
Although phenylmagnesium bromide is routinely represented as , the molecule is more complex. The compound invariably forms an adduct with two ligands from the ether or THF solvent. Thus, the Mg is tetrahedral and obeys the octet rule. The Mg–O distances are 201 and 206 pm whereas the Mg–C and Mg–Br distances are 220 pm and 244 pm, respectively.
Chemistry.
Phenylmagnesium bromide is a strong nucleophile as well as a strong base. It can abstract even mildly acidic protons, thus the substrate must be protected where necessary. It often adds to carbonyls, such as ketones, aldehydes. With carbon dioxide, it reacts to give benzoic acid after an acidic workup. If three equivalents are reacted with phosphorus trichloride, triphenylphosphine can be made.


Calcium glycerylphosphate

Calcium glycerylphosphate (or calcium glycerophosphate) is a mineral supplement. Formerly it was sold as a nerve tonic. It is added to some kinds of toothpaste.


Tetraethylammonium hydroxide

Tetraethylammonium hydroxide is the organic compound with the formula (C2H5)4NOH, abbreviated Et4NOH. It is the tetraethylammonium salt of hydroxide. It is used and dispensed as solutions in water or alcohols, which are colorless and highly alkaline. The compound is a common reagent in organic synthesis. It is also employed in the preparation of zeolites.
Structure.
Tetraethylammonium hydroxide is most commonly encountered as an aqueous solution. Several hydrates NEt4OH·xH2O. have been crystallized, including x = 4, 5, 9. These salts feature well separated Et4N+ cations and hydroxide anions. The hydroxide groups are linked by hydrogen bonds to the water of crystallization. Anhydrous Tetraethylammonium hydroxide has not been isolated.
Synthesis and reactions.
It is prepared from tetraethylammonium bromide by salt metathesis, using a hydroxide-loaded ion exchange column or by the action of silver oxide. Attempted isolation of Et4NOH induces Hofmann elimination, leading to triethylamine and ethylene.
Treatment of Et4NOH with a wide range of acids gives water and the other tetraethylammonium salts:


Organochromium chemistry

Organochromium chemistry is a branch of organometallic chemistry that deals with organic compounds containing a chromium to carbon bond and their reactions. The field is of some relevance to organic synthesis. The relevant oxidation states for organochromium complexes encompass the entire range of possible oxidation states from –4 (d10) in Na4[Cr–IV(CO)4] to +6 (d0) in oxo-alkyl complexes like Cp*CrVI(=O)2Me.
History.
The first organochromium compound was described in 1919 by Franz Hein. He treated phenylmagnesium bromide with chromium(III) chloride to give a new product (after hydrolysis) which he incorrectly identified as pentaphenyl chromium bromide (Ph5CrBr). Years later, in 1957 H.H. Zeiss et al. repeated Hein's experiments and correctly arrived at a cationic bisarene chromium sandwich compound (Ar2Cr+). Bis(benzene)chromium itself was discovered around the same time in 1956 by Ernst Otto Fischer by reaction of chromium(III) chloride, benzene and aluminum chloride. The related compound chromocene was discovered a few years earlier in 1953 also by Fischer.

In another development, Anet and Leblanc also in 1957 prepared a benzyl chromium solution from benzyl bromide and chromium(II) perchlorate. This reaction involves one-electron oxidative addition of the carbon-bromine bond, a process which was shown by Kochi to be a case of double single electron transfer, first to give the benzyl free radical and then to the benzyl anion.
G. Wilke et al. introduced tris-(η-allyl)chromium in 1963 as an early Ziegler–Natta catalyst (but not successful in the long run) Chromocene compounds were first employed in ethylene polymerization in 1972 by Union Carbide and continue to be used today in the industrial production of high-density polyethylene.
The organochromium compound (phenylmethoxycarbene)pentacarbonylchromium, Ph(OCH3)C=Cr(CO)5 was the first carbene complex to be crystallographically characterized by Fischer in 1967 (now called a Fischer carbene). The first ever carbyne, this one also containing chromium, made its debut in 1973.
The first example of a proposed metal-metal quintuple bond is found in a compound of the type [CrAr]2, where Ar is a bulky aryl ligand.
Applications in organic synthesis.
Although organochromium chemistry is heavily employed in industrial catalysis, relatively few reagents have been developed for applications in organic synthesis. Two are the Nozaki-Hiyama-Kishi reaction (1977) (transmetallation with organonickel intermediate) and the Takai olefination (1986)(oxidation of Cr(II) to Cr(III) while replacing halogens). In a niche exploit, certain tricarbonyl(arene)chromium complexes display benzylic activation.
Organochromium compounds.
Organochromium compounds can be divided into these broad compound classes:
Ethylene polymerization and oligomerization.
Chromium catalysts are important in ethylene polymerization. The Phillips catalyst is prepared by impregnating chromium(VI) oxide on silica followed activation in dry air at high temperatures. The bright yellow catalyst becomes reduced by the ethylene to afford a probable Cr(II) species that is catalytically active. A related catalytic systems developed by Union Carbide and DSM are also based on silica with chromocene and other chromium complexes. How these catalysts work is unclear. One model system describes it as coordination polymerization:
With two THF ligands the catalyst is stable but in dichloromethane one ligand is lost to form a 13 electron chromium intermediate. This enables side-on addition of an ethylene unit and a polymer chain can grow by migratory insertion.
Chromium compounds also catalyse the trimerization of ethylene to produce the monomer 1-hexene.


Zinc stearate

Zinc stearate is a "zinc soap" that is widely used industrially. In this context, soap is used in its formal sense, a metal salt of a fatty acid: in this case stearic acid. It is a white solid that repels water. It is insoluble in polar solvents such as alcohol and ether but soluble in aromatic hydrocarbons (e.g., benzene) and chlorinated hydrocarbons when heated. It is the most powerful mold release agent among all metal soaps. It contains no electrolyte and has a hydrophobic effect. Its main application areas are the plastics and rubber industry, where it is used as a releasing agent and lubricant which can be easily incorporated.
Zinc carboxylates, e.g. basic zinc acetate, adopt complex formulas, and are not simply dicarboxylates of zinc. Instead the formula for most zinc carboxylates is Zn4O(O2CR)6, consisting of a Zn4O6+ core with carboxylate ligands spanning the edges.
Applications.
It is widely used as a release agent for the production of many kinds of objects: rubber, polyurethane, polyester processing system, powder metallurgy. These applications exploit its "non-stick" properties. In cosmetics, zinc stearate is a lubricant and thickening agent used to improve texture.
It is an "activator" for accelerated rubber sulfur vulcanization. As discovered in the early days of vulcanization, zinc has a beneficial effect on the reaction of the sulfur with the polyolefin. The stearate is a form of zinc that is highly soluble in the nonpolar medium of the polyolefins.
Being lipophilic, it functions as a phase transfer catalyst for the saponification of fats.
Niche uses.
It is a component of some paints, imparting gloss. As a chief ingredient in "fanning powder", it is used by magicians performing card manipulation to decrease the friction between playing cards.


Uranium disilicide

Uranium disilicide is an inorganic chemical compound of uranium in oxidation state +4. It is a silicide of uranium.
There has been recent interest in using uranium disilicide as an alternative to uranium dioxide for
fuel in nuclear reactors. Advantages are higher
percentage of uranium and higher thermal conductivity. A direct replacement of UO2 with U3Si2
should enable a reactor to generate more energy from a set of fuel rods and also provide more "coping time" in the case of a LOCA (Loss of Cooling Accident).
The development of uranium disilicide, uranium nitride, or other high thermal conductivity uranium compound may be critical for the performance of "Accident Tolerant Fuel", a
development effort mandated by the US Department of Energy. This is due to zircalloy having a higher thermal
conductivity than all replacement materials being developed. In particular, SIC-SiC CMC (link), which has several superior material properties to zircalloy for this application, has about five times lower thermal conductivity (varies due to the manufacturing methods used for the fiber and for the matrix) than zircalloy.(refs on SiC-SiC and zircalloy). The lower thermal conductivity means that a reactor using fuel rods with SiC-SiC CMC cladding and conventional UO2 fuel will have to either: 1) Run at a lower power output to keep the fuel the same temperature, or 2) Run with the same power, with the fuel hotter, which means the reactor has less coping time (time to fix what is wrong before something fails). The alternative, enabled by U3Si2 which has about five times better thermal conductivity than UO2 , is expected to be a fuel rod capable of equal power output, slightly better energy output, and longer coping time.


Hemimorphite

Hemimorphite, is Zn4(Si2O7)(OH)2·H2O, a component of calamine. It is a sorosilicate mineral which has been historically mined from the upper parts of zinc and lead ores, chiefly associated with smithsonite, ZnCO3. They were assumed to be the same mineral and both were classed under the same name of calamine. In the second half of the 18th century it was discovered that these two different minerals were both present in calamine. They closely resemble each other.
The silicate was the rarer of the two, and was named hemimorphite, because of the hemimorph development of its crystals. This unusual form, which is typical of only a few minerals, means that the crystals are terminated by dissimilar faces. Hemimorphite most commonly forms crystalline crusts and layers, also massive, granular, rounded and reniform aggregates, concentrically striated, or finely needle-shaped, fibrous or stalactitic, and rarely fan-shaped clusters of crystals.
Some specimens show strong green fluorescence in shortwave ultraviolet light (253.7 nm) and weak light pink fluorescence in longwave UV.
Occurrence.
Hemimorphite most frequently occurs as the product of the oxidation of the upper parts of sphalerite bearing ore bodies, accompanied by other secondary minerals which form the so-called "iron cap" or "gossan". Hemimorphite is an important ore of zinc and contains up to 54.2% of the metal, together with silicon, oxygen and hydrogen. The crystals are blunt at one end and sharp at the other.
The regions on the Belgian-German border are well known for their deposits of hemimorphite of metasomatic origin, especially Vieille Montagne in Belgium and Aachen in Germany. Other deposits are in Tarnowskie Góry area in Upper Silesia, Poland; near Phoenixville, Pennsylvania; the Missouri lead-zinc district; Elkhorn, Montana; Leadville, Colorado; and Organ Mountains, New Mexico in the United States; and in several localities in North Africa. Further hemimorphite occurrences are the Padaeng deposit near Mae Sod in western Thailand; Sardinia; Nerchinsk, Siberia; Cave del Predil, Italy; Bleiberg, Carinthia, Austria; Matlock, Derbyshire, England.


Matthiessen and Hegeler Zinc Company

At one time, the family-owned Mattheissen and Hegeler (M&H) Zinc Company was the largest zinc manufacturing plant in the United States. The company brought zinc ore from Wisconsin and Missouri to the coal fields of Northern Illinois. The company and its founders had a large influence in the development of LaSalle & Peru, Illinois.
History.
Frederick Matthiessen and Edward Hegeler met while studying at the School of Mines in Freiberg, Saxony. They identified that the United States did not have a strong zinc industry and felt it was an opportunity. They were also interested in iron, coal, and zinc exporting. The pair sailed from Germany to Boston in 1856 and studied several sites in Pennsylvania, Missouri, and Wisconsin.
Zinc ore had been discovered at Mineral Point, Wisconsin and they selected LaSalle, Illinois as the location for their plant due to its coal fields, since it required about two tons of coal to smelt one ton of zinc ore. LaSalle was also connected to the zinc mining region of Wisconsin by the Illinois Central Railroad. LaSalle's mayor, Alexander Campbell, encouraged the enterprise and helped them get a deal with the railroad. The plant started in 1858 and occupied a broad flat site above the Little Vermilion River. The Civil War interrupted production, but by 1863, they began supplying zinc for arms and cartridges. They started with zinc smelting and added a rolling mill in 1866.
The company incorporated in 1871. While not the only US Zinc producer, M&H commanded a large part of the early market. Between 1881 and 1910, M&H solidified their position as the largest spelter producer in the nation and perhaps the world. In 1880, the Census of Manufacturing recorded 570 employees. Competition increased during this period, especially in LaSalle-Peru, with the arrival of Illinois Zinc Company in Peru, the Mineral Point Zinc Company in DePue, and the Collinsville Zinc Company near St. Louis. Lanyon and Company and Thomas Kinsman also ran small smelters in LaSalle. The zinc industry in LaSalle-Peru created related businesses such as sheetmetal, zinc dipping, and weather stripping companies.
Natural gas discovered in Kansas near sources of zinc ore in 1896 made Kansas and Oklahoma attractive to the zinc industry and signaled the beginning of the end for M&H. They held out through diligent work and technical expertise. In 1913 they built a zinc smelter in Joplin, Missouri. By the 1950s they had subsidiary interests throughout eastern North America.
Edward's daughter, Mary Hegeler Carus, ran the company from 1903 until her death in 1936. She was one of the first female students at the University of Michigan, and in 1882 the first woman engineering graduate. In 1885 -1886 she studied at the Bergakademie Freiberg in Germany, the first woman to be legally registered as a student there. She was ahead of her time in running a company. In 1924, the Hegeler-Carus family bought out the Matthiessen family interests in the company. Around the early 1900s, Julius and Herman Hegeler left to start their own zinc works in Danville, Illinois.
Coal mining stopped in 1936. In 1950, the company purchased Meadowbrook Corporation, a zinc smelter near Clarksburg, West Virginia, and New Castle Chemical Company near Pittsburgh from Dow Chemicals. They gave M&H problems in the 1960s due to strikes and inefficient production. The company stopped smelting zinc in 1961. Sulfuric acid discontinued in 1968 and only rolling operations remained. Today the LaSalle factory's location is a superfund site.
Hegeler furnace.
They improved upon the Belgian and Silesian furnaces and patented the new Hegleler furnace in 1881. It decreased coal usage and increased ore efficiency, eliminated some expensive labor, and created a marketable residual product, sulfuric acid. The furnace gave them an advantage over other zinc producers.
Employee relations.
M&H had an good relationship with their employees well into the 1920s. The company was noted for its safety records and employee involvement. It had low wages and reduced the daily work day to eight hours in 1885, nearly a decade before it became law. It gave loans to employees for the purpose of buying homes and was structured in a way to give its workers a significant voice. From the 1880s to 1910s when mines and factories throughout America experienced labor unrest, M&H did not have a single strike. After Matthiessen's death in 1918, contact with employees decreased. A union formed and the company had its first strike in 1936.
Zinc uses.
During WWI, zinc was used for ammunition, weapons, sheet metal, and boiler plates. After WWI, zinc was used for auto parts, brass, linoleum, batteries, semiconductors, pharmaceuticals, and rubber tires.
Subsidiaries.
M&H had a number of subsidiaries over the years:
Family impact.
Frederick Matthiessen became mayor of LaSalle and focused on civic improvements, helping bring about the LaSalle-Peru Horse and Dummy Railway and personally paying for the city waterworks, electrical light station, first library, high school buildings, and hygienic institute. He eventually donated his Deer Park estate as a public park, today Matthiessen State Park. He had interests in the LaSalle Pressed Brick Company, the LaSalle Machine and Tool Company, and re-organized a bankrupt Clock Company to become the Western Clock Manufacturing Company, also known as Westclox. His children donated funds for the Matthiessen Memorial Auditorium at the highschool in 1918.
Edward Hegeler interest in philosophy led him to establish the Open Court Press, which was instrumental in popularizing philosophy within the United States, particularly in translating Eastern religious works. He also had an interest in the DeSteiger Glass Works in LaSalle. He completed his home in 1876, now the Hegeler-Carus Mansion, a National Historic Landmark. He also donated four parks to the City of LaSalle.
Across the street from the Hegeler-Carus Mansion is the Julius W. Hegeler Mansion.
Edward Hegeler's grandson, Edward Carus founded the Carus Chemical Company in 1916, independent of M&H. It still exists today.


Autunite

Autunite (hydrated calcium uranyl phosphate), with formula Ca(UO2)2(PO4)2·10–12H2O, is a yellow-greenish fluorescent phosphate mineral with a hardness of 2–. Autunite crystallizes in the orthorhombic system and often occurs as tabular square crystals, commonly in small crusts or in fan-like masses. Due to the moderate uranium content of 48.27% it is radioactive and also used as uranium ore. Autunite fluoresces bright green to lime green under UV light. The mineral is also called calco-uranite, but this name is rarely used and effectively outdated.
Autunite was discovered in 1852 near Autun, France, which is also autunite's namesake. It occurs as an oxidation product of uranium minerals in granite pegmatites and hydrothermal deposits. Associate minerals include metaautunite, torbernite, phosphuranylite, saleeite, uranophane and sabugalite.
Etymology.
Autunite was named after the town of Autun, France, where the mineral was initially found. The mineral was named by Henry J. Brooke and William H. Miller in 1854. The archaic name "calco-uranite" pairs with the similarly composed mineral, torbernite, which was named "cupro-uranite". Both minerals are named due to their uranium content, with autunite's name referring to its calcium ("calco"), and torbernite's copper ("cupro") quantity.
Locations and mining.
Mount Kit Carson, Washington.
Researchers found autunite inside the Daybreak Mine on Mount Kit Carson, Spokane, Washington (or sometimes referred to as "near Mount Spokane"), in "vugs, fractures, and shear zones in granitic rock". These areas showed signs of another phosphate, apatite, which may have helped lead to the formation of autunite, by providing a source of phosphate and lime. The formation may have occurred with the interaction of uranium leached from a separate deposit.
90,000 lbs of U3O8 were produced from nine properties, although most of the ore came from the Daybreak Mine.
Other locations.
One of the other locations of autunite includes Autun, France, the type locality and namesake of the mineral. The mineral was formed there as an alteration of uraninite and other uranium bearing minerals. Autunite is also found in Cornwall, Saxony, and North and South Dakota.
Meta-autunite.
If the mineral dries out, it can lose its water content and convert to meta-autunite-I, which can turn into meta-autunite-II after heating. These two subsequent minerals are very rare in nature. For scientific studies it is recommended to store the mineral in a sealed container to minimize the water loss. Museums are known to have covered the mineral with lacquer to avoid drying of the mineral.


Silver arsenate

Silver arsenate is an inorganic compound with the formula Ag3AsO4. It has been used in qualitative analysis to distinguish between phosphate (Ag3PO4 is yellow) and arsenate(V) solutions.


Total gaseous mercury

Total gaseous mercury is a measurement of mercury concentration used in toxicology.


Titanyl

In inorganic chemistry, titanyl refers to the functional group TiIVO, sometimes written TiO2+. The term titanyl is used loosely to describe many titanium(IV) oxide compounds and complexes. For example, titanyl sulfate and potassium titanyl phosphate contain TiIVO centers with the connectivity O-Ti-O-Ti. In heterogeneous catalysis, titanyl refers to a terminal oxo ligand on a surface titanium(IV) center. There are a few molecular titanyl complexes where the oxo ligand is terminal, not bridging. In these cases the titanyl group is described as having a triple bond, i.e., Ti≡O.


Sodium metatitanate

Sodium metatitanate is a chemical compound with the chemical formula Na2Ti3O7. It is one of the very few sodium salts that is insoluble in water.


Organogallium chemistry

Organogallium chemistry is the chemistry of organometallic compounds containing a carbon to gallium (Ga) chemical bond. Despite their high toxicity , organogallium compounds have some use in organic synthesis. The compound trimethylgallium is of some relevance to MOCVD as a precursor to gallium arsenide via its reaction with arsine at 700 °C:
Gallium trichloride is an important reagent for the introduction of gallium into organic compounds.
The main gallium oxidation state is Ga(III), as in all lower group 13 elements (such as aluminium).
Organogallium(III) chemistry.
Compounds of the type R3Ga are monomeric. Lewis acidity decreases in the order Al > Ga > In and as a result organogallium compounds do not form bridged dimers as organoaluminum compounds do. Organogallium compounds are also less reactive than organoaluminum compounds. They do form stable peroxides.
Organogallium compounds can be synthesized by transmetallation, for example the reaction of gallium metal with dimethylmercury:
or via organolithium compounds or Grignards:
The electron-deficient nature of gallium can be removed by complex formation, for example
Pi complex formation with alkynes is also known.
Organogallium compounds are reagents or intermediates in several classes of organic reactions:
Higher group 13 organometallic chemistry.
Organoindium chemistry and organothallium (Tl) chemistry parallel that of organogallium in many regards. Indium and thallium in oxidation state +1 are more common, for example the metallocenes cyclopentadienylindium(I) and cyclopentadienylthallium. Trimethylindium is important in the semiconductor industry. A special thallium feature is electrophilic thallation of arene compounds, reminiscent of mercuration (the group 12 neighbor). A common reagent for this purpose is thallium(III) trifluoroacetate. The intermediate arylthallium bisfluoroacetate can be isolated and converted to an aryl halide, aryl cyanide, aryl thiol or nitroarene. An example is the iodination of para-xylene.
A specific niche indium research topic is indium mediated allylation.


Antigorite

Antigorite is a lamellated, monoclinic mineral in the phylosilicate serpentine subgroup with the ideal chemical formula of (Mg,Fe2+)3Si2O5(OH)4. It is the high-pressure polymorph of serpentine and is commonly found in metamorphosed serpentinites. Antigorite, and its serpentine polymorphs, play an important role in subduction zone dynamics due to their relative weakness and high weight percent of water (up to 13 weight % H2O). It is named after its type locality, the Geisspfad serpentinite, Valle Antigorio in the border region of Italy/Switzerland and is commonly used as a gemstone in jewelry and carvings.
Geologic occurrences.
Antigorite is found in low-temperature, high-pressure (or high-deformation) environments, including both extensional and compressional tectonic regimes. Serpentines are commonly found in the ultramafic greenschist facies of subduction zones, and are visible on the Earth's surface through secondary exhumation. Serpentinites that contain antigorite are usually highly deformed and show distinct textures, indicative of the dynamic region where they were formed. Antigorite serpentinites commonly have associated minerals of magnetite, chlorite, and carbonates.
Physical properties.
Lamellated antigorite occurs in tough, pleated masses. It is usually dark green in color, but may also be yellowish, gray, brown or black. It has a Mohs scale hardness of 3.5–4 and its lustre is vitreous to greasy. Antigorite has a specific gravity of 2.5-2.6. The monoclinic crystals show micaceous cleavage, a distinguished property of phylosilicates, and fuse with difficulty. Serpentinite rocks that consist of mostly antigorite are commonly mylonites. The antigorite grains that make up these rocks are very fine (on the order of 1 to 10 microns) and are fibrous, which defines a texture in the rock caused by lattice preferred orientation.
Gemstone properties.
Antigorite is used as gemstones or for carvings when it appears pure and translucent, although many crystals have black specks of magnetite suspended within. The gem types of antigorite are Bowenite and Williamsite. Bowenite, known for George T. Bowen from Rhode Island (the variety's type locality), who first analyzed the mineral, is translucent and light to dark green, often mottled with cloudy white patches and darker veining. It is the serpentine most frequently encountered in carving and jewelry, and is the state mineral of Rhode Island, United States. A bowenite cabochon featured as part of the "Our Mineral Heritage Brooch", was presented to U.S. First Lady Mrs. Lady Bird Johnson in 1967. Williamsite is very translucent and has a medium to deep apple-green color. Somewhat resembling jade, Williamsite is often cut into cabochons and beads.
Crystal structure.
The magnesian serpentines (antigorite, lizardite, chrysotile) are trioctahedral hydrous phyllosilictes. Their structure is based on 1:1 octahedral-tetrahedral layer structures. Antigorite is monoclinic in the space group Pm. Although the magensian serpentines have similar compositions, they have significantly different crystallographic structures, which are dependent on how the SiO4 tetrahedra sheets fit in with the octahedral sheets. Antigorite's basic composition has a smaller ratio of octahedral to tetrahedral cations (relative to lizardite and chrysotile), allowing the structure to compensate for the misfit of sheets through periodic flipping of the curved tetrahedra layers, and subsequently their polarity. Polysomes of antigorite are defined by the number of individual tetrahedra (denoted as the value "m") which span a wavelength of the direction of curvature. The sheets of tetrahedra allow the platy, fibrous crystals to separate parallel to the 001 (basal) plane, giving antigorite its perfect cleavage.


Tropodithietic acid

Tropodithietic acid is a tropolone derivative produced by the marine bacteria "Phaeobacter piscinae", "Phaeobacter inhibens" and "Phaeobacter gallaeciensis". Its structure is composed by a dithiete moiety fused to tropone-2-carboxylic acid.


Hydrogen energy vision and technology roadmap

The Hydrogen energy vision and technology roadmap is the roadmap of China initiated by the Ministry of Science and Technology, it makes hydrogen and fuel cell technologies important thematic priorities of the S&T development plan.


Adamite

Adamite is a zinc arsenate hydroxide mineral, Zn2AsO4OH. It is a mineral that typically occurs in the oxidized or weathered zone above zinc ore occurrences. Pure adamite is colorless, but usually it possess yellow color due to Fe compounds admixture. Tints of green also occur and are connected with copper substitutions in the mineral structure. Olivenite is a copper arsenate that is isostructural with adamite and there is considerable substitution between zinc and copper resulting in an intermediate called "cuproadamite". Zincolivenite is a recently discovered mineral being an intermediate mineral with formula CuZn(AsO4)(OH). Manganese, cobalt, and nickel also substitute in the structure. An analogous zinc phosphate, "tarbuttite", is known.
Occurrence.
Adamite occurs as a secondary mineral in the oxidized zone of zinc- and arsenic-bearing hydrothermal mineral deposits. It occurs in association with smithsonite, hemimorphite, scorodite, olivenite, calcite, quartz and iron and manganese oxides.
The yellow to bright lime-green colored crystals and druze along with its distinctive fluorescence make adamite a favorite among mineral collectors. Found in Mapimí, Durango, Mexico; Greece; and California and Utah in the United States.
Adamite was named after the French mineralogist Gilbert-Joseph Adam (1795–1881). It was first described in 1866 for an occurrence at the type locality of Chañarcillo, Copiapó Province, Atacama Region, Chile.


Substrate-level phosphorylation

Substrate-level phosphorylation is a metabolism reaction that results in the production of ATP or GTP by the transfer of a phosphate group from a substrate directly to ADP or GDP. Transferring from a higher energy (whether phosphate group attached or not) into a lower energy product. This process uses some of the released chemical energy, the Gibbs free energy, to transfer a phosphoryl (PO3) group to ADP or GDP from another phosphorylated compound. Occurs in glycolysis and in the citric acid cycle.
Unlike oxidative phosphorylation, oxidation and phosphorylation are not coupled in the process of substrate-level phosphorylation, and reactive intermediates are most often gained in the course of oxidation processes in catabolism. Most ATP is generated by oxidative phosphorylation in aerobic or anaerobic respiration while substrate-level phosphorylation provides a quicker, less efficient source of ATP, independent of external electron acceptors. This is the case in human erythrocytes, which have no mitochondria, and in oxygen-depleted muscle.
Overview.
Adenosine triphosphate is a major "energy currency" of the cell. The high energy bonds between the phosphate groups can be broken to power a variety of reactions used in all aspects of cell function.
Substrate-level phosphorylation occurs in the cytoplasm of cells during glycolysis and in mitochondria either during the Krebs cycle or by MTHFD1L (EC 6.3.4.3), an enzyme interconverting ADP + phosphate + 10-formyltetrahydrofolate to ATP + formate + tetrahydrofolate (reversibly), under both aerobic and anaerobic conditions. In the pay-off phase of glycolysis, a net of 2 ATP are produced by substrate-level phosphorylation.
Glycolysis.
The first substrate-level phosphorylation occurs after the conversion of 3-phosphoglyceraldehyde and Pi and NAD+ to 1,3-bisphosphoglycerate via glyceraldehyde 3-phosphate dehydrogenase. 1,3-bisphosphoglycerate is then dephosphorylated via phosphoglycerate kinase, producing 3-phosphoglycerate and ATP through a substrate-level phosphorylation.
The second substrate-level phosphorylation occurs by dephosphorylating phosphoenolpyruvate, catalyzed by pyruvate kinase, producing pyruvate and ATP.
During the preparatory phase, each 6-carbon glucose molecule is broken into two 3-carbon molecules. Thus, in glycolysis dephosphorylation results in the production of 4 ATP. However, the prior preparatory phase consumes 2 ATP, so the net yield in glycolysis is 2 ATP. 2 molecules of NADH are also produced and can be used in oxidative phosphorylation to generate more ATP.
Mitochondria.
ATP can be generated by substrate-level phosphorylation in mitochondria in a pathway that is independent from the proton motive force. In the matrix there are three reactions capable of substrate-level phosphorylation, utilizing either phosphoenolpyruvate carboxykinase or succinate-CoA ligase, or monofunctional C1-tetrahydrofolate synthase.
Phosphoenolpyruvate carboxykinase.
Mitochondrial phosphoenolpyruvate carboxykinase is thought to participate in the transfer of the phosphorylation potential from the matrix to the cytosol and vice versa. However, it is strongly favored towards GTP hydrolysis, thus it is not really considered as an important source of intra-mitochondrial substrate-level phosphorylation.
Succinate-CoA ligase.
Succinate-CoA ligase is a heterodimer composed of an invariant α-subunit and a substrate-specific ß-subunit, encoded by either SUCLA2 or SUCLG2. This combination results in either an ADP-forming succinate-CoA ligase (A-SUCL, EC 6.2.1.5) or a GDP-forming succinate-CoA ligase (G-SUCL, EC 6.2.1.4). The ADP-forming succinate-CoA ligase is potentially the only matrix enzyme generating ATP in the absence of a proton motive force, capable of maintaining matrix ATP levels under energy-limited conditions, such as transient hypoxia.
Monofunctional C1-tetrahydrofolate synthase.
This enzyme is encoded by MTHFD1L and reversibly interconverts ADP + phosphate + 10-formyltetrahydrofolate to ATP + formate + tetrahydrofolate.
Other mechanisms.
In working skeletal muscles and the brain, Phosphocreatine is stored as a readily available high-energy phosphate supply, and the enzyme creatine phosphokinase transfers a phosphate from phosphocreatine to ADP to produce ATP. Then the ATP releases giving chemical energy. This is sometimes erroneously considered to be substrate-level phosphorylation, although it is a transphosphorylation.
Importance of substrate-level phosphorylation in anoxia.
During anoxia, provision of ATP by substrate-level phosphorylation in the matrix is important not only as a mere means of energy, but also to prevent mitochondria from straining glycolytic ATP reserves by maintaining the adenine nucleotide translocator in ‘forward mode’ carrying ATP towards the cytosol.
Oxidative phosphorylation.
An alternative method used to create ATP is through oxidative phosphorylation, which takes place during cellular respiration. This process utilizes the oxidation of NADH to NAD+, yielding 3 ATP, and of FADH2 to FAD, yielding 2 ATP. The potential energy stored as an electrochemical gradient of protons (H+) across the inner mitochondrial membrane is required to generate ATP from ADP and Pi (inorganic phosphate molecule), a key difference from substrate-level phosphorylation. This gradient is exploited by ATP synthase acting as a pore, allowing H+ from the mitochondrial intermembrane space to move down its electrochemical gradient into the matrix and coupling the release of free energy to ATP synthesis. Conversely, electron transfer provides the energy required to actively pump H+ out of the matrix.


Zinc-activated ion channel

Zinc-activated ion channel (ZAC), is a human protein encoded by the gene. ZAC forms a cation-permeable ligand-gated ion channel of the "Cys-loop" superfamily. The ZAC gene is present in humans and dogs, but no ortholog is thought to exist in the rat or mouse genomes.
ZAC mRNA is expressed in prostate, thyroid, trachea, lung, brain (adult and fetal), spinal cord, skeletal muscle, heart, placenta, pancreas, liver, kidney and stomach. The endogenous ligand for ZAC is thought to be Zn2+, although ZAC has also been found to activate spontaneously. The function of spontaneous ZAC activation is unknown.


Greifensteinite

Greifensteinite is beryllium phosphate mineral with formula: Ca2Fe2+5Be4(PO4)6(OH)4·6H2O. It is the Fe2+ dominant member of the roscherite group. It crystallizes in the monoclinic crystal system and typically forms prismatic dark olive green crystals.
It was first described in Germany at Greifenstein Rocks, Ehrenfriedersdorf, and was named for the location. At the type locality, it occurs within a lithium-rich pegmatite in miarolitic cavities. It was approved by the International Mineralogical Association in 2002.


Untripentium


Lead fluoride

Lead fluoride may refer to:


Canavesite

Canavesite, Mg2(HBO3)(CO3)∙5H2O, is a rare carboborate mineral from the abandoned Brosso mine in Italy. Canavesite is a secondary mineral that occurs due to the weathering of ludwigite-magnetite skarn on the surface of mine walls. The physical properties consist of milky-white rosette-like aggregates of elongated transparent fibers shown in figure one. It has the crystal symmetry of a monoclinic with a diffraction symbol of "2/mP-/-".
History.
Giussani & Vighi saw and described boron mineralization in 1964 in a pyrite mine, Brosso, in Italy. They noticed secondary minerals forming between szaibelyite and ludwigite at the point where they come into contact with each other inside magnetite skarns. Years later in 1972 a few members from the Gruppo Mineralogico Lombardo found the first sample of a new unknown mineral around the same location. Since then canavesite has been traded as nesquehonite and added to geologic collections. Canavestie is one of few known carboborates. The name canavesite originates from the district of Canavese where the pyrite mine is located.
Composition.
Canavesite has the chemical formula Mg2(HBO3)(CO3)∙5H2O, which makes it a carboborate mineral. Having two complete chemical analyses and one partial gives MgO (31.06),B2O3(12.70),CO2(18.57)H2O(37.44) and MgO (31.18),B2O3(13.47),CO2(17.02)H2O(38.33) in weight % oxides. Taking eleven oxygen ions, one can calculate from this data the empirical formula is as followed Mg2.03(HBO3)1.09(CO3)∙4.91H2O. A canavestie infrared spectrum is unlike that of well known families. It has bands near 3500, 1450, and 1000 cm-1 looks almost like OH−. With a dilute HCl at or near room temperature canavestie will produce tiny bubbles of gas and foam with hissing sounds. The molecular formula has not been obtained for canavesite. The approximate density was obtained using the heavy liquid method, which shows a density of 1.8 g cm−3.
Occurrence.
The Vola Gera tunnel in the Brosso mine is located in a village northwest of Ivrea, Piedmont, Italy in Canavese district. The mine was initially a pyrite mine but it was abandoned. It is a well known mine for fine minerals that build up between the zone where monzonite stock and marbles interbedded in the mica schists of the Sesia-Lanzo zone. Mineralization of boron within magnetite skarns were first reported by Giussani and Vighi in 1964. There has since been more canavesite discovered in New Jersey at the Sterling Mine. They are the only two places that have been known to produce the secondary mineral.
Physical properties.
Canavesite occurs as a milky rosette-like aggregates of very thin [010] elongated fibers on surfaces and along fractures of ludwigite and magnetite skarns. Canavesite is an exceptionally rare secondary mineral that forms on the walls of the abandoned mine of Brosso, Italy and has formed since the mine was abandoned due to weathering of these ludwigite and magnetite skarns. Canvesite has a monoclinic P 2/m crystal structure and does not fluoresce under an ultraviolent light. It also has thin and slightly flexible vitreous fibers that are elongated parallel to [010]. Some crystal may show pseudohexagonal prismatic habit, but these cases are rare. The X-ray powder diffraction pattern strongest lines are 9.54(100)202, 8.12(40)(201), 7.80(18)(102,301), 4.56(21)(501,503), and 3.11(19)(407).


Sinhalite

Sinhalite is a borate mineral with formula MgAl(BO4).
Sinhalite was first found in Sri Lanka (Ceylon) in 1952, and was named from Sinhala - the Sanskrit name for Sri Lanka.
Gemstone quality Sinhalite can also be found in Madagascar, Tanzania and Myanmar (Burma). The most commonly seen color of Sinhalite are white to gray, grayish-blue or a shade of brown ranging from pale to dark. Pale pink and brownish pink crystals can be found in Tanzania.


Lithium superoxide

Lithium superoxide (LiO2) is an inorganic compound which has only been isolated in matrix isolation experiments at 15-40 K. It is an unstable free radical that has been analyzed using infrared (IR), Raman, electronic, electron spin resonance, soft X-ray spectroscopies, and a variety of theoretical methods.
Structure.
Experimental studies indicate that the LiO2 molecule contains highly ionic bonds. Eighteen different values were attained using six isotopic species. This indicated that the force constant between the two oxygen atoms corresponds with the constant found for the O2− ion. Studies indicate that there is little to no covalent character in the LiO2 molecule.
The bond length for the O-O bond was determined to be 1.34 Å. Using a simple crystal structure optimization, the Li-O bond was calculated to be approximately 2.10 Å.
Lithium superoxide is extremely reactive because of the odd electron present in the π* molecular orbital.
There have been quite a few studies regarding the clusters formed by LiO2 molecules. The most common dimer has been found to be the cage isomer. Second to it is the singlet bypyramidal structure. Studies have also been done on the chair complex and the planar ring, but these two are less favorable, though not necessarily impossible.
Reactions.
In a lithium-air battery, when there is a one electron reduction during discharge, lithium superoxide is formed as seen in the following reaction:
This product will then react and proceed to form lithium peroxide, Li2O2:
The mechanism for this last reaction has not been confirmed and chemists are having difficulties developing a theory of what may be happening. Another significant challenge of these batteries is finding an ideal solvent in which to perform these reactions; ether- and amide-based solvents are currently used, but these compounds readily react with oxygen and decompose. A suitable solvent would need to be able to resist autoxidation to enable a long life cycle for the battery.
Presence of the Compound.
The predominant use of lithium superoxide is in rechargeable lithium batteries. As portrayed in the reactions above, this lithium compound is a major component as an intermediate, an area where there is much research to be done. Researchers have much anticipation for the potential energy that may be provided by these batteries—-some say it is comparable to the internal combustion engine.
One study claims that alkali superoxides affect the function of the alkyl metals in the atmosphere as well. The alkali metals are found predominantly in the mesosphere and the superoxides are found just below this where the metal reacts with the excess oxygen. Rarely are superoxides stable for any significant amount of time as they exist merely as transition states. In January 2016, researchers from Argonne National Laboratory, US, claimed that crystalline LiO2 could be stabilized in a Li-O2 battery by using a suitable graphene-based cathode.


Manganese arsenide

Manganese arsenide (MnAs) is an intermetallic compound, an arsenide of manganese. It forms ferromagnetic crystals with hexagonal (NiAs-type) crystal structure, which convert to the paramagnetic orthorhombic β-phase upon heating to . MnAs has potential applications in spintronics, for electrical spin injection into GaAs and Si based devices.


Babefphite

Babefphite is a rare phosphate mineral with the general formula BaBe(PO4)(F,OH). The name is given for its composition (Ba meaning barium, Be meaning beryllium, F meaning fluorine, and P for phosphorus).
Crystallography.
Babefphite is tetragonal, which means crystallographically, it contains two horizontal and interchangeable equal axes (a1 and 2) and a longitudinal axis (c). It belongs to the 4/m 2/m 2/m point group; meaning it has four horizontal axes of twofold symmetry, two of which coincide with the crystallographic axes (a1) and (a2) and the remaining two at 45° to the (a1) and (a2).
Optical properties.
Babefphite is an anisotropic mineral meaning that the velocity of light passing through it will vary depending on the direction of travel. When measured against Canada Balsam it will display moderate positive relief. Relief is an optical property that relates the index of refraction of a mineral versus the index of refraction of another material. Positive relief denotes light being bent toward the mineral. Negative relief means the light is bent toward the mounting material. In the case of optical mineralogy, this other material is most often whatever material is used to mount the mineral in question in thin section. Based on the white color of babefphite in hand sample it is likely that it appears clear in thin section under plane polarized light. Due to its anisotropic nature, it has the possibility of displaying very weak pleochroism under the same conditions. Pleochroism is defined as a change of color, on rotation, of plane light. It occurs when the extraordinary and ordinary rays of light are absorbed differently, thereby displaying different colors. The wider the difference in the rays the more dramatic the color change. Babefphite is also weakly birefringent. Birefringence is an optical property that is determined by the direction that light passes through a mineral.
Occurrence.
Babefphite was first described in 1966 for an occurrence at the Aunik fluorite-rare metals deposit, Buryatia, Eastern-Siberian Region, Russia. It has also been reported from the Rožná pegmatite in the Vysočina Region, Moravia, Czech Republic.
In the Siberian occurrence it occurs in residual material above rare-earth metal bearing skarns associated with alkaline intrusive bodies. It occurs with zircon, ilmenorutile, fluorite, phenakite, scheelite, bertrandite, albite, microcline and quartz.


Feynmanium


Pewee Valley Confederate Cemetery

Pewee Valley Confederate Cemetery is one mile from the site of the old Kentucky Confederate Home. The cemetery is not only on the National Register of Historic Places, but an individual monument within it, the Confederate Memorial in Pewee Valley, is separately on it as part of the Civil War Monuments of Kentucky MPS. It is the only cemetery for Confederate veterans, 313 in total, that is an official state burying ground in Kentucky.
History.
In May 1871, several local citizens of Pewee Valley, Kentucky, northeast of Louisville, Kentucky, decided they wanted a local public cemetery. By August 1871 the land was purchased. The following spring, in 1872, the cemetery was organized and plants were placed to beautify the area. When they were finished, they asked for a charter from the Kentucky state government.
Close to the cemetery, the Villa Ridge Inn was built in 1889. The site was a four-story summer resort was to be popular with Louisville entrepreneurs and their families. However, despite being a $90,000 construction with "majestic architecture", it never attracted enough visitors to stay in business. Between the Inn's closure and 1902, it had temporarily become a private high school, the successor to the Kentucky College for Young Ladies.
When the approval for the Confederate home was gained in 1904, the cemetery was divided into three parts; one for Confederate veterans, one for whites, and one for blacks.
In 1902 the Kentucky state assembly unanimously approved the building of a veterans home specifically for Confederate veterans of the American Civil War by the cemetery. It was done at the behest of former Confederate officer Bennett H. Young, who had long desired such a facility as he saw many former Confederate veterans were no longer able to take care of themselves. After Young's group acquired $16,000, the legislation was approved. Young's group of Confederate veterans and the Daughters of the Confederacy chose the Villa Ridge site just outside the cemetery, as it was well-sited and very inexpensive compared to what it had cost to build in 1889. The cemetery was established shortly after the hospital was opened, with the special monument built soon afterwards. The total area of the home and cemetery was 11,275 square feet.
In the years it was an active veterans facility, it provided a hospital, nursing care, food, entertainment, and religious services for up to 350 veterans at a time, providing a home for 700 former Confederate soldiers in its years of operations. The requirements to be a resident of the home was to not only be a former Confederate soldier, but to have been a resident of Kentucky for the past six months, be mentally stable, and to have no problems with alcoholism. Many of the veterans once served under John Hunt Morgan. A fire on March 25, 1920 destroyed the main building, as well as an infirmary ward and laundry. No lives were lost, and the rest of the facility was enough to house those residents still using the home. Eventually, the number of veterans who could be served dwindled, and by 1934 the hospital was no more, with the remaining five residents being transferred to the nearby Pewee Valley Sanitorium.
Present day.
All that remains of the hospital is its main gate, which was moved to become the entrance arch for the cemetery one mile away, and part of the walkway from the house to the railroad. A sign is placed along the pathway to mark it. The source of fresh water for the facility, a reservoir, was filled in during the 1990s.
The monument is unique for Kentucky Civil War monuments in that it is built of zinc, whereas most are made of marble or limestone. Another oddity is that the obelisk and base are separated by an inscribed Gothic altar that acts as an arch on the face of the monument. Confederate flags are also crafted upon the monument.


Calcium pangamate

Calcium pangamate is a mineral supplement, and is sometimes used as a synonym for "pangamic acid."


Magnesium pidolate

Magnesium pidolate, the magnesium salt of pidolic acid (pyroglutamic acid), is a mineral supplement, which contains 8.664% magnesium w/w.


Iodine monobromide

Iodine monobromide is an interhalogen compound with the chemical symbol IBr. It is a dark red solid that melts near room temperature. Like iodine monochloride, IBr is used in some types of iodometry. It serves as a source of I+. Its Lewis acid properties are compared with those of ICl and I2 in the ECW model. It can form CT adducts with Lewis donors.
Synthesis.
Iodine monobromide is formed when iodine and bromine are combined in a chemical reaction:


Deuterated DMSO

Deuterated DMSO, also known as dimethyl sulfoxide-d6, is an isotopologue of dimethyl sulfoxide (DMSO, (CH3)2S=O)) with chemical formula ((CD3)2S=O) in which the hydrogen atoms ("H") are replaced with their isotope deuterium ("D"). Deuterated DMSO is a common solvent used in NMR spectroscopy.
Production.
Deuterated DMSO is produced by heating DMSO in heavy water (D2O) with a basic catalyst such as calcium oxide. The reaction does not give complete conversion to the d6 product, and the water produced must be removed and replaced with D2O several times to drive the equilibrium to the fully deuterated product.
Use in NMR spectroscopy.
Pure deuterated DMSO shows no peaks in 1H NMR spectroscopy and as a result is commonly used as an NMR solvent. However commercially available samples are not 100% pure and a residual DMSO-d5 1H NMR signal is observed at 2.50ppm (quintet, JHD=1.9Hz). The 13C chemical shift of DMSO-d6 is 39.52ppm (septet).


C28 fullerene


Ionic radius

Ionic radius, "r"ion, is the radius of a monatomic ion in an ionic crystal structure. Although neither atoms nor ions have sharp boundaries, they are sometimes treated as if they were hard spheres with radii such that the sum of ionic radii of the cation and anion gives the distance between the ions in a crystal lattice. Ionic radii are typically given in units of either picometers (pm) or angstroms (Å), with 1 Å = 100 pm. Typical values range from 31 pm (0.3 Å) to over 200 pm (2 Å).
The concept can be extended to solvated ions in liquid solutions taking into consideration the solvation shell.
Trends.
Ions may be larger or smaller than the neutral atom, depending on the ion's electric charge. When an atom loses an electron to form a cation, the other electrons are more attracted to the nucleus, and the radius of the ion gets smaller. Similarly, when an electron is added to an atom, forming an anion, the added electron increases the size of the electron cloud by interelectronic repulsion.
The ionic radius is not a fixed property of a given ion, but varies with coordination number, spin state and other parameters. Nevertheless, ionic radius values are sufficiently transferable to allow periodic trends to be recognized. As with other types of atomic radius, ionic radii increase on descending a group. Ionic size (for the same ion) also increases with increasing coordination number, and an ion in a high-spin state will be larger than the same ion in a low-spin state. In general, ionic radius decreases with increasing positive charge and increases with increasing negative charge.
An "anomalous" ionic radius in a crystal is often a sign of significant covalent character in the bonding. No bond is "completely" ionic, and some supposedly "ionic" compounds, especially of the transition metals, are particularly covalent in character. This is illustrated by the unit cell parameters for sodium and silver halides in the table. On the basis of the fluorides, one would say that Ag+ is larger than Na+, but on the basis of the chlorides and bromides the opposite appears to be true. This is because the greater covalent character of the bonds in AgCl and AgBr reduces the bond length and hence the apparent ionic radius of Ag+, an effect which is not present in the halides of the more electropositive sodium, nor in silver fluoride in which the fluoride ion is relatively unpolarizable.
Determination.
The distance between two ions in an ionic crystal can be determined by X-ray crystallography, which gives the lengths of the sides of the unit cell of a crystal. For example, the length of each edge of the unit cell of sodium chloride is found to be 564.02 pm. Each edge of the unit cell of sodium chloride may be considered to have the atoms arranged as Na+∙∙∙Cl−∙∙∙Na+, so the edge is twice the Na-Cl separation. Therefore, the distance between the Na+ and Cl− ions is half of 564.02 pm, which is 282.01 pm. However, although X-ray crystallography gives the distance between ions, it doesn't indicate where the boundary is between those ions, so it doesn't directly give ionic radii.
Landé estimated ionic radii by considering crystals in which the anion and cation have a large difference in size, such as LiI. The lithium ions are so much smaller than the iodide ions that the lithium fits into holes within the crystal lattice, allowing the iodide ions to touch. That is, the distance between two neighboring iodides in the crystal is assumed to be twice the radius of the iodide ion, which was deduced to be 214 pm. This value can be used to determine other radii. For example, the inter-ionic distance in RbI is 356 pm, giving 142 pm for the ionic radius of Rb+. In this way values for the radii of 8 ions were determined.
Wasastjerna estimated ionic radii by considering the relative volumes of ions as determined from electrical polarizability as determined by measurements of refractive index. These results were extended by Victor Goldschmidt. Both Wasastjerna and Goldschmidt used a value of 132 pm for the O2− ion.
Pauling used effective nuclear charge to proportion the distance between ions into anionic and a cationic radii. His data gives the O2− ion a radius of 140 pm.
A major review of crystallographic data led to the publication of revised ionic radii by Shannon. Shannon gives different radii for different coordination numbers, and for high and low spin states of the ions. To be consistent with Pauling's radii, Shannon has used a value of "r"ion(O2−) = 140 pm; data using that value are referred to as "effective" ionic radii. However, Shannon also includes data based on "r"ion(O2−) = 126 pm; data using that value are referred to as "crystal" ionic radii. Shannon states that "it is felt that crystal radii correspond more closely to the physical size of ions in a solid." The two sets of data are listed in the two tables below.
Soft-sphere model.
For many compounds, the model of ions as hard spheres does not reproduce the distance between ions, formula_1, to the accuracy with which it can be measured in crystals. One approach to improving the calculated accuracy is to model ions as "soft spheres" that overlap in the crystal. Because the ions overlap, their separation in the crystal will be less than the sum of their soft-sphere radii.
The relation between soft-sphere ionic radii, formula_2 and formula_3, and formula_1, is given by
formula_5,
where formula_6 is an exponent that varies with the type of crystal structure. In the hard-sphere model, formula_6 would be 1, giving formula_8.
In the soft-sphere model, formula_6 has a value between 1 and 2.
For example, for crystals of group 1 halides with the sodium chloride structure, a value of 1.6667 gives good agreement with experiment.
Some soft-sphere ionic radii are in the table.
These radii are larger than the crystal radii given above (Li+, 90 pm; Cl−, 167 pm). Inter-ionic separations calculated with these radii give remarkably good agreement with experimental values. Some data are given in the table. Curiously, no theoretical justification for the equation containing formula_6 has been given.
Non-spherical ions.
The concept of ionic radii is based on the assumption of a spherical ion shape. However, from a group-theoretical point of view the assumption is only justified for ions that reside on high-symmetry crystal lattice sites like Na and Cl in halite or Zn and S in sphalerite. A clear distinction can be made, when the point symmetry group of the respective lattice site is considered, which are the cubic groups "O""h" and "T""d" in NaCl and ZnS. For ions on lower-symmetry sites significant deviations of their electron density from a spherical shape may occur. This holds in particular for ions on lattice sites of polar symmetry, which are the crystallographic point groups "C"1, "C"1"h", "C""n" or "C""nv", "n" = 2, 3, 4 or 6. A thorough analysis of the bonding geometry was recently carried out for pyrite-type compounds, where monovalent chalcogen ions reside on "C"3 lattice sites. It was found that chalcogen ions have to be modeled by ellipsoidal charge distributions with different radii along the symmetry axis and perpendicular to it.


Pyrope

The mineral pyrope is a member of the garnet group. Pyrope is the only member of the garnet family to always display red colouration in natural samples, and it is from this characteristic that it gets its name: from the Greek for "fire" and "eye". Despite being less common than most garnets, it is a widely used gemstone with numerous alternative names, some of which are misnomers. "Chrome pyrope", and "Bohemian garnet" are two alternative names, the usage of the latter being discouraged by the Gemological Institute of America. Misnomers include "Colorado ruby", "Arizona ruby", "California ruby", "Rocky Mountain ruby", "Elie Ruby", "Bohemian carbuncle", and "Cape ruby".
The composition of pure pyrope is Mg3Al2(SiO4)3, although typically other elements are present in at least minor proportions—these other elements include Ca, Cr, Fe and Mn. Pyrope forms a solid solution series with almandine and spessartine, which are collectively known as the "pyralspite" garnets (pyrope, almandine, spessartine). Iron and manganese substitute for the magnesium in the pyrope structure. The resultant, mixed composition garnets are defined according to their pyrope-almandine ratio. The semi-precious stone rhodolite is a garnet of ~70% pyrope composition.
The origin of most pyrope is in ultramafic rocks, typically peridotite from the Earth's mantle: these mantle-derived peridotites can be attributed both to igneous and metamorphic processes. Pyrope also occurs in ultrahigh-pressure (UHP) metamorphic rocks, as in the Dora-Maira massif in the western Alps. In that massif, nearly pure pyrope occurs in crystals to almost 12 cm in diameter; some of that pyrope has inclusions of coesite, and some has inclusions of enstatite and sapphirine.
Pyrope is common in peridotite xenoliths from kimberlite pipes, some of which are diamond-bearing. Pyrope found in association with diamond commonly has a Cr2O3 content of 3–8%, which imparts a distinctive violet to deep purple colouration (often with a greenish tinge) and because of this is often used as a kimberlite indicator mineral in areas where erosive activity makes pin pointing the origin of the pipe difficult. These varieties are known as "chrome-pyrope", or "G9"/"G10" garnets.
Mineral identification.
In hand specimens, pyrope is very tricky to distinguish from almandine; however, it is likely to display fewer flaws and inclusions. Other distinguishing criteria are listed in the adjacent table. Care should be taken when using these properties as many of those listed have been determined from synthetically grown, pure-composition pyrope. Others, such as pyrope's high specific gravity, may be of little use when studying a small crystal embedded in a matrix of other silicate minerals. In these cases, mineral association with other mafic and ultramafic minerals may be the best indication that the garnet you are studying is pyrope.
In petrographic thin section, the most distinguishing features of pyrope are those shared with the other common garnets: high relief and isotropy. Garnets tend to be less strongly coloured than other silicate minerals in thin section, although pyrope may show a pale pinkish-purple hue in plane-polarized light. The lack of cleavage, commonly euhedral crystal morphology, and mineral associations should also be used in identification of pyrope under the microscope.


Teflic acid

Teflic acid is the chemical compound with the formula HOTeF5. This strong acid is related to orthotelluric acid, Te(OH)6. Teflic acid has a slightly distorted octahedral geometry.
Preparation.
Teflic acid was accidentally discovered by Engelbrecht and Sladky. Their synthesis did not yield the anticipated telluryl fluoride TeO2F2, but a mixture of volatile telluric compounds, "i.a." HOTeF5:
Teflic acid can also be prepared from fluorosulfonic acid and barium tellurate:
It is also the first hydrolysis product of tellurium hexafluoride:
Teflates.
The conjugate base of teflic acid is called the teflate anion, F5TeO− (not to be confused with triflate). Many teflates are known, examples being B(OTeF5)3 and the acid anhydride O(TeF5)2. Pyrolysis of the boron compound gives the dimer (TeF4O)2
The teflate anion is known to resist oxidation. This property has allowed the preparation several highly unusual species such as the hexateflates M(OTeF5)6− (M = As, Sb, Bi). Xenon forms the cation Xe(OTeF5)+.


Titanium ethoxide

Titanium ethoxide is a chemical compound with the formula Ti4(OCH2CH3)16. It is a colorless liquid that is soluble in organic solvents but hydrolyzes readily. It is sold commercially as a colorless solution. Alkoxides of titanium(IV) and zirconium(IV) are used in organic synthesis and materials science. They adopt more complex structures than suggested by their empirical formulas.
Syntheses.
Titanium ethoxide is prepared by treating titanium tetrachloride with ethanol in the presence of an amine:
The purity of titanium ethoxide is commonly assayed by proton NMR spectroscopy. Ti(OEt)4 1H NMR (90 MHz, chloroform-d, ppm): 4.36 (quartet, 8H, CH2), 1.27 (triplet, 12H, CH3).
Structure.
Both Ti(OEt)4 exist mainly as tetramers with an octahedral coordination environment around the metal centers. There are two types of titanium centers, depending on the number of terminal vs bridging alkoxide ligands. Zr(OEt)4 is structurally similar. The virtual symmetry of the M4O16 core structure for the tetramer structures of these compounds is C2h.
Related compounds.
Titanium methoxide.
Like the ethoxide, titanium methoxide Ti(OMe)4 exists as a tetramer with each of the TiIV metal centers having an octahedral coordination environment.
Titanium isopropoxide.
With bulky alky groups, Ti(O"i"Pr)4 in contrast exist as a monomer with a tetrahedral environment around the Ti center. This lower degree of coordination to the metal center is attributed to the steric bulk of the "i"Pr groups versus the "n"-alkyl groups, this serves to prevent bridging interactions between the metal centers.
Zirconium ethoxide.
Zirconium ethoxide can be prepared in a manner similar but not identical to the titanium compound:
A more common synthesis for zirconium ethoxide is to treat zirconium tetrachloride with the desired alcohol and ammonia:
Zirconium ethoxide can also be prepared with zirconocene dichloride:
Zirconium propoxide.
Zr(O"n"Pr)4 also adopts the titanium ethoxide structure.
Reactions.
Both Ti and Zr alkoxides can be used to deposit microstructured films of TiO2 or ZrO2:
These films form via a hydrolysis of the alkoxide. The resulting oxiides are chemically resilient. The structure of the metal oxide films grown in this matter is affected by the presence of base or acid catalysts for the hydrolysis. Generally acid-catalysis yields a sol where the polymer chains are randomly oriented and linear. In the base-mediated case bushy clusters or crosslinked networks are produced, these structures can trap solvent and reaction byproducts and form a gel coating. TiIV and ZrIV alkoxides are also potential starting materials for Ziegler–Natta catalysts used in alkene polymerization. Intermediates in the hydrolysis have been crystallized. They feature interior oxides in addition to the ethoxide on the exterior of the clusters.


Activated charcoal


Unbitrium


Vanadyl perchlorate

Vanadyl perchlorate or vanadyl triperchlorate is a golden yellow coloured liquid or crystalline compound of vanadium, oxygen and perchlorate group. The substance consists of molecules covalently bound and is quite volatile.
Formation.
Vanadyl perchlorate can be made by reacting vanadium pentoxide with dichlorine heptoxide at 5 °C. It is purified by distillation under a vacuum and recrystallisation at 21 °C.
A solution of vanadium(V) perchlorate can be made by dissolving vanadium pentoxide in perchloric acid.
The reaction of vanadium pentoxide and dichlorine hexoxide could produce VO(ClO4)3:
Properties.
It can react with vanadium oxychloride to form another vanadyl perchlorate (VO2ClO4):
Related.
Other perchlorates include pervanadyl perchlorate, also known as dioxovanadium perchlorate, which contains VO2+ ions, vanadyl diperchlorate, oxovanadium perchlorate or vanadium(IV) perchlorate, and VO(ClO4)2, which dissolves in water. Vanadic perchlorate, also known as vanadium(III) perchlorate solution in water, is a green-tinged blue colour, significantly different to most other V(III) solutions, which are complexed.


Manganese(II) molybdate

Managnese(II) molybdate is an inorganic compound with the chemical formula MnMoO4. α-MnMoO4 has a monoclinic crystal structure. It is also antiferromagnetic at low temperatures.
Applications.
MnMoO4 has been used as active material in electrodes for aqueous supercapacitors due to fast pseudocapacitive redox reactions, and as catalyst for hydrogen evolution.


Clinohumite

Clinohumite is an uncommon member of the humite group, a magnesium silicate according to the chemical formula (Mg, Fe)9(SiO4)4(F,OH)2. The formula can be thought of as four olivine (Mg2SiO4), plus one brucite (Mg(OH)2). Indeed, the mineral is essentially a hydrated olivine and occurs in altered ultramafic rocks and carbonatites. Most commonly found as tiny indistinct grains, large euhedral clinohumite crystals are sought by collectors and occasionally fashioned into bright, yellow-orange gemstones. Only two sources of gem-quality material are known: the Pamir Mountains of Tajikistan, and the Taymyr region of northern Siberia. It is one of two humite group minerals that have been cut into gems, the other being the much more common chondrodite.
Properties.
A monoclinic mineral, clinohumite is typically a dark to light brownish or orangy yellow, somewhat resembling the hessonite variety of grossular. Clinohumite's crystal habit is usually granular, but may also be prismatic; crystals are almost always small. Simple and multiple crystal twinning (on {001}) is common, resulting in a highly variable habit. Clinohumite is brittle with a hardness of 6 and a poor basal cleavage. Its specific gravity is 3.2–3.4, and its fracture is conchoidal to uneven; its streak is white.
Clinohumite's transparency ranges from transparent to translucent; its luster ranges from a dull vitreous to resinous. Its refractive index (as measured via sodium light, 589.3 nm) is as follows: α 1.631; β 1.638–1.647; γ 1.668;, with a maximum birefringence of 0.028 (biaxial positive). Under shortwave ultraviolet light, some clinohumite may fluoresce an orangy yellow; there is little to no response under longwave UV.
The Taymyr material is reported to be a dark reddish brown while the Pamir material is a bright yellow to orange or brownish orange. The Pamir material also has a hardness slightly greater than 6, a lower specific gravity (3.18), and higher maximum birefringence (0.036). Phillip Youngman, master faceter of Los Osos, California, noticed not only that Pamir material is harder than expected, but also that it is less brittle than expected. Youngman observed that clinohumite reacted like beryl to cutting and polishing, and that it reminded him of polishing diopside.
Like other members of the humite group, the relative amounts of hydroxyl and fluorine vary in clinohumite, and iron commonly substitutes for some of the magnesium, bringing about changes in physical and optical properties. Titanium substitution also causes pronounced changes in optical properties, producing the variety "titanclinohumite". Consequently, it is relatively easy to determine that a stone is a humite group mineral, but difficult to determine exactly which member. Other common impurities of clinohumite include aluminium, manganese, and calcium.
Formation and occurrence.
Clinohumite is a product of contact metamorphism and is commonly found as indistinct grains embedded in limestone. Its type occurrence is within the limestone ejecta of the Mount Vesuvius volcano complex near Naples, Italy, where clinohumite was discovered in 1876. The aforementioned gem-quality occurrences of Pamir and Taymyr were discovered only recently: the former in the early 1980s, and the latter in 2000. These deposits are scarce and only sporadically mined, so clinohumite remains one of the rarest gemstones with only a few thousand carats known to exist in private collections.
Other (non-gem quality) occurrences of clinohumite include: the Sør Rondane and Balchen Mountains of Antarctica; Mount Bischoff, Waratah, Tasmania; the Saualpe Mountains of Carinthia, the Koralpe mountains of Styria, and the Vals, Virgen, and Ziller valleys of the Tyrol, Austria; the Jacupiranga mine of Cajati, São Paulo State, Southeast Region, Brazil; the Pirin Mountains of Bulgaria; Bancroft, Ontario, Notre Dame du Laus, Wakefield, and Villedieu Township, Quebec, Canada; Southern and Western Finland; Bavaria and Saxony, Germany; eastern Greenland; Ambasamudram in Tamil Nadu, India; Honshū, Japan; Suan, North Korea; Nordland, Norway; KwaZulu-Natal and Northern Cape Province, South Africa; Andalusia, Spain; Värmland and Västmanland, Sweden; Isle of Skye, Scotland; and the states of California, Colorado, Massachusetts, New Jersey, New Mexico, New York, Oklahoma, Utah, and Washington, US.
Clinohumite also occurs as a minor component of some masses of peridotite from the Earth's mantle emplaced into the Earth's crust and as a very rare component of peridotite xenoliths. These occurrences and implications have been summarized by Luth (2003) in a discussion of the possible importance of the mineral as a significant reservoir of water in the Earth's mantle. Titanium is a minor constituent of clinohumite in most such occurrences. Clinohumite is stable throughout the upper mantle to depths of at least and is a potential host phase for H (water) in this region of the Earth's interior.
Minerals associated with humite include grossular, wollastonite, forsterite, monticellite, cuspidine, fluoborite, ludwigite, dolomite, calcite, talc, biotite, spinel, vesuvianite, sanidine, meionite and nepheline.
Crystal structure.
The structure is monoclinic with space group "P"21/"b" ("a"-unique). The unit cell has "a" = 4.7488 Å; "b" = 10.2875 Å; "c" = 13.6967 Å; and alpha = 100.63°; V = 667.65 Å3; Z = 2 for pure Mg hydroxyl-clinohumite. The odd setting of space group "P"21/"c" is chosen to preserve the "a" and "b" axes of olivine. The structure is closely related to that of olivine as well as the other humite minerals. Mg and Fe are in octahedral coordination with oxygen and silicon (Si) is in tetrahedral coordination. There are five distinct octahedral sites and two different tetrahedral sites. One of the octahedral sites is bonded to two OH,F atoms and is the site where Ti is partitioned. Clinohumite is a nesosilicate with no oxygen atoms shared between two silicons.


Sulfoxylic acid

Sulfoxylic acid (H2SO2) (also known as hyposulfurous acid or sulfur dihydroxide) is an unstable oxoacid of sulfur in an intermediate oxidation state between hydrogen sulfide and dithionous acid. It consists of two hydroxy groups attached to a sulfur atom. Sulfoxylic acid contains sulfur in an oxidation state of +2. Sulfur monoxide (SO) can be considered as a theoretical anhydride for sulfoxylic acid, but it is not actually known to react with water.
The complementary base is the sulfoxylate anion which is much more stable. In between these states is the ion, also somewhat stable. Sulfoxylate ions can be made by decomposing thiourea dioxide in an alkaline solution. To do this, thiourea dioxide first forms a tautomer, aminoiminomethanesulfinic acid (H2NC(=NH)SO2H, abbreviated AIMS) which breaks apart. Sulfoxylate reacts with formaldehyde to yield a hydroxymethanesulfinate called rongalite:
Formation.
Sulfoxylic acid has been detected in the gas phase. It is likely to be formed as an intermediate when hydrogen sulfide is oxidised by living organisms, or in the atmosphere, or anywhere else in the natural environment. It may also exist in circumstellar disks. When H2S is oxidised it starts from oxidation state −2, and should then pass through intermediate values of 0 and +2 before getting to well known sulfite at +4 and sulfate at +6. When sulfide in alkaline conditions is oxidised by air in the presence of nickel ions, sulfoxylate concentration first increases to around 5% and then decreases over several days. Polysulfide concentration also grows and then shrinks on a slower timescale reaching about 25% of the sulfide. The sulfur ends up forming thiosulfate.
Sulfoxylic acid has been made by ultraviolet irradiation of a mixture of solid H2S and H2O, followed by warming. This is a possible natural process in comets or circumstellar disks.
Fender "et al." claimed to make "sulfinic acid" (an isomer of sulfoxylic acid} by ultraviolet irradiation on solid sulfur dioxide and hydrogen sulfide in a solid argon matrix, measuring the infrared vibrational spectrum. However the assignment of the lines in the spectrum is doubtful, so this may not be the substance produced.
Sulfoxylic acid can be made in the gas phase in an electric discharge through a neon, H2, SO2 mixture. This also yields some sulfhydryl hydroperoxide.
Properties.
Sulfoxylic acid is an isomer of sulfinic acid, which has a hydrogen atom bonded to the sulfur, and the oxygen connected with a double bond (HS(O)OH). Other isomers are thiadioxyrane (a ring of two oxygen atoms and a sulfur), dihydrogen sulfone (a sulfur atom linked to two hydrogen and two oxygen atoms), sulfhydryl hydroperoxide (HSOOH), and dihydrogen persulfoxide H2SOO. Sulfoxylic acid has the lowest energy of any of these isomers.
The p"K"a1 of sulfoxylic acid is 7.97. The p"K"a2 of bisulfoxylate () is 13.55.
Calculations of the molecule suggest there may be two alignments termed "C"2, and "C"s. The H−O distance is 96.22 (or 96.16) pm, S−O distance is 163.64 (or 163.67) pm, ∠HOS = 108.14° (108.59°), ∠OSO = 103.28° (103.64°) HOSO twist is 84.34° (+90.56 and −90.56) ("C"s dimensions in parentheses).
The microwave spectrum has absorption lines at 10.419265, 12.2441259, 14.0223698, 16.3161779 GHz and many others for the "C"s and 12.8910254, 19.4509030, 21.4709035, 24.7588445, 29.5065050, 29.5848250, 32.8772050 GHz for the "C"2 form.
The sulfoxylate ion apparently has an X-ray absorption near edge structure at 2476.1 eV. With sulfur the X-ray absorption edge changes with oxidation state as per Kunzl's law. The edge corresponds to the energy needed to excite and inner 1S electron to a 3P orbital. Sulfoxylate has an infrared absorption peak at 918.2 cm−1.
Reactions.
Sulfoxylic acid disproportionates into sulfur and hydrogensulfite . Some of this in turn reacts to form thiosulfate .
Sulfoxylates are sensitive to air, and will be oxidised by the oxygen in it.
Sulfoxylate is oxidised to sulfur dioxide radical anion and then to sulfur dioxide.
The known sulfoxylate salts include cobalt sulfoxylate CoSO2·3H2O. This can dissolve in an ammoniacal solution. However cobalt sulfide will precipitate if sulfide is formed during a reaction.
Sulfoxylate in solution reacts with thiosulfate to form sulfides and sulfites.
Sulfoxylate reduces nitrite to hydronitrite radical dianion . This in turn reacts with water forming hydroxide ions and nitric oxide (NO). Nitric oxide and nitrous oxide N2O in turn are further reduced by sulfoxylate.
When sulfoxylate reacts with hypochlorite, bromine or chlorine dioxide it forms hydrogen sulfite and sulfates.
Dithionite is unstable in a pH 4 solution, decomposing to sulfoxylic acid and hydrogen sulfite. This sulfoxylic acid reacts with more dithionite to yield more hydrogen sulfite, and some kind of sulfur, and a small amount of thiosulfate.
By reducing sulfur dioxide, hydrogen sulfoxylate forms as an intermediate, and this is much more reactive. Hydrogen sulfoxylate reacts with organic compounds with a double bond (vinyls) to make an organic sulfinate. Hydrogen sulfoxylate reacts with divinyl sulfone to make 1,4-dithiane 1,1,4,4-tetroxide. Perfluorophenyl iodide is reduced to pentafluorobenzene.
The reaction of sulfoxylic acid with sulfite yields trithionate () and with thiosulfate yields pentathionate ().
Salts.
Salts of sulfoxylic acid that have been claimed to have been made include cobalt, thallium and zinc. Cobalt sulfoxylate is made from sodium hyposulfite, cobalt chloride and ammonia. Zinc sulfoxylate is produced by reacting zinc metal with sulfuryl chloride. Thallous sulfoxylate is made by allowing oxygen onto thallous sulfide. This is olive brown in colour. When heated to 250 °C it recrystallizes to another yellow form. Cuprous sufoxylate Cu2SO2 can be made as a solid or liquid by heating cuprous sulfide and copper sulfate. Cu2SO2 melts at and is stable as a liquid phase to over 680K. There are also solid phase transitions at and .
Derivatives.
Organic derivatives of sulfoxylic acid exist, including 1-hydroxysulfanyloxy-4-methylbenzene, and S(OCH3)2 and S(OCH2CH3)2. Naming for substances with the −OSOH group, can use suffixes -oxysulfanol (preferred), -hydrogen sulfoxylate, or -oxysulfenic acid; or prefix hydroxysulfanyloxy- (preferred) or sulfenooxy-. The ionic group -OSO− can use the preferred suffix -oxysulfanolate or -sulfoxylate; or preferred prefix oxysulfanolato- or sulfenatooxy-. The R-OSO-R' can be suffixed with -dioxysulfane or -sulfoxylate; or prefixed with oxysulfanyloxy- or sulfenooxy-.
A reaction with sulfur dichloride with saturated primary or secondary alcohols yields a diarylsulfoxylate ROSOR compound. With 1,2 diols, SCl2 forms polymeric sulfoxylates. This method also produces dialkoxydisulfes and dialylsulfites. The yield of sulfoxylates is maximised by doing the reaction at low temperatures around -75°C, and diluting the reactants, reducing free chlorine and disulfur dichloride production. Diethylsulfoxylate can be made by reacting diethoxydisulfide with sodium ethoxide. Other sulfoxylate esters include propyl, isopropyl, "n"-butyl, "n"-pentyl and cholesterol. Unsaturated or aromatic alcohols do not react with SCl2 to form sulfoxylates. For 1,3-diols SCl2 reaction can result in a six member 1,3,2-dioxathiane ring or a twelve membered cyclic dimer sulfoxylate ring or even larger rings and polymers.
Sulfoxylic acid dimethyl ester (also called dimethylsulfoxylate or dimethoxysulfane) (S(OCH3)2) has been studied using electron diffraction, X-ray crystallography, Raman spectroscopy, and infrared spectroscopy. Studying this molecule is much easier than the unstable sulfoxylic acid. It is a liquid at standard conditions that boils at 74°C and freezes at -67° In the gas state the molecular bond angle ∠OSO is 103°, with the oxygen-sulfur distance 1.625 Å. The oxygen-carbon distance is 1.426 Å and carbon-hydrogen distance 1.105 Å with ∠COS 115.9° and ∠COS 109°. The molecule as "C"2 symmetry with one methyl group rotated above the plane determined by the OSO atoms, and the other below. The C−O−S−O dihedral angle is about 84°. The energy barrier to moving one methyl group to the other side of the plane is 37 kJ/mol and it would be in a 12 kJ/mol higher energy state in this "C"s symmetry state.


Peroxomonosulfate

The peroxomonosulfate ion, , is a sulfur oxoanion. It is sometimes referred to as the persulfate ion, but this term also refers to the peroxydisulfate ion, .
Its other IUPAC names are sulfuroperoxoate and trioxidoperoxidosulfate(2−).


Cheralite

Cheralite is an anhydrous phosphate mineral with the ideal chemical formula CaTh(PO4)2. It is isomorphous with huttonite and monazite. It can be regarded as the product of the complete cationic substitution in the system:
2 LREE3+ ↔ Ca2+ + Th4+.
It was previously known under the name brabantite.


Scandium oxide

Scandium(III) oxide or scandia is a inorganic compound with formula Sc2O3. It is one of several oxides of rare earth elements with a high melting point. It is used in the preparation of other scandium compounds as well as in high-temperature systems (for its resistance to heat and thermal shock), electronic ceramics, and glass composition (as a helper material).
Structure and physical properties.
Scandium(III) oxide adopts a cubic crystal structure (point group: tetrahedral (Th), space group: Ia) containing 6-coordinate metal centres. Powder diffraction analysis shows Sc−O bond distances of 2.159–2.071 Å.
Scandium oxide is an insulator with a band gap of 6.0 eV.
Production.
Scandium oxide is the primary form of refined scandium produced by the mining industry. Scandium-rich ores, such as thortveitite (Sc,Y)2(Si2O7) and kolbeckite ScPO4·2H2O are rare, however trace amounts of scandium are present in many other minerals. Scandium oxide is therefore predominantly produced as a by-product from the extraction of other elements.
Reactions.
Scandium oxide is the primary form of refined scandium produced by the mining industry, making it the start point for all scandium chemistry.
Scandium oxide reacts with most acids upon heating, to produce the expected hydrated product. For example, heating in excess aqueous HCl produces hydrated ScCl3·"n"H2O. This can be rendered anhydrous by evaporation to dryness in the presence of NH4Cl, with the mixture then being purified by removal of NH4Cl by sublimation at 300-500 °C. The presence of NH4Cl is required, as the hydrated ScCl3·"n"H2O would otherwise form a mixed oxychloride upon drying.
Likewise, it is converted into hydrated scandium(III) triflate (Sc(OTf)3·"n"H2O) by a reaction with triflic acid.
Metallic scandium is produced industrially by the reduction of scandium oxide; this proceeds via conversion to scandium fluoride followed by a reduction with metallic calcium. This process is in some ways similar to the Kroll process for the production of metallic titanium.
Scandium oxide forms scandate salts with alkalis, unlike its higher homologues yttrium oxide and lanthanum oxide (but like lutetium oxide), for example forming K3Sc(OH)6 with KOH. In this, scandium oxide shows more similarity with aluminium oxide.
Natural occurrence.
Natural scandia, although impure, occurs as mineral kangite.