NOx

In atmospheric chemistry, is a generic term for the nitrogen oxides that are most relevant for air pollution, namely nitric oxide (NO) and nitrogen dioxide (). These gases contribute to the formation of smog and acid rain, as well as affecting tropospheric ozone.
 gases are usually produced from the reaction among nitrogen and oxygen during combustion of fuels, such as hydrocarbons, in air; especially at high temperatures, such as in car engines. In areas of high motor vehicle traffic, such as in large cities, the nitrogen oxides emitted can be a significant source of air pollution. gases are also produced naturally by lightning.
The term is chemistry shorthand for molecules containing one nitrogen and one or more oxygen atom. It is generally meant to include nitrous oxide (N2O), although nitrous oxide is a fairly inert oxide of nitrogen that has many uses as an oxidizer for rockets and car engines, an anesthetic, and a propellant for aerosol sprays and whipped cream. Nitrous oxide plays hardly any role in air pollution, although it may have a significant impact on the ozone layer, and is a significant greenhouse gas.
' is defined as the sum of plus the ' compounds produced from the oxidation of which include nitric acid, nitrous acid (HONO), dinitrogen pentoxide (N2O5), peroxyacetyl nitrate (PAN), alkyl nitrates (RONO2), peroxyalkyl nitrates (ROONO2), the nitrate radical (NO3), and peroxynitric acid (HNO4).
Formation and reactions.
Because of energy limitations, oxygen and nitrogen do not react at ambient temperatures. But at high temperatures, they undergo an endothermic reaction producing various oxides of nitrogen. Such temperatures arise inside an internal combustion engine or a power station boiler, during the combustion of a mixture of air and fuel, and naturally in a lightning flash.
In atmospheric chemistry, the term denotes the total concentration of NO and since the conversion between these two species is rapid in the stratosphere and troposphere. During daylight hours, these concentrations together with that of ozone are in steady state, also known as photostationary state(PSS); the ratio of NO to is determined by the intensity of sunshine (which converts to NO) and the concentration of ozone (which reacts with NO to again form ).
In other words, the concentration of ozone in the atmosphere is determined by the ratio of these two species.
This relationship between and ozone is also known as the Leighton relationship.
The time τ that is needed to reach a steady state among and ozone is dominated by reaction (), which reverses reactions ()+():
for mixing ratio of NO, [NO] = 10 part per billion (ppb), the time constant is 40 minutes; for [NO] = 1 ppb, 4 minutes.
Formation of smog.
When and volatile organic compounds (VOCs) react in the presence of sunlight, they form photochemical smog, a significant form of air pollution. The presence of photochemical smog increases during the summer when the incident solar radiation is higher. The emitted hydrocarbons from industrial activities and transportation react with NOx quickly and increase the concentration of ozone and peroxide compounds, especially peroxyacetyl nitrate (PAN).
Children, people with lung diseases such as asthma, and people who work or exercise outside are particularly susceptible to adverse effects of smog such as damage to lung tissue and reduction in lung function.
Formation of nitric acid and acid rain.
NO2 is further oxidized in the gas phase during daytime by reaction with OH
where M denotes a third molecule required to stabilize the addition product. Nitric acid (HNO3) is highly soluble in liquid water in aerosol particles or cloud drops.
NO2 also reacts with ozone to form nitrate radical
During the daytime, NO3 is quickly photolyzed back to NO2, but at night it can react with a second NO2 to form dinitrogen pentoxide.
N2O5 reacts rapidly with liquid water (in aerosol particles or cloud drops, but not in the gas phase) to form HNO3,
These are thought to be the principal pathways for formation of nitric acid in the atmosphere. This nitric acid contributes to acid rain or may deposit to soil, where it makes nitrate, which is of use to growing plants. The aqueous phase reaction
is too slow to be of any significance in the atmosphere.
Sources.
Natural sources.
Nitric oxide is produced during thunderstorms due to the extreme heating and cooling within a lightning strike. This causes stable molecules such as N2 and O2 to convert into significant amounts of NO similar to the process that occurs during high temperature fuel combustion. NOx from lightning can become oxidized to produce nitric acid (HNO3), this can be precipitated out as acid rain or deposited onto particles in the air. Elevated production of NOx from lightning depends on the season and geographic location. The occurrence of lightning is more common over land near the equator in the inter-tropical convergence zone (ITCZ) during summer months. This area migrates slightly as seasons change. NOx production from lightning can be observed through satellite observations.
Scientists Ott et al. estimated that each flash of lightning on average in the several mid-latitude and subtropical thunderstorms studied turned of nitrogen into chemically reactive . With 1.4 billion lightning flashes per year, multiplied by 7 kilograms per lightning strike, they estimated the total amount of produced by lightning per year is 8.6 million tonnes. However, emissions resulting from fossil fuel combustion are estimated at 28.5 million tonnes.
A recent discovery indicated that cosmic ray and solar flares can significantly influence the number of lightning strikes occurring on Earth. Therefore, space weather can be a major driving force of lightning-produced atmospheric . Atmospheric constituents such as nitrogen oxides can be stratified vertically in the atmosphere. Ott noted that the lightning-produced is typically found at altitudes greater than 5 km, while combustion and biogenic (soil) are typically found near the sources at near surface elevation (where it can cause the most significant health effects).
Biogenic sources.
Agricultural fertilization and the use of nitrogen fixing plants also contribute to atmospheric , by promoting nitrogen fixation by microorganisms. The nitrification process transforms ammonia into nitrate. And the denitrification is basically the reverse process of nitrification. During the denitrification, nitrate is reduced to nitrite then NO then N2O and finally nitrogen. Through these processes, NOx is emitted to the atmosphere.
A recent study conducted by the University of California Davis, found that adding nitrogen fertilizer to soil in California is contributing 25 percent or more to state-wide NOx pollution levels. When nitrogen fertilizer is added to the soil, excess ammonium and nitrate not used by plants can be converted to NO by microorganism in the soil, which escapes into the air. NOx is a precursor for smog formation which is already a known issue for the state of California. In addition to contributing to smog, when nitrogen fertilizer is added to the soil and the excess is released in the form of NO, or leached as nitrate this can be a costly process for the farming industry.
A 2018 study by the Indiana University determined that forests in the eastern United States can expect to see increases in NOx, as a result to changes in the types of trees which predominate. Due to human activity and climate change, the maples, sassafras, and tulip poplar are pushing out the beneficial oak, beech, and hickory. The team determined that the first three tree species, maples, sassafras, and tulip poplar, are associated with ammonia-oxidizing bacteria known to "emit reactive nitrogen from soil." By contrast, the second three tree species, oak, beech and hickory, are associated with microbes that "absorb reactive nitrogen oxides," and thus can have a positive impact on the nitrogen oxide component of air quality. Nitrogen oxide release from forest soils is expected to be highest in Indiana, Illinois, Michigan, Kentucky and Ohio.
Industrial sources (anthropogenic sources).
The three primary sources of in combustion processes:
Thermal formation, which is highly temperature dependent, is recognized as the most relevant source when combusting natural gas. Fuel tends to dominate during the combustion of fuels, such as coal, which have a significant nitrogen content, particularly when burned in combustors designed to minimise thermal . The contribution of prompt is normally considered negligible. A fourth source, called "feed" is associated with the combustion of nitrogen present in the feed material of cement rotary kilns, at between 300 °C and 800 °C, where it is considered a minor contributor.
Thermal.
Thermal refers to formed through high temperature oxidation of the diatomic nitrogen found in combustion air. The formation rate is primarily a function of temperature and the residence time of nitrogen at that temperature. At high temperatures, usually above 1300 °C (2600 °F), molecular nitrogen (N2) and oxygen (O2) in the combustion air dissociate into their atomic states and participate in a series of reactions.
The three principal reactions (the extended Zel'dovich mechanism) producing thermal are:
All three reactions are reversible. Zeldovich was the first to suggest the importance of the first two reactions. The last reaction of atomic nitrogen with the hydroxyl radical, •HO, was added by Lavoie, Heywood and Keck to the mechanism and makes a significant contribution to the formation of thermal .
Fuel.
It is estimated that transportation fuels cause 54% of the anthropogenic (i.e. human-caused) . The major source of production from nitrogen-bearing fuels such as certain coals and oil, is the conversion of fuel bound nitrogen to during combustion. During combustion, the nitrogen bound in the fuel is released as a free radical and ultimately forms free N2, or NO. Fuel can contribute as much as 50% of total emissions through the combusting oil and as much as 80% through the combusting of coal.
Although the complete mechanism is not fully understood, there are two primary pathways of formation. The first involves the oxidation of volatile nitrogen species during the initial stages of combustion. During the release and before the oxidation of the volatiles, nitrogen reacts to form several intermediaries which are then oxidized into NO. If the volatiles evolve into a reducing atmosphere, the nitrogen evolved can readily be made to form nitrogen gas, rather than . The second pathway involves the combustion of nitrogen contained in the char matrix during the combustion of the char portion of the fuels. This reaction occurs much more slowly than the volatile phase. Only around 20% of the char nitrogen is ultimately emitted as , since much of the that forms during this process is reduced to nitrogen by the char, which is nearly pure carbon.
Prompt.
Nitrogen oxides are released during manufacturing of nitrogen fertilizers. Though nitrous oxide is emitted during its application, it is then reacted in atmosphere to form nitrogen oxides. This third source is attributed to the reaction of atmospheric nitrogen, N2, with radicals such as C, CH, and CH2 fragments derived from fuel, rather than thermal or fuel processes. Occurring in the earliest stage of combustion, this results in the formation of fixed species of nitrogen such as NH (nitrogen monohydride), NCN (diradical cyanonitrene), HCN (hydrogen cyanide), •H2CN (dihydrogen cyanide) and •CN (cyano radical) which can oxidize to NO. In fuels that contain nitrogen, the incidence of prompt is comparatively small and it is generally only of interest for the most exacting emission targets.
Health and environment effects.
There is strong evidence that respiratory exposure can trigger and exacerbate existing asthma symptoms, and may even lead to the development of asthma over longer periods of time. It has also been associated with heart disease, diabetes, birth outcomes, and all-cause mortality, but these nonrespiratory effects are less well-established.
 reacts with ammonia, moisture, and other compounds to form nitric acid vapor and related particles.
 reacts with volatile organic compounds in the presence of sunlight to form ozone. Ozone can cause adverse effects such as damage to lung tissue and reduction in lung function mostly in susceptible populations (children, elderly, asthmatics). Ozone can be transported by wind currents and cause health impacts far from the original sources. The American Lung Association estimates that nearly 50 percent of United States inhabitants live in counties that are not in ozone compliance. In South East England, ground level ozone pollution tends to be highest in the countryside and in suburbs, while in central London and on major roads NO emissions are able to "mop up" ozone to form and oxygen.
 also readily reacts with common organic chemicals, and even ozone, to form a wide variety of toxic products: nitroarenes, nitrosamines and also the nitrate radical some of which may cause DNA mutations. Recently another pathway, via , to ozone has been found that predominantly occurs in coastal areas via formation of nitryl chloride when comes into contact with salt mist.
The direct effect of the emission of has positive contribution to the greenhouse effect. Instead of reacting with ozone in Reaction 3, NO can also react with HO2· and organic peroxyradicals (RO2·) and thus increase the concentration of ozone. Once the concentration of exceeds a certain level, atmospheric reactions result in net ozone formation. Since tropospheric ozone can absorb infrared radiation, this indirect effect of is intensifying global warming.
There are also other indirect effects of that can either increase or decrease the greenhouse effect. First of all, through the reaction of NO with HO2 radicals, •OH radicals are recycled, which oxidize methane molecules, meaning emissions can counter the effect of greenhouse gases. For instance, ship traffic emits a great amount of NOx which provides a source of NOx over the ocean. Then, photolysis of NO2 leads to the formation of ozone and the further formation of hydroxyl radicals (·OH) through ozone photolysis. Since the major sink of methane in the atmosphere is by reaction with •OH radicals, the NOx emissions from ship travel may lead to a net global cooling. However, in the atmosphere may undergo dry or wet deposition and return to land in the form of HNO3/NO3−. Through this way, the deposition leads to nitrogen fertilization and the subsequent formation of nitrous oxide (N2O) in soil, which is another greenhouse gas. In conclusion, considering several direct and indirect effects, emissions have a negative contribution to global warming.
 in the atmosphere is removed through several pathways. During daytime, NO2 reacts with hydroxyl radicals (·OH) and forms nitric acid (HNO3), which can easily be removed by dry and wet deposition. Organic peroxyradicals (RO2·) can also react with NO and NO2 and result in the formation of organic nitrates. These are ultimately broken down to inorganic nitrate, which is a useful nutrient for plants. During nighttime, NO2 and NO can form nitrous acid (HONO) through surface-catalyzed reaction. Although the reaction is relatively slow, it is an important reaction in urban areas. In addition, the nitrate radical (NO3) is formed by the reaction between NO2 and ozone. At night, NO3 further reacts with NO2 and establishes an equilibrium reaction with dinitrogen pentoxide (N2O5). Via heterogeneous reaction, N2O5 reacts with water vapor or liquid water and forms nitric acid (HNO3). As mentioned above, nitric acid can be removed through wet and dry deposition and this results in the removal of from the atmosphere.
Biodiesel and.
Biodiesel and its blends in general are known to reduce harmful tailpipe emissions such as: carbon monoxide; particulate matter (PM), otherwise known as soot; and unburned hydrocarbon emissions. While earlier studies suggested biodiesel could sometimes decrease NOx and sometimes increase NOx emissions, subsequent investigation has shown that blends of up to 20% biodiesel in USEPA-approved diesel fuel have no significant impact on NOx emissions compared with regular diesel. The state of California uses a special formulation of diesel fuel to produce less NOx relative to diesel fuel used in the other 49 states. This has been deemed necessary by the California Air Resources Board (CARB) to offset the combination of vehicle congestion, warm temperatures, extensive sunlight, PM, and topography that all contribute to the formation of ozone and smog. CARB has established a special regulation for Alternative Diesel Fuels to ensure that any new fuels, including biodiesel, coming into the market do not substantially increase NOx emissions. The reduction of emissions is one of the most important challenges for advances in vehicle technology. While diesel vehicles sold in the US since 2010 are dramatically cleaner than previous diesel vehicles, urban areas continue to seek more ways to reduce the formation of smog and ozone. formation during combustion is associated with a number of factors such as combustion temperature. As such, it can be observed that the vehicle drive cycle, or the load on the engine have more significant impact on NOx emissions than the type of fuel used. This may be especially true for modern, clean diesel vehicles that continuously monitor engine operation electronically and actively control engine parameters and exhaust system operations to limit NOx emission to less than 0.2 g/km. Low-temperature combustion or LTC technology may help reduce thermal formation of during combustion, however a tradeoff exists as high temperature combustion produces less PM or soot and results in greater power and fuel efficiency.
Regulation and emission control technologies.
Selective catalytic reduction (SCR) and selective non-catalytic reduction (SNCR) reduce post combustion by reacting the exhaust with urea or ammonia to produce nitrogen and water. SCR is now being used in ships, diesel trucks and in some diesel cars. The use of exhaust gas recirculation and catalytic converters in motor vehicle engines have significantly reduced vehicular emissions. was the main focus of the Volkswagen emissions violations.
Other technologies such as flameless oxidation (FLOX) and staged combustion significantly reduce thermal in industrial processes. Bowin low technology is a hybrid of staged-premixed-radiant combustion technology with a major surface combustion preceded by a minor radiant combustion. In the Bowin burner, air and fuel gas are premixed at a ratio greater than or equal to the stoichiometric combustion requirement. Water Injection technology, whereby water is introduced into the combustion chamber, is also becoming an important means of reduction through increased efficiency in the overall combustion process. Alternatively, the water (e.g. 10 to 50%) is emulsified into the fuel oil before the injection and combustion. This emulsification can either be made in-line (unstabilized) just before the injection or as a drop-in fuel with chemical additives for long term emulsion stability (stabilized).


Rhodium(III) nitrate

Rhodium(III) nitrate is a nitrate of rhodium. It is a yellow, water-soluble salt.
Uses.
Rhodium(III) nitrate is used as a precursor to synthesize rhodium.
References.


Vinyl tributyltin

Vinyl tributyltin is an organotin compound with the formula Bu3SnCH=CH2 (Bu = butyl). It is a white, air-stable solid. It is used as a source of vinyl anion equivalent in Stille coupling reactions. As a source of vinyltin reagents, early work used vinyl trimethyltin, but trimethyltin compounds are avoided nowadays owing to their toxicity.
Preparation.
The compound is prepared by the reaction of vinylmagnesium bromide with tributyltin chloride. It can be synthesized in the laboratory by hydrostannylation of acetylene with tributyltin hydride. It is commercially available.


Bismuthine

Bismuthine (IUPAC name: bismuthane) is the chemical compound with the formula BiH3. As the heaviest analogue of ammonia (a pnictogen hydride), BiH3 is unstable, decomposing to bismuth metal well below 0 °C. This compound adopts the expected pyramidal structure with H-Bi-H angles of around 90°.
The term "bismuthine" may also refer to a member of the family of organobismuth(III) species having the general formula , where R is an organic substituent. For example, Bi(CH3)3 is "trimethylbismuthine".
Preparation and properties.
BiH3 is prepared by the redistribution of methylbismuthine (BiH2Me):
The required BiH2Me, which is also thermally unstable, is generated by reduction of methylbismuth dichloride, BiCl2Me with LiAlH4.
As suggested by the behavior of SbH3, BiH3 is unstable and decomposes to its constituent elements according to the following equation:
The methodology used for detection of arsenic ("Marsh test") can also be used to detect BiH3. This test relies on the thermal decomposition of these trihydrides to the metallic mirrors of reduced As, Sb, and Bi. These deposits can be further distinguished by their distinctive solubility characteristics: arsenic dissolves in NaOCl, antimony dissolves in ammonium polysulfide, and bismuth resists both reagents.
Uses and safety considerations.
The low stability of BiH3 precludes significant health effects, it decomposes rapidly well below room temperature.


Zinc peroxide

Zinc peroxide (ZnO2) appears as a bright yellow powder at room temperature. It was historically used as a surgical antiseptic. More recently zinc peroxide has also been used as an oxidant in explosives and pyrotechnic mixtures. Its properties have been described as a transition between ionic and covalent peroxides.
Zinc peroxide can be synthesized through the reaction of zinc chloride and hydrogen peroxide.
Preparation.
Zinc hydroxide is reacted with a mixture of hydrochloric acid and hydrogen peroxide and precipitated with sodium hydroxide also containing hydrogen peroxide to ensure a higher yield of zinc peroxide. Unlike in the preparation of copper peroxide, the zinc ion does not cause the peroxide to decompose.
Applications.
Since the 1930s zinc peroxide has been applied in a variety of settings, from medicine to aesthetics and even fireworks.
Medical Use.
The treatment of burrowing ulcers in the abdominal wall with zinc peroxide was first recorded in 1933 and throughout the 1940s ZnO2 was used as a disinfectant in surgical infections.
Zinc peroxide was, however, deemed ineffective against certain bacterial strains, such as "Streptococcus viridans", "staphylococcus aureus", "E. coli", "B. proteus", and "B. pyocyoneus". One aspect of the compound’s microorganism toxicity is the resultant stagnation of microbial populations upon administration. This effect was hypothesized to depend upon the compound’s penchant for oxygen donation. It has been suggested that the increase in oxygen concentration associated with the presence of ZnO2 interferes with the replicative processes of anaerobic and micro-aerophillic organisms, both of which require low oxygen environments for their survival.
While this mechanism was sufficient to explain the stagnation of microbe populations, it did not account for the active reduction in colony size. As to the microbicidal function, the Zinc ion itself has been postulated to have antibacterial properties, facilitated by the binding of the Zn ion to the bacterial cell wall, which allows for the exertion of cytotoxic effects. Zinc has been observed to be more effective in the elimination of gram-positive bacteria than gram-negative bacteria. This difference has been attributed to a difference in the protein composition of the respective cell walls, with the gram-positive wall providing a composition more conducive to binding.
Mineral stain.
Recently the compound has found use as a mineral stain for wood and other substances. The mechanism of this action involves the application of a metal salt (such as iron(II) chloride) and the zinc peroxide to the substrate material (wood or wood-like material, i.e. bamboo, paper, cloths, and cellulose products). The metal salt is applied in solution and allowed to dry for up to 30 minutes. Next, zinc peroxide is applied, also in solution. Color change is immediately visible. The two solutions soak into the material, and react, thereby becoming ingrained in the matrix of the substrate. While these stains can produce a variety of colors ranging from a reddish brown to a yellow hue, they are generally used to mimic the look of endangered wood species in cheaper and more commonly available stock.
Pyrotechnics.
In the 1980s the discovery of zinc peroxide’s ability to complement pyrotechnical mixtures was discovered. It was noted that ZnO2 was preferable to the usage of barium compounds, as it was deemed to be less toxic. The zinc compound proves an effective component in explosives because of its oxidizing properties. Many chemical explosives rely on rapid oxidation reactions, for this reason ZnO2 is an ideal candidate for use in pyrotechnics. Another advantage of ZnO2 was, as compared to the barium and strontium pyroingredients, it effects less corrosion in the metallic materials which contain the compounds in the pyrotechnic instrument.
In one modality it is necessary for zinc peroxide to act in conjunction with a reducer like calcium silicide, to create the necessary red/ox reaction. In another modality, a ‘secondary’ explosive is mixed with zinc peroxide. Secondary explosives include nitrocellulose, pentaerythritol tetranitrate (PETN), as well as a variety of compounds, like trinitrobenzene, which provide powerful negative charge balance.
These secondary explosives are relatively insensitive to stimuli such as physical impact, heat or charge. The explosive mixture would be comprised, in bulk, of the secondary explosive with a much smaller fraction of zinc peroxide, present to initiate the reaction.
Safety.
Zinc peroxide is very hazardous in case of skin contact, of eye contact, ingestion, or inhalation. It has been shown to be corrosive to skin. Prolonged exposure may result in skin burns and ulcerations. Over-exposure by inhalation may cause respiratory irritation. Skin inflammation is characterized by itching, scaling, reddening, or, occasionally, blistering. Zinc peroxide is toxic to lungs and mucous membranes. Repeated or prolonged exposure can produce organ damage. Repeated or prolonged inhalation of vapors may lead to chronic respiratory irritation.


Mercury(I) hydride

Mercury(I) hydride (systematically named mercury hydride) is an inorganic compound with the empirical chemical formula HgH. It has not yet been obtained in bulk, hence its bulk properties remain unknown. However, molecular mercury(I) hydrides with the formulae HgH and have been isolated in solid gas matrices. The molecular hydrides are very unstable toward thermal decomposition. As such the compound is not well characterised, although many of its properties have been calculated via computational chemistry.
Molecular forms.
History.
In 1979 and 1985, Swiss chemical physicists, Egger and Gerber, and Soviet chemical physicists, Kolbycheva and Kolbychev, independently, theoretically determined that it is feasible to develop a mercury(I) hydride molecular laser.
Chemical properties.
Mercury(I) hydride is an unstable gas and is the heaviest group 12 monohydride. The composition of mercury(I) hydride is 0.50% hydrogen and 99.50% mercury. In mercury(I) hydride, the formal oxidation states of hydrogen and mercury are −1 and +1, respectively, because of the electronegativity of mercury is lower than that of hydrogen. The stability of metal hydrides with the formula MH (M = Zn-Hg) increases as the atomic number of M increases.
The Hg-H bond is very weak and therefore the compound has only been matrix isolated at temperatures up to 6 K. The dihydride, HgH2, has also been detected this way.
A related compound is bis(hydridomercury)("Hg"—"Hg") with the formula , which can be considered to be dimeric mercury(I) hydride. It spontaneously decomposes into the monomeric form.
Radicality.
The mercury centre in mercury complexes such as hydridomercury can accept or donate a single electron by association:
Because of this acceptance or donation of the electron, hydridomercury has radical character. It is a moderately reactive monoradical.


Lanthanocene

A lanthanocene is a type of metallocene compound that contains an element from the lanthanide series. The most common lanthanocene complexes contain two cyclopentadienyl anions and an X type ligand, usually hydride or alkyl ligand.
History.
In 1954, Wilkinson and Birmingham described the tris(cyclopentadienyl)lanthanide complex, Ln(C5H5)3 (Ln = La, Ce, Pr, Nd, Sm). However, due to the highly moisture and oxygen sensitive character of organolanthanide compounds, as well as the incapability of the separation of simple alkyl and aryl derivatives of the type LnR3, this area of organomemtallic chemistry experienced a period of relative stagnation for two decades before lanthanocene complexes were prepared for some of the later lanthanides (Ln = Gd, Er, Yb, Lu).
Synthesis.
The synthesis part will focus on lanthanide(III) metallocene complexes that contain Ln-C bonds, which are widely prepared from the corresponding Ln-Cl precursors as shown.
The synthetic route leading to lanthanocene chlorides are summarized:
Reactions.
With the large 4f orbitals, lanthanide elements display properties significantly different from the common d-block transition metals. The large ionic radii limits the extent to which 4f orbitals can overlap with ligands, but at the same time allows the organolanthanide complexes to attain higher coordination numbers. As such, cyclopentadienyl anions (abbreviated Cp) are typically used to occupy the unsaturated site as well as stabilize the metal complex. Some of the alkyl and hydride lanthanocene complexes exhibit unique activities towards C-H bond activation, alkene functionalization, and carbonyl activation.
In 1983, Watson reported one of the first lanthanocene catalyzed C-H bond activation reactions. The active catalysts are lutetium-methyl or lutetium-hydride complexes, which react at room temperature in hydrocarbon solvents with benzene, pyridine, and the ylide CH2PPh3 to give stable, isolatable products.
Studies have shown that organolanthanides are extraordinary catalysts for hydrofunctionalization reactions including hydrogenation, hydrosilylation, hydroboration, hydroamination, etc. Examples for each type have shown below.
Mechanism for hydrogenation is shown below, where the active catalyst is generated by sigma-bond metathesis, followed by olefin insertion and another sigma-bond metathesis to regenerate the catalyst. This is also the mechanism for other hydrofunctionalization.
Yet, as the large cyclopentadienyl ligand hinders the metal center, reactions with substituted alkenes are inhibited. Two general methods are used to overcome this difficulty. One is to increase the size of metal by incorporating lanthanides with larger ionic radii. Due to the lanthanide contraction, this means replacing the late lanthanides with the early lanthanides. Another method is to decrease the size of the ligand by manipulating the geometry of ligands and substitutions on ligands. For example, a hinged or ansa-bridged cyclopentadienyl ligand could be used to pull ligands closer to each other, and hence creating more open access to the metal center as shown on the right.


Murdochite

Murdochite is a mineral combining lead and copper oxides with the chemical formula  (x ≤ 0.5).
It was first discovered in 1953 in the Mammoth-Saint Anthony Mine in Pinal County, Arizona by Percy W. Porter, a mining engineer who handpicked a 401.5-mg sample. Porter would later submit for analysis and it was then that Fred A. Hildebrand suggested that the sample was a new mineral after taking a powder x-ray picture. It was named for Joseph Murdoch (1890–1973), American mineralogist. Murdochite was first suggested to be of a cubic structure. After this suggestion, the term “murdochite-type structure” began to be used when describing a structure that is similar to that of murdochite. Murdochite was later found to be octahedral.
Composition.
Various studies have examined the composition of murdochite. When murdochite was first discovered, its chemical formula was determined through stoichiometry to be Cu6•1Pb0•9O8•O. The ideal formula for murdochite was originally thought to be . This formula failed to take into consideration the significant amounts of chlorine and bromine present in murdochite. Using electron microprobe analysis, a new chemical composition for murdochite was determined that included chlorine and bromine. This new formula was and was created with regard to the Cu/Pb variance observed in samples of murdochite from both Hansonburg, New Mexico, and T. Khuni mine, Iran. This new composition was also proven incorrect and the formula that is widely accepted today,  (x ≤ 0.5), was proposed.
Structure.
The structure originally proposed for murdochite was a simple cubic structure. This structure was later proven incorrect and determined to be octahedral. The structure of murdochite is described as edges of [PbO8] cubes spanned by Cu2+ ions that give a square-planar CuO4 arrangement in a three-dimensional network. Halogen ions complete a (4 + 2_ elongated octahedron.
Physical properties.
Murdochite is an opaque mineral that exhibits a black color, adamantine luster, black streak, and isotropic optical class. Samples from T. Khuni mine indicate that polished sections of murdochite resemble the color and reflectivity of magnetite but vary because of zoning. The lighter zones in murdochite are caused because they are richer in lead than the darker zones. The mineral is known to be brittle, have a hardness of 4, and have {111} cleavage. Murdochite also has a pink tinge that is distinctive and its edges seem to be more reflective than the center because samples are commonly porous in the center and compact on the sides with clean-cut faces. The space group of murdochite has been confirmed to be Fm3m. Murdochite can be cubic, with a point group of 4/m 3 2/m, or octahedral and cubo-octahedral whose point group is 2 mm. Murdochite can occur in the forms (100) and (111) with crystal twinning being very common. The calculated density of murdochite was determined to be 6.1 g/cm3 and have a measured density of 6.47 g/cm3.
Geologic occurrence.
Naturally occurring murdochite,  (x ≤ 0.5), was discovered in Mammoth Mine, Arizona. The mine, about 46 miles northeast of Tucson, contains deposits of molybdenum, lead, gold, zinc, vanadium, sphalerite, and galena. In this mine, murdochite was found embedded in plates of wulfenite as well as on the surface of crystals of fluorite. Other crystals found in Mammoth Mine include hemimorphite, willemite, and quartz. Larger masses of murdochite have also been found at the T. Kguni Mine in Anarak, Iran. Here murdochite was found in conjunction with khuniite (iranite), chrominium (phoenicochroite), plumangite and an unknown mineral with the formula Pb9O16. All these minerals are lead oxide minerals, meaning that murdochite can be found in oxidized Pb-Cu deposits. More recently, small murdochite crystals were found encrusting aurichalcite crystals in the Granite Gap Mines in New Mexico.
Origin of the name.
The mineral murdochite is named after Joseph Murdoch, who was once a professor of geology at the University of California, Los Angeles and a past president of the Mineralogical Society of America. Professor Murdoch was born February 19, 1890 and died December 31, 1973 at 83 years of age. Murdoch obtained a baccalaureate degree, an M.S. and a Ph.D in geologic science by 1915 from Harvard University and would first be involved in a non-geological business career before joining the faculty at the department of geology at the University of California, Los Angeles in 1928. It was here that Murdoch returned to work in the geologic field that would lead to the publishing of numerous papers which would include 23 publications ranging from the years 1913-1968, along with ten published abstracts and many book reviews.


Vanadyl isopropoxide

Vanadyl isopropoxide is the metal alkoxide with the formula VO(O-iPr)3 (iPr = CH(CH3)2). A yellow volatile liquid, it is a common alkoxide of vanadium. It is used as a reagent and as a precursor to vanadium oxides. The compound is diamagnetic. It is prepared by alcoholysis of vanadyl trichloride:
The related cyclopentanoxide VO(O-CH(CH2)4)3 is a dimer, one pair of alkoxide ligands bind weakly trans to the vanadyl oxygens.


C84 fullerene


Scandium(III) hydride

Scandium trihydride is an unstable molecular chemical compound with the chemical formula ScH3. It has been formed as one of a number of other molecular scandium hydride products at low temperature using laser ablation and identified by infrared spectroscopy.
Scandium trihydride has recently been the subject of Dirac–Hartree–Fock relativistic calculation studies, which investigate the stabilities, geometries, and relative energies of hydrides of the formula MH3, MH2, or MH.
Properties and bonding.
Scandium trihydride is a quastrigonal planar molecule with three equivalent Sc-H bonds. (C3v) structure an equilibrium distance between Sc and hydrogen of 182.0 pm, the bond angle is 119.2 degrees. By weight percent, the composition of scandium trihydride is 6.30% hydrogen and 93.70% scandium. In scandium trihydride, the formal oxidation states of hydrogen and scandium are -1 and +3 respectively, because of the electronegativity of scandium is lower than that of hydrogen. The stability of metal hydrides with the formula MH3 (M = Sc-Lu) increases as the atomic number of M increases.
Early theoretical studies of ScH3 revealed that the molecule is unstable, the bulk substance is likely to be a colourless gas with a low activation energy toward the conversion into trimeric clusters due to the electron deficient nature of the monomer, not unlike the group 13 hydrides. One major difference, is that the dimer is the most stable cluster for group 13 hydrides. This can be attributed to the distortion caused by the "d"-orbitals. It cannot be made by methods used to synthesise BH3 or AlH3.


Rutile

Rutile is a mineral composed primarily of titanium dioxide (TiO2), and is the most common natural form of TiO2. Other rarer polymorphs of TiO2 are known, including anatase, akaogiite, and brookite.
Rutile has one of the highest refractive indices at visible wavelengths of any known crystal and also exhibits a particularly large birefringence and high dispersion. Owing to these properties, it is useful for the manufacture of certain optical elements, especially polarization optics, for longer visible and infrared wavelengths up to about 4.5 μm.
Natural rutile may contain up to 10% iron and significant amounts of niobium and tantalum. Rutile derives its name from the Latin "rutilus", red, in reference to the deep red color observed in some specimens when viewed by transmitted light. Rutile was first described in 1803 by Abraham Gottlob Werner.
Occurrence.
Rutile is a common accessory mineral in high-temperature and high-pressure metamorphic rocks and in igneous rocks.
Thermodynamically, rutile is the most stable polymorph of TiO2 at all temperatures, exhibiting lower total free energy than metastable phases of anatase or brookite. Consequently, the transformation of the metastable TiO2 polymorphs to rutile is irreversible. As it has the lowest molecular volume of the three main polymorphs, it is generally the primary titanium bearing phase in most high-pressure metamorphic rocks, chiefly eclogites.
Within the igneous environment, rutile is a common accessory mineral in plutonic igneous rocks, though it is also found occasionally in extrusive igneous rocks, particularly those such as kimberlites and lamproites that have deep mantle sources. Anatase and brookite are found in the igneous environment particularly as products of autogenic alteration during the cooling of plutonic rocks; anatase is also found in placer deposits sourced from primary rutile.
The occurrence of large specimen crystals is most common in pegmatites, skarns, and granite greisens. Rutile is found as an accessory mineral in some altered igneous rocks, and in certain gneisses and schists. In groups of acicular crystals it is frequently seen penetrating quartz as in the "fléches d'amour" from Graubünden, Switzerland. In 2005 the Republic of Sierra Leone in West Africa had a production capacity of 23% of the world's annual rutile supply, which rose to approximately 30% in 2008.
Crystal structure.
Rutile has a tetragonal unit cell, with unit cell parameters "a" = "b" = 4.584 Å, and "c" = 2.953 Å. The titanium cations have a coordination number of 6, meaning they are surrounded by an octahedron of 6 oxygen atoms. The oxygen anions have a coordination number of 3, resulting in a trigonal planar coordination. Rutile also shows a screw axis when its octahedra are viewed sequentially.
Rutile crystals are most commonly observed to exhibit a prismatic or acicular growth habit with preferential orientation along their "c" axis, the [001] direction. This growth habit is favored as the {110} facets of rutile exhibit the lowest surface free energy and are therefore thermodynamically most stable. The "c"-axis oriented growth of rutile appears clearly in nanorods, nanowires and abnormal grain growth phenomena of this phase.
Uses and economic importance.
In large enough quantities in beach sands, rutile forms an important constituent of heavy minerals and ore deposits. Miners extract and separate the valuable minerals – e.g., rutile, zircon, and ilmenite. The main uses for rutile are the manufacture of refractory ceramic, as a pigment, and for the production of titanium metal.
Finely powdered rutile is a brilliant white pigment and is used in paints, plastics, paper, foods, and other applications that call for a bright white color. Titanium dioxide pigment is the single greatest use of titanium worldwide. Nanoscale particles of rutile are transparent to visible light but are highly effective in the absorption of ultraviolet radiation. The UV absorption of nano-sized rutile particles is blue-shifted compared to bulk rutile, so that higher-energy UV light is absorbed by the nanoparticles. Hence, they are used in sunscreens to protect against UV-induced skin damage.
Small rutile needles present in gems are responsible for an optical phenomenon known as asterism. Asteriated gems are known as "star" gems. Star sapphires, star rubies, and other "star" gems are highly sought after and are generally more valuable than their normal counterparts.
Rutile is widely used as a welding electrode covering. It is also used as a part of the ZTR index, which classifies highly weathered sediments.
Rutile, as a large band-gap semiconductor, has in recent decades been the subject of significant research towards applications as a functional oxide for applications in photocatalysis and dilute magnetism. Research efforts typically utilize small quantities of synthetic rutile rather than mineral-deposit derived materials.
Synthetic rutile.
Synthetic rutile was first produced in 1948 and is sold under a variety of names. It can be produced from the titanium ore ilmenite through the Becher process. Very pure synthetic rutile is transparent and almost colorless, being slightly yellow, in large pieces. Synthetic rutile can be made in a variety of colors by doping. The high refractive index gives an adamantine luster and strong refraction that leads to a diamond-like appearance. The near-colorless diamond substitute is sold as "Titania", which is the old-fashioned chemical name for this oxide. However, rutile is seldom used in jewellery because it is not very hard (scratch-resistant), measuring only about 6 on the Mohs hardness scale.
As the result of growing research interest in the photocatalytic activity of titanium dioxide, in both anatase and rutile phases (as well as biphasic mixtures of the two phases), rutile TiO2 in powder and thin film form is frequently fabricated in laboratory conditions through solution based routes using inorganic precursors (typically TiCl4) or organometallic precursors (typically alkoxides such as titanium isopropoxide, also known as TTIP). Depending on synthesis conditions, the first phase to crystallize may be the metastable anatase phase, which can then be converted to the equilibrium rutile phase through thermal treatment. The physical properties of rutile are often modified using dopants to impart improved photocatalytic activity through improved photo-generated charge carrier separation, altered electronic band structures and improved surface reactivity.


Molybdenum dioxide

Molybdenum dioxide is the chemical compound with the formula MoO2. It is a violet-colored solid and is a metallic conductor. The mineralogical form of this compound is called tugarinovite, and is only very rarely found.
Structure.
It crystallizes in a monoclinic cell, and has a distorted rutile, (TiO2) crystal structure. In TiO2 the oxide anions are close packed and titanium atoms occupy half of the octahedral interstices (holes). In MoO2 the octahedra are distorted, the Mo atoms are off-centre, leading to alternating short and long Mo – Mo distances and Mo-Mo bonding. The short Mo – Mo distance is 251 pm which is less than the Mo – Mo distance in the metal, 272.5 pm. The bond length is shorter than would be expected for a single bond. The bonding is complex and involves a delocalisation of some of the Mo electrons in a conductance band accounting for the metallic conductivity.
Preparation.
MoO2 can be prepared :
Single crystals are obtained by chemical transport using iodine. Iodine reversibly converts MoO2 into the volatile species MoO2I2.
Uses.
Molybdenum oxide is a constituent of "technical molybdenum oxide" produced during the industrial processing of MoS2:
MoO2 has been reported as catalysing the dehydrogenation of alcohols, the reformation of hydrocarbons and biodiesel. Molybdenum nano-wires have been produced by reducing MoO2 deposited on graphite. Molybdenum oxide has also been suggested as possible anode material for Li-ion batteries.


Unpentbium


Unhexquadium


Lithium acetate

Lithium acetate (CH3COOLi) is a salt of lithium and acetic acid.
Uses.
Lithium acetate is used in the laboratory as buffer for gel electrophoresis of DNA and RNA. It has a lower electrical conductivity and can be run at higher speeds than can gels made from TAE buffer (5-30V/cm as compared to 5-10V/cm). At a given voltage, the heat generation and thus the gel temperature is much lower than with TAE buffers, therefore the voltage can be increased to speed up electrophoresis so that a gel run takes only a fraction of the usual time. Downstream applications, such as isolation of DNA from a gel slice or Southern blot analysis, work as expected when using lithium acetate gels.
Lithium boric acid or sodium boric acid are usually preferable to lithium acetate or TAE when analyzing smaller fragments of DNA (less than 500 bp) due to the higher resolution of borate-based buffers in this size range as compared to acetate buffers.
Lithium acetate is also used to permeabilize the cell wall of yeast for use in DNA transformation. It is believed that the beneficial effect of LiOAc is caused by its chaotropic effect; denaturing DNA, RNA and proteins.


Dibromoacetylene

Dibromoacetylene is a molecular chemical compound containing acetylene, with its hydrogen substituted by bromine.
Production.
Dibromoacetylene can be made by reacting 1,1,2-tribromoethylene with potassium hydroxide. This method has a danger of causing explosions.
Another way is to react acetylene with phenyl lithium (at −50 °C) to make lithium acetylide, which then reacts with bromine to yield the product.
Yet another way is to react acetylene with sodium hypobromite NaOBr.
Properties.
Dibromoacetylene is explosive and sensitive to air. Its appearance is a clear water like liquid. It has a sweetish smell, but makes a white fume in air that then smells like ozone, possibly because it forms ozone. Dibromoacetylene is lachrymatory.
Dibromoacetylene can be polymerised to polydibromoacetylene using catalysts like titanium tetrachloride and triethylaluminium. Polydibromoacetylene is black, electrically conducting, and stable in air to over 200 °C.
The infrared spectrum has absorption at 2185 cm−1 due to symmetrical stretching on the C≡C bond, 832 cm−1 asymmetrical stretch on C-Br bond, 311 cm−1 bending on Z shape BrCC, 267 symmetrical stretch on all bonds, and 167 cm−1 for C shaped bending.
Reactions.
In air dibromoacetylene spontaneously inflames producing black sooty smoke and a red flame. When heated, it explodes, yielding carbon and other substances. A slower reaction with oxygen and water, yields oxalic acid and hydrobromic acid and other bromine containing substances. A reaction with hydrogen iodide yields dibromoiodoethylene.
Dibromoacetylene reacts with bromine to yield tetrabromoethene.
Dibromoacetylene reacts with Apiezon L, which is used as a vacuum grease in distillation, so its use is unsuitable with this chemical.


Malayaite

Malayaite is a calcium tin silicate mineral with formula CaSnOSiO4. It is a member of the titanite group.
Discovery.
Malayaite was originally found in Perak (a state in Malaysia) and was first described in literature in 1961, though it was not yet given a name. In 1965, the mineral was named and recognized by the International Mineralogical Association. It was named for the locality in which it was discovered; which is the Malay Peninsula in Malaysia.
Crystal structure and symmetry.
This mineral is classified under the nesosilicate group for silicate minerals because the titanite group falls under this category as well. Minerals in the nesosilicate group have isolated SiO4 tetrahedra connected to cations. In malayaite, the tetrahedra are connected to the chain of distorted SnO6 octahedra, in which the octahedra are linked by vertex ["trans" corners] sharing and form chains parallel to the miller index of [100]. Within the SnO6-SiO4 framework, the CaO7 polyhedra form chains parallel to [101].
Malayaite belongs to the monoclinic crystal system and has a 2/m (prismatic) crystal class. According to the Hermann–Mauguin notation, the '2' refers to the two-fold axis while the 'm' refers to the presence of a single mirror plane. The '/' symbol indicates that the twofold axis is perpendicular to the mirror plane.
The space group for malayaite is A2/a. According to the Bravais lattice symbol, 'A' refers to single face centering of the motif. This means that there is one more point on the center of one face. '2' refers to the two-fold rotation axis while the '/' indicates that a mirror plane is perpendicular to the "a"-axis, which is the last symbol on that space group notation.
Optical properties.
Malayaite exhibits anisotropy so it is doubly refractive and breaks up the light that travels through it into two distinct rays of different speeds. This mineral is known to have very high optical relief and has three indices of refraction. Since it is a monoclinic mineral, Malayaite exhibits two different colors under plane polarized light making it a pleochroic mineral. Malayaite is biaxial birefringent (trirefringent).
Functions and purposes.
Malayaite has excellent thermal properties and stability making its host lattice ideal for ceramic pigmentation unlike its analogous mineral, titanite that undergoes phase transition at a higher temperature. This makes chromium-doped malayaite very desirable as it provides a consistent pigment for ceramics. Malayaite contains the element tin which is sometimes doped with chromium or nickel. These chromophore elements allows the mineral to produce different pigmentation. If malayaite is chromium-doped, it produces a pink-red color while with nickel, it produces a purple color. Pink chromium-doped malayaite is among a few important minerals used in producing pigments for the ceramic industry, specifically for coloring glazes as it produces a resilient color.
Occurrence.
This mineral is generally found in tin-rich contact metamorphic skarn deposits. There is a possibility that malayaite is a hydrothermal altered form of cassiterite-quartz assemblage or some tin-bearing mineral. There have been specimens that were found with malayaite as a coating mineral on cassiterite.


Ti-sapphire laser

Ti:sapphire lasers (also known as Ti:Al2O3 lasers, titanium-sapphire lasers, or Ti:sapphs) are tunable lasers which emit red and near-infrared light in the range from 650 to 1100 nanometers. These lasers are mainly used in scientific research because of their tunability and their ability to generate ultrashort pulses. Lasers based on Ti:sapphire were first constructed and invented in June 1982 by Peter Moulton at the MIT Lincoln Laboratory.
Titanium-sapphire refers to the lasing medium, a crystal of sapphire (Al2O3) that is doped with Ti3+ ions. A Ti:sapphire laser is usually pumped with another laser with a wavelength of 514 to 532 nm, for which argon-ion lasers (514.5 nm) and frequency-doubled , , and Nd:YVO lasers (527-532 nm) are used. They are capable of laser operation from 670 nm to  nm wavelength. Ti:sapphire lasers operate most efficiently at wavelengths near 800 nm.
Types of Ti:sapphire lasers.
Mode-locked oscillators.
Mode-locked oscillators generate ultrashort pulses with a typical duration between a few picoseconds and 10 femtoseconds, in special cases even around 5 femtoseconds. The pulse repetition frequency is in most cases around 70 to 90 MHz. Ti:sapphire oscillators are normally pumped with a continuous-wave laser beam from an argon or frequency-doubled laser. Typically, such an oscillator has an average output power of 0.4 to 2.5 watts.
Chirped-pulse amplifiers.
These devices generate ultrashort, ultra-high-intensity pulses with a duration of 20 to 100 femtoseconds. A typical one stage amplifier can produce pulses of up to 5 millijoules in energy at a repetition frequency of 1000 hertz, while a larger, multistage facility can produce pulses up to several joules, with a repetition rate of up to 10 Hz. Usually, amplifiers crystals are pumped with a pulsed frequency-doubled Nd:YLF laser at 527 nm and operate at 800 nm. Two different designs exist for the amplifier: regenerative amplifier and multi-pass amplifier.
Regenerative amplifiers operate by amplifying single pulses from an oscillator (see above). Instead of a normal cavity with a partially reflective mirror, they contain high-speed optical switches that insert a pulse into a cavity and take the pulse out of the cavity exactly at the right moment when it has been amplified to a high intensity.
The term 'chirped-pulse' refers to a special construction that is necessary to prevent the pulse from damaging the components in the laser. The pulse is stretched in time so that the energy is not all located at the same point in time and space. This prevents damage to the optics in the amplifier. Then the pulse is optically amplified and recompressed in time to form a short, localized pulse. All optics after this point should be chosen to take the high energy density into consideration.
In a multi-pass amplifier, there are no optical switches. Instead, mirrors guide the beam a fixed number of times (two or more) through the Ti:sapphire crystal with slightly different directions. A pulsed pump beam can also be multi-passed through the crystal, so that more and more passes pump the crystal. First the pump beam pumps a spot in the gain medium. Then the signal beam first passes through the center for maximal amplification, but in later passes the diameter is increased to stay below the damage-threshold, to avoid amplification the outer parts of the beam, thus increasing beam quality and cutting off some amplified spontaneous emission and to completely deplete the inversion in the gain medium.
The pulses from chirped-pulse amplifiers are often converted to other wavelengths by means of various nonlinear optical processes.
At 5 mJ in 100 femtoseconds, the peak power of such a laser is 50 gigawatts. When focused by a lens, these laser pulses will ionise any material placed in the focus, including air molecules.
Tunable continuous wave lasers.
Titanium-sapphire is especially suitable for pulsed lasers since an ultrashort pulse inherently contains a wide spectrum of frequency components. This is due to the inverse relationship between the frequency bandwidth of a pulse and its time duration, due to their being conjugate variables. However, with an appropriate design, titanium-sapphire can also be used in continuous wave lasers with extremely narrow linewidths tunable over a wide range.
History and applications.
The Ti:sapphire laser was invented by Peter Moulton in June 1982 at MIT Lincoln Laboratory in its continuous wave version. Subsequently, these lasers were shown to generate ultrashort pulses through Kerr-lens modelocking. Strickland and Mourou, in addition to others, working at the University of Rochester, showed chirped pulse amplification of this laser within a few years, for which these two shared in the 2018 Nobel Prize in physics (along with Arthur Ashkin for optical tweezers). The cumulative product sales of the Ti:sapphire laser has amounted to more than $600 million, making it a big commercial success that has sustained the solid state laser industry for more than three decades.
The ultrashort pulses generated by Ti:sapphire lasers in the time domain corresponds to mode-locked optical frequency combs in the spectral domain. Both the temporal and spectral properties of these lasers make them highly desirable for frequency metrology, spectroscopy, or for pumping nonlinear optical processes. One half of the Nobel prize for physics in 2005 was awarded to the development of the optical frequency comb technique, which heavily relied on the Ti:sapphire laser and its self-modelocking properties. The continuous wave versions of these lasers can be designed to have nearly quantum limited performance, resulting in a low noise and a narrow linewidth, making them attractive for quantum optics experiments. The reduced amplified spontaneous emission noise in the radiation of Ti:sapphire lasers lends great strength in their application as optical lattices for the operation of state-of-the-art atomic clocks. Apart from fundamental science applications in the laboratory, this laser has found biological applications such as deep-tissue multiphoton imaging and industrial applications cold micromachining. When operated in the chirped pulse amplification mode, they can be used to generate extremely high peak powers in the terawatt range, which finds use in nuclear fusion research.


Desloge Consolidated Lead Company

Desloge Consolidated Lead Company was a lead mining company in the Southeast Missouri Lead District that was operated by the Desloge family in the 19th and early 20th century. The Desloge lead operations in the "Old Lead Belt", in the eastern Ozark Mountains, helped Missouri become the world's premier lead mining area.
History.
Desloge Lead Company.
The businesses that would become the Desloge Consolidated Lead Company began around 1824, when Firmin Rene Desloge — founder of the Desloge Family in America — built a smelting furnace as an extension of his mercantile business in Potosi, Missouri.
His son, Firmin V. Desloge, expanded mining operations and moved management to Bonne Terre, Missouri. On June 5, 1874, a charter was granted to the Missouri Lead and Smelting Company, which was renamed the Desloge Lead Company on February 21, 1876. In 1876 and 1877, the company sank three shafts and built a new mill in Bonne Terre.
A fire in March 1886 destroyed the concentrating mill plant and damaged the rest of the surface plant. The following year, the company's remaining assets were sold to St. Joseph Lead Company, which made Firmin V. Desloge a trustee on its board, a post he would hold until his death. This merger helped the St. Joe company become the "greatest lead-mining and smelting company in the world.
Desloge Consolidated Lead Company.
After the fire, Desloge and his family took an option on a piece of land owned by the Bogy family which was called "Mine-a-Joe", and started another mining operation under the name Desloge Consolidated Lead Company. The new company cleared the land and built company houses for its workers just west of present-day Desloge, Missouri.
In 1893, the Desloges opened a new mine in St. Francois County, Missouri, just north of the St. Joe Lead Company property in Bonne Terre, on a tract of land originally granted to Jean Bte. Pratte and designated as U. S. Survey No. 3099. The mill and smelting plant could produce 500 tons of lead per day.
Firmin Desloge II expanded lead mining operations by buying the Bogy Lead Mine Company and the St. Francois Mining Company. Along with his partners, he organized a new company called Desloge Consolidated Lead Company. The company sank mine shafts and built mills, smelting furnaces, and power stations.
Ultimately, the company's lead operations required a massive real estate effort involving hundreds of land leases, purchases, options, rights of refusals, mineral rights, chattel mortgages, deed transfers, quit-claims, trust deeds, judgment sales, sheriff’s sales, bankruptcy sales, grants, bonds, notes, and various claims from area miners covering thousands of acres and hundreds of parcels of land.
Railroads.
To serve the Desloge and St. Joe mines, Firmin Desloge II built the first railroads to penetrate the disseminated lead field of St. Francois County: the Desloge Railway, the Mississippi River and Bonne Terre Railway and the Valley Railroad.
Desloge also helped develop the St. Louis, Iron Mountain and Southern Railway (aka the Iron Mountain Railroad), which opened in 1874 from St. Louis to Texarkana, Arkansas. When the St. Joseph Lead Company built a 13.5-mile narrow-gauge railroad from its mines to the Iron Mountain tracks at Summit in Washington County, the Desloge Company paid one-third of the costs.
Late-century lobbying.
The Revenue Act of 1894, also known as the Wilson-Gorman Tariff, lowered protective tariffs on lead (although it raised rates on sugar and other materials). The prospect of cheaper imported lead threatened the Desloge company and other southeast Missouri lead businesses, so Firmin Desloge II immediately went to Washington, D.C., to lobby for repeal. In the meantime, he spent part of his own fortune to continue operations, retained his employees, and stockpiled instead of sold his pig lead. The lead industry's lobbying efforts paid off in 1897, when the Dingley Act once again raised barriers to lead imports.
1900s.
Around 1916, the Desloge Consolidated Lead Company moved its corporate offices from Desloge to the 1892 Rialto Building in downtown St. Louis. Despite the city's history as a fur-trading center, "more money passed through St. Louis as a result of the lead business in Missouri,” Robert E. McHenry wrote in his 2006 book "Chat Dumps of The Missouri Lead Belt".
Sale.
In 1929, the company sold its plant to the St. Joe Lead Company for $18 million ($ today). A trade newspaper wrote, “With the absorption of the Desloge concern by the St. Joseph Lead Company, one of the oldest mining companies of the district goes out of existence as a company.”
The Desloge Consolidated Lead Company, and specifically Firmin Rene Desloge and Firmin Desloge II, were described in 1932 as "the most distinguished of American mining engineers".
The modern Desloge Consolidated Lead Company.
The name Desloge Consolidated Lead Company was used for a company formed in 2004 as an investment holding company by a Desloge family descendant, Christopher Desloge. The name was changed in 2005 to Madaket Growth, LLC.


Gallium(III) oxide

Gallium(III) trioxide is an inorganic compound with the formula Ga2O3. It exists as several polymorphs, all of which are white, water-insoluble solids. Although no commercial applications exist, Ga2O3 is an intermediate in the purification of gallium, which is consumed almost exclusively as gallium arsenide. The thermal conductivity of β-Ga2O3 is at least one order of magnitude lower than the other wide bandgap semiconductors, such as GaN and SiC. It is further reduced for related nanostructures which are usually used in electronic devices. Heterogeneous integration with high thermal conductivity substrates such as diamond and SiC helps heat dissipation of β-Ga2O3 electronics.
Preparation.
Gallium trioxide is precipitated in hydrated form upon neutralization of acidic or basic solution of gallium salt. Also, it is formed on heating gallium in air or by thermally decomposing gallium nitrate at 200–250 ˚C. It can occur in five different modifications, α, β, γ, δ, and ε. Of these modifications β-Ga2O3 is the most stable form.
Reactions.
Gallium(III) trioxide is amphoteric. It reacts with alkali metal oxides at high temperature to form, e.g., NaGaO2, and with Mg, Zn, Co, Ni, Cu oxides to form spinels, e.g. MgGa2O4.
It dissolves in strong alkali to form a solution of the gallate ion, .
With HCl, it forms gallium trichloride GaCl3.
It can be reduced to gallium suboxide (gallium(I) oxide) Ga2O by H2. or by reaction with gallium metal:
Structure.
β-Ga2O3, with a melting point of 1900 ˚C, is the most stable crystalline modification. The oxide ions are in a distorted cubic closest packing arrangement, and the gallium (III) ions occupy distorted tetrahedral and octahedral sites, with Ga-O bond distances of 1.83 and 2.00 Å respectively.
α-Ga2O3 has the same structure (corundum) as α-Al2O3, wherein Ga ions are 6-coordinate. γ-Ga2O3 has a defect spinel structure similar to that of γ-Al2O3.
ε-Ga2O3 films deposited by metalorganic vapour phase epitaxy show a columnar structure with orthorhombic crystal symmetry. Macroscopically, this structure is seen by X-ray crystallography as hexagonal close packed.
Potential applications.
Gallium(III) oxide has been studied in the use of lasers, phosphors, and luminescent materials. It has also been used as an insulating barrier in tight junctions. Monoclinic β-Ga2O3 is used in gas sensors and luminescent phosphors and can be applied to dielectric coatings for solar cells. This stable oxide has also shown potential for deep-ultraviolet transparent conductive oxides, and transistor applications.
ε-Ga2O3 thin films deposited on sapphire show potential applications as solar-blind UV photodetector.
Thin Ga2O3 films are of commercial interest as gas sensitive materials and Ga2O3. Ellipsometry is a procedure that can be used to determine optical functions of the β-Ga2O3.
β-Ga2O3 is used in the production of Ga2O3-Al2O3 catalyst.


Isopropylmagnesium chloride

Isopropylmagnesium chloride is an organometallic compound with the general formula (CH3)2HCMgCl. This highly flammable, colorless, and moisture sensitive material is the Grignard reagent derived from isopropyl chloride. It is commercially available, usually as a solution in tetrahydrofuran. This reagent is used to prepare Grignard reagents by transmetalation reactions as well as installing isopropyl groups. An illustrative generic reaction involves the generation of the Grignard reagent derived from bromo-3,5-bis(trifluoromethyl)benzene:
Isopropylmagnesium chloride is also used to prepare other isopropyl compounds, such as chlorodiisopropylphosphine:
This reaction exploits the bulky nature of the isopropyl substituent.


Tricarbon monosulfide

Tricarbon monosulfide (C3S) or tricarbon sulfur is a reactive molecular substance found in space. Tricarbon monosulfide is a heterocumulene or thiocumulene, consisting of a straight chain of three carbon atoms and a terminal sulfur atom.
Properties.
The dipole moment of tricarbon monosulfide is 3.704 Debye. The bond-lengths are 1.275 Å, for terminal C=C bond, 1.292 Å for internal bond, and 1.535 Å for the C=S bond. The similar bond lengths between the carbon atoms indicate they each have a double bond nature. The rotational constants for 12C12C12C32S are B0 = 2890.38000 MHz and D0 = 0.00022416.
There is a characteristic infrared absorption band at 2047.5 cm−1 due to stretching of a C=C bond.
Formation.
Along with the related dicarbon monosulfide (CCS), tricarbon monosulfide was made by a glow discharge though carbon disulfide vapour in helium. Microwave emission lines from rotational transitions matched up with previously unknown molecular lines from the Taurus Molecular Cloud 1. Maximal concentrations occurred with a carbon disulfide pressure of 0.02 torr.
In molecular clouds, the formation mechanism is speculated to be CCS + CH → CCCS + H.
On dust grains, in space the formation mechanism is theorised to be: CCC + H2S → C3•HSH → CCCS + H2 when irradiated with visible or UV light. THis reaction has been reproduced in a solid argon matrix.
Natural occurrence.
Tricarbon monosulfide has been detected in molecular clouds in space, such as Taurus Molecular Cloud 1 and Orion KL. The ratio of tricarbon monoxide to tricarbon monosulfide reflects the ratio of sulfur to oxygen in the cloud. The ratio of concentration of sulfur to oxygen analogues follows the same pattern. The clouds can be cold and dark, or warm. CCCS has also been found in the stellar envelope of carbon rich AGB stars, including in IRC+10216.


Transatomic Power

Transatomic Power was an American company that designed Generation IV nuclear reactors based on molten salt reactor (MSR) technology.
MIT alumni Dr. Leslie Dewan and Mark Massie founded Transatomic Power in 2011, and its board directors included E Ink Corporation co-founder Russ Wilcox. Among its backers were the venture capital outfit Founders Fund, of which Peter Thiel is a partner. In 2013, the U.S. Department of Energy awarded Transatomic first prize in the ARPA-E Future Energy innovation contest.
In 2018 the company announced that it would be winding down and open source its intellectual property. The company discovered that in 2016 it had made errors in its early analysis and realized that the design couldn't consume nuclear waste. Transatomic Power ceased operation on September 25, 2018.
Reactor concept.
Transatomic initial concept was that of a "Waste Annihilating Molten Salt Reactor" (WAMSR) designed to digest spent nuclear fuel.
The concept was based on the Molten-Salt Reactor Experiment reactor that ran at Oak Ridge National Laboratory (ORNL) from 1964 to 1969.
It was open to the use of thorium or uranium as a fuel in its reactors.
This design was later updated and corrected, and the claim that the reactor is able to use nuclear waste as fuel was dropped.
The latest design is based on a low pressure, high temperature molten salt reactor.
Features include:


Zinc mining

Zinc mining is the process by which mineral forms of the metal zinc are extracted from the earth through mining. A zinc mine is a mine that produces zinc minerals in ore as its primary product. Common co-products in zinc ores include minerals of lead and silver. Other mines may produce zinc minerals as a by-product of the production of ores containing more valuable minerals or metals, such as gold, silver or copper. Mined ore is processed, usually on site, to produce one or more metal-rich concentrates, then transported to a zinc smelter for production of zinc metal.
Global zinc mine production in 2019 was estimated to be 12.9 million tonnes. The largest producers were China (34%), Peru (11%), Australia (10%), United States (6.1%), India (5.5%), and Mexico (5.4%), with Australia having the largest reserves.
The world's largest zinc mine is the Red Dog open-pit zinc-lead-silver mine in Alaska, with 4.2% of world production. Major zinc mine operators include Vedanta Resources, Glencore, BHP, Teck Resources, Sumitomo, Nexa Resources, Boliden AB, and China Minmetals.
History.
Zinc deposits have been exploited for thousands of years, with the oldest zinc mine, located in Rajasthan, India established nearly 2000 years BP.
Pure zinc production occurred in the 9th century AD while, earlier in antiquity zinc was primarily utilized in the alloying of copper to produce Brass. This is because the isolation of zinc metal from its ore poses a unique challenge. This is because at the temperature zinc is released from its ore it vaporizes into a gas, and if the furnace is not air tight, the gaseous zinc reacts with the air to form zinc oxide.
Metallic zinc smelting occurred in 9th century BC in India, followed soon by China 300 years later, and In Europe by 1738 AD. The methods of smelting in China and India were most likely independently developed, while the method of smelting developed in Europe was likely derived by the Indian method.
The primary modern use for zinc is for coating iron and steel in order to prevent its corrosion, with nearly half of worldwide zinc production going towards that purpose. Approximately 20% of the world's zinc is used in the production of brass, where zinc is alloyed with copper in between ratios of 20-40% zinc. Of the remaining 30% of the global production of zinc, half is used in the production of zinc alloys, where zinc is combined with varying amounts of aluminum, and magnesium. The remaining zinc is used in various other industries from agriculture as a fertilizer and human consumption as a supplement.
Methods of extraction.
Zinc is mined both at the surface and at depth. Surface mining of zinc, typically used for oxide ores, while underground mining yields zinc sulfide ores. Some of the common methods of zinc mining are open pit mining, open stope, and cut and fill mining:
Open-pit mining: Surface mining involves the removal of waste rock from above an ore deposit before it can be extracted. Once the waste overburden is removed, ore and waste are then mined in parallel, primarily using track-mounted excavators and rubber-tired trucks. In smaller scale operations, front loaders may be used.
Open Stope mining: This is a method of underground mining where ore bodies are completely removed leaving sizeable caverns (stopes) within the mine. Open stope mining leaves these caverns with no additional bracing or external support. What is used to support the cavern walls, are random pillars of ore which have not removed.
Cut and Fill stoping: A method of underground mining which removes ore from below the deposit. The stope is then filled with waste rock to replace the mined out ore to support the stope walls, and to provide an elevated floor for the miners and equipment to further extract ore from the deposit.
Production.
Global mine production of zinc in 2019 was 12.9 million tonnes, a 0.9% increase from 2018, with the increase primarily attributed to increased output from zinc mines located in Australia and South Africa. In 2020 production of zinc is expected to rise 3.7% to 13.99 million tonnes, with the increase due to increased production of zinc by China and India.
In 2019 global demand for refined zinc exceeded supply and resulted in a deficit of 0.178 million tonnes, while in 2020 there is an expected surplus of 0.192 million tonnes.
Major zinc producing countries ranked by their output for 2019 are as follows:
Environmental impact.
Research conducted in the health of the benthic macroinvertebrate populations in the mining areas of southeastern Missouri, a US state, have yielded a wealth of information on the effect of zinc mining and its effect on its local environment. Fish and Crayfish populations in localities near mining sites have been observed to be much lower that other populations found in reference sites; with the crayfish possessing metal concentrations within their tissues at a much greater concentration than their reference counterparts. Other investigation into the effect of the health of mussel populations that reside near lead-zinc mining areas have found that the populations residing near mining areas possessed reduced biomass, and were less speciose than those found in their reference sites. Plant tissue have been reported to possess concentrations of metals 10-60% higher than reference. Macroinvertebrate assessments of localities immediately downstream of mining activity have observed a reduction in biotic condition 10-58% and with the locality possessing an impaired ability to support its populations when compared to other reference sites.
Benthic macro-invertebrates such as crayfish and mussels represent a pathway for biomagnification, where the concentration of noxious materials within organisms at higher trophic levels as a result of consuming contaminated prey items. In addition, benthic macroinvertebrate populations are frequently used as indicators of overall ecosystem health.
Assessment of soil samples from agricultural areas near a lead-zinc mining region in Guangxi, China have observed a "Serious pollution level" of zinc in the soils of the paddy fields relatively close to the mining area and a "Moderate pollution level" in the aerated fields relatively further away from the mining area. The research also indicated that as a result of their Nemerow synthetic index assessment, the region under study is not fit for agricultural purposes. Other investigation into the effect of zinc mining on agricultural soils in the Heilongjiang Province of china has found that the soils were "moderately contaminated" and a significant reduction in the population and diversity of the bacterial assemblages within the soils and reduced activity of soil enzymes. The activity of the bacteria and enzymes aid plant matter in the uptake of nutrients, decompose decaying matter, and other ecosystem interactions. Their reduction and impaired effectiveness result in poorer agricultural productivity.
Zinc mines.
The world's ten largest zinc producing mines (by tonnes of zinc) are:


Sarabauite

Sarabauite (sar-a-bau'-ite) is a red monoclinic sulfide mineral with the chemical formula: CaSb10O10S6.
Origin of name and type locality.
Sarabauite was first described in 1977 and named for its type locality, the Sarabau Mine (Lucky Hill mine) in Sarawak, Malaysia. It has also been reported from the Castelo Branco District of Portugal.
Crystallography.
Sarabauite is a monoclinic mineral. Monoclinic minerals have three crystallographic axes of unequal length. Two of the inter-axial angles are 90 degrees, while the third angle, between the optic axis and the plane containing the other two axes, is unequal to the other inter-axial angles. The monoclinic system includes crystal classes with a single twofold rotation axis and/or a mirror plane. Sarabauite is of the 2/m crystal class; it has a 2 fold rotation axis perpendicular to a mirror plane.
Optical properties.
Sarabauite is a biaxial negative mineral. All minerals of the monoclinic, orthorhombic, and triclinic systems are biaxial minerals. Unlike uniaxial minerals, biaxial crystals have 2 optic axes. This means they have two directions in which light can travel without birefringence, three principal axes, and accordingly three different indices of refraction. Biaxial minerals are further subdivided by optic sign. The optic angle, 2V, is the angle between the two optic axes that is split by the direction of the lowest index of refraction. If 2V is acute, the mineral is biaxial positive. In sarabauite crystals this angle is obtuse, thus the mineral is biaxial negative.
Geologic setting.
The mining region that contains the Sarabau Mine is part of the Sumatra Orogen. The region consists of a system of faults and joints through which hydrothermal fluids responsible for the mineralization of sarabauite rise into the overlying carbonate rock.
Mineral importance.
Sarabauite is a notable mineral for its antimony content, fluid inclusions, accompaniment of gold, and its two-stage hydrothermal formation process. The Sarabau Mine is a gold and antimony mine. Sarabauite can be found there in mineral veins through altered limestone containing quartz, wollastonite, calcite, stibnite and senarmonite. Both temperature and fluids are responsible for its formation through a two-stage hydrothermal mineralization process. First, wollastonite, diopside, and epidote form at temperatures over 400 °C. In the second stage, at over 377 °C, sarabauite and native gold develop. Calcite, stibnite, and senarmontite also form during the second stage as the mineralization cools further to 377–194 °C. Sarabauite's formation makes it the only hypogene antimony mineral whose fluid inclusions can be studied in thin section under normal light.


Haloxazolam

Haloxazolam (marketed in Japan under the brand name Somelin), is a drug which is a benzodiazepine derivative. It has similar hypnotic properties as the benzodiazepine drugs triazolam, temazepam, and flunitrazepam and as such is indicated for the treatment insomnia. A study in cats comparing estazolam and haloxazolam found that haloxazolam only affects gamma motor neurons, whereas estazolam affects both alpha and gamma motor neurons.


Lithium sulfite

Lithium sulfite, or lithium sulphite, is an ionic compound with the formula Li2SO3.


Heinrich Barbl

Heinrich Barbl (born 3 March 1900, Sarleinsbach, Austria; date of death unknown, not before 1965) was an Austrian-born SS-"Rottenführer". He participated in the T-4 euthanasia program in Nazi Germany and, after the invasion of Poland, in Operation Reinhard phase of the Holocaust.
Life.
Barbl was a tinsmith and plumber employed at the nitrogen works in Linz as a gas and waterpipe fitter. He joined the Nazi Party and the SS after the Anschluss of Austria. He was posted to Hartheim Euthanasia Centre at the onset of the Action T4 forced euthanasia program, and served at Grafeneck Euthanasia Centre as well. At both institutions he was a "Stanzer", stamping sheet metal with names of the dead to form nameplates. He would then attach them to small urns which he used to fill up with ash indistinctively. The urns were then sent to victims' relatives correctly named, but invariably with the wrong contents, as he filled them up with ashes shoveled indiscriminately from the adjacent crematorium.
Operation Reinhard.
In 1942 Barbl was posted to Bełżec extermination camp, the first of the Nazi German extermination camps created for the purpose of implementing the secretive Operation Reinhard. Barbl was often drunk on duty. Not particularly intelligent (which might also explain his low rank), he was frequently made fun of by his SS colleagues. Camp commandant, "SS-Hauptsturmführer" Gottlieb Hering did not allow him to participate in the execution of sick and elderly arrivals because, "he is so daft that he would shoot us, not the Jews". Consequently, he was the only person to be excused from this duty. However, no one was certain whether Barbl was genuinely stupid, or if he was merely acting a fool just to avoid more demanding tasks.
Barbl was not exempt from his commandant's punishments. He was frequently whipped by Christian Wirth for his drunken behavior, and for an unknown reason in the winter of 1942–43, Hering allegedly had him imprisoned in a cell for several days. Barbl was sent to Sobibor extermination camp by Wirth for his plumbing skills, to help fit the exhaust piping for the gas chambers. Barbl boasted that he had made the gas chambers look like neat shower rooms.
After the war Barbl was interrogated by Austrian police but was never put on trial. He submitted his testimony in Linz on September 6, 1965 at the Austrian Ministry of Internal Affairs. No further details are known of his life and death.


Sodium pyrosulfate

Sodium pyrosulfate is an inorganic compound with the chemical formula of Na2S2O7. It is a colorless salt.
Preparation.
Sodium pyrosulfate is obtained by the dehydration of sodium bisulfate:
Temperatures above 460 °C further decompose the compound, producing sodium sulfate and sulfur trioxide:
Applications.
Sodium pyrosulfate was used in analytical chemistry. Samples are fused with sodium pyrosulfate to ensure complete dissolution before a quantitative analysis.


Reactive magnesia

Reactive magnesia is also variously known as caustic calcined magnesia, caustic magnesia or CCM. The temperature of firing has a greater influence on reactivity than grind size as excess energy goes into lattice energy.
Crystalline magnesium oxide, or periclase, has a calculated lattice energy of 3795 kJ mol-1 which must be overcome for it to go into solution or for reaction to occur.
Reactive magnesia is essentially amorphous magnesia with low lattice energy and is made at low temperatures and finely ground.
CCM is manufactured worldwide. China supplies low-cost, low-purity (~90% MgO) CCM. Turkey (KUMAS) and Australia (Causmag) supply high-purity (~95% MgO) CCM. Low-purity CCM is used in industrial applications such as waste water treatment. High-purity CCM is used in pharmaceuticals, cattle feed supplement, fertilizer, adhesive, Tyre, Magnetic Material etc.


Methylammonium lead halide

Methylammonium lead halides (MALHs) are solid compounds with perovskite structure and a chemical formula of CH3NH3PbX3, where X = I, Br or Cl. They have potential applications in solar cells, lasers, light-emitting diodes, photodetectors, radiation detectors, scintillator, magneto-optical data storage and hydrogen production.
Properties and synthesis.
In the CH3NH3PbX3 cubic crystal structure the methylammonium cation (CH3NH3+) is surrounded by PbX6 octahedra. The X ions are not fixed and can migrate through the crystal with an activation energy of 0.6 eV; the migration is vacancy assisted. The methylammonium cations can rotate within their cages. At room temperature the ions have the CN axis aligned towards the face directions of the unit cells and the molecules randomly change to another of the six face directions on a 3 ps time scale.
The solubility of MALHs strongly decreases with increased temperature: from 0.8 g/mL at 20 °C to 0.3 g/mL at 80 °C for CH3NH3PbI3 in dimethylformamide. This property is used in the growth of MALH single crystals and films from solution, using a mixture of CH3NH3X and PbX2 powders as the precursor. The growth rates are 3–20 mm3/hour for CH3NH3PbI3 and reach 38 mm3/hour for CH3NH3PbBr3 crystals.
The resulting crystals are metastable and dissolve in the growth solution when cooled to room temperature. They have bandgaps of 2.18 eV for CH3NH3PbBr3 and 1.51 eV for CH3NH3PbI3, while their respective carrier mobilities are 24 and 67 cm2/(V·s). Their thermal conductivity is exceptionally low, ~0.5 W/(K·m) at room temperature for CH3NH3PbI3.
Photodecomposition and thermal decomposition of CH3NH3PbX3.
Initially, a proposed decomposition pathway mechanism for CH3NH3PbI3 in presence of water
releasing CH3NH2 and HI gases was broadly adopted by researchers in perovskite solar cell. Later, it was found that the major gases released during high temperature (> 360 °C) thermal degradation of CH3NH3PbI3 are methyl iodide (CH3I) and ammonia (NH3).
In 2017, it has been inferred using in situ XPS measurements that in the presence of water vapour, CH3NH3I salt can not be a product of the degradation of CH3NH3PbI3 perovskite.
Similar high temperature degradation reaction has been confirmed for CH3NH3PbBr3
Furthermore, high resolution mass spectrometry measurements at low temperature conditions (< 100 °C) compatible with photovoltaic operation found that CH3NH3PbI3 undergoes reversible,
and irreversible chemical decomposition reactions under vacuum when illumination or heat pulses are applied.
Recently, a method to quantify the intrinsic chemical stability of arbitrarily mixed hybrid halide perovskites has been proposed.
Applications.
MALHs have potential applications in solar cells, lasers, light-emitting diodes, photodetectors, radiation detectors, scintillator and hydrogen production. The power conversion efficiency of MALH solar cells exceeds 19%.


Sodium tetrasulfide

Sodium tetrasulfide is an inorganic compound with the formula Na2S4. It is a yellow-orange solid that dissolved with hydrolysis in water. They are precursors to some specialty polymers and intermediates in prototypes of the sodium-sulfur battery.
Synthesis and structure.
It is produced through the reaction between elemental sulfur and sodium hydrosulfide in alcoholic solution:
The polysulfide anions adopt zig-zag chains of sulfur atoms. The S-S distances are about 2.05 Å and the S-S-S-S dihedral angles are around 90°.
Reactions and applications.
Upon treatment with acid, it is converted to hydrogen sulfide and elemental sulfur. Treatment with alkylating agents gives organic polysulfides. In one commercial application, it is used to produce the cross-linking agent bis(triethoxysilylpropyl)tetrasulfide:
Sometimes as a mixture with other polysulfides, sodium tetrasulfide is used to produce the polymer called thiokol. The reaction involves alkylation with ethylene chloride:
These materials, which have the approximate formula (C2H4)Sx]n (x ~ 4), are highly resistant to degradation by solvents and acids.


List of countries by chromium production

Chromium is a chemical element that is designated by the symbol Cr and has an atomic number of 24. It is usually found as the mineral chromite, from which ferrochrome is produced in a smelting process.
List of Mining Countries.
This is a list of countries by chromium ore mining in 2019, based on United States Geological Survey accessed in 2020.


Bronopol

Bronopol (INN; chemical name 2-bromo-2-nitro-1,3-propanediol) is an organic compound that is used as an antimicrobial. It is a white solid although commercial samples appear yellow.
The first reported synthesis of bronopol was in 1897.
Bronopol was invented by The Boots Company PLC in the early 1960s and first applications were as a preservative for pharmaceuticals. Due to its low mammalian toxicity at in-use levels and high activity against bacteria, especially Gram-negative species, bronopol became popular as a preservative in many consumer products such as shampoos and cosmetics. It was subsequently adopted as an antimicrobial in other industrial environments such as paper mills, oil exploration, and production facilities, as well as cooling water disinfection plants.
Production.
Bronopol is produced by the bromination of di(hydroxymethyl)nitromethane, which is derived from nitromethane by a nitroaldol reaction. World production increased from the tens of tonnes in the late 1970s to current estimates in excess of 5,000 tonnes. Manufacturing today is the business of low cost producers, mainly in China.
Applications.
Bronopol is used in consumer products as an effective preservative agent, as well as a wide variety of industrial applications (almost any industrial water system is a potential environment for bacterial growth, leading to slime and corrosion problems - in many of these systems bronopol can be a highly effective treatment).
The use of bronopol in personal care products (cosmetics, toiletries) has declined since the late 1980s due to the potential formation of nitrosamines. While bronopol is not in itself a nitrosating agent, under conditions where it decomposes "(alkaline solution and/or elevated temperatures)" it can liberate nitrite and low levels of formaldehyde and these decomposition products can react with any contaminant secondary amines or amides in a personal care formulation to produce significant levels of nitrosamines (due to the toxicity of these substances, the term 'significant' means levels as low as 10s of parts per billion).
Manufacturers of personal care products are therefore instructed by regulatory authorities to avoid the formation of nitrosamines which might mean removing amines or amides from the formulation, removing bronopol from a formulation, or using nitrosamine inhibitors.
Bronopol has been restricted for use in cosmetics in Canada.
Physical and chemical properties.
Appearance.
Bronopol is supplied as crystals or crystalline powder, which may vary from white to pale yellow in colour depending on the grade. The yellow coloration is due to chelation of iron during the manufacturing process.
Melting point.
As a pure material, bronopol has a melting point of about 130 °C. However, due to its polymorphic characteristics, bronopol undergoes a lattice rearrangement at 100 to 105 °C and this can often be wrongly interpreted as the melting point.
At temperatures above 140 °C, bronopol decomposes exothermically releasing hydrogen bromide and oxides of nitrogen.
Solubility.
Bronopol is readily soluble in water; the dissolution process is endothermic. Solutions containing up to 28% w/v are possible at ambient temperature.
Bronopol is poorly soluble in non-polar solvents but shows a high affinity for polar organic solvents.
Partition coefficient.
Study of the solubility data shows that bronopol has a high affinity for polar rather than non-polar environments. In two-phase systems, bronopol partitions preferentially into the polar (usually aqueous) phase.
Stability in aqueous solution.
In aqueous solutions, bronopol is most stable when the pH of the system is on the acid side of neutral. Temperature also has a significant effect on stability in alkaline systems.
Degradation.
Under extreme alkaline conditions, bronopol decomposes in aqueous solution and very low levels of formaldehyde are produced. Liberated formaldehyde is not responsible for the biological activity associated with bronopol. Other decomposition products detected after bronopol breakdown are bromide ion, nitrite ion, bromonitroethanol and 2-hydroxymethyl-2-nitropropane-1,3-diol.
Allergy.
In 2005–2006, it was the 15th-most-prevalent allergen in patch tests (3.4%) of people with suspected allergic contact dermatitis. It is used as a substitute for formaldehyde, a disinfectant and preservative, in solvents. It is prevalent in skin and personal care products and topical medications.


Trichlorophenylsilane

Trichlorophenylsilane is a compound with formula Si(C6H5)Cl3.
Similarly to other alkylchlorosilanes, trichlorophenylsilane is a possible precursor to silicone. It hydrolyses in water to give HCl and phenylsilantriol, with the latter condensating to a polymeric substance.


Ciprofibrate

Ciprofibrate is a fibrate that was developed as a lipid-lowering agent.
It was patented in 1972 and approved for medical use in 1985.