2 2 azino bis

 2′-AZINO-BIS (3-ETHYLBENZTHIAZOLINE -6

2 2′-AZINO-BIS (3-ETHYLBENZTHIAZOLINE -6
CAS No.
Chemical Name: 2 2′-AZINO-BIS (3-ETHYLBENZTHIAZOLINE -6
Synonyms: 2 2′-AZINO-BIS (3-ETHYLBENZTHIAZOLINE -6
CBNumber: CB6552899
Molecular Formula:
Formula Weight: 0
MOL File: Mol file
2 2′-AZINO-BIS (3-ETHYLBENZTHIAZOLINE -6 Property
storage temp. : 2-8°C
Safety

2 2′-AZINO-BIS (3-ETHYLBENZTHIAZOLINE -6 Chemical Properties,Usage,Production

2 2′-AZINO-BIS (3-ETHYLBENZTHIAZOLINE -6 Preparation Products And Raw materials
Raw materials
Preparation Products

fluoride suppliers 16984-48-8

Fluoride is the anion F, the reduced form of fluorine when as an ion and when bonded to another element. Inorganic fluorine containing compounds are called fluorides. Fluoride, like other halides, is a monovalent ion (−1 charge). Its compounds often have properties that are distinct relative to other halides. Structurally, and to some extent chemically, the fluoride ion resembles the hydroxide ion.

Fluoride
Identifiers
CAS number 16984-48-8 Yes
PubChem 28179
ChemSpider 26214 Yes
KEGG C00742 Yes
MeSH Fluoride
ChEBI CHEBI:17051
ChEMBL CHEMBL1362 Yes
Gmelin Reference 14905
Jmol-3D images Image 1
Properties
Molecular formula F
Molar mass 18.9984032 g mol−1
Thermochemistry
Std enthalpy of
formationΔfHo298
−333 kJ mol−1
Standard molar
entropy So298
145.58 J/mol K (gaseous)[2]
Related compounds
Other anions
  • Iodide
  • Bromide
  • Chloride

Occurrence

The mineral fluorite, a common mineral and chief source of fluoride for commercial applications.[3][4]

Solutions of inorganic fluorides in water contain F and bifluoride HF−
2.[5] Few inorganic fluorides are soluble in water without undergoing significant hydrolysis. In terms of its reactivity, fluoride differs significantly from chloride and other halides, and is more strongly solvated due to its smaller radius/charge ratio. Its closest chemical relative is hydroxide. When relatively unsolvated, fluoride anions are called “naked”. Naked fluoride is a very strong lewis base.[6] The presence of fluoride and its compounds can be detected by 19F NMR spectroscopy.

[edit]Natural occurrence

Many fluoride minerals are known, but of paramount commercial importance are fluorite and fluorapatite.[3]

Fluoride is usually found naturally in low concentration in drinking water and foods. The concentration in seawater averages 1.3 parts per million(ppm). Fresh water supplies generally contain between 0.01–0.3 ppm, whereas the ocean contains between 1.2 and 1.5 ppm.[7] In some locations, the fresh water contains dangerously high levels of fluoride, leading to serious health problems.

[edit]Applications

Fluorides are pervasive in modern technology. Hydrofluoric acid is the fluoride synthesized on the largest scale. It is produced by treating fluoride minerals with sulfuric acid. Hydrofluoric acid and its anhydrous form hydrogen fluoride are used in the production of fluorocarbons and aluminium fluorides. Hydrofluoric acid has a variety of specialized applications, including its ability to dissolve glass.[3]

[edit]Inorganic chemicals

Fluoride salts are used in the manufacture of many inorganic chemicals, many of which contain fluoride covalently bonded to the metal or nonmetal in question. Some examples of these are:

  • Cryolite (Na3AlF6) is a pesticide that can leave fluoride on agricultural commodities.[8][9] Cryolite was originally utilized in the preparation ofaluminium.
  • Sulfuryl fluoride (SO2F2) is used as a pesticide and fumigant on agricultural crops. In 2010, the United States Environmental Protection Agencyproposed to withdraw the use of sulfuryl fluoride on food. Sulfuryl fluoride releases fluoride when metabolized.[10][11]
  • Sulfur hexafluoride is an inert, nontoxic insulator gas that is used in electrical transformers and as a tracer gas in indoor air quality investigations.
  • Uranium hexafluoride, although not ionic, is prepared from fluoride reagents. It is utilized in the separation of isotopes of uranium between thefissile isotope U-235 and the non-fissile isotope U-238 in preparation of nuclear reactor fuel and atomic bombs. This is due to the volatility of fluoridesof uranium.

[edit]Organic chemicals

Main article: Organofluorine chemistry

Fluoride reagents are significant in synthetic organic chemistry. Organofluorine chemistry has produced many useful compounds over the last 50 years. Included in this area are polytetrafluorethylene (Teflon), polychlorotrifluoroethylene (moisture barriers), efavirenz (pharmaceutical used for treatment of HIV), fluoxetine (an antidepressant), 5-fluorouracil (an anticancer drug), hydrochlorofluorocarbons and hydrofluorcarbons (refrigerants, blowing agents and propellants).

Due to the affinity of silicon for fluoride, and the ability of silicon to expand its coordination number, silyl ether protecting groups can be easily removed by the fluoride sources such as sodium fluoride and tetra-n-butylammonium fluoride (TBAF). This is quite useful for organic synthesis and the production of fine chemicals. The Si-F linkage is one of the strongest single bonds. In contrast, other silyl halides are easily hydrolyzed.

[edit]Cavity prevention

Main articles: fluoride therapy and water fluoridation

Fluoride-containing compounds are used in topical and systemic fluoride therapy for preventing tooth decay. They are used for water fluoridation and in many products associated with oral hygiene.[12] Originally, sodium fluoride was used to fluoridate water; hexafluorosilicic acid (H2SiF6) and its salt sodium hexafluorosilicate (Na2SiF6) are more commonly used additives, especially in the United States. The fluoridation of water is known to prevent tooth decay[13][14] and is considered by the U.S. Centers for Disease Control and Prevention as “one of 10 great public health achievements of the 20th century”.[15][16] In some countries where large, centralized water systems are uncommon, fluoride is delivered to the populace by fluoridating table salt. Fluoridation of water has its critics (see Water fluoridation controversy).[17]

Structure ofhalothane.

[edit]Biomedical applications

Positron emission tomography is commonly carried out using fluoride-containing pharmaceuticals such as fluorodeoxyglucose, which is labelled with theradioactive isotope fluorine-18, which emits positrons when it decays into 18O.

Numerous drugs contain fluorine including antipsychotics such as fluphenazine, HIV protease inhibitors such as tipranavir, antibiotics such as ofloxacin andtrovafloxacin, and anesthetics such as halothane.[18] Fluorine is incorporated in the drug structures to reduce drug metabolism, as the strong C-F bond resists deactivation in the liver by cytochrome P450 oxidases.[19]

Fluoride salts are commonly used to inhibit the activity of phosphatases, such as serine/threonine phosphatases.[20] Fluoride mimics the nucleophilic hydroxyl ion in these enzymes’ active sites.[21] Beryllium fluoride and aluminium fluoride are also used as phosphatase inhibitors, since these compounds are structural mimics of thephosphate group and can act as analogues of the transition state of the reaction.[22][23]

[edit]Toxicology

Main article: Fluoride toxicity

Reaction of the irreversible inhibitordiisopropylfluorophosphatewith a serine protease

Fluoride-containing compounds are so diverse that it is not possible to generalize on their toxicity, which depends on their reactivity and structure, and in the case of salts, their solubility and ability to release fluoride ions.

Soluble fluoride salts, of which sodium fluoride is the most common, are mildly toxic but have resulted in both accidental and suicidal deaths from acute poisoning.[3] While the minimum fatal dose in humans is not known, the lethal dose for most adult humans is estimated at 5 to 10 g (which is equivalent to 32 to 64 mg/kg elemental fluoride/kg body weight).[24][25][26] However, a case of a fatal poisoning of an adult with 4 grams of sodium fluoride is documented,[27]while a dose of 120 g sodium fluoride has been survived.[28] A toxic dose that may lead to adverse health effects is estimated at 3 to 5 mg/kg of elemental fluoride.[29] For Sodium fluorosilicate (Na2SiF6), the median lethal dose (LD50) orally in rats is 0.125 g/kg, corresponding to 12.5 g for a 100 kg adult.[30] The fatal period ranges from 5 min to 12 hours.[27] The mechanism of toxicity involves the combination of the fluoride anion with the calcium ions in the blood to form insoluble calcium fluoride, resulting in hypocalcemia; calcium is indispensable for the function of the nervous system, and the condition can be fatal. Treatment may involve oral administration of dilute calcium hydroxide or calcium chloride to prevent further absorption, and injection of calcium gluconateto increase the calcium levels in the blood.[27] Hydrogen fluoride is more dangerous than salts such as NaF because it is corrosive and volatile, and can result in fatal exposure through inhalation or upon contact with the skin; calcium gluconate gel is the usual antidote.[31]

In the higher doses used to treat osteoporosis, sodium fluoride can cause pain in the legs and incomplete stress fractures when the doses are too high; it also irritates the stomach, sometimes so severely as to cause ulcers. Slow-release and enteric-coated versions of sodium fluoride do not have gastric side effects in any significant way, and have milder and less frequent complications in the bones.[32] In the lower doses used for water fluoridation, the only clear adverse effect is dental fluorosis, which can alter the appearance of children’s teeth during tooth development; this is mostly mild and is unlikely to represent any real effect on aesthetic appearance or on public health.[33]

krypton fluoride 13773-81-4

Krypton difluoride, KrF2 is a chemical compound of krypton and fluorine. It was the first compound of krypton discovered.[2] It is a volatile, colourless solid. The structure of the KrF2 molecule is linear, with Kr−F distances of 188.9 pm. It reacts with strong Lewis acids to form salts of the KrF+and Kr2F3+ cations.

Krypton difluoride
Identifiers
CAS number 13773-81-4 
PubChem 83721
ChemSpider 75543 Yes
Jmol-3D images Image 1
Properties
Molecular formula F2Kr
Molar mass 121.79 g mol−1
Appearance Colourless crystals (solid)
Density 3.24 g cm−3 (solid)
Solubility in water Reacts
Structure
Crystal structure Body-centered tetragonal[1]
Space group P42/mnm, No. 136
Lattice constant a = 0.4585 nm,c = 0.5827 nm
Molecular shape Linear
Dipole moment 0 D
Related compounds
Related compounds Xenon difluoride

Synthesis

Krypton difluoride can be synthesized using many different methods including electrical discharge, photochemical, hot wire, and proton bombardment. It can also be prepared by irradiating krypton with ultraviolet rays in a fluorine-argon gas mixture at liquid helium temperature. The product can be stored at −78 °C without decomposition.[4]

[edit]Electrical discharge

The first method used to make krypton difluoride and the only one ever reported to produce krypton tetrafluoride (although the identification of krypton tetrafluoride was later shown to be mistaken) was the electrical discharge method. The electrical discharge method involves having 1:1 to 2:1 mixtures of F2 to Kr at a pressure of 40 to 60 torr and then arcing large amounts of energy between it. Rates of almost 0.25 g/h can be achieved. The problem with this method is that it is unreliable with respect to yield.[3][5]

[edit]Proton bombardment

Using proton bombardment for the production of KrF2 has a maximum production rate of about 1 g/h. This is achieved by bombarding mixtures of Kr and F2 with a proton beam that is operating at an energy level of 10 MeV and at a temperature of about 133 K. It is a fast method of producing relatively large amounts of KrF2, but requires a source of α-particles which usually would come from a cyclotron.[3][6]

[edit]Photochemical

The photochemical process for the production of KrF2 involves the use of UV light and can produce under ideal circumstances 1.22 g/h. The ideal wavelengths to use are in the range of 303–313 nm. It is important to note that harder UV radiation is detrimental to the production of KrF2. In order to avoid the harder wavelengths, simply using Pyrex glass or Vycor or quartz will significantly increase yield because they all block harder UV light. In a series of experiments performed by S. A Kinkead et al., it was shown that a quartz insert (UV cut off of 170 nm) produced on average 158 mg/h, Vycor 7913 (UV cut off of 210 nm) produced on average 204 mg/h and Pyrex 7740 (UV cut off of 280 nm) produced on average 507 mg/h. It is clear from these results that higher energy ultra violet light reduces the yield significantly. The ideal circumstances for the production KrF2 by a photochemical process appear to occur when krypton is a solid and fluorine is a liquid which occur at 77K. The biggest problem with this method is that it requires the handling of liquid F2 and the potential of it being released if it becomes over pressurized.[3][5]

[edit]Hot wire

The hot wire method for the production of KrF2 involves having the krypton in a solid state with a hot wire running a few centimeters away from it as fluorine gas is then run past the wire. The wire has a large current, causing it to reach temperatures around 680 °C. This causes the fluorine gas to split into its radicals which then can react with the solid krypton. Under ideal conditions, it has been known to reach a maximum yield of 6 g/h. In order to achieve optimal yields the gap between the wire and the solid krypton should be 1 cm, giving rise to a temperature gradient of about 900 °C/cm. The only major downside to this method is the amount of electricity that has to be passed through the wire thus making it dangerous if not properly set up.[3][5]

[edit]Structure

α-KrF2

Krypton difluoride can exist in one of two possible crystallographic morphologies: α-phase and β-phase. β-KrF2 generally exists at above −80 °C, while the α-KrF2 is more stable at lower temperatures.[3] The unit cell of α-KrF2 is body-centred tetragonal.

[edit]Chemistry

Krypton difluoride is primarily a powerful oxidising and fluorinating agent. It can oxidise gold to its highest-known oxidation state, +5:[7]

7 KrF2 (g) + 2 Au (s) → 2 KrF+
AuF−
6 (s) + 5 Kr (g)

KrF+
AuF−
6 decomposes at 60°C into gold(V) fluoride and krypton and fluorine gases:[8]

KrF+
AuF−
6 → AuF5 (s) + Kr (g) + F2 (g)

KrF2 can also directly oxidise xenon to xenon hexafluoride:[7]

3 KrF2 + Xe → XeF6 + 3 Kr

KrF2 is used to synthesize the highly reactive BrF+
6 cation.[4] KrF2 reacts with SbF5 to form the salt KrF+
SbF−
6; the KrF+
cation is capable of oxidising both BrF5 and ClF5 to BrF+
6 and ClF+
6, respectively.[9]

soy bean products

The soybean (US) or soya bean (UK) (Glycine max)[2] is a species of legume native to East Asia, widely grown for its edible bean which has numerous uses. The plant is classed as an oilseed rather than a pulse by the UN Food and Agricultural Organization (FAO).

Scientific classification
Kingdom: Plantae
(unranked): Angiosperms
(unranked): Eudicots
(unranked): Rosids
Order: Fabales
Family: Fabaceae
Subfamily: Faboideae
Genus: Glycine
Species: G. max

Fat-free (defatted) soybean meal is a significant and cheap source of protein for animal feeds and many prepackaged meals;[3] soy vegetable oil is another product of processing the soybean crop. For example, soybean products such as textured vegetable protein (TVP) are ingredients in many meat and dairy analogues.[4] Soybeans produce significantly more protein per acre than most other uses of land.[5]

Traditional nonfermented food uses of soybeans include soy milk, and from the latter tofu and tofu skin. Fermented foods include soy sauce,fermented bean paste, natto, and tempeh, among others. The oil is used in many industrial applications. The main producers of soy are the United States (35%), Brazil (27%), Argentina (19%), China (6%) and India (4%).[6][7] Today, the United States is also the world’s largest consumer of soybeans, with an average annual consumption of 45,313 TMT.[8] The beans contain significant amounts of phytic acid, alpha-linolenic acid, and isoflavones

Name

The plant is sometimes referred to as greater bean (大豆 – Chinese dàdòu and Japanese daizu). Both the immature soybean and its dish are called edamame in Japan,[9][10] but in English, edamame refers only to a specific dish.[citation needed]

[edit]Classification

Varieties of soybeans are used for many purposes.

The genus name Glycine was originally introduced by Carl Linnaeus (1737) in his first edition of Genera Plantarum. The word glycine is derived from the Greek – glykys (sweet) and likely refers to the sweetness of the pear-shaped (apios in Greek) edible tubers produced by the native North American twining or climbing herbaceous yambean legume, Glycine apios, now known as Apios americana. The cultivated soybean first appeared in Species Plantarum, by Linnaeus, under the name Phaseolus max L. The combination Glycine max (L.) Merr., as proposed by Merrill in 1917, has become the valid name for this useful plant.

The genus Glycine Willd. is divided into two subgenera, Glycine and Soja. The subgenus Soja (Moench) F.J. Herm. includes the cultivated soybean,Glycine max (L.) Merr., and the wild soybean, Glycine soja Sieb. & Zucc. Both species are annuals. Glycine soja is the wild ancestor of Glycine max, and grows wild in China, Japan, Korea, Taiwan and Russia.[11] The subgenus Glycine consists of at least 16 wild perennial species: for example, Glycine canescens F.J. Herm. and G. tomentella Hayata, both found in Australia and Papua New Guinea.[12][13]

Like some other crops of long domestication, the relationship of the modern soybean to wild-growing species can no longer be traced with any degree of certainty. It is a cultural variety with a very large number of cultivars.

[edit]Description and physical characteristics

Soy varies in growth and habit. The height of the plant varies from less than 0.2 to 2.0 m (0.66 to 6.6 ft).

The pods, stems, and leaves are covered with fine brown or gray hairs. The leaves are trifoliolate, having three to four leaflets per leaf, and the leaflets are 6–15 cm (2.4–5.9 in) long and 2–7 cm (0.79–2.8 in) broad. The leaves fall before the seeds are mature. The inconspicuous, self-fertile flowers are borne in the axil of the leaf and are white, pink or purple.

Small, purple soybean flowers

The fruit is a hairy pod that grows in clusters of three to five, each pod is 3–8 cm long (1–3 in) and usually contains two to four (rarely more) seeds 5–11 mm in diameter.

Soybeans occur in various sizes, and in many hull or seed coat colors, including black, brown, blue, yellow, green and mottled. The hull of the mature bean is hard, water-resistant, and protects the cotyledon and hypocotyl (or “germ”) from damage. If the seed coat is cracked, the seed will notgerminate. The scar, visible on the seed coat, is called the hilum (colors include black, brown, buff, gray and yellow) and at one end of the hilum is the micropyle, or small opening in the seed coat which can allow the absorption of water for sprouting.

Remarkably, seeds such as soybeans containing very high levels of protein can undergo desiccation, yet survive and revive after water absorption. A. Carl Leopold, son of Aldo Leopold, began studying this capability at the Boyce Thompson Institute for Plant Research at Cornell University in the mid 1980s. He found soybeans and corn to have a range of soluble carbohydrates protecting the seed’s cell viability.[14] Patents were awarded to him in the early 1990s on techniques for protecting “biological membranes” and proteins in the dry state. Compare to tardigrades.

[edit]Nitrogen-fixing abilty

Many legumes (alfalfa, clover, peas, beans, lentils, soybeans, peanuts and others) contain symbiotic bacteria called Rhizobia within nodules of their root systems. These bacteria have the special ability of fixing nitrogen from atmospheric, molecular nitrogen (N2) into ammonia (NH3).[15] The chemical reaction is:

N2 + 8 H+ + 8 e → 2 NH3 + H2

Ammonia is then converted to another form, ammonium (NH4+), usable by (some) plants by the following reaction:

NH3 + H+ → NH4+

This arrangement means that the root nodules are sources of nitrogen for legumes, making them relatively rich in plant proteins.

[edit]Chemical composition of the seed

Soybean, mature seeds, raw
Nutritional value per 100 g (3.5 oz)
Energy 1,866 kJ (446 kcal)
Carbohydrates 30.16 g
- Sugars 7.33 g
- Dietary fiber 9.3 g
Fat 19.94 g
- saturated 2.884 g
- monounsaturated 4.404 g
- polyunsaturated 11.255 g
Protein 36.49 g
- Tryptophan 0.591 g
- Threonine 1.766 g
- Isoleucine 1.971 g
- Leucine 3.309 g
- Lysine 2.706 g
- Methionine 0.547 g
- Cystine 0.655 g
- Phenylalanine 2.122 g
- Tyrosine 1.539 g
- Valine 2.029 g
- Arginine 3.153 g
- Histidine 1.097 g
- Alanine 1.915 g
- Aspartic acid 5.112 g
- Glutamic acid 7.874 g
- Glycine 1.880 g
- Proline 2.379 g
- Serine 2.357 g
Water 8.54 g
Vitamin A equiv. 1 μg (0%)
Vitamin B6 0.377 mg (29%)
Vitamin B12 0 μg (0%)
Choline 115.9 mg (24%)
Vitamin C 6.0 mg (7%)
Vitamin K 47 μg (45%)
Calcium 277 mg (28%)
Iron 15.70 mg (121%)
Magnesium 280 mg (79%)
Phosphorus 704 mg (101%)
Potassium 1797 mg (38%)
Sodium 2 mg (0%)
Zinc 4.89 mg (51%)
Percentages are relative to
US recommendations for adults.
Source: USDA Nutrient Database

Together, soybean oil and protein content account for about 60% of dry soybeans by weight (protein at 40% and oil at 20%). The remainder consists of 35% carbohydrate and about 5% ash. Soybean cultivars comprise approximately 8% seed coat or hull, 90% cotyledons and 2%hypocotyl axis or germ.

Most soy protein is a relatively heat-stable storage protein. This heat stability enables soy food products requiring high temperature cooking, such as tofu, soy milk and textured vegetable protein (soy flour) to be made.

The principal soluble carbohydrates of mature soybeans are the disaccharide sucrose (range 2.5–8.2%), the trisaccharide raffinose (0.1–1.0%) composed of one sucrose molecule connected to one molecule of galactose, and the tetrasaccharide stachyose (1.4 to 4.1%) composed of one sucrose connected to two molecules of galactose. While the oligosaccharides raffinose and stachyose protect the viability of the soybean seed from desiccation (see above section on physical characteristics) they are not digestible sugars, so contribute to flatulence and abdominal discomfort in humans and other monogastric animals, comparable to the disaccharide trehalose. Undigested oligosaccharides are broken down in the intestine by native microbes, producing gases such as carbon dioxide, hydrogen, and methane.

Since soluble soy carbohydrates are found in the whey and are broken down during fermentation, soy concentrate, soy protein isolates, tofu, soy sauce, and sprouted soybeans are without flatus activity. On the other hand, there may be some beneficial effects to ingesting oligosaccharides such as raffinose and stachyose, namely, encouraging indigenous bifidobacteria in the colon against putrefactive bacteria.

The insoluble carbohydrates in soybeans consist of the complex polysaccharides cellulose, hemicellulose, and pectin. The majority of soybean carbohydrates can be classed as belonging to dietary fiber.

Within soybean oil or the lipid portion of the seed is contained the phytosterols: stigmasterol (17–21%), sitosterol(53–56%) and campesterol (20–23%) accounting for 2.5% of the lipid fraction.

Saponins, a class of natural surfactants (soaps), are sterols that are present naturally in a wide variety of food-plants: vegetables, legumes, and cereals–ranging from beans and spinach to tomatoes, potatoes and oats. Whole soybeans contain from 0.17 to 6.16% saponins, 0.35 to 2.3% in defatted soy flour and 0.06 to 1.9% in tofu. Legumes such as soybean and chickpeas are the major source of saponins in the human diet. Sources of non-dietary saponins include alfalfa, sunflower, herbs and barbasco. Recent studies have shown that saponins are potential functional food ingredients because of their physiological properties.[16]

Soy contains isoflavones like genistein and daidzein. It also contains glycitein, an O-methylated isoflavone which accounts for 5–10% of the total isoflavones in soy food products. Glycitein is a phytoestrogen with weak estrogenic activity, comparable to that of the other soy isoflavones.[17]

[edit]Nutrition

Further information: Soy protein

For human consumption, soybeans must be cooked with “wet” heat to destroy the trypsin inhibitors (serine protease inhibitors). Raw soybeans, including the immature green form, are toxic to humans, swine, chickens, and in fact, all monogastric animals.[18]

Soybeans are considered by many agencies to be a source of complete protein.[19] A complete protein is one that contains significant amounts of all the essential amino acids that must be provided to the human body because of the body’s inability to synthesize them. For this reason, soy is a good source of protein, amongst many others, for vegetarians and vegans or for people who want to reduce the amount of meat they eat. According to the US Food and Drug Administration:

Soy protein products can be good substitutes for animal products because, unlike some other beans, soy offers a ‘complete’ protein profile. … Soy protein products can replace animal-based foods—which also have complete proteins but tend to contain more fat, especially saturated fat—without requiring major adjustments elsewhere in the diet.[19]

The gold standard for measuring protein quality, since 1990, is the Protein Digestibility Corrected Amino Acid Score (PDCAAS) and by this criterion soy protein is the nutritional equivalent of meat, eggs, and casein for human growth and health. Soybean protein isolate has a biological valueof 74, whole soybeans 96, soybean milk 91, and eggs 97.[20]

Soy protein is essentially identical to the protein of other legume seeds and pulses.[21][22][23] Moreover, soybeans can produce at least twice as much protein per acre than any other major vegetable or grain crop besides hemp, five to 10 times more protein per acre than land set aside for grazing animals to make milk, and up to 15 times more protein per acre than land set aside for meat production.[5]

Consumption of soy may also reduce the risk of colon cancer, possibly due to the presence of sphingolipids.[24]

[edit]Comparison of Soybean to other major staple foods

The following table shows the nutrient content of green soybean and other major staple foods, each in respective raw form. Raw staples, however, aren’t edible and can not be digested. These must be sprouted, or prepared and cooked for human consumption. In sprouted and cooked form, the relative nutritional and anti-nutritional contents of each of these grains is remarkably different from that of raw form of these grains reported in this table. The nutritional value of soybean and each cooked staple depends on the pre-processing and the method of cooking: boiling, frying, roasting, baking, e

dodecyl sulfate de sodium 151-21-3

Sodium acetate, CH3COONa, also abbreviated NaOAc,[1] also sodium ethanoate, is the sodium salt of acetic acid. This colourless salt has a wide range of uses.

 

odium lauryl sulfate
Identifiers
CAS number 151-21-3 Yes
PubChem 3423265
ChemSpider 8677 Yes
DrugBank DB00815
ChEBI CHEBI:8984 Yes
ChEMBL CHEMBL23393 Yes
ATC code A06AG11
Jmol-3D images Image 1
Properties
Molecular formula NaC12H25SO4
Molar mass 288.372 g/mol
Appearance white or cream-colored solid
Odor odorless
Density 1.01 g/cm³
Melting point 206 °C, 479 K, 403 °F
Refractive index (nD) 1.461
Hazards
LD50 1288 mg/Kg (rat, oral)

Applications

SDS is synthesized by treating lauryl alcohol with sulfur trioxide gas, oleum, or chlorosulfuric acid to produce hydrogen lauryl sulfate. The industrially practiced method typically uses sulfur trioxide gas. The resulting product is then neutralized through the addition of sodium hydroxide or sodium carbonate. Lauryl alcohol is in turn usually derived from either coconut or palm kernel oil by hydrolysis, which liberates their fatty acids, followed by hydrogenation.

Due to this synthesis method, commercial samples of SDS are often a mixture of other alkyl sulfates, dodecyl sulfate being the main component.[3]

SDS is available commercially in powder and pellet forms. It seems the pellet form dissolves faster than the powder form in water.[4]

[edit]Applications

SDS is mainly used in detergents for laundry with many cleaning applications.[5] SDS is a highly effective surfactant and is used in any task requiring the removal of oily stains and residues. For example, it is found in higher concentrations with industrial products including engine degreasers, floor cleaners, and car wash soaps. It is found in toothpastes, shampoos, shaving foams, and bubble bath formulations in part for its thickening effect and its ability to create a lather.[6]

[edit]Laboratory applications

Bottle of solution of sodium dodecyl sulfate for use in the laboratory.

It can be used to aid in lysing cells during DNA extraction and for unraveling proteins in SDS-PAGE. Sodium lauryl sulfate, in science referred to as sodium dodecyl sulfate (SDS) or Duponol, is commonly used in preparing proteins forelectrophoresis in the SDS-PAGE technique.[7] This compound works by disrupting non-covalent bonds in the proteins, denaturing them, and causing the molecules to lose their native shape (conformation).

This new negative charge is significantly greater than the original charge of that protein. The electrostatic repulsionthat is created by binding of SDS causes proteins to unfold into a rod-like shape thereby eliminating differences in shape as a factor for separation in the gel.

Sodium lauryl sulfate is probably the most researched anionic surfactant compound. Like all detergent surfactants, sodium lauryl sulfate removes oils from the skin, and can cause skin and eye irritation. The critical micelle concentration (CMC) in pure water at 25°C is 0.0082 M,[8] and the aggregation number at this concentration is usually considered to be about 62.[9] The micelle ionization fraction (α) is around 0.3 (or 30%).[10]

Aqueous solutions of SDS are also popular for dispersing or suspending nanotubes.

Sodium lauryl sulfate is also used in the analysis of hemoglobin. The hydrophobic group of SLS acts upon the globin subunit, causing a conformational change. The hydrophillic group of SLS then binds with the oxidized iron subunit, producing a stable reaction product which can then be analyzed, giving a hemoglobin value which is used as part of a complete blood count.

[edit]Niche uses

[edit]Biocide

SDS represent a potentially effective topical microbicide, which can also inhibit and possibly prevent infection by various enveloped and non-enveloped viruses such as the Herpes simplex viruses, HIV, and the Semliki Forest Virus.[11][12]

[edit]Medicinal applications

In medicine, sodium lauryl sulfate is used rectally as a laxative in enemas, and as an excipient on some dissolvable aspirins and other fiber therapy caplets.

[edit]Shark repellant

Evidence suggests that surfactants such as sodium lauryl sulfate can act as a shark repellent at concentrations on the order of 100 parts per million. However, this does not meet the desired “cloud” deterrence level of 0.1 parts per million.[13] [14]

[edit]Taste alteration

Sodium lauryl sulfate temporarily diminishes perception of sweetness,[15] an effect commonly observed after recent use of toothpaste containing this ingredient.[16]

[edit]Toxicology

[edit]Carcinogenicity

SDS is not carcinogenic when either applied directly to skin or consumed.[17] A review of the scientific literature stated “SLS [SDS] was negative in an Ames (bacterial mutation) test, a gene mutation and sister chromatid exchange test in mammalian cells, as well as in an in vivo micronucleus assay in mice. The negative results from in vitro and in vivo studies indicate SDS does not interact with DNA.”[18] The same review also stated “In the only carcinogenicity study available, SLS was not carcinogenic in Beagle dogs, though the short study duration and limited details provided limit the significance that can be attached to the data.”

[edit]Sensitivity

It has been shown to irritate the skin of the face with prolonged and constant exposure (more than an hour) in young adults.[19] SDS may worsen skin problems in individuals with chronic skin hypersensitivity, with some people being affected more than others.[20][21][22] In animal studies SDS appears to cause skin and eye irritation.[18]

[edit]Aphthous ulcers

There have been several studies on how SDS in toothpaste affects the recurrence of aphthous ulcers, commonly referred to in some countries as canker sores or white sores. The results of these studies have been inconsistent. In 1994, a preliminary crossover study showed patients had a significantly higher number of aphthous ulcers after using SLS-containing toothpaste, compared with an SLS-free toothpaste.[23] A follow-up double-blind crossover study in 1996 further supported these results,[24] as did a separate study in 1997.[25] However, a double-blind crossover study published in 1999 failed to find any statistically significant difference.[26] A double-blind crossover study in 2012 also failed to find a significant difference in number of ulcers, but did find a significant difference in ulcer duration and pain scores.[27] According to the 2012 study, patients using an SLS-free toothpaste experienced faster healing of ulcers and less ulcer-related pain on average than patients using SLS-containing toothpaste.[27]

[edit]Interaction with fluoride

Some studies have suggested that SLS in toothpaste may decrease the effectiveness of fluoride at preventing dental caries (cavities). This may be due to SLS interacting with the deposition of fluoride on tooth enamel.[28]

aluminum potassium sulphate 10043-67-1

Potassium alumpotash alum or tawas is the potassium double sulfate of aluminium. Its chemical formula is KAl(SO4)2 and it is commonly found in its dodecahydrate form as KAl(SO4)2·12(H2O). Alum is the common name for this chemical compound, given the nomenclature of potassium aluminum sulfate dodecahydrate. It is commonly used in water purification, leather tanning, dyeing, fireproof textiles, and baking powder[citation needed]. It also has cosmetic uses as a deodorant, as an aftershave treatment and as a styptic for minor bleeding from shaving.

Identifiers
CAS number 10043-67-1 Yes,
7784-24-9 (dodecahydrate)
PubChem 24856
Jmol-3D images Image 1
Properties
Molecular formula KAl(SO4)2
Molar mass 258.205 g/mol
Odor odorless
Density 1.725 g/cm3
Melting point 92–93 °C
Boiling point 200 °C
Solubility in water 14.00 g/100 mL (20 °C)
36.80 g/100 mL (50 °C)
Solubility insoluble in acetone
Refractive index(nD) 1.4564

Characteristics

Potassium alum crystallizes in regular octahedra with flattened corners, and is very soluble in water. The solution reddens litmus and is anastringent. When heated to nearly a red heat it gives a porous, friable mass which is known as “burnt alum.” It fuses at 92 °C in its own water of crystallization. “Neutral alum” is obtained by the addition of as much sodium carbonate to a solution of alum as will begin to cause the separation of alumina. Alum finds application as a mordant, in the preparation of lakes for sizing handmade paper and in the clarifying of turbid liquids.It can also be used as fire proof material and in preparation of many fire proof clothing .

[edit]Mineral form and occurrence

Potassium alum or alum-(K) is a naturally occurring sulfate mineral which typically occurs as encrustations on rocks in areas of weathering andoxidation of sulfide minerals and potassium-bearing minerals. In the past, alum was obtained from alunite, a mineral mined from sulfur-containing volcanic sediments source.[2] Alunite is an associate and likely potassium and aluminium source.[3][4] It has been reported at Vesuvius, Italy, east of Springsure, Queensland, Alum Cave, Tennessee, Alum Gulch, Santa Cruz County, Arizona and the Philippine island of Cebu. A related mineral is kalinite, a fibrous mineral with formula KAl(SO4)2·11(H2O).[5]

[edit]Uses

Potassium alum is an astringent/styptic and antiseptic. For this reason, it can be used as a natural deodorant by inhibiting the growth of thebacteria responsible for body odor. Use of mineral salts in such a fashion does not prevent perspiration. Its astringent/styptic properties are often employed after shaving and to reduce bleeding in minor cuts and abrasions, nosebleeds, and hemorrhoids. It is frequently used topically and internally in traditional systems of medicine including Ayurveda, where it is called phitkari or saurashtri, and Traditional Chinese Medicine, where it is called ming fan.[6] It is also used as a hardener for photographic emulsions (films and papers), usually as part of the fixer, although modern materials are adequately hardened and this practice has fallen out of favor. It is also used in tanning of leather.

[edit]Toxicology and safety

Deodorant crystals containing synthetically made potassium alum are a weak irritant to the skin

properties periodic table

periodic table is a tabular display of the chemical elements, organized on the basis of their atomic numbers, electron configurations, and recurring chemical properties. Elements are presented in order of increasing atomic number (number of protons). The standard form of table comprises an 18 × 7 grid or main body of elements, positioned above a smaller double row of elements. The table can also be deconstructed into four rectangular blocks: the s-block to the left, the p-block to the right, the d-block in the middle, and the f-blockbelow that. The rows of the table are called periods; the columns of the s-, d-, and p-blocks are called groups, with some of these having names such as the halogens or the noble gases. Since, by definition, a periodic table incorporates recurring trends, any such table can be used to derive relationships between the properties of the elements and predict the properties of new, yet to be discovered or synthesized, elements. As a result, a periodic table—whether in the standard form or some other variant—provides a useful framework for analyzing chemical behavior, and such tables are widely use in chemistry and other sciences.

Although precursors exist, Dmitri Mendeleev is generally credited with the publication, in 1869, of the first widely recognized periodic table. He developed his table to illustrate periodic trends in the properties of the then-known elements. Mendeleev also predicted some properties of then-unknown elements that would be expected to fill gaps in this table. Most of his predictions were proved correct when the elements in question were subsequently discovered. Mendeleev’s periodic table has since been expanded and refined with the discovery or synthesis of further new elements and the development of new theoretical models to explain chemical behavior.

All elements from atomic numbers 1 (hydrogen) to 118 (ununoctium) have been discovered or synthesized. Of these, all up to and including californium exist naturally; the rest have only been synthesised in laboratories. Production of elements beyond ununoctium is being pursued, with the question of how the periodic table may need to be modified to accommodate any such additions being a matter of ongoing debate. Numerous synthetic radionuclides of naturally occurring elements have also been produced in laboratories.

properties of titanium dioxide 13463-67-7

Titanium dioxide, also known as titanium(IV) oxide or titania, is the naturally occurring oxide of titanium, chemical formula TiO2. When used as apigment, it is called titanium whitePigment White 6, or CI 77891. Generally it is sourced from ilminite, rutile and anatase. It has a wide range of applications, from paint to sunscreen to food colouring. When used as a food colouring, it has E number E171.

Titanium dioxide
Identifiers
CAS number 13463-67-7 Yes
PubChem 26042
ChemSpider 24256 Yes
UNII 15FIX9V2JP Yes
KEGG C13409 
ChEBI CHEBI:32234 Yes
ChEMBL CHEMBL1201136 
RTECS number XR2775000
Jmol-3D images Image 1
Properties
Molecular formula TiO2
Molar mass 79.866 g/mol
Appearance White solid
Odor odorless
Density 4.23 g/cm3
Melting point 1843 °C
Boiling point 2972 °C
Solubility in water insoluble
Refractive index (nD) 2.488 (anatase)
2.583 (brookite)
2.609 (rutile)
Thermochemistry
Std enthalpy of
formation ΔfHo298
−945 kJ·mol−1[1]
Standard molar
entropy So298
50 J·mol−1·K−1[1]
Hazards
MSDS ICSC 0338
EU classification Not listed
NFPA 704
NFPA 704.svg
0
1
0
Flash point Non-flammable
Related compounds
Other cations Zirconium dioxide
Hafnium dioxide
Related titanium oxides Titanium(II) oxide
Titanium(III) oxide
Titanium(III,IV) oxide
Related compounds Titanic acid

Occurrence

Titanium dioxide occurs in nature as well-known minerals rutile, anatase and brookite, and additionally as two high pressure forms, a monoclinicbaddeleyite-like form and an orthorhombic α-PbO2-like form, both found recently at the Ries crater in Bavaria.[2][3] It is mainly sourced fromilmenite ore. This is the most wide spread form of titanium dioxide-bearing ore around the world. Rutile is the next most abundant and contains around 98% titanium dioxide in the ore.The metastable anatase and brookite phases convert to rutile upon heating.[4]

Titanium dioxide has eight modifications – in addition to rutile, anatase, and brookite three metastable phases can be produced synthetically (monoclinic, tetragonal and orthorombic), and five high pressure forms (α-PbO2-like, baddeleyite-like, cotunnite-like, orthorhombic OI, and cubic phases):

Form Crystal system Synthesis
rutile tetragonal
anatase tetragonal
brookite orthorhombic
TiO2(B)[5] monoclinic Hydrolysis of K2Ti4O9 followed by heating
TiO2(H), hollandite-like form[6] tetragonal Oxidation of the related potassium titanate bronze, K0.25TiO2
TiO2(R), ramsdellite-like form[7] orthorhombic Oxidation of the related lithium titanate bronze Li0.5TiO2
TiO2(II)-(α-PbO2-like form)[8] orthorhombic
baddeleyite-like form, (7 coordinated Ti)[9] monoclinic
TiO2 -OI[10] orthorhombic
cubic form[11] cubic P > 40 GPa, T > 1600 °C
TiO2 -OII, cotunnite(PbCl2)-like[12] orthorhombic P > 40 GPa, T > 700 °C

The cotunnite-type phase was claimed by L. Dubrovinsky and co-authors to be the hardest known oxide with the Vickers hardness of 38 GPa and the bulk modulus of 431 GPa (i.e. close to diamond’s value of 446 GPa) at atmospheric pressure.[12] However, later studies came to different conclusions with much lower values for both the hardness (7–20 GPa, which makes it softer than common oxides like corundum Al2O3 and rutile TiO2)[13] and bulk modulus (~300 GPa).[14][15]

The oxides are commercially important ores of titanium. The metal can also be mined from other minerals such as ilmenite or leucoxene ores, or one of the purest forms, rutile beach sand. Star sapphires and rubies get their asterism from rutile impurities present in them.[16]

Titanium dioxide (B) is found as a mineral in magmatic rocks and hydrothermal veins, as well as weathering rims on perovskite. TiO2 also formslamellae in other minerals.[17]

Spectral lines from titanium oxide are prominent in class M stars, which are cool enough to allow molecules of this chemical to form.

[edit]Production

The production method depends on the feedstock. The most common method for the production of titanium dioxide utilizes ilmenite. Ilmenite is mixed with sulfuric acid. This reacts to remove the iron oxide group in the ilmenite. The by-product iron(II) sulfate is crystallized and filtered-off to yield only the titanium salt in the digestion solution. This product is called synthetic rutile. This is further processed in a similar way to rutile to give the titanium dioxide product. Synthetic rutile and titanium slags are made especially for titanium dioxide production.[18] The use of ilminite ore usually only produces pigment grade titanium dioxide. Another method for the production of synthetic rutile from ilminite utilizes the Becher Process.

Rutile is the second most abundant mineral sand. Rutile found in primary rock cannot be extracted hence the deposits containing rutile sand can be mined meaning a reduced availability to the high concentration ore. Crude titanium dioxide (in the form of rutile or sunthetic rutile) is purified via converting to titanium tetrachloride in the chloride process. In this process, the crude ore (containing at least 70% TiO2) is reduced withcarbon, oxidized with chlorine to give titanium tetrachloride; i.e., carbothermal chlorination. This titanium tetrachloride is distilled, and re-oxidized in a pure oxygen flame or plasma at 1500–2000 K to give pure titanium dioxide while also regenerating chlorine.[19] Aluminium chloride is often added to the process as a rutile promotor; the product is mostly anatase in its absence. The preferred raw material for the chloride process is natural rutile because of its high titanium dioxide content.[20]

One method for the production of titanium dioxide with relevance to nanotechnology is solvothermal Synthesis of titanium dioxide.

Titanium oxide nanotubes, SEMimage.

[edit]Nanotubes

Anatase can be converted by hydrothermal synthesis to delaminated anatase inorganic nanotubes[21] and titanate nanoribbons which are of potential interest as catalytic supports and photocatalysts. In the synthesis, anatase is mixed with 10 M sodium hydroxide and heated at 130 °C for 72 hours. The reaction product is washed with dilute hydrochloric acid and heated at 400 °C for another 15 hours. The yield of nanotubes is quantitative and the tubes have an outer diameter of 10 to 20 nm and an inner diameter of 5 to 8 nm and have a length of 1 μm. A higher reaction temperature (170 °C) and less reaction volume gives the corresponding nanowires.[22]

Another process for synthesizing TiO2 is through anodization in an electrolytic solution. When anodized in a 0.5 weight percent HF solution for 20 minutes, well-aligned titanium oxide nanotube arrays can be fabricated an average tube diameter of 60 nm and length of 250 nm. Based on X-ray Diffraction, nanotubes grown through anodization are amorphous.[23]

[edit]Applications

[edit]Pigment

Titanium dioxide is the most widely used white pigment because of its brightness and very high refractive index, in which it is surpassed only by a few other materials. Approximately 4.6 million tons of pigmentary TiO2 are consumed annually worldwide, and this number is expected to increase as consumption continues to rise. [24] When deposited as a thin film, its refractive index and colour make it an excellent reflective optical coating fordielectric mirrors and some gemstones like “mystic fire topaz”. TiO2 is also an effective opacifier in powder form, where it is employed as a pigment to provide whiteness and opacity to products such as paints, coatings, plastics, papers, inks, foods, medicines (i.e. pills and tablets) as well as mosttoothpastes. In paint, it is often referred to offhandedly as “the perfect white”, “the whitest white”, or other similar terms. Opacity is improved by optimal sizing of the titanium dioxide particles.

In ceramic glazes titanium dioxide acts as an opacifier and seeds crystal formation.

Titanium dioxide has been shown statistically to increase skimmed milk’s whiteness, increasing skimmed milk’s sensory acceptance score.[25]

Titanium dioxide is used to mark the white lines of some tennis courts.[26]

The exterior of the Saturn V rocket was painted with titanium dioxide; this later allowed astronomers to determine that J002E3 was the S-IVB stage from Apollo 12 and not an asteroid.

[edit]Sunscreen and UV absorber

In cosmetic and skin care products, titanium dioxide is used as a pigment, sunscreen and a thickener. It is also used as a tattoo pigment and in styptic pencils. Titanium dioxide is produced in varying particle sizes, oil and water dispersible, and with varying coatings for the cosmetic industry. This pigment is used extensively in plastics and other applications for its UV resistant properties where it acts as a UV absorber, efficiently transforming destructive UV light energy into heat.

Titanium dioxide is found in almost every sunscreen with a physical blocker because of its high refractive index, its strong UV light absorbing capabilities and its resistance to discolouration under ultraviolet light. This advantage enhances its stability and ability to protect the skin from ultraviolet light. Nano-scaled titanium dioxide particles are primarily used in sun screen lotion because they scatter visible light less than titanium dioxide pigments while still providing UV protection.[27] Sunscreens designed for infants or people with sensitive skin are often based on titanium dioxide and/or zinc oxide, as these mineral UV blockers are believed to cause less skin irritation than other UV absorbing chemicals. The titanium dioxide particles used in sunscreens have to be coated with silica or alumina, because titanium dioxide creates radicals in the photocatalytic reaction. These radicals are carcinogenic, and could damage the skin.

[edit]Photocatalyst

TiO2 fibers and spirals.

Titanium dioxide, particularly in the anatase form, is a photocatalyst under ultraviolet (UV) light. Recently it has been found that titanium dioxide, when spiked with nitrogen ions or doped with metal oxide like tungsten trioxide, is also a photocatalyst under either visible or UV light.[28] The strongoxidative potential of the positive holes oxidizes water to create hydroxyl radicals. It can also oxidize oxygen or organic materials directly. Titanium dioxide is thus added to paints, cements, windows, tiles, or other products for its sterilizing, deodorizing and anti-fouling properties and is used as ahydrolysis catalyst. It is also used in dye-sensitized solar cells, which are a type of chemical solar cell (also known as a Graetzel cell).

The photocatalytic properties of titanium dioxide were discovered by Akira Fujishima in 1967[29] and published in 1972.[30] The process on the surface of the titanium dioxide was called the Honda-Fujishima effect.[29] Titanium dioxide has potential for use in energy production: as a photocatalyst, it can carry out hydrolysis; i.e., break water into hydrogen and oxygen. Were the hydrogen collected, it could be used as a fuel. The efficiency of this process can be greatly improved by doping the oxide with carbon.[31] Further efficiency and durability has been obtained by introducing disorder to the lattice structure of the surface layer of titanium dioxide nanocrystals, permitting infrared absorption.[32]

Titanium dioxide can also produce electricity when in nanoparticle form. Research suggests that by using these nanoparticles to form the pixels of a screen, they generate electricity when transparent and under the influence of light. If subjected to electricity on the other hand, the nanoparticles blacken, forming the basic characteristics of a LCD screen. According to creator Zoran Radivojevic, Nokia has already built a functional 200-by-200-pixel monochromatic screen which is energetically self-sufficient.

In 1995 Fujishima and his group discovered the superhydrophilicity phenomenon for titanium dioxide coated glass exposed to sun light.[29] This resulted in the development of self-cleaning glass and anti-fogging coatings.

TiO2 incorporated into outdoor building materials, such as paving stones in noxer blocks or paints, can substantially reduce concentrations of airborne pollutants such as volatile organic compounds and nitrogen oxides.[33]

A photocatalytic cement that uses titanium dioxide as a primary component, produced by Italcementi Group, was included in Time’s Top 50 Inventions of 2008.[34]

Attempts have been made to photocatalytically mineralize pollutants (to convert into CO2 and H2O) in waste water.[35] TiO2 offers great potential as an industrial technology for detoxification or remediation of wastewater due to several factors:[36]

  1. The process uses natural oxygen and sunlight and thus occurs under ambient conditions; it is wavelength selective and is accelerated by UV light.
  2. The photocatalyst is inexpensive, readily available, non-toxic, chemically and mechanically stable, and has a high turnover.
  3. The formation of photocyclized intermediate products, unlike direct photolysis techniques, is avoided.
  4. Oxidation of the substrates to CO2 is complete.
  5. TiO2 can be supported on suitable reactor substrates.

[edit]Electronic data storage medium

In 2010, researchers at the University of Tokyo, Japan have created a crystal form of titanium oxide with particles 5 to 20 nanometers that can be switched between two states with light. Use of the 5 nm particles could theoretically lead to a 25 TB storage disc.[37]

[edit]Other applications

Synthetic single crystals of TiO2, ca. 2–3 mm in size, cut from a larger plate.

  • Titanium dioxide in solution or suspension can be used to cleave protein that contains the amino acid proline at the site where proline is present. This breakthrough in cost-effective protein splitting took place at Arizona State University in 2006.[38]
  • Titanium dioxide is also used as a material in the memristor, a new electronic circuit element. It can be employed for solar energy conversion based on dye, polymer, or quantum dot sensitized nanocrystalline TiO2 solar cells using conjugated polymers as solid electrolytes.[39]
  • Synthetic single crystals and films of TiO2 are used as a semiconductor,[40] and also in Bragg-stack style dielectric mirrors due to the high refractive index of TiO2 (2.5–2.9).[41][42]

[edit]Health and safety

Ambox scales.svg
This article has been nominated to be checked for its neutrality. Discussion of this nomination can be found on thetalk page. (March 2011)

Titanium dioxide is incompatible with strong reducing agents and strong acids.[43] Violent or incandescent reactions occur with molten metals that are very electropositive, e.g. aluminium, calcium, magnesium, potassium, sodium, zinc and lithium.[44]

Titanium dioxide accounts for 70% of the total production volume of pigments worldwide. It is widely used to provide whiteness and opacity to products such as paints, plastics, papers, inks, foods, and toothpastes. It is also used in cosmetic and skin care products, and it is present in almost every sunblock, where it helps protect the skin from ultraviolet light.

Many sunscreens use nanoparticle titanium dioxide (along with nanoparticle zinc oxide) which does get absorbed into the skin.[45][46] The effects on human health are not yet well understood.[47] Nevertheless, allergy to topical application has been confirmed.[48]

Titanium dioxide dust, when inhaled, has been classified by the International Agency for Research on Cancer (IARC) as an IARC Group 2B carcinogen, meaning it is possibly carcinogenic to humans.[49] The findings of the IARC are based on the discovery that high concentrations of pigment-grade (powdered) and ultrafine titanium dioxide dust caused respiratory tract cancer in rats exposed by inhalation and intratracheal instillation.[50] The series of biological events or steps that produce the rat lung cancers (e.g. particle deposition, impaired lung clearance, cell injury, fibrosis, mutations and ultimately cancer) have also been seen in people working in dusty environments. Therefore, the observations of cancer in animals were considered, by IARC, as relevant to people doing jobs with exposures to titanium dioxide dust. For example, titanium dioxide production workers may be exposed to high dust concentrations during packing, milling, site cleaning and maintenance, if there are insufficient dust control measures in place. However, the human studies conducted so far do not suggest an association between occupational exposure to titanium dioxide and an increased risk for cancer. The safety of the use of nano-particle sized titanium dioxide, which can penetrate the body and reach internal organs, has been criticized.[51] Studies have also found that titanium dioxide nanoparticles cause inflammatory response and genetic damage in mice.[52][53] There is some evidence the rare diseaseYellow nail syndrome may be caused by titanium, either implanted for medical reasons or through eating various foods containing titanium dioxide. [54]

Bundschuh et al. from the University Koblenz-Landau have found that waterflea (or daphnia) (a common test for mutagenicity) is damaged into the second generation when swimming in Water containing TiO2 nanoparticles.[55]

The molecular mechanism by which TiO2 induces cancer is unclear. In 2012, Yanglong Zhu, John W. Eaton and Chi Li, on the basis of molecular research, suggested that cell cytotoxicity due to TiO2 results from the interaction between TiO2 nanoparticles and the lysosomal compartment, i.e., independently of the known apoptotic signalling pathways

sodium acetate tri hydrate 6131-90-4

Sodium acetate, CH3COONa, also abbreviated NaOAc,[1] also sodium ethanoate, is the sodium salt of acetic acid. This colourless salt has a wide range of uses.

Sodium acetate
Identifiers
CAS number 127-09-3 Yes, 6131-90-4 (trihydrate)
PubChem 517045
ChemSpider 29105 Yes
UNII NVG71ZZ7P0 Yes
ChEBI CHEBI:32954 Yes
ChEMBL CHEMBL1354 Yes
RTECS number AJ4300010 (anhydrous)
AJ4580000
ATC code B05XA08
Jmol-3D images Image 1
Properties
Molecular formula C2H3NaO2
Molar mass 82.03 g mol−1
Appearance White deliquescent powder
Odor odorless
Density 1.528 g/cm3
1.45 g/cm3 (trihydrate)
Melting point 324 °C (anhydrous)
58 °C (trihydrate)
Boiling point 881.4 °C (anhydrous)
122 °C (trihydrate)(decomposes)
Solubility in water 36.2 g/100 ml (0°C)
46.4 g/100 mL (20°C)
139 g/100 mL (60°C)
170.15 g/100 mL (100°C)
Solubility soluble in ethanol (5.3 g/100 mL (trihydrate)
Basicity (pKb) 9.25
Refractive index(nD) 1.464
Structure
Crystal structure monoclinic
Hazards
MSDS External MSDS
Main hazards Irritant
NFPA 704
NFPA 704.svg
1
1
0
Flash point 250 °C
Autoignition
temperature
607 °C
Related compounds
Other anions Sodium formate
Sodium propionate
Other cations Potassium acetate
Calcium acetate
 Yes (verify) (what is: Yes/?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Sodium acetate, CH3COONa, also abbreviated NaOAc,[1] also sodium ethanoate, is the sodium salt of acetic acid. This colourless salt has a wide range of uses.

Contents

[hide]

  • 1 Applications
    • 1.1 Industrial
    • 1.2 Concrete longevity
    • 1.3 Food
    • 1.4 Buffer solution
    • 1.5 Heating pad
  • 2 Preparation
  • 3 Reactions
  • 4 References
  • 5 External links

[edit]Applications

[edit]Industrial

Sodium acetate is used in the textile industry to neutralize sulfuric acid waste streams, and as a photoresist while using aniline dyes. It is also apickling agent in chrome tanning, and it helps to retard vulcanization of chloroprene in synthetic rubber production. In processing cotton for disposable cotton pads, sodium acetate is used to eliminate the buildup of static electricity.

[edit]Concrete longevity

Sodium acetate is used to reduce the damage water can potentially do to concrete by acting as a Concrete sealant, while also being environmentally benign and cheaper than the epoxy alternative that is usually employed for sealing concrete against water permeation .[2]

[edit]Food

Sodium acetate may be added to foods as a seasoning. It may be used in the form of sodium diacetate — a 1:1 complex of sodium acetate and acetic acid,[3] given the E-number E262. A frequent use of this form is in salt and vinegar chips in the United States. Many US brands, including national manufacturer Frito-Lay, sell “salt and vinegar flavored” chips that use this chemical, with lactose and smaller percentages of other chemicals, in lieu of a real salt and vinegar preparation.[4]

[edit]Buffer solution

As the conjugate base of a weak acid, a solution of sodium acetate and acetic acid can act as a buffer to keep a relatively constant pH. This is useful especially in biochemical applications where reactions are pH dependent.

[edit]Heating pad

Sodium acetate is also used in consumer heating pads or hand warmers and is also used in hot ice. Sodium acetate trihydrate crystals melt at 54°C,[5] (to 58°C [6]) dissolving in their water of crystallization. When they are heated to around 100°C, and subsequently allowed to cool, the aqueous solution becomes supersaturated. This solution is capable of cooling to room temperature without forming crystals. By clicking on a metal disc in the heating pad, a nucleation centre is formed which causes the solution to crystallize into solid sodium acetate trihydrate again. The bond-forming process of crystallization is exothermic.[7][8][9] The latent heat of fusion is about 264–289 kJ/kg.[6] Unlike some other types of heat packs that depend on irreversible chemical reactions, sodium acetate heat packs can be easily recharged by placing in boiling water for a few minutes until all crystals are dissolved; they can be reused many times.

An inexpensive hand warmercontaining a supersaturated solutionof sodium acetate which releases heat on crystallization

[edit]Preparation

Applications

[edit]Industrial

Sodium acetate is used in the textile industry to neutralize sulfuric acid waste streams, and as a photoresist while using aniline dyes. It is also apickling agent in chrome tanning, and it helps to retard vulcanization of chloroprene in synthetic rubber production. In processing cotton for disposable cotton pads, sodium acetate is used to eliminate the buildup of static electricity.

[edit]Concrete longevity

Sodium acetate is used to reduce the damage water can potentially do to concrete by acting as a Concrete sealant, while also being environmentally benign and cheaper than the epoxy alternative that is usually employed for sealing concrete against water permeation .[2]

[edit]Food

Sodium acetate may be added to foods as a seasoning. It may be used in the form of sodium diacetate — a 1:1 complex of sodium acetate and acetic acid,[3] given the E-number E262. A frequent use of this form is in salt and vinegar chips in the United States. Many US brands, including national manufacturer Frito-Lay, sell “salt and vinegar flavored” chips that use this chemical, with lactose and smaller percentages of other chemicals, in lieu of a real salt and vinegar preparation.[4]

[edit]Buffer solution

As the conjugate base of a weak acid, a solution of sodium acetate and acetic acid can act as a buffer to keep a relatively constant pH. This is useful especially in biochemical applications where reactions are pH dependent.

[edit]Heating pad

Sodium acetate is also used in consumer heating pads or hand warmers and is also used in hot ice. Sodium acetate trihydrate crystals melt at 54°C,[5] (to 58°C [6]) dissolving in their water of crystallization. When they are heated to around 100°C, and subsequently allowed to cool, the aqueous solution becomes supersaturated. This solution is capable of cooling to room temperature without forming crystals. By clicking on a metal disc in the heating pad, a nucleation centre is formed which causes the solution to crystallize into solid sodium acetate trihydrate again. The bond-forming process of crystallization is exothermic.[7][8][9] The latent heat of fusion is about 264–289 kJ/kg.[6] Unlike some other types of heat packs that depend on irreversible chemical reactions, sodium acetate heat packs can be easily recharged by placing in boiling water for a few minutes until all crystals are dissolved; they can be reused many times.

An inexpensive hand warmercontaining a supersaturated solutionof sodium acetate which releases heat on crystallization

[edit]Preparation

a crystal of sodium acetate trihydrate (length 1.7 centimetre)

For laboratory use, sodium acetate is very inexpensive, and is usually purchased instead of being synthesized. It is sometimes produced in a laboratory experiment by the reaction of acetic acid (ethanoic acid) with sodium carbonate, sodium bicarbonate, or sodium hydroxide. These reactions produce aqueous sodium acetate and water. Carbon dioxide is produced in the reaction with sodium carbonate and bicarbonate, and it leaves the reaction vessel as a gas (unless the reaction vessel is pressurized). This is the well-known “volcano” reaction between baking soda (sodium bicarbonate) and vinegar.

CH3COOH + NaHCO3 → CH3COONa + H2O + CO2

Industrially, sodium acetate is prepared from glacial acetic acid and sodium hydroxide.

CH3COOH + NaOH → CH3COONa + H2O

[edit]Reactions

Sodium acetate can be used to form an ester with an alkyl halide such as bromoethane:

CH3COONa + BrCH2CH3 → CH3COOCH2CH3 + NaBr

Caesium salts catalyze this reaction.

Magnesium Sulfate 7487-88-9

Magnesium sulfate (or magnesium sulphate) is an inorganic salt (chemical compound) containing magnesium, sulfur and oxygen, with the formula MgSO4. It is often encountered as the heptahydrate sulfate mineral epsomite (MgSO4·7H2O), commonly called Epsom salt, named for a bitter saline spring from the town of Epsom in Surrey, England, where the salt was produced from the springs that arise where the porous chalk of the North Downs meets non-porous London clay. Epsom salt occurs naturally as a pure mineral. Another hydrate form is kieserite.

Anhydrous magnesium sulfate

Epsomite (heptahydrate)
Identifiers
CAS number 7487-88-Yes9, 14168-73-1 (monohydrate), 24378-31-2 (tetrahydrate), 15553-21-6 (pentahydrate), 13778-97-7 (hexahydrate), 10034-99-8 (heptahydrate)
PubChem 24083
ChemSpider 22515 Yes
UNII ML30MJ2U7I Yes
DrugBank DB00653
ChEBI CHEBI:32599 Yes
ChEMBL CHEMBL1200456 
RTECS number OM4500000
ATC code A06AD04,A12CC02B05XA05 D11AX05 V04CC02
Jmol-3D images Image 1
Properties
Molecular formula MgSO4
Molar mass 120.366 g/mol (anhydrous)
246.47 g/mol (heptahydrate)
Appearance white crystalline solid
Odor odorless
Density 2.66 g/cm3 (anhydrous)
2.445 g/cm3 (monohydrate)
1.68 g/cm3 (heptahydrate)
1.512 g/cm3 (11-hydrate)
Melting point 1124 °C (anhydrous, decomp)
200 °C (monohydrate, decomp)
150 °C (heptahydrate, decomp)
2 °C (11-hydrate, decomp)
Solubility inwater anhydrous
26.9 g/100 mL (0 °C)
25.5 g/100 mL (20 °C)


heptahydrate
71 g/100 mL (20 °C)

Solubility 1.16 g/100 mL (18 °C, ether)
slightly soluble in alcohol,glycerol
insoluble in acetone
Refractive index (nD) 1.523 (monohydrate)
1.433 (heptahydrate)
Structure
Crystal structure monoclinic (hydrate)
Hazards
MSDS External MSDS
EU Index Not listed
Related compounds
Othercations Beryllium sulfate
Calcium sulfate
Strontium sulfate
Barium sulfate

Anhydrous magnesium sulfate is used as a drying agent. Since the anhydrous form is hygroscopic (readily absorbs water from the air) and is therefore difficult to weigh accurately, the hydrate is often preferred when preparing solutions, for example in medical preparations. Epsom salt has been traditionally used as a component of bath salts. Epsom salt can also be used as a beauty product. Athletes use it to soothe sore muscles, while gardeners use it to improve crops. It has a variety of other uses

Magnesium sulfate is highly soluble in water. The anhydrous form is strongly hygroscopic, and can be used as a desiccant.

Epsom salts have medicinal properties when used both externally and internally.

Magnesium sulfate is the primary substance that causes the absorption of sound in seawater[2] (acoustic energy is converted to thermal energy). Absorption is strongly dependent on frequency: lower frequencies are less absorbed by the salt, so that the sound travels much farther in the ocean. Boric acid also contributes to absorption, but the most abundant salt in seawater, sodium chloride, has negligible sound absorption.

[edit]Hydrates

Almost all known mineralogical forms of MgSO4 occur as hydrates. Epsomite is the natural analogue of “Epsom salt”. Another heptahydrate, thecopper-containing mineral alpersite (Mg,Cu)SO4·7H2O,[3] was recently recognized. Both are, however, not the highest known hydrates of MgSO4, due to the recent terrestrial find of meridianiite, MgSO4·11H2O, which is thought to also occur on Mars. Hexahydrite is the next lower (6) hydrate. Three next lower hydrates — pentahydrite (5), starkeyite (4) and especially sanderite (2) — are more rarely found. Kieserite is a monohydrate and is common among evaporitic deposits. Anhydrous magnesium sulfate was reported from some burning coal dumps but never treated as a mineral.

The pH of hydrates is average 6.0 (5.5 to 6.5). Magnesium hydrates have, like copper(II) sulfate, coordinated water.[4]

[edit]Manufacturing

The heptahydrate can be prepared by neutralizing sulfuric acid with magnesium carbonate or oxide, but it is usually obtained directly from natural sources.

Anhydrous magnesium sulfate is prepared only by the dehydration of a hydrate.

[edit]Occurrence

Magnesium sulfates are common minerals in geological environments. Their occurrence is mostly connected with supergene processes. Some of them are also important constituents of evaporitic potassium-magnesium (K-Mg) salts deposits.

[edit]Applications

Anhydrous magnesium sulfate is commonly used as a desiccant in organic synthesis due to its affinity for water. During work-up, an organic phase is saturated with magnesium sulfate until it no longer forms clumps. The hydrated solid is then removed with filtration or decantation. Other inorganic sulfate salts such as sodium sulfate and calcium sulfate may also be used in the same way.

Magnesium sulfate is used in bath salts, particularly in flotation therapy where high concentrations raise the bath water’s specific gravity, effectively making the body more buoyant. Traditionally, it is also used to prepare foot baths, intended to soothe sore feet. The reason for the inclusion of the salt is partially cosmetic: the increase in ionic strength prevents some of the temporary skin wrinkling (partial maceration) which is caused by prolonged immersion of extremities in pure water. However, magnesium sulfate can also be absorbed into the skin, reducing inflammation.[citation needed] It is naturally present in some mineral waters.[citation needed]

It may also be used as a coagulant for making tofu.[5]

Magnesium sulfate heptahydrate is also used to maintain the magnesium concentration in marine aquaria which contain large amounts of stony corals as it is slowly depleted in their calcification process. In a magnesium-deficient marine aquarium calcium and alkalinity concentrations are very difficult to control because not enough magnesium is present to stabilize these ions in the saltwater and prevent their spontaneous precipitation into calcium carbonate.[6]

Magnesium sulfate is used as the electrolyte to prepare copper sulfate. A magnesium sulfate solution is electrolyzed with a copper anode to form copper sulfate, magnesium hydroxide, and hydrogen:

Cu + MgSO4 + 2 H2O → H2 + CuSO4 + Mg(OH)2.[citation needed]

Magnesium sulfate is used as a brewing salt in beer production to adjust the ion content of the brewing water and enhance enzyme action in the mash or promote a desired flavor profile in the beer.

[edit]Agriculture

In gardening and other agriculture, magnesium sulfate is used to correct a magnesium or sulfur deficiency in soil; magnesium is an essential element in the chlorophyll molecule, and sulfur is another important micronutrient. It is most commonly applied to potted plants, or to magnesium-hungry crops, such as potatoes, roses, tomatoes, lemon trees, peppers and cannabis. The advantage of magnesium sulfate over other magnesium soil amendments (such as dolomitic lime) is its high solubility, which also allows the option of foliar feeding. Solutions of magnesium sulfate are also nearly neutral, as compared to alkaline salts of magnesium, as found in limestone; therefore the use of magnesium sulfate as a magnesium source for soil does not significantly change the soil pH.[citation needed]

[edit]Medical use

Magnesium sulfate is a common pharmaceutical preparation of magnesium, commonly known as Epsom salts, used both externally and internally. Epsom salts are used as bath salts. The sulfate is supplied in a gel preparation for topical application in treating aches and pains. Oral magnesium sulfate is commonly used as a saline laxative or osmotic purgative. Magnesium sulfate is the main preparation of intravenous magnesium.

Indications for internal use are:

  • Replacement therapy for hypomagnesemia.[7]
  • Magnesium sulfate is the first-line antiarrhythmic agent for torsades de pointes in cardiac arrest under the 2005 ECC guidelines and for managing quinidine-induced arrhythmias.[8]
  • As a bronchodilator after beta-agonist and anticholinergic agents have been tried, e.g. in severe exacerbations of asthma.[9] Recent studies have revealed that magnesium sulfate can be nebulized to reduce the symptoms of acute asthma.[9] It is commonly administered via theintravenous route for the management of severe asthma attacks.
  • Magnesium sulfate can be used to treat eclampsia in pregnant women.[10]
  • Magnesium sulfate can also delay labor (tocolysis) by inhibiting uterine muscle contraction in the case of premature labor, to delay preterm birth.[11][12]
  • Intravenous magnesium sulfate has been shown to prevent cerebral palsy in preterm babies.[13] A recent systematic review suggests that antenatal intravenous magnesium sulphate can reduce the risk of cerebral palsy and gross motor dysfunction in preterm infants by on average 30%.[14]
  • Magnesium sulfate is administered intravenously as a bolus and as an infusion due to its inhibition of skeletal muscle contraction and vasodilatory properties for the management ofChironex fleckeri or Box Jellyfish envenomation that is unresponsive to antivenom therapy, or for treatment of Irukandji syndrome.
  • Solutions of sulfate salts such as Epsom salt may be given as first aid for barium chloride poisoning.[15]

An overdose of magnesium causes hypermagnesemia.