Antimony telluride is an inorganic compound with the chemical formula Sb2Te3. It is a grey crystalline solid with layered structure. Layers consist of two atomic sheets of antimony and three atomic sheets of tellurium and are held together by weak van der Waals forces. Sb2Te3 is a narrow-gap semiconductor with a band gap 0.21 eV. Antimony telluride can be formed by the reaction of antimony with tellurium at 500–900 °C. 2 Sb + 3 Te → Sb2Te3 Like other binary chalcogenides of antimony and bismuth, Sb2Te3 has been investigated for its semiconductor properties. It can be transformed into both n-type and p-type semiconductors by doping with an appropriate dopant. Sb2Te3 forms the pseudobinary intermetallic system germanium-antimony-tellurium with germanium telluride, GeTe. Like bismuth telluride, Bi2Te3, antimony telluride has a large thermoelectric effect and is therefore used in solid state refrigerators
Encyclopædia Britannica, Eleventh Edition
The Encyclopædia Britannica, Eleventh Edition is a 29-volume reference work, an edition of the Encyclopædia Britannica. It was developed during the encyclopaedia's transition from a British to an American publication; some of its articles were written by the best-known scholars of the time. This edition of the encyclopedia, containing 40,000 entries, is now in the public domain, many of its articles have been used as a basis for articles in Wikipedia. However, the outdated nature of some of its content makes its use as a source for modern scholarship problematic; some articles have special value and interest to modern scholars as cultural artifacts of the 19th and early 20th centuries. The 1911 eleventh edition was assembled with the management of American publisher Horace Everett Hooper. Hugh Chisholm, who had edited the previous edition, was appointed editor in chief, with Walter Alison Phillips as his principal assistant editor. Hooper bought the rights to the 25-volume 9th edition and persuaded the British newspaper The Times to issue its reprint, with eleven additional volumes as the tenth edition, published in 1902.
Hooper's association with The Times ceased in 1909, he negotiated with the Cambridge University Press to publish the 29-volume eleventh edition. Though it is perceived as a quintessentially British work, the eleventh edition had substantial American influences, not only in the increased amount of American and Canadian content, but in the efforts made to make it more popular. American marketing methods assisted sales; some 14% of the contributors were from North America, a New York office was established to coordinate their work. The initials of the encyclopedia's contributors appear at the end of selected articles or at the end of a section in the case of longer articles, such as that on China, a key is given in each volume to these initials; some articles were written by the best-known scholars of the time, such as Edmund Gosse, J. B. Bury, Algernon Charles Swinburne, John Muir, Peter Kropotkin, T. H. Huxley, James Hopwood Jeans and William Michael Rossetti. Among the lesser-known contributors were some who would become distinguished, such as Ernest Rutherford and Bertrand Russell.
Many articles were carried over from some with minimal updating. Some of the book-length articles were divided into smaller parts for easier reference, yet others much abridged; the best-known authors contributed only a single article or part of an article. Most of the work was done by British Museum scholars and other scholars; the 1911 edition was the first edition of the encyclopædia to include more than just a handful of female contributors, with 34 women contributing articles to the edition. The eleventh edition introduced a number of changes of the format of the Britannica, it was the first to be published complete, instead of the previous method of volumes being released as they were ready. The print type was subject to continual updating until publication, it was the first edition of Britannica to be issued with a comprehensive index volume in, added a categorical index, where like topics were listed. It was the first not to include long treatise-length articles. Though the overall length of the work was about the same as that of its predecessor, the number of articles had increased from 17,000 to 40,000.
It was the first edition of Britannica to include biographies of living people. Sixteen maps of the famous 9th edition of Stielers Handatlas were translated to English, converted to Imperial units, printed in Gotha, Germany by Justus Perthes and became part this edition. Editions only included Perthes' great maps as low quality reproductions. According to Coleman and Simmons, the content of the encyclopedia was distributed as follows: Hooper sold the rights to Sears Roebuck of Chicago in 1920, completing the Britannica's transition to becoming a American publication. In 1922, an additional three volumes, were published, covering the events of the intervening years, including World War I. These, together with a reprint of the eleventh edition, formed the twelfth edition of the work. A similar thirteenth edition, consisting of three volumes plus a reprint of the twelfth edition, was published in 1926, so the twelfth and thirteenth editions were related to the eleventh edition and shared much of the same content.
However, it became apparent that a more thorough update of the work was required. The fourteenth edition, published in 1929, was revised, with much text eliminated or abridged to make room for new topics; the eleventh edition was the basis of every version of the Encyclopædia Britannica until the new fifteenth edition was published in 1974, using modern information presentation. The eleventh edition's articles are still of value and interest to modern readers and scholars as a cultural artifact: the British Empire was at its maximum, imperialism was unchallenged, much of the world was still ruled by monarchs, the tragedy of the modern world wars was still in the future, they are an invaluable resource for topics omitted from modern encyclopedias for biography and the history of science and technology. As a literary text, the encyclopedia has value as an example of early 20th-century prose. For example, it employs literary devices, such as pathetic fallacy, which are not as common in modern reference texts.
In 1917, using the pseudonym of S. S. Van Dine, the US art critic and author Willard Huntington Wright published Misinforming a Nation, a 200+
Water of crystallization
In chemistry, water of crystallization or water of hydration are water molecules that are present inside crystals. Water is incorporated in the formation of crystals from aqueous solutions. In some contexts, water of crystallization is the total mass of water in a substance at a given temperature and is present in a definite ratio. Classically, "water of crystallization" refers to water, found in the crystalline framework of a metal complex or a salt, not directly bonded to the metal cation. Upon crystallization from water or moist solvents, many compounds incorporate water molecules in their crystalline frameworks. Water of crystallization can be removed by heating a sample but the crystalline properties are lost. For example, in the case of sodium chloride, the dihydrate is unstable at room temperature. Compared to inorganic salts, proteins crystallize with large amounts of water in the crystal lattice. A water content of 50% is not uncommon for proteins. In molecular formulas water of crystallization can be denoted in differential: "hydrated compound · nH2O" or "hydrated compound×nH2O"This notation is used when the compound only contains lattice water or when the crystal structure is undetermined.
For example Calcium chloride: CaCl2 · 2H2O"hydrated compoundn"A hydrate with coordinated water. For example Zinc chloride: ZnCl24Both notations can be combined as for example in copper sulfate: SO4 · H2O A salt with associated water of crystallization is known as a hydrate; the structure of hydrates can be quite elaborate, because of the existence of hydrogen bonds that define polymeric structures. The structures of many hydrates were unknown, the dot in the formula of a hydrate was employed to specify the composition without indicating how the water is bound. Examples: CuSO4 · 5H2O - copper sulfate pentahydrate CoCl2 · 6H2O - cobalt chloride hexahydrate SnCl2 · 2H2O - tin chloride dihydrateFor many salts, the exact bonding of the water is unimportant because the water molecules are labilized upon dissolution. For example, an aqueous solution prepared from CuSO4 · 5H2O and anhydrous CuSO4 behave identically. Therefore, knowledge of the degree of hydration is important only for determining the equivalent weight: one mole of CuSO4 · 5H2O weighs more than one mole of CuSO4.
In some cases, the degree of hydration can be critical to the resulting chemical properties. For example, anhydrous RhCl3 is not soluble in water and is useless in organometallic chemistry whereas RhCl3 · 3H2O is versatile. Hydrated AlCl3 is a poor Lewis acid and thus inactive as a catalyst for Friedel-Crafts reactions. Samples of AlCl3 must therefore be protected from atmospheric moisture to preclude the formation of hydrates. Crystals of hydrated copper sulfate consist of 2+ centers linked to SO42− ions. Copper is surrounded by six oxygen atoms, provided by two different sulfate groups and four molecules of water. A fifth water does not bind directly to copper; the cobalt chloride mentioned above occurs as 2+ and Cl−. In tin chloride, each Sn center is pyramidal being bound to one water; the second water in the formula unit is hydrogen-bonded to the chloride and to the coordinated water molecule. Water of crystallization is stabilized by electrostatic attractions hydrates are common for salts that contain +2 and +3 cations as well as −2 anions.
In some cases, the majority of the weight of a compound arises from water. Glauber's salt, Na2SO410, is a white crystalline solid with greater than 50% water by weight. Consider the case of nickel chloride hexahydrate; this species has the formula NiCl26. Crystallographic analysis reveals that the solid consists of subunits that are hydrogen bonded to each other as well as two additional molecules of H2O, thus 1/3 of the water molecules in the crystal are not directly bonded to Ni2+, these might be termed "water of crystallization". The water content of most compounds can be determined with a knowledge of its formula. An unknown sample can be determined through thermogravimetric analysis where the sample is heated and the accurate weight of a sample is plotted against the temperature; the amount of water driven off is divided by the molar mass of water to obtain the number of molecules of water bound to the salt. Water is common solvent to be found in crystals because it is small and polar, but all solvents can be found in some host crystals.
Water is noteworthy because it is reactive, whereas other solvents such as benzene are considered to be chemically innocuous. More than one solvent is found in a crystal, the stoichiometry is variable, reflected in the crystallographic concept of "partial occupancy." It is common and conventional for a chemist to "dry" a sample with a combination of vacuum and heat "to constant weight." For other solvents of crystallization, analysis is conveniently accomplished by dissolving the sample in a deuterated solvent and analyzing the sample for solvent signals by NMR spectroscopy. Single crystal X-ray crystallography is able to detect the presence of these solvents of crystallization as well. Other methods may be available. In the table below are indicated the number of molecules of water per metal in various salts. Transition metal sulfates form mono-, tetra-, pentahydrates, each of which crystallizes in only one form; the water in these salts is coordinated, together with sulfate to the metal center.
The sulfates of these same metals crystallize as both tetragonal and monoclinic hexahydrates, wherein all water
Sodium benzoate is a substance which has the chemical formula C6H5COONa. It is a used food preservative, with an E number of E211, it exists in this form when dissolved in water. It can be produced by reacting sodium hydroxide with benzoic acid. Sodium benzoate is produced by the neutralization of benzoic acid, itself produced commercially by partial oxidation of toluene with oxygen. Sodium benzoate occurs along with benzoic acid and its esters, in many foods. Fruits and vegetables can be rich sources berries such as cranberry and bilberry. Other sources include seafood, such as prawns, dairy products like milk and yogurt. Sodium benzoate is a preservative, with the E number E211, it is most used in acidic foods such as salad dressings, carbonated drinks and fruit juices, pickles and frogurt toppings. It is used as a preservative in medicines and cosmetics. Under these conditions it is converted into benzoic acid, bacteriostatic and fungistatic. Benzoic acid is not used directly due to its poor water solubility.
Concentration as a food preservative is limited by the FDA in the U. S. to 0.1% by weight. Sodium benzoate is allowed as an animal food additive at up to 0.1%, according to AFCO's official publication. Sodium benzoate has been replaced by potassium sorbate in the majority of soft drinks in the United Kingdom. Sodium benzoate is used as a treatment for urea cycle disorders due to its ability to bind amino acids; this leads to excretion of a decrease in ammonia levels. Recent research shows. Total Positive and Negative Syndrome Scale scores dropped by 21% compared to placebo. Sodium benzoate, along with phenylbutyrate, is used to treat hyperammonemia. Sodium benzoate is used in fireworks as a fuel in whistle mix, a powder that emits a whistling noise when compressed into a tube and ignited; the mechanism starts with the absorption of benzoic acid into the cell. If the intracellular pH falls to 5 or lower, the anaerobic fermentation of glucose through phosphofructokinase decreases which inhibits the growth and survival of microorganisms that cause food spoilage.
In the United States, sodium benzoate is designated as recognized as safe by the Food and Drug Administration. The International Programme on Chemical Safety found no adverse effects in humans at doses of 647–825 mg/kg of body weight per day. Cats have a lower tolerance against benzoic acid and its salts than rats and mice; the human body clears sodium benzoate by combining it with glycine to form hippuric acid, excreted. The metabolic pathway for this begins with the conversion of benzoate by butyrate-CoA ligase into an intermediate product, benzoyl-CoA, metabolized by glycine N-acyltransferase into hippuric acid. In combination with ascorbic acid, sodium benzoate and potassium benzoate may form benzene; when tested by the FDA, most beverages that contained both ascorbic acid and benzoate had benzene levels that were below those considered dangerous for consumption by the World Health Organization. Most of the beverages that tested higher have been reformulated and subsequently tested below the safety limit.
Heat and shelf life can increase the rate at which benzene is formed. Research published in 2007 for the UK's Food Standards Agency suggests that certain artificial colors, when paired with sodium benzoate, may be linked to hyperactive behavior; the results were inconsistent regarding sodium benzoate, so the FSA recommended further study. The Food Standards Agency concluded that the observed increases in hyperactive behavior, if real, were more to be linked to the artificial colors than to sodium benzoate; the report's author, Jim Stevenson from Southampton University, said: "The results suggest that consumption of certain mixtures of artificial food colours and sodium benzoate preservative are associated with increases in hyperactive behaviour in children.... Many other influences are at work but this at least is one a child can avoid." British Pharmacopoeia European Pharmacopoeia Food Chemicals Codex Japanese Pharmacopoeia United States Pharmacopeia Acceptable daily intake List of investigational antipsychotics Potassium benzoate International Programme on Chemical Safety - Benzoic Acid and Sodium Benzoate report Kubota K, Ishizaki T. "Dose-dependent pharmacokinetics of benzoic acid following oral administration of sodium benzoate to humans".
Eur. J. Clin. Pharmacol. 41: 363–8. Doi:10.1007/BF00314969. PMID 1804654. Although the maximum rate of biotransformation of benzoic acid to hippuric acid varied between 17.2 and 28.8 mg.kg-1.h-1 among the six individuals, the mean value was close to that provided by daily maximum dose recommended in the treatment of hyperammonaemia in patients with inborn errors of ureagenesis Safety data for sodium benzoate The Ketchup Conundrum
Potassium is a chemical element with symbol K and atomic number 19. It was first isolated from the ashes of plants, from which its name derives. In the periodic table, potassium is one of the alkali metals. All of the alkali metals have a single valence electron in the outer electron shell, removed to create an ion with a positive charge – a cation, which combines with anions to form salts. Potassium in nature occurs only in ionic salts. Elemental potassium is a soft silvery-white alkali metal that oxidizes in air and reacts vigorously with water, generating sufficient heat to ignite hydrogen emitted in the reaction, burning with a lilac-colored flame, it is found dissolved in sea water, is part of many minerals. Potassium is chemically similar to sodium, the previous element in group 1 of the periodic table, they have a similar first ionization energy, which allows for each atom to give up its sole outer electron. That they are different elements that combine with the same anions to make similar salts was suspected in 1702, was proven in 1807 using electrolysis.
Occurring potassium is composed of three isotopes, of which 40K is radioactive. Traces of 40K are found in all potassium, it is the most common radioisotope in the human body. Potassium ions are vital for the functioning of all living cells; the transfer of potassium ions across nerve cell membranes is necessary for normal nerve transmission. Fresh fruits and vegetables are good dietary sources of potassium; the body responds to the influx of dietary potassium, which raises serum potassium levels, with a shift of potassium from outside to inside cells and an increase in potassium excretion by the kidneys. Most industrial applications of potassium exploit the high solubility in water of potassium compounds, such as potassium soaps. Heavy crop production depletes the soil of potassium, this can be remedied with agricultural fertilizers containing potassium, accounting for 95% of global potassium chemical production; the English name for the element potassium comes from the word "potash", which refers to an early method of extracting various potassium salts: placing in a pot the ash of burnt wood or tree leaves, adding water and evaporating the solution.
When Humphry Davy first isolated the pure element using electrolysis in 1807, he named it potassium, which he derived from the word potash. The symbol "K" stems from kali, itself from the root word alkali, which in turn comes from Arabic: القَلْيَه al-qalyah "plant ashes". In 1797, the German chemist Martin Klaproth discovered "potash" in the minerals leucite and lepidolite, realized that "potash" was not a product of plant growth but contained a new element, which he proposed to call kali. In 1807, Humphry Davy produced the element via electrolysis: in 1809, Ludwig Wilhelm Gilbert proposed the name Kalium for Davy's "potassium". In 1814, the Swedish chemist Berzelius advocated the name kalium for potassium, with the chemical symbol "K"; the English and French speaking countries adopted Davy and Gay-Lussac/Thénard's name Potassium, while the Germanic countries adopted Gilbert/Klaproth's name Kalium. The "Gold Book" of the International Union of Physical and Applied Chemistry has designated the official chemical symbol as K.
Potassium is the second least dense metal after lithium. It is a soft solid with a low melting point, can be cut with a knife. Freshly cut potassium is silvery in appearance, but it begins to tarnish toward gray on exposure to air. In a flame test and its compounds emit a lilac color with a peak emission wavelength of 766.5 nanometers. Neutral potassium atoms have 19 electrons, one more than the stable configuration of the noble gas argon; because of this and its low first ionization energy of 418.8 kJ/mol, the potassium atom is much more to lose the last electron and acquire a positive charge than to gain one and acquire a negative charge. This process requires so little energy that potassium is oxidized by atmospheric oxygen. In contrast, the second ionization energy is high, because removal of two electrons breaks the stable noble gas electronic configuration. Potassium therefore does not form compounds with the oxidation state of higher. Potassium is an active metal that reacts violently with oxygen in water and air.
With oxygen it forms potassium peroxide, with water potassium forms potassium hydroxide. The reaction of potassium with water is dangerous because of its violent exothermic character and the production of hydrogen gas. Hydrogen reacts again with atmospheric oxygen, producing water, which reacts with the remaining potassium; this reaction requires only traces of water. Because of the sensitivity of potassium to water and air, reactions with other elements are possible only in an inert atmosphere such as argon gas using air-free techniques. Potassium does not react with most hydrocarbons such as mineral kerosene, it dissolves in liquid ammonia, up to 480 g per 1000 g of ammonia at 0 °C. Depending on the concentration, the ammonia solutions are blue to yellow, their electrical conductivity is similar to that of liquid metals. In a pure solution, potassium reacts with ammonia to form KNH2, but this reaction is accelerated by minute amounts of transition metal s
The melting point of a substance is the temperature at which it changes state from solid to liquid. At the melting point the solid and liquid phase exist in equilibrium; the melting point of a substance depends on pressure and is specified at a standard pressure such as 1 atmosphere or 100 kPa. When considered as the temperature of the reverse change from liquid to solid, it is referred to as the freezing point or crystallization point; because of the ability of some substances to supercool, the freezing point is not considered as a characteristic property of a substance. When the "characteristic freezing point" of a substance is determined, in fact the actual methodology is always "the principle of observing the disappearance rather than the formation of ice", that is, the melting point. For most substances and freezing points are equal. For example, the melting point and freezing point of mercury is 234.32 kelvins. However, certain substances possess differing solid-liquid transition temperatures.
For example, agar melts at 85 °C and solidifies from 31 °C. The melting point of ice at 1 atmosphere of pressure is close to 0 °C. In the presence of nucleating substances, the freezing point of water is not always the same as the melting point. In the absence of nucleators water can exist as a supercooled liquid down to −48.3 °C before freezing. The chemical element with the highest melting point is tungsten, at 3,414 °C; the often-cited carbon does not melt at ambient pressure but sublimes at about 3,726.85 °C. Tantalum hafnium carbide is a refractory compound with a high melting point of 4215 K. At the other end of the scale, helium does not freeze at all at normal pressure at temperatures arbitrarily close to absolute zero. Many laboratory techniques exist for the determination of melting points. A Kofler bench is a metal strip with a temperature gradient. Any substance can be placed on a section of the strip, revealing its thermal behaviour at the temperature at that point. Differential scanning calorimetry gives information on melting point together with its enthalpy of fusion.
A basic melting point apparatus for the analysis of crystalline solids consists of an oil bath with a transparent window and a simple magnifier. The several grains of a solid are placed in a thin glass tube and immersed in the oil bath; the oil bath is heated and with the aid of the magnifier melting of the individual crystals at a certain temperature can be observed. In large/small devices, the sample is placed in a heating block, optical detection is automated; the measurement can be made continuously with an operating process. For instance, oil refineries measure the freeze point of diesel fuel online, meaning that the sample is taken from the process and measured automatically; this allows for more frequent measurements as the sample does not have to be manually collected and taken to a remote laboratory. For refractory materials the high melting point may be determined by heating the material in a black body furnace and measuring the black-body temperature with an optical pyrometer. For the highest melting materials, this may require extrapolation by several hundred degrees.
The spectral radiance from an incandescent body is known to be a function of its temperature. An optical pyrometer matches the radiance of a body under study to the radiance of a source, calibrated as a function of temperature. In this way, the measurement of the absolute magnitude of the intensity of radiation is unnecessary. However, known temperatures must be used to determine the calibration of the pyrometer. For temperatures above the calibration range of the source, an extrapolation technique must be employed; this extrapolation is accomplished by using Planck's law of radiation. The constants in this equation are not known with sufficient accuracy, causing errors in the extrapolation to become larger at higher temperatures. However, standard techniques have been developed to perform this extrapolation. Consider the case of using gold as the source. In this technique, the current through the filament of the pyrometer is adjusted until the light intensity of the filament matches that of a black-body at the melting point of gold.
This establishes the primary calibration temperature and can be expressed in terms of current through the pyrometer lamp. With the same current setting, the pyrometer is sighted on another black-body at a higher temperature. An absorbing medium of known transmission is inserted between this black-body; the temperature of the black-body is adjusted until a match exists between its intensity and that of the pyrometer filament. The true higher temperature of the black-body is determined from Planck's Law; the absorbing medium is removed and the current through the filament is adjusted to match the filament intensity to that of the black-body. This establishes a second calibration point for the pyrometer; this step is repeated to carry the calibration to hi
Sodium bromide is an inorganic compound with the formula NaBr. It is a high-melting crystalline solid that resembles sodium chloride, it is a used source of the bromide ion and has many applications. NaBr crystallizes in the same cubic motif as NaCl, NaF and NaI; the anhydrous salt crystallizes above 50.7 °C. Dihydrate salts crystallize out of water solution below 50.7 °C. NaBr is produced by treating sodium hydroxide with hydrogen bromide. Sodium bromide can be used as a source of the chemical element bromine; this can be accomplished by treating an aqueous solution of NaBr with chlorine gas: 2 NaBr + Cl2 → Br2 + 2 NaCl Sodium bromide is the most useful inorganic bromide in industry. It is used as a catalyst in TEMPO-mediated oxidation reactions. Known as Sedoneural, sodium bromide has been used as a hypnotic and sedative in medicine used as an anticonvulsant and a sedative in the late 19th and early 20th centuries, its action is due to the bromide ion, for this reason potassium bromide is effective.
In 1975, bromides were removed from drugs in the U. S. such as Bromo-Seltzer due to toxicity. Sodium bromide is used for the preparation of other bromides in organic synthesis and other areas, it is a source of the bromide nucleophile to convert alkyl chlorides to more reactive alkyl bromides by the Finkelstein reaction: NaBr + RCl → RBr + NaCl Once a large need in photography, but now shrinking, the photosensitive salt silver bromide is prepared using NaBr. Sodium bromide is used in conjunction with chlorine as a disinfectant for hot tubs and swimming pools. Sodium bromide is used to prepare dense fluids used in oil wells. NaBr has a low toxicity with an oral LD50 estimated at 3.5 g/kg for rats. However, this is a single-dose value. Bromide ion is a cumulative toxin with a long half life: see potassium bromide. Information about NaBr. Bromide Poisoning in Angola