Cobalt(II) chloride
Cobalt chloride is an inorganic compound of cobalt and chlorine, with the formula CoCl2. It is a sky blue crystalline solid; the compound forms several hydrates CoCl2•nH2O, for n = 1, 2, 6, 9. Claims of the formation of tri- and tetrahydrates have not been confirmed; the dihydrate is purple and hexahydrate is pink. It is supplied as the hexahydrate CoCl2·6H2O, one of the most used cobalt compounds in the lab; because of the ease of the hydration/dehydration reaction, the resulting color change, cobalt chloride is used as an indicator for water in desiccants. Niche uses of cobalt chloride include its role in organic synthesis and electroplating objects with cobalt metal. Cobalt chloride has been classified as a substance of high concern by the European Chemicals Agency as it is a suspected carcinogen. At room temperature, anhydrous cobalt chloride has the CdCl2 structure in which the cobalt ions are octahedrally coordinated. At about 706 C, the coordination is believed to change to tetrahedral; the vapor pressure has been reported as 7.6 mm Hg at the melting point.
Cobalt chloride is soluble in water. Under atmospheric pressure, the mass concentration of a saturated solution of CoCl2 in water is about 54% at the boiling point, 120.2 °C. Diluted aqueous solutions of CoCl2 contain the species 2+, besides chloride ions. Concentrated solutions are red at room temperature but become blue at higher temperatures; the crystal unit of the solid hexahydrate CoCl2•6H2O contains the neutral molecule trans-CoCl24 and two molecules of water of crystallization. This species dissolves in water and alcohol; the anhydrous salt is hygroscopic and the hexahydrate is deliquescent. Cobalt chloride can be prepared in aqueous solution from cobalt hydroxide or cobalt carbonate and hydrochloric acid: CoCO3 + 2 HCl → CoCl2 + CO2 Co2 + 2 HCl → CoCl2 + 2H2OThe solid dihydrate and hexahydrate can be obtained by evaporation. Cooling saturated aqueous solutions yields the dihydrate between 120.2 °C and 51.25 °C, the hexahydrate below 51.25 °C. Water ice, rather than cobalt chloride, will crystallize from solutions with concentration below 29%.
The monohydrate and the anhydrous forms can be obtained by cooling solutions only under high pressure, above 206 °C and 335 °C, respectively. The anhydrous compound can be prepared by heating the hydrates. On rapid heating or in a closed container, each of the 6-, 2-, 1- hydrates melts into a mixture of the next lower hydrate and a saturated solution -- at 51.25 °C, 206 °C, 335 °C, respectively. On slow heating in an open container, water evaporates out of each of the solid 6-, 2-, 1- hydrates, leaving the next lower hydrate -- at about 40 °C, 89 °C, 126 °C, respectively. Dehydration can be effected with trimethylsilyl chloride: CoCl2•6H2O + 12 3SiCl → CoCl2 + 62O + 12 HClThe anhydrous compound can be purified by sublimation in vacuum. In the laboratory, cobalt chloride serves as a common precursor to other cobalt compounds. Aqueous solutions of the salt behave like other cobalt salts since these solutions consist of the 2+ ion regardless of the anion. For example, such solutions give a precipitate of Cobalt sulfide CoS upon treatment with hydrogen sulfide H2S.
The hexahydrate and the anhydrous salt are weak Lewis acids. The adducts are either octahedral or tetrahedral. With pyridine, one obtains an octahedral complex: CoCl2·6H2O + 4 C5H5N → CoCl24 + 6 H2OWith triphenylphosphine, a tetrahedral complex results: CoCl2·6H2O + 2 P3 → CoCl22 + 6 H2OSalts of the anionic complex CoCl42− can be prepared using tetraethylammonium chloride: CoCl2 + 2 Cl → 2The 2− ion is the blue ion that forms upon addition of hydrochloric acid to aqueous solutions of hydrated cobalt chloride, which are pink. Reaction of the anhydrous compound with sodium cyclopentadienide gives cobaltocene Co2; this 19-electron species is a good reducing agent, being oxidised to the yellow 18-electron cobaltacenium cation +. Compounds of cobalt in the +3 oxidation state exist, such as cobalt fluoride CoF3, nitrate Co3, sulfate Co23. On the other hand, cobalt chlorides can be obtained if the cobalt is bound to other ligands of greater Lewis basicity than chloride, such as amines. For example, in the presence of ammonia, cobalt chloride is oxidised by atmospheric oxygen to hexamminecobalt chloride: 4 CoCl2·6H2O + 4 NH4Cl + 20 NH3 + O2 → 4 Cl3 + 26 H2OSimilar reactions occur with other amines.
These reactions are performed in the presence of charcoal as a catalyst, or with hydrogen peroxide H2O2 substituted for atmospheric oxygen. Other basic ligands, including carbonate and oxalate, induce the formation of Co derivatives. Simple carboxylates and halides do not. Unlike Co complexes, Co complexes are slow to exchange ligands, so they are said to be kinetically inert; the German chemist Alfred Werner was awarded the Nobel prize in 1913 for his studies on a series of these cobalt compounds, work that led to an understanding of the structures of such coordination compounds. Reaction of 1-norbonyllithium with the CoCl2·THF in pentane produces the brown, thermally stable cobalt tetralkyl — a rare example of a stable transition metal/saturated alkane compound, different products are obtained in other solvents. In 2005–06, cobalt chloride was the eighth-most-prevalent allergen in p
Crystal
A crystal or crystalline solid is a solid material whose constituents are arranged in a ordered microscopic structure, forming a crystal lattice that extends in all directions. In addition, macroscopic single crystals are identifiable by their geometrical shape, consisting of flat faces with specific, characteristic orientations; the scientific study of crystals and crystal formation is known as crystallography. The process of crystal formation via mechanisms of crystal growth is called crystallization or solidification; the word crystal derives from the Ancient Greek word κρύσταλλος, meaning both "ice" and "rock crystal", from κρύος, "icy cold, frost". Examples of large crystals include snowflakes and table salt. Most inorganic solids are not crystals but polycrystals, i.e. many microscopic crystals fused together into a single solid. Examples of polycrystals include most metals, rocks and ice. A third category of solids is amorphous solids, where the atoms have no periodic structure whatsoever.
Examples of amorphous solids include glass and many plastics. Despite the name, lead crystal, crystal glass, related products are not crystals, but rather types of glass, i.e. amorphous solids. Crystals are used in pseudoscientific practices such as crystal therapy, along with gemstones, are sometimes associated with spellwork in Wiccan beliefs and related religious movements; the scientific definition of a "crystal" is based on the microscopic arrangement of atoms inside it, called the crystal structure. A crystal is a solid where the atoms form a periodic arrangement.. Not all solids are crystals. For example, when liquid water starts freezing, the phase change begins with small ice crystals that grow until they fuse, forming a polycrystalline structure. In the final block of ice, each of the small crystals is a true crystal with a periodic arrangement of atoms, but the whole polycrystal does not have a periodic arrangement of atoms, because the periodic pattern is broken at the grain boundaries.
Most macroscopic inorganic solids are polycrystalline, including all metals, ice, etc. Solids that are neither crystalline nor polycrystalline, such as glass, are called amorphous solids called glassy, vitreous, or noncrystalline; these have no periodic order microscopically. There are distinct differences between crystalline solids and amorphous solids: most notably, the process of forming a glass does not release the latent heat of fusion, but forming a crystal does. A crystal structure is characterized by its unit cell, a small imaginary box containing one or more atoms in a specific spatial arrangement; the unit cells are stacked in three-dimensional space to form the crystal. The symmetry of a crystal is constrained by the requirement that the unit cells stack with no gaps. There are 219 possible crystal symmetries, called crystallographic space groups; these are grouped into 7 crystal systems, such as hexagonal crystal system. Crystals are recognized by their shape, consisting of flat faces with sharp angles.
These shape characteristics are not necessary for a crystal—a crystal is scientifically defined by its microscopic atomic arrangement, not its macroscopic shape—but the characteristic macroscopic shape is present and easy to see. Euhedral crystals are those with well-formed flat faces. Anhedral crystals do not because the crystal is one grain in a polycrystalline solid; the flat faces of a euhedral crystal are oriented in a specific way relative to the underlying atomic arrangement of the crystal: they are planes of low Miller index. This occurs; as a crystal grows, new atoms attach to the rougher and less stable parts of the surface, but less to the flat, stable surfaces. Therefore, the flat surfaces tend to grow larger and smoother, until the whole crystal surface consists of these plane surfaces. One of the oldest techniques in the science of crystallography consists of measuring the three-dimensional orientations of the faces of a crystal, using them to infer the underlying crystal symmetry.
A crystal's habit is its visible external shape. This is determined by the crystal structure, the specific crystal chemistry and bonding, the conditions under which the crystal formed. By volume and weight, the largest concentrations of crystals in the Earth are part of its solid bedrock. Crystals found in rocks range in size from a fraction of a millimetre to several centimetres across, although exceptionally large crystals are found; as of 1999, the world's largest known occurring crystal is a crystal of beryl from Malakialina, Madagascar, 18 m long and 3.5 m in diameter, weighing 380,000 kg. Some crystals have formed by magmatic and metamorphic processes, giving origin to large masses of crystalline rock; the vast majority of igneous rocks are formed from molten magma and the degree of crystallization depends on the conditions under which they solidified. Such rocks as granite, which have cooled slowly and under great pressures, have crystallized.
Drinking
Drinking is the act of ingesting water or other liquids into the body through the mouth. Water is required for many of life’s physiological processes. Both excessive and inadequate water intake are associated with health problems; when a liquid is poured into an open human mouth, the swallowing process is completed by peristalsis which delivers the liquid to the stomach. The liquid may be poured from the hands or drinkware may be used as vessels. Drinking can be performed by acts of inhalation when imbibing hot liquids or drinking from a spoon. Infants employ a method of suction wherein the lips are pressed tight around a source, as in breastfeeding: a combination of breath and tongue movement creates a vacuum which draws in liquid. Amphibians and aquatic animals which live in freshwater do not need to drink: they absorb water through the skin by osmosis. Saltwater fish, drink through the mouth as they swim, purge the excess salt through the gills. By necessity, terrestrial animals in captivity become accustomed to drinking water, but most free-roaming animals stay hydrated through the fluids and moisture in fresh food.
When conditions impel them to drink from bodies of water, the methods and motions differ among species. Many desert animals do not drink if water becomes available, but rely on eating succulent plants. Cats and ruminants all lower the neck and lap in water with their powerful tongues. Cats and canines lap up water with the tongue in a spoon-like shape. Ruminants and most other herbivores submerge the tip of the mouth in order to draw in water by means of a plunging action with the tongue held straight. Cats drink at a slower pace than ruminants, who face greater natural predation hazards. Uniquely, elephants squirt it into their mouths. Most birds scoop or draw water into the buccal areas of their bills and tilting their heads back to drink. An exception is the common pigeon. Like nearly all other life forms, humans require water for tissue hydration. Lack of hydration causes thirst, a desire to drink, regulated by the hypothalamus in response to subtle changes in the body's electrolyte levels and blood volume.
A decline in total body water is called dehydration and will lead to death by hypernatremia. Methods used in the management of dehydration include oral rehydration therapy. An overconsumption of water can lead to water intoxication, which can dangerously dilute the concentration of salts in the body. Overhydration sometimes occurs among athletes and outdoor laborers, but it can be a sign of disease or damage to the hypothalamus. A persistent desire to drink inordinate quantities of water is a psychological condition termed polydipsia, it is accompanied by polyuria and may itself be a symptom of Diabetes mellitus or Diabetes insipidus. A daily intake of water is required for the normal physiological functioning of the human body; the USDA recommends a daily intake of total water: not by drinking but by consumption of water contained in other beverages and foods. The recommended intake is 3.7 liters per day for an adult male, 2.7 liters for an adult female. Other sources, claim that a high intake of fresh drinking water and distinct from other sources of moisture, is necessary for good health – eight servings per day of eight fluid ounces is the amount recommended by many nutritionists, although there is no scientific evidence supporting this recommendation.
The term “drinking” is used metonymically for the consumption of alcoholic beverages. Most cultures throughout history have incorporated some number of the wide variety of "strong drinks" into their meals, ceremonies and other occasions. Evidence of fermented drinks in human culture goes back as early as the Neolithic Period, the first pictorial evidence can be found in Egypt around 4,000 BC. Alcohol consumption has developed into a variety of well-established drinking cultures around the world. Despite its popularity, alcohol consumption poses significant health risks. Alcohol abuse and the addiction of alcoholism are common maladies in developed countries worldwide. A high rate of consumption can lead to cirrhosis, gout, hypertension, various forms of cancer, numerous other illnesses. Eating Hydration BibliographyBroom, Donald M.. Biology of Behaviour: Mechanisms and Applications. Cambridge: Cambridge University Press. ISBN 0-521-29906-3. Retrieved 31 August 2013. Curtis, Helena. Invitation to Biology.
Macmillan. ISBN 0879016795. Retrieved 31 August 2013. Fiebach, Nicholas H. ed.. Principles of Ambulatory Medicine. Lippincott Williams & Wilkins. ISBN 0-7817-6227-8. Retrieved 31 August 2013. Flint, Austin; the Physiology of Man. New York: D. Appleton and Co. OCLC 5357686. Retrieved 31 August 2013. Gately, Iain. Drink: A Cultural History of Alcohol. New York: Penguin. Pp. 1–14. ISBN 1-59240-464-2. Retrieved 31 August 2013. Mayer, William. Physiological Mammalogy. II. Elsevier. ISBN 9780323155250. Retrieved 31 August 2013. Provan, Drew. Oxford Handbook of Clinical and Laboratory Investigation. Oxford: Oxford University Press. ISBN 0-19-923371-3. Retrieved 31 August 2013. Smith, Robert Meade; the Physiology of the Domestic Animals. Philadelphia, London: F. A. Davis. Retrieved 31 August 2013. "Are You Drinking Enough?", recommendations by the European Hydration Institute
Efflorescence
In chemistry, efflorescence is the migration of a salt to the surface of a porous material, where it forms a coating. The essential process involves the dissolving of an internally held salt in water, or in another solvent; the water, with the salt now held in solution, migrates to the surface evaporates, leaving a coating of the salt. In what has been described as "primary efflorescence," the water is the invader and the salt was present internally; some people describe a reverse process, where the salt is present externally and is carried inside in solution, as "secondary efflorescence." However, others would give this latter phenomenon another name entirely. Efflorescences can occur in natural and built environments. On porous construction materials it may present a cosmetic outer problem only, but can sometimes indicate internal structural weakness. Efflorescence may clog the pores of porous materials, resulting the destruction of those materials by internal water pressure, as seen in the spalling of brick.
A 5 molar concentration aqueous droplet of NaCl will spontaneously crystallize at 45% relative humidity to form an NaCl cube by the mechanism of homogeneous nucleation. The original water is released to the gas phase. Gypsum is a hydrate solid that, in a sufficiently dry environment, will give up its water to the gas phase and form anhydrite. Copper sulfate is a blue crystalline solid that when exposed to air loses water of crystallization from its surface to form a white layer of anhydrous copper sulfate. Sodium carbonate deca hydrate will lose water. Primary efflorescence is named such, as it occurs during the initial cure of a cementitious product, it occurs on masonry construction brick, as well as some firestop mortars, when water moving through a wall or other structure, or water being driven out as a result of the heat of hydration as cement stone is being formed, brings salts to the surface that are not bound as part of the cement stone. As the water evaporates, it leaves the salt behind, which forms a white, fluffy deposit, that can be brushed off.
The resulting white deposits are referred to as "efflorescence" in this instance. In this context efflorescence is sometimes referred to as "saltpetering." Since primary efflorescence brings out salts that are not ordinarily part of the cement stone, it is not a structural, rather, an aesthetic concern. For controlling primary efflorescence, formulations containing liquid fatty acid mixtures have been used; the oily liquid admixture is introduced into the batch mix at an early stage by coating onto the sand particles prior to the introduction of any mix water, so that the oily admixture is distributed uniformly throughout the concrete batch mix. Secondary efflorescence is named such as it does not occur as a result of the forming of the cement stone or its accompanying hydration products. Rather, it is due to the external influence of concrete poisons, such as chlorides. A common example of where secondary efflorescence occurs is steel-reinforced concrete bridges as well as parking garages. Saline solutions are formed due to the presence of road salt in the winter.
This saline solution is absorbed into the concrete, where it can begin to dissolve cement stone, of primary structural importance. Virtual stalactites can be formed in some cases as a result of dissolved cement stone, hanging off cracks in concrete structures. Where this process has taken hold, the structural integrity of a concrete element is at risk; this is a common traffic building maintenance concern. Secondary efflorescence is akin to osteoporosis of the concrete. For controlling secondary efflorescence, admixtures containing aqueous-based calcium stearate dispersion are added at a stage of the batching process with the mix water. In a typical batching process, sand is first charged into the mixer the oil-based primary anti-efflorescence admixture is added with constant mixing to allow the oil to coat the sand. Coarse aggregates and cement are added, followed by water. If CSD is used, it is introduced at this point during or after the addition of the mix water. CSD is an aqueous dispersion wherein fine solid particles of calcium stearate are suspended in the water uniformly.
Commercially available CSD has an average particle size of about 1 to 10 micrometres. The uniform distribution of CSD in the mix may render the resulting concrete masonry unit water repellent, as CSD particles are well distributed in the pores of the unit to interfere with the capillary movement of water. Calthemite is a secondary deposit derived from concrete, mortar or lime, which can be mistakenly assumed to be efflorescence. Calthemites are deposited as calcite, the most stable polymorph of calcium carbonate; the only way to and permanently prevent efflorescence in cementitious materials is by using special admixtures that chemically react with and bind the salt-based impurities in the concrete when hydrogen is present. The chemical reaction in these special additives fuses the sodium chloride on a nanomolecular level, converting it into non-sodium chemicals and other harmless matter that will not leach out or migrate to the surface. In fact, the nanotechnology in these additives can be up to 100,000 times smaller than the smallest cement particles, allowing their molecules to pass through cement minerals or sand particles and be
Heavy water
Heavy water is a form of water that contains a larger than normal amount of the hydrogen isotope deuterium, rather than the common hydrogen-1 isotope that makes up most of the hydrogen in normal water. The presence of deuterium gives the water different nuclear properties, the increase of mass gives it different physical and chemical properties when compared to normal water. Deuterium is a hydrogen isotope with a nucleus containing a proton; the additional neutron makes a deuterium atom twice as heavy as a protium atom. A molecule of heavy water has two deuterium atoms in place of the two protium atoms of ordinary "light" water; the weight of a heavy water molecule, however, is not different from that of a normal water molecule, because about 89% of the molecular weight of water comes from the single oxygen atom rather than the two hydrogen atoms. The colloquial term'heavy water' refers to a enriched water mixture that contains deuterium oxide D2O, but some hydrogen-deuterium oxide and a smaller amount of ordinary hydrogen oxide H2O.
For instance, the heavy water used in CANDU reactors is 99.75% enriched by hydrogen atom-fraction—meaning that 99.75% of the hydrogen atoms are of the heavy type. For comparison, ordinary water contains only about 156 deuterium atoms per million hydrogen atoms, meaning that 0.0156% of the hydrogen atoms are of the heavy type. Heavy water is not radioactive. In its pure form, it has a density about 11% greater than water, but is otherwise physically and chemically similar; the various differences in deuterium-containing water are larger than in any other occurring isotope-substituted compound because deuterium is unique among heavy stable isotopes in being twice as heavy as the lightest isotope. This difference increases the strength of water's hydrogen-oxygen bonds, this in turn is enough to cause differences that are important to some biochemical reactions; the human body contains deuterium equivalent to about five grams of heavy water, harmless. When a large fraction of water in higher organisms is replaced by heavy water, the result is cell dysfunction and death.
Heavy water was first produced in a few months after the discovery of deuterium. With the discovery of nuclear fission in late 1938, the need for a neutron moderator that captured few neutrons, heavy water became a component of early nuclear energy research. Since heavy water has been an essential component in some types of reactors, both those that generate power and those designed to produce isotopes for nuclear weapons; these heavy water reactors have the advantage of being able to run on natural uranium without using graphite moderators that pose radiological and dust explosion hazards in the decommissioning phase. Most modern reactors use enriched uranium with ordinary water as the moderator. Semiheavy water, HDO, exists whenever there is water with light deuterium in the mix; this is because hydrogen atoms are exchanged between water molecules. Water containing 50% H and 50% D in its hydrogen contains about 50% HDO and 25% each of H2O and D2O, in dynamic equilibrium. In normal water, about 1 molecule in 3,200 is HDO, heavy water molecules only occur in a proportion of about 1 molecule in 41 million.
Thus semiheavy water molecules are far more common than "pure" heavy water molecules. Water enriched in the heavier oxygen isotopes 17O and 18O is commercially available, e.g. for use as a non-radioactive isotopic tracer. It is "heavy water" as it is denser than normal water —but is called heavy water, since it does not contain the deuterium that gives D2O its unusual nuclear and biological properties, it is more expensive than D2O due to the more difficult separation of 17O and 18O. H218O is used for production of fluorine-18 for radiopharmaceuticals and radiotracers and for positron emission tomography. Tritiated water contains tritium in place of protium or deuterium, therefore it is radioactive; the physical properties of water and heavy water differ in several respects. Heavy water is less dissociated than light water at given temperature, the true concentration of D+ ions is less than H+ ions would be for a light water sample at the same temperature; the same is true of OD OH − ions. For heavy water Kw D2O = 1.35 × 10−15, must equal for neutral water.
Thus pKw D2O = p + p = 7.44 + 7.44 = 14.87, the p of neutral heavy water at 25.0 °C is 7.44. The pD of heavy water is measured using pH electrodes giving a pH value, or pHa, at various temperatures a true acidic pD can be estimated from the directly pH meter measured pHa, such that pD+ = pHa + 0.41. The electrode correction for alkaline conditions is 0.456 for heavy water. The alkaline correction is pD+ = pHa + 0.456. These corrections are different from the differences in p and p of 0.44 from the corresponding ones in heavy water. Heavy water is 10.6% denser than ordinary water, heavy water's physically different properties can be seen without equipment if a frozen sample is dropped into normal water, as it will sink. If the water is ice-cold the higher melting tem
Hydrogen bond
A hydrogen bond is a electrostatic force of attraction between a hydrogen atom, covalently bound to a more electronegative atom or group the second-row elements nitrogen, oxygen, or fluorine —the hydrogen bond donor —and another electronegative atom bearing a lone pair of electrons—the hydrogen bond acceptor. Such an interacting system is denoted Dn–H···Ac, where the solid line denotes a covalent bond, the dotted line indicates the hydrogen bond. There is general agreement that there is a minor covalent component to hydrogen bonding for moderate to strong hydrogen bonds, although the importance of covalency in hydrogen bonding is debated. At the opposite end of the scale, there is no clear boundary between a weak hydrogen bond and a van der Waals interaction. Weaker hydrogen bonds are known for hydrogen atoms bound to elements such as chlorine; the hydrogen bond is responsible for many of the anomalous physical and chemical properties of compounds of N, O, F. Hydrogen bonds can be intramolecular.
Depending on the nature of the donor and acceptor atoms which constitute the bond, their geometry, environment, the energy of a hydrogen bond can vary between 1 and 40 kcal/mol. This makes them somewhat stronger than a van der Waals interaction, weaker than covalent or ionic bonds; this type of bond can occur in inorganic molecules such as water and in organic molecules like DNA and proteins. Intermolecular hydrogen bonding is responsible for the high boiling point of water compared to the other group 16 hydrides that have much weaker hydrogen bonds. Intramolecular hydrogen bonding is responsible for the secondary and tertiary structures of proteins and nucleic acids, it plays an important role in the structure of polymers, both synthetic and natural. It was recognized that there are many examples of weaker hydrogen bonding involving donor Dn other than N, O, or F and/or acceptor Ac with close to or the same electronegativity as hydrogen. Though they are quite weak, they are ubiquitous and are recognized as important control elements in receptor-ligand interactions in medicinal chemistry or intra-/intermolecular interactions in materials sciences.
Thus, there is a trend of gradual broadening for the definition of hydrogen bonding. In 2011, an IUPAC Task Group recommended a modern evidence-based definition of hydrogen bonding, published in the IUPAC journal Pure and Applied Chemistry; this definition specifies: The hydrogen bond is an attractive interaction between a hydrogen atom from a molecule or a molecular fragment X–H in which X is more electronegative than H, an atom or a group of atoms in the same or a different molecule, in which there is evidence of bond formation. Most introductory textbooks still restrict the definition of hydrogen bond to the "classical" type of hydrogen bond characterized in the opening paragraph. A hydrogen atom attached to a electronegative atom is the hydrogen bond donor. C-H bonds only participate in hydrogen bonding when the carbon atom is bound to electronegative substituents, as is the case in chloroform, CHCl3. In a hydrogen bond, the electronegative atom not covalently attached to the hydrogen is named proton acceptor, whereas the one covalently bound to the hydrogen is named the proton donor.
In the donor molecule, the H center is protic. The donor is a Lewis base. Hydrogen bonds are represented as H · · · Y system. Liquids that display hydrogen bonding are called associated liquids; the hydrogen bond is described as an electrostatic dipole-dipole interaction. However, it has some features of covalent bonding: it is directional and strong, produces interatomic distances shorter than the sum of the van der Waals radii, involves a limited number of interaction partners, which can be interpreted as a type of valence; these covalent features are more substantial when acceptors bind hydrogens from more electronegative donors. Hydrogen bonds can vary in strength from weak to strong. Typical enthalpies in vapor include: F−H···:F, illustrated uniquely by HF2−, bifluoride O−H···:N, illustrated water-ammonia O−H···:O, illustrated water-water, alcohol-alcohol N−H···:N, illustrated by ammonia-ammonia N−H···:O, illustrated water-amide HO−H···:OH+3 The strength of intermolecular hydrogen bonds is most evaluated by measurements of equilibria between molecules containing donor and/or acceptor units, most in solution.
The strength of intramolecular hydrogen bonds can be studied with equilibria between conformers with and without hydrogen bonds. The most important method for the identification of hydrogen bonds in complicated molecules is crystallography, sometimes NMR-spectroscopy. Structural details, in particular distances between donor and acceptor which are smaller than the sum of the van der Waals radii can be taken as indication of the hydrogen bond strength. One scheme gives the following somewhat arbitrary classification: those that are 15 to 40 kcal/mol, 5 to 15 kcal/mol, >0 to 5 kcal/mol are considered strong, moder
Tetrahydrofuran
Tetrahydrofuran is an organic compound with the formula 4O. The compound is classified as heterocyclic compound a cyclic ether, it is a water-miscible organic liquid with low viscosity. It is used as a precursor to polymers. Being polar and having a wide liquid range, THF is a versatile solvent. About 200,000 tonnes of tetrahydrofuran are produced annually; the most used industrial process involves the acid-catalyzed dehydration of 1,4-butanediol. Ashland/ISP is one the biggest producers of this chemical route; the method is similar to the production of diethyl ether from ethanol. The butanediol is derived from condensation of acetylene with formaldehyde followed by hydrogenation. DuPont developed a process for producing THF by oxidizing n-butane to crude maleic anhydride, followed by catalytic hydrogenation. A third major industrial route entails hydroformylation of allyl alcohol followed by hydrogenation to 1,4-butanediol. THF can be synthesized by catalytic hydrogenation of furan. Certain sugars can be converted to THF, although this method is not practiced.
Furan is thus derivable from renewable resources. In the presence of strong acids, THF converts to a linear polymer called poly glycol known as polytetramethylene oxide: n C4H8O → −n−This polymer is used to make elastomeric polyurethane fibers like Spandex; the other main application of THF is as an industrial solvent for polyvinyl chloride and in varnishes. It is an aprotic solvent with a dielectric constant of 7.6. It is a moderately polar solvent and can dissolve a wide range of nonpolar and polar chemical compounds. THF is water-miscible and can form solid clathrate hydrate structures with water at low temperatures. THF has been explored as a miscible co-solvent in aqueous solution to aid in the liquefaction and delignification of plant lignocellulosic biomass for production of renewable platform chemicals and sugars as potential precursors to biofuels. Aqueous THF augments the hydrolysis of glycans from biomass and dissolves the majority of biomass lignin making it a suitable solvent for biomass pretreatment.
THF is used in polymer science. For example, it can be used to dissolve polymers prior to determining their molecular mass using gel permeation chromatography. THF dissolves PVC as well, thus it is the main ingredient in PVC adhesives, it can be used to liquefy old PVC cement and is used industrially to degrease metal parts. THF is used as a component in mobile phases for reversed-phase liquid chromatography, it has a greater elution strength than methanol or acetonitrile, but is less used than these solvents. THF is used as a solvent in 3D printing, it can be used to clean clogged 3D printer parts, as well as when finishing prints to remove extruder lines and add a shine to the finished product. In the laboratory, THF is a popular solvent, it is more basic than diethyl ether and forms stronger complexes with Li+, Mg2+, boranes. It is a popular solvent for hydroboration reactions and for organometallic compounds such as organolithium and Grignard reagents. Although similar to diethyl ether, THF is a stronger base.
Thus, while diethyl ether remains the solvent of choice for some reactions, THF fills that role in many others, where strong coordination is desirable and the precise properties of ethereal solvents such as these allows fine-tuning modern chemical reactions. Commercial THF contains substantial water that must be removed for sensitive operations, e.g. those involving organometallic compounds. Although THF is traditionally dried by distillation from an aggressive desiccant, molecular sieves are superior. THF is a weak Lewis base. Typical complexes are of the stoichiometry MCl33; such compounds are used reagents. In the presence of a solid acid catalyst, THF reacts with hydrogen sulfide to give tetrahydrothiophene. THF is a nontoxic solvent, with the median lethal dose comparable to that for acetone. Reflecting its remarkable solvent properties, it penetrates the skin. THF dissolves latex and is handled with nitrile or neoprene rubber gloves, it is flammable. One danger posed by THF follows from its tendency to form explosive peroxides on storage in air.
To minimize this problem, commercial samples of THF are inhibited with butylated hydroxytoluene. Distillation of THF to dryness is avoided because the explosive peroxides concentrate in the residue. Polytetrahydrofuran 2-Methyltetrahydrofuran Trapp mixture Other cyclic ethers: oxirane, oxane Loudon, G. Mark. Organic Chemistry. New York: Oxford University Press. P. 318. ISBN 9780981519432. International Chemical Safety Card 0578 NIOSH Pocket Guide to Chemical Hazards U. S. OSHA info on THF "2-Methyltetrahydrofuran, An alternative to Tetrahydrofuran and Dichloromethane". Sigma-Aldrich. Retrieved 2007-05-23
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