Supersaturation is a solution that contains more of the dissolved material than could be dissolved by the solvent under normal circumstances. It can refer to a vapor of a compound that has a higher pressure than the vapor pressure of that compound. Special conditions need to be met in order to generate a supersaturated solution. One of the easiest ways to do this relies on the temperature dependence of solubility; as a general rule, the more heat is added to a system, the more soluble a substance becomes.. Therefore, at high temperatures, more solute can be dissolved than at lower temperatures. If this solution were to be cooled at a rate faster than the rate of precipitation, the solution will become supersaturated until the solute precipitates to the temperature-determined saturation point; the precipitation or crystallization of the solute takes longer than the actual cooling time because the molecules need to meet up and form the precipitate without being knocked apart by water. Thus, the larger the molecule, the longer the solute will take to crystallize due to the principles of Brownian motion.
The condition of supersaturation does not have to be reached through the manipulation of heat. The ideal gas law suggests that pressure and volume can be changed to force a system into a supersaturated state. If the volume of solvent is decreased, the concentration of the solute can be above the saturation point and thus create a supersaturated solution; the decrease in volume is most generated through evaporation. An increase in pressure can drive a solution to a supersaturated state. All three of these mechanisms rely on the fact that the conditions of the solution can be changed quicker than the solute can precipitate or crystallize out. Supersaturated solutions will undergo crystallization under specific conditions. In a normal solution, once the maximum amount of solute is dissolved, adding more solute would either cause the dissolved solute to precipitate out and/or for the solute to not dissolve at all. There are cases wherein solubility of a saturated solution is decreased by manipulating temperature, pressure, or volume but a supersaturated state does not occur.
In these cases, the solute will precipitate out. This is. A supersaturated solution of gases in a liquid may form bubbles. Supersaturation may be defined as a sum of all gas partial pressures in the liquid which exceeds the ambient pressure in the liquid. Crystallization will occur to allow the solution to reach a lower energy state.. The activation energy comes in the form of a nuclei crystal being added to the liquid solution; this nuclei can be either added from another source, known as seeding, or can spontaneously form within the solution due to ion and molecule interactions. This process is known as primary nucleation, it is necessary for the nuclei to be identical to the solute, crystallizing. This will allow for the dissolved ions to build up on the nuclei and each other in the process of crystal growth or secondary nucleation. There are a multitude of factors that will affect the rate and order of magnitude with which crystallization proceeds as well as the difference in formation of crystallites and single crystals.
A crystallization phase diagram shows where undersaturation and supersaturation occur at certain concentrations. Concentrations below the solubility curve result in an undersaturation solution. Saturation occurs. If the concentrations are above the solubility curve, the solution is considered supersaturated. There are three mechanisms with which supersaturation occurs: precipitation and metastable. In the precipitation zone, the molecules in a solution are in excess and will separate from the solution to form amorphous aggregates; the excess of molecules aggregate to form a crystalline structure. In the metastable zone, the solution takes time to nucleate. In order to grow crystals while in the metastable zone, the conditions would require the formation of one nucleus while in the nucleation zone, just past the metastable region; the supersaturated solution can return to the metastable region. Supersaturation in vapor phase is related to the surface tension of liquids through the Kelvin equation, the Gibbs–Thomson effect and the Poynting effect.
The International Association for the Properties of Water and Steam provides a special equation for the Gibbs free energy in the metastable-vapor region of water in its Revised Release on the IAPWS Industrial Formulation 1997 for the Thermodynamic Properties of Water and Steam. All thermodynamic properties for the metastable-vapor region of water can be derived from this equation by means of the appropriate relations of thermodynamic properties to the Gibbs free energy. Table 1. Supersaturation measurement methods. Supersaturation has been a frequent topic of research throughout history. Early studies of these solutions were conducted with sodium sulfate known as Glauber’s Salt, due to the stability of the crystal and the rising role it had in industry. Through the use of this salt, an important scientific discovery was made by Jean-Baptiste Ziz, a botanist from Mayence, in 1809, his experiments allowed him to conclude that the crystallization of a supersaturated solution does not come from its agitation, but from solid matter entering and acting as a “starting” site for crystals to form
Sulfuryl chloride is an inorganic compound with the formula SO2Cl2. At room temperature, it is a colorless liquid with a pungent odor. Sulfuryl chloride is not found in nature. Sulfuryl chloride is confused with thionyl chloride, SOCl2; the properties of these two sulfur oxychlorides are quite different: sulfuryl chloride is a source of chlorine whereas thionyl chloride is a source of chloride ions. An alternative IUPAC name is sulfuroyl dichloride. Sulfur is tetrahedral in SO2Cl2 and the oxidation state of the sulfur atom is +6, as in sulfuric acid. SO2Cl2 is prepared by the reaction of sulfur dioxide and chlorine in the presence of a catalyst, such as activated carbon. SO2 + Cl2 → SO2Cl2The crude product can be purified by fractional distillation, it is uncommon to prepare SO2Cl2 in the laboratory. Sulfuryl chloride can be considered a derivative of sulfuric acid. Sulfuryl chloride was first prepared in 1838 by the French chemist Henri Victor Regnault. Sulfuryl chloride reacts with water, releasing hydrogen chloride gas and sulfuric acid: 2 H2O + SO2Cl2 → 2 HCl + H2SO4SO2Cl2 will decompose when heated to or above 100 °C, about 30 °C above its boiling point.
Upon standing, SO2Cl2 decomposes to sulfur dioxide and chlorine, which gives the older samples a yellowish color. Sulfuryl chloride is used as a source of Cl2; because it is a pourable liquid, it is considered more convenient than Cl2 to measure and dispense. SO2Cl2 is used as a reagent in the conversion of C−H to C−Cl adjacent to activating substituents such as carbonyls and sulfoxides, it chlorinates alkanes, alkynes, aromatics and epoxides. Such reactions occur under free radical conditions using an initiator such as AIBN, it can be used to convert thiols or disulfides into their corresponding sulfenyl chlorides, though sulfinyl chlorides result from thiols in some cases. SO2Cl2 can convert alcohols to alkyl chlorides. In industry, sulfuryl chloride is most used in producing pesticides. SO2Cl2 can be used to treat wool to prevent shrinking. SO2Cl2 is toxic and acts as a lachrymator, it can form fuming mixtures with water, as well as donor solvents such as dimethyl sulfoxide and dimethylformamide
Xenon is a chemical element with symbol Xe and atomic number 54. It is a colorless, odorless noble gas found in the Earth's atmosphere in trace amounts. Although unreactive, xenon can undergo a few chemical reactions such as the formation of xenon hexafluoroplatinate, the first noble gas compound to be synthesized. Xenon is used in flash lamps and arc lamps, as a general anesthetic; the first excimer laser design used a xenon dimer molecule as the lasing medium, the earliest laser designs used xenon flash lamps as pumps. Xenon is used to search for hypothetical weakly interacting massive particles and as the propellant for ion thrusters in spacecraft. Occurring xenon consists of eight stable isotopes. More than 40 unstable xenon isotopes undergo radioactive decay, the isotope ratios of xenon are an important tool for studying the early history of the Solar System. Radioactive xenon-135 is produced by beta decay from iodine-135, is the most significant neutron absorber in nuclear reactors. Xenon was discovered in England by the Scottish chemist William Ramsay and English chemist Morris Travers in September 1898, shortly after their discovery of the elements krypton and neon.
They found xenon in the residue left over from evaporating components of liquid air. Ramsay suggested the name xenon for this gas from the Greek word ξένον, neuter singular form of ξένος, meaning'foreign','strange', or'guest'. In 1902, Ramsay estimated the proportion of xenon in the Earth's atmosphere to be one part in 20 million. During the 1930s, American engineer Harold Edgerton began exploring strobe light technology for high speed photography; this led him to the invention of the xenon flash lamp in which light is generated by passing brief electric current through a tube filled with xenon gas. In 1934, Edgerton was able to generate flashes as brief as one microsecond with this method. In 1939, American physician Albert R. Behnke Jr. began exploring the causes of "drunkenness" in deep-sea divers. He tested the effects of varying the breathing mixtures on his subjects, discovered that this caused the divers to perceive a change in depth. From his results, he deduced. Although Russian toxicologist Nikolay V. Lazarev studied xenon anesthesia in 1941, the first published report confirming xenon anesthesia was in 1946 by American medical researcher John H. Lawrence, who experimented on mice.
Xenon was first used as a surgical anesthetic in 1951 by American anesthesiologist Stuart C. Cullen, who used it with two patients. Xenon and the other noble gases were for a long time considered to be chemically inert and not able to form compounds. However, while teaching at the University of British Columbia, Neil Bartlett discovered that the gas platinum hexafluoride was a powerful oxidizing agent that could oxidize oxygen gas to form dioxygenyl hexafluoroplatinate. Since O2 and xenon have the same first ionization potential, Bartlett realized that platinum hexafluoride might be able to oxidize xenon. On March 23, 1962, he mixed the two gases and produced the first known compound of a noble gas, xenon hexafluoroplatinate. Bartlett thought its composition to be Xe+−, but work revealed that it was a mixture of various xenon-containing salts. Since many other xenon compounds have been discovered, in addition to some compounds of the noble gases argon and radon, including argon fluorohydride, krypton difluoride, radon fluoride.
By 1971, more than 80 xenon compounds were known. In November 1989, IBM scientists demonstrated a technology capable of manipulating individual atoms; the program, called IBM in atoms, used a scanning tunneling microscope to arrange 35 individual xenon atoms on a substrate of chilled crystal of nickel to spell out the three letter company initialism. It was the first time atoms had been positioned on a flat surface. Xenon has atomic number 54. At standard temperature and pressure, pure xenon gas has a density of 5.761 kg/m3, about 4.5 times the density of the Earth's atmosphere at sea level, 1.217 kg/m3. As a liquid, xenon has a density of up to 3.100 g/mL, with the density maximum occurring at the triple point. Liquid xenon has a high polarizability due to its large atomic volume, thus is an excellent solvent, it can dissolve hydrocarbons, biological molecules, water. Under the same conditions, the density of solid xenon, 3.640 g/cm3, is greater than the average density of granite, 2.75 g/cm3.
Under gigapascals of pressure, xenon forms a metallic phase. Solid xenon changes from face-centered cubic to hexagonal close packed crystal phase under pressure and begins to turn metallic at about 140 GPa, with no noticeable volume change in the hcp phase, it is metallic at 155 GPa. When metallized, xenon appears sky blue because it absorbs red light and transmits other visible frequencies; such behavior is unusual for a metal and is explained by the small width of the electron bands in that state. Liquid or solid xenon nanoparticles can be formed at room temperature by implanting Xe+ ions into a solid matrix. Many solids have lattice constants smaller than solid Xe; this results in compression of the implanted Xe to pressures that may be sufficient for its liquefaction or solidification. Xenon is a member of the zero-valence elements that are called inert gases, it is inert to most common chemical reactions because the outer valence shell contains eight electrons. This produces a stable, minimum energy configuration in which the outer electrons are bound.
In a gas-filled tube, xenon em
In chemistry, a solution is a special type of homogeneous mixture composed of two or more substances. In such a mixture, a solute is a substance dissolved in another substance, known as a solvent; the mixing process of a solution happens at a scale where the effects of chemical polarity are involved, resulting in interactions that are specific to solvation. The solution assumes the phase of the solvent when the solvent is the larger fraction of the mixture, as is the case; the concentration of a solute in a solution is the mass of that solute expressed as a percentage of the mass of the whole solution. The term aqueous solution is. A solution is a homogeneous mixture of two or more substances; the particles of solute in a solution cannot be seen by the naked eye. A solution does not allow beams of light to scatter. A solution is stable; the solute from a solution cannot be separated by filtration. It is composed of only one phase. Homogeneous means. Heterogeneous means; the properties of the mixture can be uniformly distributed through the volume but only in absence of diffusion phenomena or after their completion.
The substance present in the greatest amount is considered the solvent. Solvents can be liquids or solids. One or more components present in the solution other; the solution has the same physical state as the solvent. If the solvent is a gas, only gases are dissolved under a given set of conditions. An example of a gaseous solution is air. Since interactions between molecules play no role, dilute gases form rather trivial solutions. In part of the literature, they are not classified as solutions, but addressed as mixtures. If the solvent is a liquid almost all gases and solids can be dissolved. Here are some examples: Gas in liquid: Oxygen in water Carbon dioxide in water – a less simple example, because the solution is accompanied by a chemical reaction. Note that the visible bubbles in carbonated water are not the dissolved gas, but only an effervescence of carbon dioxide that has come out of solution. Liquid in liquid: The mixing of two or more substances of the same chemistry but different concentrations to form a constant.
Alcoholic beverages are solutions of ethanol in water. Solid in liquid: Sucrose in water Sodium chloride or any other salt in water, which forms an electrolyte: When dissolving, salt dissociates into ions. Solutions in water are common, are called aqueous solutions. Non-aqueous solutions are. Counter examples are provided by liquid mixtures that are not homogeneous: colloids, emulsions are not considered solutions. Body fluids are examples for complex liquid solutions. Many of these are electrolytes. Furthermore, they contain solute molecules like urea. Oxygen and carbon dioxide are essential components of blood chemistry, where significant changes in their concentrations may be a sign of severe illness or injury. If the solvent is a solid gases and solids can be dissolved. Gas in solids: Hydrogen dissolves rather well in metals in palladium. Liquid in solid: Mercury in gold, forming an amalgam Water in solid salt or sugar, forming moist solids Hexane in paraffin wax Solid in solid: Steel a solution of carbon atoms in a crystalline matrix of iron atoms Alloys like bronze and many others Polymers containing plasticizers The ability of one compound to dissolve in another compound is called solubility.
When a liquid can dissolve in another liquid the two liquids are miscible. Two substances that can never mix to form a solution are said to be immiscible. All solutions have a positive entropy of mixing; the interactions between different molecules or ions may be energetically favored or not. If interactions are unfavorable the free energy decreases with increasing solute concentration. At some point the energy loss outweighs the entropy gain, no more solute particles can be dissolved. However, the point at which a solution can become saturated can change with different environmental factors, such as temperature and contamination. For some solute-solvent combinations a supersaturated solution can be prepared by raising the solubility to dissolve more solute, lowering it; the greater the temperature of the solvent, the more of a given solid solute it can dissolve. However, most gases and some compounds exhibit solubilities that decrease with increased temperature; such behavior is a result of an exothermic enthalpy of solution.
Some surfactants exhibit this behaviour. The solubility of liquids in liquids is less temperature-sensitive than that of solids or gases; the physical properties of compounds such as melting point and boiling point change when other compounds are added. Together they are called colligative properties. There are several ways to quantify the amount of one compound dissolved in the other compounds collectively called concentration. Examples include molarity, volume fraction, mole fraction; the properties of ideal solutions can be calculated by the linear combination of the properties of
Phosphoryl chloride is a colourless liquid with the formula POCl3. It hydrolyses in moist air releasing phosphoric acid and fumes of hydrogen chloride, it is manufactured industrially on a large scale from phosphorus trichloride and oxygen or phosphorus pentoxide. It is used to make phosphate esters such as tricresyl phosphate. Like phosphate, phosphoryl chloride is tetrahedral in shape, it features three P−Cl bonds and one strong P=O double bond, with an estimated bond dissociation energy of 533.5 kJ/mol. On the basis of bond length and electronegativity, the Schomaker-Stevenson rule suggests that the double bond form is dominant, in contrast with the case of POF3; the P=O bond involves the donation of the lone pair electrons on oxygen p-orbitals to the antibonding combinations associated with phosphorus-chlorine bonds, thus constituting π bonding. With a freezing point of 1 °C and boiling point of 106 °C, the liquid range of POCl3 is rather similar to water. Like water, POCl3 autoionizes, owing to the reversible formation of POCl2+,Cl−.
POCl3 reacts with water to give hydrogen chloride and phosphoric acid: O=PCl3 + 3 H2O → O=P3 + 3 HClIntermediates in the conversion have been isolated, including pyrophosphoryl chloride, P2O3Cl4. Upon treatment with excess alcohols and phenols, POCl3 gives phosphate esters: O=PCl3 + 3 ROH → O=P3 + 3 HClSuch reactions are performed in the presence of an HCl acceptor such as pyridine or an amine. POCl3 can act as a Lewis base, forming adducts with a variety of Lewis acids such as titanium tetrachloride: Cl3PO + TiCl4 → Cl3POTiCl4The aluminium chloride adduct is quite stable, so POCl3 can be used to remove AlCl3 from reaction mixtures, for example at the end of a Friedel-Crafts reaction. POCl3 reacts with hydrogen bromide in the presence of Lewis-acidic catalysts to produce POBr3. Phosphoryl chloride can be prepared by many methods. Phosphoryl chloride was first reported in 1847 by the French chemist Adolphe Wurtz by reacting phosphorus pentachloride with water; the commercial method involves oxidation of phosphorus trichloride with oxygen: 2 PCl3 + O2 → 2 POCl3A related reaction include the oxidation of phosphorus trichloride with potassium chlorate: 3 PCl3 + KClO3 → 3 POCl3 + KCl The reaction of phosphorus pentachloride with phosphorus pentoxide.
6 PCl5 + P4O10 → 10 POCl3The reaction can be simplified by chlorinating a mixture of PCl3 and P4O10, generating the PCl5 in situ. The reaction of phosphorus pentachloride with boric acid or oxalic acid: 3 PCl5 + 2 B3 → 3 POCl3 + B2O3 + 6 HCl PCl5 + 2 → POCl3 + CO + CO2 + 2 HCl Reduction of tricalcium phosphate with carbon in the presence of chlorine gas: Ca32 + 6 C + 6 Cl2 → 3 CaCl2 + 6 CO + 2 POCl3The reaction of phosphorus pentoxide with sodium chloride is reported: 2 P2O5 + 3 NaCl → 3 NaPO3 + POCl3. In one commercial application, phosphoryl chloride is used in the manufacture of phosphate esters. Triarylphosphates such as triphenyl phosphate and tricresyl phosphate are used as flame retardants and plasticisers for PVC. Trialkylphosphates such as tributyl phosphate are used as liquid–liquid extraction solvents in nuclear reprocessing and elsewhere. In the semiconductor industry, POCl3 is used as a safe liquid phosphorus source in diffusion processes; the phosphorus acts. In the laboratory, POCl3 is a reagent in dehydrations.
One example involves conversion of primary amides to nitriles: RCNH2 + POCl3 → RCN + "PO2Cl" + 2 HClIn a related reaction, certain aryl-substituted amides can be cyclised using the Bischler-Napieralski reaction. Such reactions are believed to proceed via an imidoyl chloride. In certain cases, the imidoyl chloride is the final product. For example and pyrimidones can be converted to chloro- derivatives such as 2-chloropyridines and 2-chloropyrimidines, which are intermediates in the pharmaceutical industry. In the Vilsmeier-Haack reaction, POCl3 reacts with amides to produce a "Vilsmeier reagent", a chloro-iminium salt, which subsequently reacts with electron-rich aromatic compounds to produce aromatic aldehydes upon aqueous work-up
An alloy is a combination of metals and of a metal or another element. Alloys are defined by a metallic bonding character. An alloy may be a mixture of metallic phases. Intermetallic compounds are alloys with a defined crystal structure. Zintl phases are sometimes considered alloys depending on bond types. Alloys are used in a wide variety of applications. In some cases, a combination of metals may reduce the overall cost of the material while preserving important properties. In other cases, the combination of metals imparts synergistic properties to the constituent metal elements such as corrosion resistance or mechanical strength. Examples of alloys are steel, brass, duralumin and amalgams; the alloy constituents are measured by mass percentage for practical applications, in atomic fraction for basic science studies. Alloys are classified as substitutional or interstitial alloys, depending on the atomic arrangement that forms the alloy, they can be heterogeneous or intermetallic. An alloy is a mixture of chemical elements, which forms an impure substance that retains the characteristics of a metal.
An alloy is distinct from an impure metal in that, with an alloy, the added elements are well controlled to produce desirable properties, while impure metals such as wrought iron are less controlled, but are considered useful. Alloys are made by mixing two or more elements, at least one of, a metal; this is called the primary metal or the base metal, the name of this metal may be the name of the alloy. The other constituents may or may not be metals but, when mixed with the molten base, they will be soluble and dissolve into the mixture; the mechanical properties of alloys will be quite different from those of its individual constituents. A metal, very soft, such as aluminium, can be altered by alloying it with another soft metal, such as copper. Although both metals are soft and ductile, the resulting aluminium alloy will have much greater strength. Adding a small amount of non-metallic carbon to iron trades its great ductility for the greater strength of an alloy called steel. Due to its very-high strength, but still substantial toughness, its ability to be altered by heat treatment, steel is one of the most useful and common alloys in modern use.
By adding chromium to steel, its resistance to corrosion can be enhanced, creating stainless steel, while adding silicon will alter its electrical characteristics, producing silicon steel. Like oil and water, a molten metal may not always mix with another element. For example, pure iron is completely insoluble with copper; when the constituents are soluble, each will have a saturation point, beyond which no more of the constituent can be added. Iron, for example, can hold a maximum of 6.67% carbon. Although the elements of an alloy must be soluble in the liquid state, they may not always be soluble in the solid state. If the metals remain soluble when solid, the alloy forms a solid solution, becoming a homogeneous structure consisting of identical crystals, called a phase. If as the mixture cools the constituents become insoluble, they may separate to form two or more different types of crystals, creating a heterogeneous microstructure of different phases, some with more of one constituent than the other phase has.
However, in other alloys, the insoluble elements may not separate until after crystallization occurs. If cooled quickly, they first crystallize as a homogeneous phase, but they are supersaturated with the secondary constituents; as time passes, the atoms of these supersaturated alloys can separate from the crystal lattice, becoming more stable, form a second phase that serve to reinforce the crystals internally. Some alloys, such as electrum, an alloy consisting of silver and gold, occur naturally. Meteorites are sometimes made of occurring alloys of iron and nickel, but are not native to the Earth. One of the first alloys made by humans was bronze, a mixture of the metals tin and copper. Bronze was an useful alloy to the ancients, because it is much stronger and harder than either of its components. Steel was another common alloy. However, in ancient times, it could only be created as an accidental byproduct from the heating of iron ore in fires during the manufacture of iron. Other ancient alloys include pewter and pig iron.
In the modern age, steel can be created in many forms. Carbon steel can be made by varying only the carbon content, producing soft alloys like mild steel or hard alloys like spring steel. Alloy steels can be made by adding other elements, such as chromium, vanadium or nickel, resulting in alloys such as high-speed steel or tool steel. Small amounts of manganese are alloyed with most modern steels because of its ability to remove unwanted impurities, like phosphorus and oxygen, which can have detrimental effects on the alloy. However, most alloys were not created until the 1900s, such as various aluminium, titanium and magnesium alloys; some modern superalloys, such as incoloy and hastelloy, may consist of a multitude of different elements. As a noun, the term alloy is used to describe a mixture of atoms in which the primary constituent is a metal; when used as a verb, the term refers to the act of mixing a metal with other elements. The primary metal is called the matrix, or the solvent; the secondary constituents are called s
Hydrogen fluoride is a chemical compound with the chemical formula HF. This colorless gas or liquid is the principal industrial source of fluorine as an aqueous solution called hydrofluoric acid, it is an important feedstock in the preparation of many important compounds including pharmaceuticals and polymers. HF is used in the petrochemical industry as a component of superacids. Hydrogen fluoride boils near room temperature, much higher than other hydrogen halides. Hydrogen fluoride is a dangerous gas, forming corrosive and penetrating hydrofluoric acid upon contact with moisture; the gas can cause blindness by rapid destruction of the corneas. French chemist Edmond Frémy is credited with discovering anhydrous hydrogen fluoride while trying to isolate fluorine. Although Carl Wilhelm Scheele prepared hydrofluoric acid in large quantities in 1771, this acid was known in the glass industry before then. Although a diatomic molecule, HF forms strong intermolecular hydrogen bonds. Solid HF consists of zigzag chains of HF molecules.
The HF molecules, with a short H–F bond of 95 pm, are linked to neighboring molecules by intermolecular H–F distances of 155 pm. Liquid HF consists of chains of HF molecules, but the chains are shorter, consisting on average of only five or six molecules. Hydrogen fluoride does not boil until 20 °C in contrast to the heavier hydrogen halides which boil between −85 °C and −35 °C; this hydrogen bonding between HF molecules gives rise to high viscosity in the liquid phase and lower than expected pressure in the gas phase. Hydrogen fluoride is miscible with water, whereas the other hydrogen halides have large solubility gaps with water. Hydrogen fluoride and water form several compounds in the solid state, most notably a 1:1 compound that does not melt until −40 °C, 44 °C above the melting point of pure HF. Unlike other hydrohalic acids, such as hydrochloric acid, hydrogen fluoride is only a weak acid in dilute aqueous solution; this is in part a result of the strength of the hydrogen–fluorine bond, but of other factors such as the tendency of HF, H2O, F− anions to form clusters.
At high concentrations, HF molecules undergo homoassociation to form polyatomic ions and protons, thus increasing the acidity. This leads to protonation of strong acids like hydrochloric, sulfuric, or nitric when using concentrated hydrofluoric acid solutions. Although hydrofluoric acid is regarded as a weak acid, it is corrosive attacking glass when hydrated; the acidity of hydrofluoric acid solutions varies with concentration owing to hydrogen-bond interactions of the fluoride ion. Dilute solutions are weakly acidic with an acid ionization constant Ka = 6.6×10−4, in contrast to corresponding solutions of the other hydrogen halides, which are strong acids. Concentrated solutions of hydrogen fluoride are much more acidic than implied by this value, as shown by measurements of the Hammett acidity function H0; the H0 for 100% HF is estimated to be between −10.2 and −11, comparable to the value −12 for sulfuric acid. In thermodynamic terms, HF solutions are non-ideal, with the activity of HF increasing much more than its concentration.
The weak acidity in dilute solution is sometimes attributed to the high H—F bond strength, which combines with the high dissolution enthalpy of HF to outweigh the more negative enthalpy of hydration of the fluoride ion. However, Paul Giguère and Sylvia Turrell have shown by infrared spectroscopy that the predominant solute species is the hydrogen-bonded ion pair, which suggests that the ionization can be described as a pair of successive equilibria: The first equilibrium lies well to the right and the second to the left, meaning that HF is extensively dissociated, but that the tight ion pairs reduce the thermodynamic activity coefficient of H3O+, so that the solution is less acidic. In concentrated solution, the additional HF causes the ion pair to dissociate with formation of the hydrogen-bonded hydrogen difluoride ion. + HF ⇌ H3O+ + HF−2The increase in free H3O+ due to this reaction accounts for the rapid increase in acidity, while fluoride ions are stabilized by strong hydrogen bonding to HF to form HF−2.
This interaction between the acid and its own conjugate base is an example of homoassociation. At the limit of 100% liquid HF, there is self-ionization 3 HF ⇌ H2F+ + HF−2which forms an acidic solution; the acidity of anhydrous HF can be increased further by the addition of Lewis acids such as SbF5, which can reduce H0 to −21. Dry hydrogen fluoride dissolves low-valent metal fluorides, as well as several molecular fluorides. Many proteins and carbohydrates can be recovered from it. In contrast, most non-fluoride inorganic chemicals react with HF rather than dissolving. Hydrogen fluoride is produced by the action of sulfuric acid on pure grades of the mineral fluorite and as a side-product of the extraction of the fertilizer precursor phosphoric acid from various minerals. See hydrofluoric acid; the anhydrous compound hydrogen fluoride is more used than its aqueous solution, hydrofluoric acid. HF serves. A component of high-octane petrol called "alkylate" is generated in alkylation units that combine C3 and C4 olefins and iso-butane to generate petrol.
HF is a reactive solvent in the electrochemical fluorination of organic compounds. In this approach, HF is oxidized in the presence of a hydrocarbon and the fluorine replaces C–H bonds with C–F bonds. P