The density, or more the volumetric mass density, of a substance is its mass per unit volume. The symbol most used for density is ρ, although the Latin letter D can be used. Mathematically, density is defined as mass divided by volume: ρ = m V where ρ is the density, m is the mass, V is the volume. In some cases, density is loosely defined as its weight per unit volume, although this is scientifically inaccurate – this quantity is more called specific weight. For a pure substance the density has the same numerical value as its mass concentration. Different materials have different densities, density may be relevant to buoyancy and packaging. Osmium and iridium are the densest known elements at standard conditions for temperature and pressure but certain chemical compounds may be denser. To simplify comparisons of density across different systems of units, it is sometimes replaced by the dimensionless quantity "relative density" or "specific gravity", i.e. the ratio of the density of the material to that of a standard material water.
Thus a relative density less than one means. The density of a material varies with pressure; this variation is small for solids and liquids but much greater for gases. Increasing the pressure on an object decreases the volume of the object and thus increases its density. Increasing the temperature of a substance decreases its density by increasing its volume. In most materials, heating the bottom of a fluid results in convection of the heat from the bottom to the top, due to the decrease in the density of the heated fluid; this causes it to rise relative to more dense unheated material. The reciprocal of the density of a substance is called its specific volume, a term sometimes used in thermodynamics. Density is an intensive property in that increasing the amount of a substance does not increase its density. In a well-known but apocryphal tale, Archimedes was given the task of determining whether King Hiero's goldsmith was embezzling gold during the manufacture of a golden wreath dedicated to the gods and replacing it with another, cheaper alloy.
Archimedes knew that the irregularly shaped wreath could be crushed into a cube whose volume could be calculated and compared with the mass. Baffled, Archimedes is said to have taken an immersion bath and observed from the rise of the water upon entering that he could calculate the volume of the gold wreath through the displacement of the water. Upon this discovery, he leapt from his bath and ran naked through the streets shouting, "Eureka! Eureka!". As a result, the term "eureka" entered common parlance and is used today to indicate a moment of enlightenment; the story first appeared in written form in Vitruvius' books of architecture, two centuries after it took place. Some scholars have doubted the accuracy of this tale, saying among other things that the method would have required precise measurements that would have been difficult to make at the time. From the equation for density, mass density has units of mass divided by volume; as there are many units of mass and volume covering many different magnitudes there are a large number of units for mass density in use.
The SI unit of kilogram per cubic metre and the cgs unit of gram per cubic centimetre are the most used units for density. One g/cm3 is equal to one thousand kg/m3. One cubic centimetre is equal to one millilitre. In industry, other larger or smaller units of mass and or volume are more practical and US customary units may be used. See below for a list of some of the most common units of density. A number of techniques as well as standards exist for the measurement of density of materials; such techniques include the use of a hydrometer, Hydrostatic balance, immersed body method, air comparison pycnometer, oscillating densitometer, as well as pour and tap. However, each individual method or technique measures different types of density, therefore it is necessary to have an understanding of the type of density being measured as well as the type of material in question; the density at all points of a homogeneous object equals its total mass divided by its total volume. The mass is measured with a scale or balance.
To determine the density of a liquid or a gas, a hydrometer, a dasymeter or a Coriolis flow meter may be used, respectively. Hydrostatic weighing uses the displacement of water due to a submerged object to determine the density of the object. If the body is not homogeneous its density varies between different regions of the object. In that case the density around any given location is determined by calculating the density of a small volume around that location. In the limit of an infinitesimal volume the density of an inhomogeneous object at a point becomes: ρ = d m / d V, where d V is an elementary volume at position r; the mass of the body t
Potassium bromide is a salt used as an anticonvulsant and a sedative in the late 19th and early 20th centuries, with over-the-counter use extending to 1975 in the US. Its action is due to the bromide ion. Potassium bromide is used as an antiepileptic medication for dogs. Under standard conditions, potassium bromide is a white crystalline powder, it is soluble in water. In a dilute aqueous solution, potassium bromide tastes sweet, at higher concentrations it tastes bitter, tastes salty when the concentration is higher; these effects are due to the properties of the potassium ion—sodium bromide tastes salty at any concentration. In high concentration, potassium bromide irritates the gastric mucous membrane, causing nausea and sometimes vomiting. Potassium bromide, a typical ionic salt, is dissociated and near pH 7 in aqueous solution, it serves as a source of bromide ions. This reaction is important for the manufacture of silver bromide for photographic film: KBr + AgNO3 → AgBr + KNO3Aqueous bromide Br− forms complexes when reacted with some metal halides such as copper bromide: 2 KBr + CuBr2 → K2 A traditional method for the manufacture of KBr is the reaction of potassium carbonate with an iron bromide, Fe3Br8, made by treating scrap iron under water with excess bromine: 4 K2CO3 + Fe3Br8 → 8 KBr + Fe3O4 + 4 CO2 The anticonvulsant properties of potassium bromide were first noted by Sir Charles Locock at a meeting of the Royal Medical and Chirurgical Society in 1857.
Bromide can be regarded as the first effective medication for epilepsy. At the time, it was thought that epilepsy was caused by masturbation. Locock noted that bromide calmed sexual excitement and thought this was responsible for his success in treating seizures. In the latter half of the 19th century, potassium bromide was used for the calming of seizure and nervous disorders on an enormous scale, with the use by single hospitals being as much as several tons a year. By the beginning of the 20th century the generic word had become so associated with being sedate that bromide came to mean a dull, sedate person or a boring platitude uttered by such a person. There was not a better epilepsy drug until phenobarbital in 1912; the British Army has been claimed to lace soldiers' tea with bromide to quell sexual arousal but, untrue as doing so would diminish alertness in battle. Similar stories exist about a number of substances. Bromide compounds sodium bromide, remained in over-the-counter sedatives and headache remedies in the US until 1975, when bromides were outlawed in all over-the-counter medicines, due to chronic toxicity.
Bromide's exceedingly long half life in the body made it difficult to dose without side effects. Medical use of bromides in the US was discontinued at this time, as many better and shorter-acting sedatives were known by then. Potassium bromide is used in veterinary medicine to treat epilepsy in dogs, either as first-line treatment or in addition to phenobarbital, when seizures are not adequately controlled with phenobarbital alone. Use of bromide in cats is limited because it carries a substantial risk of causing lung inflammation in them; the use of bromide as a treatment drug for animals means that veterinary medical diagnostic laboratories are able as a matter of routine to measure serum levels of bromide on order of a veterinarian, whereas human medical diagnostic labs in the US do not measure bromide as a routine test. Potassium bromide is not approved by the US Food and Drug Administration for use in humans to control seizures. In Germany, it is still approved as an antiepileptic drug for humans children and adolescents.
These indications include severe forms of generalized tonic-clonic seizures, early-childhood-related Grand-Mal-seizures, severe myoclonic seizures during childhood. Adults who have reacted positively to the drug during childhood/adolescence may continue treatment. Potassium bromide tablets are sold under the brand name Dibro-Be mono; the drug has complete bioavailability, but the bromide ion has a long half life of 12 days in the blood, making bromide salts difficult to adjust and dose. Bromide is not known to interfere with the absorption or excretion of any other anticonvulsant, though it does have strong interactions with chloride in the body, the normal body uptake and excretion of which influences bromide's excretion; the therapeutic index for bromide is small. As with other antiepileptics, sometimes therapeutic doses may give rise to intoxication. Indistinguishable from'expected' side-effects, these include: Bromism These are central nervous system reactions, they may include:depression, somnolence loss of appetite and cachexia, nausea/emesis with exicosis loss of reflexes or pathologic reflexes clonic seizures tremor ataxia loss of neural sensitivity paresis cerebral edema with associated headache and papilledema of the eyes delirium: confusion, abnormal speech, loss of concentration and memory, aggressiveness psychosisAcne-form dermatitis and other forms of skin disease may be seen, as well as mucous hypersecretion in the lungs.
Asthma and rhinitis may worsen. Tongue disorder, aphthous stomatitis, bad breath, constipation occur. Potassium bromide is transparent from the near ultraviolet to long-wave
Potassium perchlorate is the inorganic salt with the chemical formula KClO4. Like other perchlorates, this salt is a strong oxidizer although it reacts slowly with organic substances; this obtained as a colorless, crystalline solid, is a common oxidizer used in fireworks, ammunition percussion caps, explosive primers, is used variously in propellants, flash compositions and sparklers. It has been used as a solid rocket propellant, although in that application it has been replaced by the higher performance ammonium perchlorate. KClO4 has the lowest solubility of the alkali metal perchlorates. Potassium perchlorate is prepared industrially by treating an aqueous solution of sodium perchlorate with potassium chloride; this single precipitation reaction exploits the low solubility of KClO4, about 1/100 as much as the solubility of NaClO4. It can be produced by bubbling chlorine gas through a solution of potassium chlorate and potassium hydroxide, by the reaction of perchloric acid with potassium hydroxide.
Another preparation involves the electrolysis of a potassium chlorate solution, causing KClO4 to form and precipitate at the anode. This procedure is complicated by the low solubility of both potassium chlorate and potassium perchlorate, the latter of which may precipitate onto the electrodes and impede the current. KClO4 is an oxidizer in the sense that it exothermically transfers oxygen to combustible materials increasing their rate of combustion relative to that in air. Thus, with glucose it gives carbon dioxide: 3 KClO4 + C6H12O6 → 6 H2O + 6 CO2 + 3 KClThe conversion of solid glucose into hot gaseous CO2 is the basis of the explosive force of this and other such mixtures. With sugar, KClO4 yields a low explosive, provided the necessary confinement. Otherwise such mixtures deflagrate with an intense purple flame characteristic of potassium. Flash compositions used in firecrackers consist of a mixture of aluminium powder and potassium perchlorate; this mixture, sometimes called flash powder, is used in ground and air fireworks.
As an oxidizer, potassium perchlorate can be used safely in the presence of sulfur, whereas potassium chlorate cannot. The greater reactivity of chlorate is typical – perchlorates are kinetically poorer oxidants. Chlorate produces chloric acid, unstable and can lead to premature ignition of the composition. Correspondingly, perchloric acid is quite stable. In commercial use it is mixed 50/50 with potassium nitrate to create Pyrodex black powder substitute, when not compressed within a muzzle loading firearm or in a cartridge, burns at a sufficiently slow rate to reduce it from being categorized with black powder as a low explosive, to "flammable". Potassium perchlorate can be used as an antithyroid agent used to treat hyperthyroidism in combination with one other medication; this application exploits hydrophilicity of perchlorate and iodide. The administration of known goitrogen substances can be used as a prevention in reducing the bio-uptake of iodine. Perchlorate ions, a common water contaminant in the USA due to the aerospace industry, has been shown to reduce iodine uptake and thus is classified as a goitrogen.
Perchlorate ions are a competitive inhibitor of the process by which iodide, is deposited into thyroid follicular cells. Studies involving healthy adult volunteers determined that at levels above 0.007 milligrams per kilogram per day, perchlorate begins to temporarily inhibit the thyroid gland's ability to absorb iodine from the bloodstream. The reduction of the iodide pool by perchlorate has dual effects – reduction of excess hormone synthesis and hyperthyroidism, on the one hand, reduction of thyroid inhibitor synthesis and hypothyroidism on the other. Perchlorate remains useful as a single dose application in tests measuring the discharge of radioiodide accumulated in the thyroid as a result of many different disruptions in the further metabolism of iodide in the thyroid gland. Treatment of thyrotoxicosis with 600-2,000 mg potassium perchlorate daily for periods of several months or longer was once common practice in Europe, perchlorate use at lower doses to treat thryoid problems continues to this day.
Although 400 mg of potassium perchlorate divided into four or five daily doses was used and found effective, higher doses were introduced when 400 mg/d was discovered not to control thyrotoxicosis in all subjects. Current regimens for treatment of thyrotoxicosis, when a patient is exposed to additional sources of Iodine include 500 mg potassium perchlorate twice per day for 18–40 days. Prophylaxis with perchlorate containing water at concentrations of 17 ppm, which corresponds to 0.5 mg/ personal intake, if one is 70 kg and consumes 2 litres of water per day, was found to reduce baseline radioiodine uptake by 67% This is equivalent to ingesting a total of just 35 mg of Perchlorate ions per day. In another related study were subjects drank just 1 litre of perchlorate containing water per day at a concentration of 10 ppm, i.e. daily 10 mg of Perchlorate ions were ingested, an average 38% reduction in the uptake of Iodine was observed. However, when the average perchlorate absorption in perchlorate plant workers subjected to the highest exposure has been estimate
Simplified molecular-input line-entry system
The simplified molecular-input line-entry system is a specification in the form of a line notation for describing the structure of chemical species using short ASCII strings. SMILES strings can be imported by most molecule editors for conversion back into two-dimensional drawings or three-dimensional models of the molecules; the original SMILES specification was initiated in the 1980s. It has since been extended. In 2007, an open standard called. Other linear notations include the Wiswesser line notation, ROSDAL, SYBYL Line Notation; the original SMILES specification was initiated by David Weininger at the USEPA Mid-Continent Ecology Division Laboratory in Duluth in the 1980s. Acknowledged for their parts in the early development were "Gilman Veith and Rose Russo and Albert Leo and Corwin Hansch for supporting the work, Arthur Weininger and Jeremy Scofield for assistance in programming the system." The Environmental Protection Agency funded the initial project to develop SMILES. It has since been modified and extended by others, most notably by Daylight Chemical Information Systems.
In 2007, an open standard called "OpenSMILES" was developed by the Blue Obelisk open-source chemistry community. Other'linear' notations include the Wiswesser Line Notation, ROSDAL and SLN. In July 2006, the IUPAC introduced the InChI as a standard for formula representation. SMILES is considered to have the advantage of being more human-readable than InChI; the term SMILES refers to a line notation for encoding molecular structures and specific instances should be called SMILES strings. However, the term SMILES is commonly used to refer to both a single SMILES string and a number of SMILES strings; the terms "canonical" and "isomeric" can lead to some confusion when applied to SMILES. The terms are not mutually exclusive. A number of valid SMILES strings can be written for a molecule. For example, CCO, OCC and CC all specify the structure of ethanol. Algorithms have been developed to generate the same SMILES string for a given molecule; this SMILES is unique for each structure, although dependent on the canonicalization algorithm used to generate it, is termed the canonical SMILES.
These algorithms first convert the SMILES to an internal representation of the molecular structure. Various algorithms for generating canonical SMILES have been developed and include those by Daylight Chemical Information Systems, OpenEye Scientific Software, MEDIT, Chemical Computing Group, MolSoft LLC, the Chemistry Development Kit. A common application of canonical SMILES is indexing and ensuring uniqueness of molecules in a database; the original paper that described the CANGEN algorithm claimed to generate unique SMILES strings for graphs representing molecules, but the algorithm fails for a number of simple cases and cannot be considered a correct method for representing a graph canonically. There is no systematic comparison across commercial software to test if such flaws exist in those packages. SMILES notation allows the specification of configuration at tetrahedral centers, double bond geometry; these are structural features that cannot be specified by connectivity alone and SMILES which encode this information are termed isomeric SMILES.
A notable feature of these rules is. The term isomeric SMILES is applied to SMILES in which isotopes are specified. In terms of a graph-based computational procedure, SMILES is a string obtained by printing the symbol nodes encountered in a depth-first tree traversal of a chemical graph; the chemical graph is first trimmed to remove hydrogen atoms and cycles are broken to turn it into a spanning tree. Where cycles have been broken, numeric suffix labels are included to indicate the connected nodes. Parentheses are used to indicate points of branching on the tree; the resultant SMILES form depends on the choices: of the bonds chosen to break cycles, of the starting atom used for the depth-first traversal, of the order in which branches are listed when encountered. Atoms are represented by the standard abbreviation of the chemical elements, in square brackets, such as for gold. Brackets may be omitted in the common case of atoms which: are in the "organic subset" of B, C, N, O, P, S, F, Cl, Br, or I, have no formal charge, have the number of hydrogens attached implied by the SMILES valence model, are the normal isotopes, are not chiral centers.
All other elements must be enclosed in brackets, have charges and hydrogens shown explicitly. For instance, the SMILES for water may be written as either O or. Hydrogen may be written as a separate atom; when brackets are used, the symbol H is added if the atom in brackets is bonded to one or more hydrogen, followed by the number of hydrogen atoms if greater than 1 by the sign + for a positive charge or by - for a negative charge. For example, for ammonium. If there is more than one charge, it is written as digit.
A chemical compound is a chemical substance composed of many identical molecules composed of atoms from more than one element held together by chemical bonds. A chemical element bonded to an identical chemical element is not a chemical compound since only one element, not two different elements, is involved. There are four types of compounds, depending on how the constituent atoms are held together: molecules held together by covalent bonds ionic compounds held together by ionic bonds intermetallic compounds held together by metallic bonds certain complexes held together by coordinate covalent bonds. A chemical formula is a way of expressing information about the proportions of atoms that constitute a particular chemical compound, using the standard abbreviations for the chemical elements, subscripts to indicate the number of atoms involved. For example, water is composed of two hydrogen atoms bonded to one oxygen atom: the chemical formula is H2O. Many chemical compounds have a unique numerical identifier assigned by the Chemical Abstracts Service: its CAS number.
A compound can be converted to a different chemical composition by interaction with a second chemical compound via a chemical reaction. In this process, bonds between atoms are broken in both of the interacting compounds, bonds are reformed so that new associations are made between atoms. Any substance consisting of two or more different types of atoms in a fixed stoichiometric proportion can be termed a chemical compound, it follows from their being composed of fixed proportions of two or more types of atoms that chemical compounds can be converted, via chemical reaction, into compounds or substances each having fewer atoms. The ratio of each element in the compound is expressed in a ratio in its chemical formula. A chemical formula is a way of expressing information about the proportions of atoms that constitute a particular chemical compound, using the standard abbreviations for the chemical elements, subscripts to indicate the number of atoms involved. For example, water is composed of two hydrogen atoms bonded to one oxygen atom: the chemical formula is H2O.
In the case of non-stoichiometric compounds, the proportions may be reproducible with regard to their preparation, give fixed proportions of their component elements, but proportions that are not integral. Chemical compounds have a unique and defined chemical structure held together in a defined spatial arrangement by chemical bonds. Chemical compounds can be molecular compounds held together by covalent bonds, salts held together by ionic bonds, intermetallic compounds held together by metallic bonds, or the subset of chemical complexes that are held together by coordinate covalent bonds. Pure chemical elements are not considered chemical compounds, failing the two or more atom requirement, though they consist of molecules composed of multiple atoms. Many chemical compounds have a unique numerical identifier assigned by the Chemical Abstracts Service: its CAS number. There is varying and sometimes inconsistent nomenclature differentiating substances, which include non-stoichiometric examples, from chemical compounds, which require the fixed ratios.
Many solid chemical substances—for example many silicate minerals—are chemical substances, but do not have simple formulae reflecting chemically bonding of elements to one another in fixed ratios. It may be argued that they are related to, rather than being chemical compounds, insofar as the variability in their compositions is due to either the presence of foreign elements trapped within the crystal structure of an otherwise known true chemical compound, or due to perturbations in structure relative to the known compound that arise because of an excess of deficit of the constituent elements at places in its structure. Other compounds regarded as chemically identical may have varying amounts of heavy or light isotopes of the constituent elements, which changes the ratio of elements by mass slightly. Compounds are held together through a variety of different types of bonding and forces; the differences in the types of bonds in compounds differ based on the types of elements present in the compound.
London dispersion forces are the weakest force of all intermolecular forces. They are temporary attractive forces that form when the electrons in two adjacent atoms are positioned so that they create a temporary dipole. Additionally, London dispersion forces are responsible for condensing non polar substances to liquids, to further freeze to a solid state dependent on how low the temperature of the environment is. A covalent bond known as a molecular bond, involves the sharing of electrons between two atoms; this type of bond occurs between elements that fall close to each other on the periodic table of elements, yet it is observed between some metals and nonmetals. This is due to the mechanism of this type of bond. Elements that fall close to each other on the periodic table tend to have similar electronegativities, which means they have a similar affinity for electrons. Since neither element has a stronger affinity to donate or gain electrons, it causes the elements to share electrons so both elements have a more stable octet.
Ionic bonding occurs when valence electrons are transferred between elements. Opposite to covalent bonding, this chemical bond creates two oppositely charged ions; the metals in ionic bonding
The Jmol applet, among other abilities, offers an alternative to the Chime plug-in, no longer under active development. While Jmol has many features that Chime lacks, it does not claim to reproduce all Chime functions, most notably, the Sculpt mode. Chime requires plug-in installation and Internet Explorer 6.0 or Firefox 2.0 on Microsoft Windows, or Netscape Communicator 4.8 on Mac OS 9. Jmol operates on a wide variety of platforms. For example, Jmol is functional in Mozilla Firefox, Internet Explorer, Google Chrome, Safari. Chemistry Development Kit Comparison of software for molecular mechanics modeling Jmol extension for MediaWiki List of molecular graphics systems Molecular graphics Molecule editor Proteopedia PyMOL SAMSON Official website Wiki with listings of websites and moodles Willighagen, Egon. "Fast and Scriptable Molecular Graphics in Web Browsers without Java3D". Doi:10.1038/npre.2007.50.1
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