Stellar age estimation
Various methods and tools are involved in stellar age estimation, an attempt to identify within reasonable degrees of confidence what the age of a star is. These methods include stellar evolutionary models, membership in a given star cluster or system, fitting the star with the standard spectral and luminosity classification system, the presence of a protoplanetary disk, among others. Nearly all of the methods of determining age require knowledge of the mass of the star, which can be known through various methods. No individual method can provide accurate results for all types of stars; as stars grow older, their luminosity increases at an appreciable rate. Given the mass of the star, one can use this rate of increase in luminosity in order to determine the age of the star; this method only works for calculating stellar age on the main sequence, because in advanced evolutionary stages of the star, such as the red giant stage, the standard relationship for the determination of age no longer holds.
However, when one can observe a red giant star with a known mass, one can calculate the main-sequence lifetime, thus the minimum age of star is known given that it is in an advanced stage of its evolution. As the star spends only about 1% of its total lifetime as a red giant, this is an accurate method of determining age. Various properties of stars can be used to determine their age. For example, the Eta Carinae system is emitting large quantities of dust; these enormous outbursts can be used to infer that the star system is nearing the end of its life, will explode as a supernova within a short period of astronomical time. Large stars like VY Canis Majoris, one of the largest stars known, together with Mu Cephei, Betelgeuse and UY Scuti all have radii larger than that of the average orbital radius of Jupiter in the Solar System, thus showing that they are in late evolutionary stages. Betelgeuse in particular is expected to die in a supernova explosion within the next million years; as well as the scenarios of supermassive stars violently casting off their outer layers before their deaths, other examples can be found of the properties of stars which illustrate their age.
For example, Cepheid variables have a characteristic pattern in their lightcurves, the rate of repetition of, dependent on the luminosity of the star. Since Cepheid variables are a short evolutionary stage in the lifecycle of stars, knowing the mass of the star allows for the star to be tracked in its evolutionary path, one can estimate the age of the Cepheid variable. Exceptional stellar properties which allow for an estimation of age are not confined to advanced evolutionary stages; when a solar-mass star exhibitis T Tauri variability, astronomers can locate the age of the star as being before the beginning of the main sequence phase of the star's life. Additionally, more massive pre-main-sequence stars could be Herbig Ae/Be stars. If a red dwarf star is emitting immense stellar flares and x-rays, the star can be calculated to be in an early stage of its main-sequence lifetime, after which it will become less variable and become stable. Membership in a star cluster or star system permits an assignment of rough ages to a large number of stars present within.
When one can determine the age of stars through other methods, such as the ones listed above, one can identify the age of all of the bodies in a system. This is useful in clusters of stars which exhibit a large amount of variety in their stellar masses, evolutionary stages, classifications. While not independent of the properties of the stars in the cluster, system, or other reasonably-sized association of stars, an astronomer would only need a representative sample of stars to determine the age of the cluster, rather than painstakingly finding the age of every star in the cluster through other properties. In addition, knowing the age of one member of a star system can help determine the age of that system. In a star system, stars always form at the same time as each other, given the age of one star, the age of all of the others can be known. However, this method does not work for galaxies; these units are much larger, are not a one-off creation of stars which allows their age to be determined in this fashion.
The creation of stars in a galaxy takes place over billions of years though star production may long since have ceased. The oldest stars in a galaxy can only set a minimum age for the galaxy but by no means determine the actual age. Along with other factors, the presence of a protoplanetary disk sets a maximum limit on the age of stars. Stars with protoplanetary disks are young, having moved onto the main sequence only a short time ago. Over time, this disk would coalesce to form planets, with leftover material being deposited into various asteroid belts and other similar locations. However, the presence of pulsar planets complicates this method as a determinant of age. Gyro-chronology is a method used to determine the age of field stars by measuring their rotation rate, comparing this rate with the rotation rate of the Sun, which serves as a precalibrated clock for this measurement; this method has been seen as a more accurate method for the determination of stellar ages than other methods for field stars.
Gyrochronology Stellar evolution Red Dwarf Stars from UniverseToday What is a Binary Star? from UniverseToday Review Article:Protoplanetary Disks and their Evolution from Astrobites
Carbon–hydrogen bond activation
Carbon–hydrogen bond functionalization is a type of reaction in which a carbon–hydrogen bond is cleaved and replaced with a carbon-X bond. The term implies that a transition metal is involved in the C-H cleavage process. Reactions classified by the term involve the hydrocarbon first to react with a metal catalyst to create an organometallic complex in which the hydrocarbon is coordinated to the inner-sphere of a metal, either via an intermediate "alkane or arene complex" or as a transition state leading to a "M−C" intermediate; the intermediate of this first step can undergo subsequent reactions to produce the functionalized product. Important to this definition is the requirement that during the C–H cleavage event, the hydrocarbyl species remains associated in the inner-sphere and under the influence of "M". Mechanisms for C-H activations fall under three general categories: oxidative addition, in which a metal center inserts into a carbon-hydrogen bond, which cleaves the bond and oxidizes the metal, producing an intermediate that can undergo reductive elimination to yield the organometallic reactive intermediate electrophilic activation, in which a electron rich substrate undergoes an SEAr-type mechanism Sigma-bond metathesis, which proceeds through a "four-centered" transition state in which bonds break and form in a single step: the target hydrocarbon bond breaks as the carbon bonds to the metal and the hydrogen bonds to one of the metal's ligands, which causes bond breakage between the ligand and the metal.
C–H bonds, which are traditionally considered unreactive, can be cleaved by coordination. Much research has been devoted to the design and synthesis of new reagents and catalysts that can effect C–H activation. C-H activation chemistry has the potential to transform the chemical world through the development of novel synthetic methods. C-H activation could enable the conversion of cheap and abundant alkanes into valuable functionalized organic compounds and the efficient structural editing of complex molecules. Selective activation of a specific C-H bond poses a great challenge. In addition to a high bond dissociation energy, C-H bonds have low polarity because these two elements have similar electronegativities; the first C–H activation reaction is attributed to Otto Dimroth, who in 1902, reported that benzene reacted with mercury acetate, but some scholars do not view this reaction as being true C–H activation. Many electrophilic metal centers undergo this reaction. Joseph Chatt has been credited by many to be the first to perform the first C-H activation reaction in 1965 with the insertion of ruthenium, in the form of RuCl22, into the C-H bond of naphthalene.
However, in 1955, Shunsuke Murahashi reported a cobalt-catalyzed chelation-assisted C-H functionalization of 2-phenylisoindolin-1-one from -N,1-diphenylmethanimine. In 1969, A. E. Shilov reported that potassium tetrachloroplatinate induced isotope scrambling between methane and heavy water; the pathway was proposed to involve binding of methane to Pt. In 1972, the Shilov group was able to produce methanol and methyl chloride in a similar reaction involving a stoichiometric amount of potassium tetrachloroplatinate, catalytic potassium hexachloroplatinate and water. Due to the fact that Shilov worked and published in the Soviet Union during the Cold War era, his work was ignored by Western scientists; this so-called Shilov system is today one of the few true catalytic systems for alkane functionalizations. In some cases, discoveries in C-H activation were being made in conjunction with those of cross coupling. In 1969, Yuzo Fujiwara reported the synthesis of -1,2-diphenylethene from benzene and styrene with Pd2 and Cu2, a procedure similar to that of cross coupling.
On the category of oxidative addition, M. L. H. Green in 1970 reported on the photochemical insertion of tungsten in a benzene C–H bond and George M. Whitesides in 1979 was the first to carry out an intramolecular aliphatic C–H activation The next breakthrough was reported independently by two research groups in 1982. R. G. Bergman reported the first transition metal-mediated intermolecular C–H activation of unactivated and saturated hydrocarbons by oxidative addition. Using a photochemical approach, photolysis of Cp*IrH2, where Cp* is a pentamethylcyclopentadienyl ligand, led to the coordinatively unsaturated species Cp*Ir which reacted via oxidative addition with cyclohexane and neopentane to form the corresponding hydridoalkyl complexes, Cp*IrHR, where R = cyclohexyl and neopentyl, respectively. W. A. G. Graham found that the same hydrocarbons react with Cp*Ir2 upon irradiation to afford the related alkylhydrido complexes Cp*IrHR, where R = cyclohexyl and neopentyl, respectively. In the latter example, the reaction is presumed to proceed via the oxidative addition of alkane to a 16-electron iridium intermediate, Cp*Ir, formed by irradiation of Cp*Ir2.
The selective activation and functionalization of alkane C–H bonds was reported using a tungsten complex outfitted with pentamethylcyclopentadienyl, nitrosyl and neopentyl ligands, Cp*W. In one example involving this system, the alkane pentane is selectively converted to the halocarbon 1-iodopentane; this transformation was achieved via the thermolysis of Cp*W in pentane at room temperature, resulting in elimination of neopentane by a pseudo-first-order process, generating an undetectable electronically and steric
Diatomic molecules are molecules composed of only two atoms, of the same or different chemical elements. The prefix di- is of Greek origin, meaning "two". If a diatomic molecule consists of two atoms of the same element, such as hydrogen or oxygen it is said to be homonuclear. Otherwise, if a diatomic molecule consists of two different atoms, such as carbon monoxide or nitric oxide, the molecule is said to be heteronuclear; the only chemical elements that form stable homonuclear diatomic molecules at standard temperature and pressure are the gases hydrogen, oxygen and chlorine. The noble gases are gases at STP, but they are monatomic; the homonuclear diatomic gases and noble gases together are called "elemental gases" or "molecular gases", to distinguish them from other gases that are chemical compounds. At elevated temperatures, the halogens bromine and iodine form diatomic gases. All halogens have been observed as diatomic molecules, except for astatine, uncertain; the mnemonics BrINClHOF, pronounced "Brinklehof", HONClBrIF, pronounced "Honkelbrif", HOFBrINCl have been coined to aid recall of the list of diatomic elements.
Other elements form diatomic molecules when evaporated, but these diatomic species repolymerize when cooled. Heating elemental phosphorus gives diphosphorus, P2. Sulfur vapor is disulfur. Dilithium is known in the gas phase. Ditungsten and dimolybdenum form with sextuple bonds in the gas phase; the bond in a homonuclear diatomic molecule is non-polar. Dirubidium is diatomic. All other diatomic molecules are chemical compounds of two different elements. Many elements can combine to form heteronuclear diatomic molecules, depending on temperature and pressure; some examples include, gases carbon monoxide, nitric oxide, hydrogen chloride. Many 1:1 binary compounds are not considered diatomic because they are polymeric at room temperature, but they form diatomic molecules when evaporated, for example gaseous MgO, SiO, many others. Hundreds of diatomic molecules have been identified in the environment of the Earth, in the laboratory, in interstellar space. About 99% of the Earth's atmosphere is composed of two species of diatomic molecules: nitrogen and oxygen.
The natural abundance of hydrogen in the Earth's atmosphere is only of the order of parts per million, but H2 is the most abundant diatomic molecule in the universe. The interstellar medium is, dominated by hydrogen atoms. Diatomic elements played an important role in the elucidation of the concepts of element and molecule in the 19th century, because some of the most common elements, such as hydrogen and nitrogen, occur as diatomic molecules. John Dalton's original atomic hypothesis assumed that all elements were monatomic and that the atoms in compounds would have the simplest atomic ratios with respect to one another. For example, Dalton assumed water's formula to be HO, giving the atomic weight of oxygen as eight times that of hydrogen, instead of the modern value of about 16; as a consequence, confusion existed regarding atomic weights and molecular formulas for about half a century. As early as 1805, Gay-Lussac and von Humboldt showed that water is formed of two volumes of hydrogen and one volume of oxygen, by 1811 Amedeo Avogadro had arrived at the correct interpretation of water's composition, based on what is now called Avogadro's law and the assumption of diatomic elemental molecules.
However, these results were ignored until 1860 due to the belief that atoms of one element would have no chemical affinity toward atoms of the same element, partly due to apparent exceptions to Avogadro's law that were not explained until in terms of dissociating molecules. At the 1860 Karlsruhe Congress on atomic weights, Cannizzaro resurrected Avogadro's ideas and used them to produce a consistent table of atomic weights, which agree with modern values; these weights were an important prerequisite for the discovery of the periodic law by Dmitri Mendeleev and Lothar Meyer. Diatomic molecules are in their lowest or ground state, which conventionally is known as the X state; when a gas of diatomic molecules is bombarded by energetic electrons, some of the molecules may be excited to higher electronic states, as occurs, for example, in the natural aurora. Such excitation can occur when the gas absorbs light or other electromagnetic radiation; the excited states are unstable and relax back to the ground state.
Over various short time scales after the excitation, transitions occur from higher to lower electronic states and to the ground state, in each transition results a photon is emitted. This emission is known as fluorescence. Successively higher electronic states are conventionally named A, B, C, etc.. The excitation energy must be greater than or equal to the energy of the electronic state in order for the excitation to occur. In quantum theory, an electronic state of a diatomic molecule is represented by 2 S + 1 Λ ( v
In organic chemistry, an alkene is an unsaturated hydrocarbon that contains at least one carbon–carbon double bond. The words alkene and olefin are used interchangeably. Acyclic alkenes, with only one double bond and no other functional groups, known as mono-enes, form a homologous series of hydrocarbons with the general formula CnH2n. Alkenes have two hydrogen atoms fewer than the corresponding alkane; the simplest alkene, with the International Union of Pure and Applied Chemistry name ethene, is the organic compound produced on the largest scale industrially. Aromatic compounds are drawn as cyclic alkenes, but their structure and properties are different and they are not considered to be alkenes. Like a single covalent bond, double bonds can be described in terms of overlapping atomic orbitals, except that, unlike a single bond, a carbon–carbon double bond consists of one sigma bond and one pi bond; this double bond is stronger than a single covalent bond and shorter, with an average bond length of 1.33 ångströms.
Each carbon of the double bond uses its three sp2 hybrid orbitals to form sigma bonds to three atoms. The unhybridized 2p atomic orbitals, which lie perpendicular to the plane created by the axes of the three sp² hybrid orbitals, combine to form the pi bond; this bond lies outside the main C–C axis, with half of the bond on one side of the molecule and half on the other. With a strength of 65 kcal/mol, the pi bond is weaker than the sigma bond. Rotation about the carbon–carbon double bond is restricted because it incurs an energetic cost to break the alignment of the p orbitals on the two carbon atoms; as a consequence, substituted alkenes may exist as one of called cis or trans isomers. More complex alkenes may be named with the E–Z notation for molecules with three or four different substituents. For example, of the isomers of butene, the two methyl groups of -but-2-ene appear on the same side of the double bond, in -but-2-ene the methyl groups appear on opposite sides; these two isomers of butene are different in their chemical and physical properties.
Twisting to a 90° dihedral angle between two of the groups on the carbons requires less energy than the strength of a pi bond, the bond still holds. The carbons of the double bond become pyramidal, which allows preserving some p orbital alignment—and hence pi bonding; the other two attached. This contradicts a common textbook assertion that the two carbons retain their planar nature when twisting, in which case the p orbitals would rotate enough away from each other to be unable to sustain a pi bond. In a 90°-twisted alkene, the p orbitals are only misaligned by 42° and the strain energy is only around 40 kcal/mol. In contrast, a broken pi bond has an energetic cost of around 65 kcal/mol; some pyramidal alkenes are stable. For example, trans-cyclooctene is a stable strained alkene and the orbital misalignment is only 19°, despite having a significant dihedral angle of 137° and a degree of pyramidalization of 18°. Trans-cycloheptene is stable at low temperatures; as predicted by the VSEPR model of electron pair repulsion, the molecular geometry of alkenes includes bond angles about each carbon in a double bond of about 120°.
The angle may vary because of steric strain introduced by nonbonded interactions between functional groups attached to the carbons of the double bond. For example, the C–C–C bond angle in propylene is 123.9°. For bridged alkenes, Bredt's rule states that a double bond cannot occur at the bridgehead of a bridged ring system unless the rings are large enough. Following Fawcett and defining S as the total number of non-bridgehead atoms in the rings, bicyclic systems require S ≥ 7 for stability and tricyclic systems require S ≥ 11. Many of the physical properties of alkenes and alkanes are similar: they are colourless and combustable; the physical state depends on molecular mass: like the corresponding saturated hydrocarbons, the simplest alkenes, ethene and butene are gases at room temperature. Linear alkenes of five to sixteen carbons are liquids, higher alkenes are waxy solids; the melting point of the solids increases with increase in molecular mass. Alkenes have stronger smells than the corresponding alkane.
Ethylene is described to have a "sweet" odor, for example. The binding of cupric ion to the olefin in the mammalian olfactory receptor MOR244-3 is implicated in the smell of alkenes. Strained alkenes, in particular, like norbornene and trans-cyclooctene are known to have strong, unpleasant odors, a fact consistent with the stronger π complexes they form with metal ions including copper. Alkenes are stable compounds, but are more reactive than alkanes, either because of the reactivity of the carbon–carbon pi-bond or the presence of allylic CH centers. Most reactions of alkenes involve additions to this pi bond. Alkenes serve as a feedstock for the petrochemical industry because they can participate in a wide variety of reactions, prominently polymerization and alkylation. Alkenes react in many addition reactions. Most of these addition reactions follow the mechanism of electrophilic addition. Examples are hydrohalogenation, halohydrin formation, hydroboration, dichlorocarbene addition, Simmons–Smith reaction, catalytic hydrogenation, epox
In organic chemistry, a hydrocarbon is an organic compound consisting of hydrogen and carbon. Hydrocarbons are examples of group 14 hydrides. Hydrocarbons from which one hydrogen atom has been removed are functional groups called hydrocarbyls; because carbon has 4 electrons in its outermost shell carbon has four bonds to make, is only stable if all 4 of these bonds are used. Aromatic hydrocarbons, alkanes and alkyne-based compounds are different types of hydrocarbons. Most hydrocarbons found on Earth occur in crude oil, where decomposed organic matter provides an abundance of carbon and hydrogen which, when bonded, can catenate to form limitless chains; as defined by IUPAC nomenclature of organic chemistry, the classifications for hydrocarbons are: Saturated hydrocarbons are the simplest of the hydrocarbon species. They are composed of single bonds and are saturated with hydrogen; the formula for acyclic saturated hydrocarbons is CnH2n+2. The most general form of saturated hydrocarbons is CnH2n +2.
Those with one ring are the cycloalkanes. Saturated hydrocarbons are the basis of petroleum fuels and are found as either linear or branched species. Substitution reaction is their characteristics property. Hydrocarbons with the same molecular formula but different structural formulae are called structural isomers; as given in the example of 3-methylhexane and its higher homologues, branched hydrocarbons can be chiral. Chiral saturated hydrocarbons constitute the side chains of biomolecules such as chlorophyll and tocopherol. Unsaturated hydrocarbons have one or more triple bonds between carbon atoms; those with double bond are called alkenes. Those with one double bond have the formula CnH2n; those containing triple bonds are called alkyne. Those with one triple bond have the formula CnH2n−2. Aromatic hydrocarbons known as arenes, are hydrocarbons that have at least one aromatic ring. Hydrocarbons can be gases, waxes or low melting solids or polymers; because of differences in molecular structure, the empirical formula remains different between hydrocarbons.
This inherent ability of hydrocarbons to bond to themselves is known as catenation, allows hydrocarbons to form more complex molecules, such as cyclohexane, in rarer cases, arenes such as benzene. This ability comes from the fact that the bond character between carbon atoms is non-polar, in that the distribution of electrons between the two elements is somewhat due to the same electronegativity values of the elements, does not result in the formation of an electrophile. With catenation comes the loss of the total amount of bonded hydrocarbons and an increase in the amount of energy required for bond cleavage due to strain exerted upon the molecule. In simple chemistry, as per valence bond theory, the carbon atom must follow the 4-hydrogen rule, which states that the maximum number of atoms available to bond with carbon is equal to the number of electrons that are attracted into the outer shell of carbon. In terms of shells, carbon consists of an incomplete outer shell, which comprises 4 electrons, thus has 4 electrons available for covalent or dative bonding.
Hydrocarbons are hydrophobic like lipids. Some hydrocarbons are abundant in the solar system. Lakes of liquid methane and ethane have been found on Titan, Saturn's largest moon, confirmed by the Cassini-Huygens Mission. Hydrocarbons are abundant in nebulae forming polycyclic aromatic hydrocarbon compounds. Hydrocarbons are a primary energy source for current civilizations; the predominant use of hydrocarbons is as a combustible fuel source. In their solid form, hydrocarbons take the form of asphalt. Mixtures of volatile hydrocarbons are now used in preference to the chlorofluorocarbons as a propellant for aerosol sprays, due to chlorofluorocarbons' impact on the ozone layer. Methane and ethane are gaseous at ambient temperatures and cannot be liquefied by pressure alone. Propane is however liquefied, exists in'propane bottles' as a liquid. Butane is so liquefied that it provides a safe, volatile fuel for small pocket lighters. Pentane is a colorless liquid at room temperature used in chemistry and industry as a powerful nearly odorless solvent of waxes and high molecular weight organic compounds, including greases.
Hexane is a used non-polar, non-aromatic solvent, as well as a significant fraction of common gasoline. The C6 through C10 alkanes and isomeric cycloalkanes are the top components of gasoline, jet fuel and specialized industrial solvent mixtures. With the progressive addition of carbon units, the simple non-ring structured hydrocarbons have higher viscosities, lubricating indices, boiling points, solidification temperatures, deeper color. At the opposite extreme from methane lie the heavy tars that remain as the lowest fraction in a crude oil refining retort, they are collected and utilized as roofing comp
In molecular geometry, bond length or bond distance is the average distance between nuclei of two bonded atoms in a molecule. It is a transferable property of a bond between atoms of fixed types independent of the rest of the molecule. Bond length is related to bond order: when more electrons participate in bond formation the bond is shorter. Bond length is inversely related to bond strength and the bond dissociation energy: all other factors being equal, a stronger bond will be shorter. In a bond between two identical atoms, half the bond distance is equal to the covalent radius. Bond lengths are measured in the solid phase by means of X-ray diffraction, or approximated in the gas phase by microwave spectroscopy. A bond between a given pair of atoms may vary between different molecules. For example, the carbon to hydrogen bonds in methane are different from those in methyl chloride, it is however possible to make generalizations. A table with experimental single bonds for carbon to other elements is given below.
Bond lengths are given in picometers. By approximation the bond distance between two different atoms is the sum of the individual covalent radii; as a general trend, bond distances decrease across the row in the periodic table and increase down a group. This trend is identical to that of the atomic radius; the bond length between two atoms in a molecule depends not only on the atoms but on such factors as the orbital hybridization and the electronic and steric nature of the substituents. The carbon–carbon bond length in diamond is 154 pm, the largest bond length that exists for ordinary carbon covalent bonds. Since one atomic unit of length is 52.9177 pm, the C–C bond length is 2.91 atomic units, or three Bohr radii long. Unusually long bond lengths do exist. In one compound, tricyclobutabenzene, a bond length of 160 pm is reported; the current record holder is another cyclobutabenzene with length 174 pm based on X-ray crystallography. In this type of compound the cyclobutane ring would force 90° angles on the carbon atoms connected to the benzene ring where they ordinarily have angles of 120°.
The existence of a long C–C bond length of up to 290 pm is claimed in a dimer of two tetracyanoethylene dianions, although this concerns a 2-electron-4-center bond. This type of bonding has been observed in neutral phenalenyl dimers; the bond lengths of these so-called "pancake bonds" are up to 305 pm. Shorter than average C–C bond distances are possible: alkenes and alkynes have bond lengths of 133 and 120 pm due to increased s-character of the sigma bond. In benzene all bonds have the same length: 139 pm. Carbon–carbon single bonds increased s-character is notable in the central bond of diacetylene and that of a certain tetrahedrane dimer. In propionitrile the cyano group withdraws electrons resulting in a reduced bond length. Squeezing a C–C bond is possible by application of strain. An unusual organic compound exists called In-methylcyclophane with a short bond distance of 147 pm for the methyl group being squeezed between a triptycene and a phenyl group. In an in silico experiment a bond distance of 136 pm was estimated for neopentane locked up in fullerene.
The smallest theoretical C–C single bond obtained in this study is 131 pm for a hypothetical tetrahedrane derivative. The same study estimated that stretching or squeezing the C–C bond in an ethane molecule by 5 pm required 2.8 or 3.5 kJ/mol, respectively. Stretching or squeezing the same bond by 15 pm required an estimated 21.9 or 37.7 kJ/mol. Bond length tutorial
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