Francium is a chemical element with symbol Fr and atomic number 87. It used to be known as eka-caesium, it is radioactive. It is the second-most electropositive element, behind only caesium, is the second rarest occurring element; the isotopes of francium decay into astatine and radon. The electronic structure of a francium atom is 7s1, so the element is classed as an alkali metal. Bulk francium has never been viewed; because of the general appearance of the other elements in its periodic table column, it is assumed that francium would appear as a reactive metal, if enough could be collected together to be viewed as a bulk solid or liquid. Obtaining such a sample is improbable, since the extreme heat of decay caused by its short half-life would vaporize any viewable quantity of the element. Francium was discovered by Marguerite Perey in France in 1939, it was the last element first discovered in nature, rather than by synthesis. Outside the laboratory, francium is rare, with trace amounts found in uranium and thorium ores, where the isotope francium-223 continually forms and decays.
As little as 20–30 g exists at any given time throughout the Earth's crust. The largest amount produced in the laboratory was a cluster of more than 300,000 atoms. Francium is one of the most unstable of the occurring elements: its longest-lived isotope, francium-223, has a half-life of only 22 minutes; the only comparable element is astatine, whose most stable natural isotope, astatine-219, has a half-life of 56 seconds, although synthetic astatine-210 is much longer-lived with a half-life of 8.1 hours. All isotopes of francium decay into radium, or radon. Francium-223 has a shorter half-life than the longest-lived isotope of each synthetic element up to and including element 105, dubnium. Francium is an alkali metal whose chemical properties resemble those of caesium. A heavy element with a single valence electron, it has the highest equivalent weight of any element. Liquid francium—if created—should have a surface tension of 0.05092 N/m at its melting point. Francium's melting point was calculated to be around 27 °C.
The melting point is uncertain because of the element's extreme radioactivity. The estimated boiling point of 677 °C is uncertain. Linus Pauling estimated the electronegativity of francium at 0.7 on the Pauling scale, the same as caesium. Francium has a higher ionization energy than caesium, 392.811 kJ/mol as opposed to 375.7041 kJ/mol for caesium, as would be expected from relativistic effects, this would imply that caesium is the less electronegative of the two. Francium should have a higher electron affinity than caesium and the Fr− ion should be more polarizable than the Cs− ion; the CsFr molecule is predicted to have francium at the negative end of the dipole, unlike all known heterodiatomic alkali metal molecules. Francium superoxide is expected to have a more covalent character than its lighter congeners. Francium coprecipitates with several caesium salts, such as caesium perchlorate, which results in small amounts of francium perchlorate; this coprecipitation can be used to isolate francium, by adapting the radiocaesium coprecipitation method of Lawrence E. Glendenin and Nelson.
It will additionally coprecipitate with many other caesium salts, including the iodate, the picrate, the tartrate, the chloroplatinate, the silicotungstate. It coprecipitates with silicotungstic acid, with perchloric acid, without another alkali metal as a carrier, which provides other methods of separation. Nearly all francium salts are water-soluble. There are 34 known isotopes of francium ranging in atomic mass from 199 to 232. Francium has seven metastable nuclear isomers. Francium-223 and francium-221 are the only isotopes that occur in nature, though the former is far more common. Francium-223 is the most stable isotope, with a half-life of 21.8 minutes, it is unlikely that an isotope of francium with a longer half-life will be discovered or synthesized. Francium-223 is the fifth product of the actinium decay series as the daughter isotope of actinium-227. Francium-223 decays into radium-223 by beta decay, with a minor alpha decay path to astatine-219. Francium-221 has a half-life of 4.8 minutes.
It is the ninth product of the neptunium decay series as a daughter isotope of actinium-225. Francium-221 decays into astatine-217 by alpha decay; the least stable ground state isotope is francium-215, with a half-life of 0.12 μs: it undergoes a 9.54 MeV alpha decay to astatine-211. Its metastable isomer, francium-215m, is less stable still, with a half-life of only 3.5 ns. Due to its instability and rarity, there are no commercial applications for francium, it has been used of atomic structure. Its use as a potential diagnostic aid for various cancers has been explored, but this application has been deemed impractical. Francium's ability to be synthesized and cooled, along with its simple atomic structure, has made it the subject of specialized sp
Lithium aluminium hydride
Lithium aluminium hydride abbreviated to LAH, is an inorganic compound with the chemical formula LiAlH4. It was discovered by Finholt and Schlesinger in 1947; this compound is used as a reducing agent in organic synthesis for the reduction of esters, carboxylic acids, amides. The solid is dangerously reactive toward water, releasing gaseous hydrogen; some related derivatives have been discussed for hydrogen storage. LAH is a colorless solid, but commercial samples are gray due to contamination; this material can be purified by recrystallization from diethyl ether. Large-scale purifications employ a Soxhlet extractor; the impure gray material is used in synthesis, since the impurities are innocuous and can be separated from the organic products. The pure powdered material is pyrophoric, but not its large crystals; some commercial materials contain mineral oil to inhibit reactions with atmospheric moisture, but more it is packed in moisture-proof plastic sacks. LAH violently reacts with water, including atmospheric moisture.
The reaction proceeds according to the following idealized equation: LiAlH4 + 4 H2O → LiOH + Al3 + 4 H2This reaction provides a useful method to generate hydrogen in the laboratory. Aged, air-exposed samples appear white because they have absorbed enough moisture to generate a mixture of the white compounds lithium hydroxide and aluminium hydroxide. LAH crystallizes in the monoclinic space group P21/c; the unit cell has the dimensions: a = 4.82, b = 7.81, c = 7.92 Å, α = γ=90° and β=112°. In the structure, Li+ centers are surrounded by five AlH−4 tetrahedra; the Li+ centers are bonded to one hydrogen atom from each of the surrounding tetrahedra creating a bipyramid arrangement. At high pressures a phase transition may occur to give β-LAH. LiAH was first prepared from the reaction between lithium hydride and aluminium chloride: 4 LiH + AlCl3 → LiAlH4 + 3 LiClIn addition to this method, the industrial synthesis entails the initial preparation of sodium aluminium hydride from the elements under high pressure and temperature: Na + Al + 2 H2 → NaAlH4LiAlH4 is prepared by a salt metathesis reaction according to: NaAlH4 + LiCl → LiAlH4 + NaClwhich proceeds in a high yield.
LiCl is removed by filtration from an ethereal solution of LiAH, with subsequent precipitation of LiAlH4 to yield a product containing around 1% w/w LiCl. An alternative preparation starts from LiH, metal Al instead of AlCl3. Catalyzed by a small quantity of TiCl3, the reaction proceeds well using dimethylether as solvent; this method avoids the cogeneration of salt. LAH is soluble in many ethereal solutions. However, it may spontaneously decompose due to the presence of catalytic impurities, though, it appears to be more stable in tetrahydrofuran. Thus, THF is preferred over, despite the lower solubility; the table summarizes thermodynamic data for LAH and reactions involving LAH, in the form of standard enthalpy and Gibbs free energy change, respectively. LAH is metastable at room temperature. During prolonged storage it decomposes to Li3AlH6 and LiH; this process can be accelerated by the presence of catalytic elements, such as titanium, iron or vanadium. When heated LAH decomposes in a three-step reaction mechanism: R1 is initiated by the melting of LAH in the temperature range 150–170 °C followed by decomposition into solid Li3AlH6, although R1 is known to proceed below the melting point of LiAlH4 as well.
At about 200 °C, Li3AlH6 decomposes into LiH and Al which subsequently convert into LiAl above 400 °C. Reaction R1 is irreversible. R3 is reversible with an equilibrium pressure of about 0.25 bar at 500 °C. R1 and R2 can occur at room temperature with suitable catalysts. Lithium aluminium hydride is used in organic chemistry as a reducing agent, it is more powerful than the related reagent sodium borohydride owing to the weaker Al-H bond compared to the B-H bond. As a solution in diethyl ether and followed by an acid workup, it will convert esters, carboxylic acids, acyl chlorides and ketones into the corresponding alcohols, it converts amide, nitrile, imine and azide compounds into the amines. It reduces quaternary ammonium cations into the corresponding tertiary amines. Reactivity can be tuned by replacing hydride groups by alkoxy groups. Due to its pyrophoric nature, toxicity, low shelf life and handling problems associated with its reactivity, it has been replaced in the last decade, both at the small-industrial scale and for large-scale reductions by the more convenient related reagent sodium bis aluminium hydride, which exhibits similar reactivity but with higher safety, easier handling and better economics.
LAH is most used for the reduction of esters and carboxylic acids to primary alcohols. Aldehydes and ketones can be reduced to alcohols by LAH, but this is done using milder reagents such as NaBH4; when epoxides are reduced using LAH, the reagent attacks the less hindered end of the epoxide producing a secondary or tertiary alcohol. Epoxycyclohexanes are reduced to give axial alcohols preferentially. Partial reduction of acid chlorides to give the corresponding aldehyde product cannot proceed via LAH, since the latter reduces all the way to the primary alcohol. Instead, the milder lithium aluminium trihydride, which reacts faster with the acid chloride than with the aldehyde, must be used. For example, when isovaleric acid
Radium is a chemical element with symbol Ra and atomic number 88. It is the sixth element in group 2 of the periodic table known as the alkaline earth metals. Pure radium is silvery-white, but it reacts with nitrogen on exposure to air, forming a black surface layer of radium nitride. All isotopes of radium are radioactive, with the most stable isotope being radium-226, which has a half-life of 1600 years and decays into radon gas; when radium decays, ionizing radiation is a product, which can excite fluorescent chemicals and cause radioluminescence. Radium, in the form of radium chloride, was discovered by Marie and Pierre Curie in 1898, they extracted the radium compound from uraninite and published the discovery at the French Academy of Sciences five days later. Radium was isolated in its metallic state by Marie Curie and André-Louis Debierne through the electrolysis of radium chloride in 1911. In nature, radium is found in uranium and thorium ores in trace amounts as small as a seventh of a gram per ton of uraninite.
Radium is not necessary for living organisms, adverse health effects are when it is incorporated into biochemical processes because of its radioactivity and chemical reactivity. Other than its use in nuclear medicine, radium has no commercial applications. Today, these former applications are no longer in vogue because radium's toxicity has since become known, less dangerous isotopes are used instead in radioluminescent devices. Radium is the only radioactive member of its group, its physical and chemical properties most resemble its lighter congener barium. Pure radium is a volatile silvery-white metal, although its lighter congeners calcium and barium have a slight yellow tint; this tint vanishes on exposure to air, yielding a black layer of radium nitride. Its melting point is either 700 °C or 960 °C and its boiling point is 1,737 °C. Both of these values are lower than those of barium, confirming periodic trends down the group 2 elements. Like barium and the alkali metals, radium crystallizes in the body-centered cubic structure at standard temperature and pressure: the radium–radium bond distance is 514.8 picometers.
Radium has a density of 5.5 g/cm3, higher than that of barium, again confirming periodic trends. Radium has 33 known isotopes, with mass numbers from 202 to 234: all of them are radioactive. Four of these – 223Ra, 224Ra, 226Ra, 228Ra – occur in the decay chains of primordial thorium-232, uranium-235, uranium-238; these isotopes still have half-lives too short to be primordial radionuclides and only exist in nature from these decay chains. Together with the artificial 225Ra, these are the five most stable isotopes of radium. All other known radium isotopes have half-lives under two hours, the majority have half-lives under a minute. At least 12 nuclear isomers have been reported. In the early history of the study of radioactivity, the different natural isotopes of radium were given different names. In this scheme, 223Ra was named actinium X, 224Ra thorium X, 226Ra radium, 228Ra mesothorium 1; when it was realized that all of these are isotopes of the same element, many of these names fell out of use, "radium" came to refer to all isotopes, not just 226Ra.
Some of radium-226's decay products received historical names including "radium", ranging from radium A to radium G, with the letter indicating how far they were down the chain from their parent 226Ra.226Ra is the most stable isotope of radium and is the last isotope in the decay chain of uranium-238 with a half-life of over a millennium: it makes up all of natural radium. Its immediate decay product is the dense radioactive noble gas radon, responsible for much of the danger of environmental radium, it is 2.7 million times more radioactive than the same molar amount of natural uranium, due to its proportionally shorter half-life. A sample of radium metal maintains itself at a higher temperature than its surroundings because of the radiation it emits – alpha particles, beta particles, gamma rays. More natural radium emits alpha particles, but other steps in its decay chain emit alpha or beta particles, all particle emissions are accompanied by gamma rays. In 2013 it was discovered; this was the first discovery of an asymmetric nucleus.
Radium, like barium, is a reactive metal and always exhibits its group oxidation state of +2. It forms the colorless Ra2+ cation in aqueous solution, basic and does not form complexes readily. Most radium compounds are therefore simple ionic compounds, though participation from the 6s and 6p electrons is expected due to relativistic effects and would enhance the covalent character of radium compounds such as RaF2 and RaAt2. For this reason, the standard electrode potential for the half-reaction Ra2+ + 2e− →
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
Radon is a chemical element with symbol Rn and atomic number 86. It is a radioactive, odorless, tasteless noble gas, it occurs in minute quantities as an intermediate step in the normal radioactive decay chains through which thorium and uranium decay into lead and various other short-lived radioactive elements. Its most stable isotope, 222Rn, has a half-life of only 3.8 days, making radon one of the rarest elements since it decays away so quickly. However, since thorium and uranium are two of the most common radioactive elements on Earth, they have three isotopes with long half-lives, on the order of several billions of years, radon will be present on Earth long into the future in spite of its short half-life as it is continually being generated; the decay of radon produces many other short-lived nuclides known as radon daughters, ending at stable isotopes of lead. Unlike all the other intermediate elements in the aforementioned decay chains, radon is, under normal conditions and inhaled. Radon gas is considered a health hazard.
It is the single largest contributor to an individual's background radiation dose, but due to local differences in geology, the level of the radon-gas hazard differs from location to location. Despite its short lifetime, radon gas from natural sources, such as uranium-containing minerals, can accumulate in buildings due to its high density, in low areas such as basements and crawl spaces. Radon can occur in ground water – for example, in some spring waters and hot springs. Epidemiological studies have shown a clear link between breathing high concentrations of radon and incidence of lung cancer. Radon is a contaminant. According to the United States Environmental Protection Agency, radon is the second most frequent cause of lung cancer, after cigarette smoking, causing 21,000 lung cancer deaths per year in the United States. About 2,900 of these deaths occur among people. While radon is the second most frequent cause of lung cancer, it is the number one cause among non-smokers, according to EPA estimates.
As radon itself decays, it produces decay products, which are other radioactive elements called radon daughters. Unlike the gaseous radon itself, radon daughters are solids and stick to surfaces, such as dust particles in the air. If such contaminated dust is inhaled, these particles can cause lung cancer. Radon is a colorless and tasteless gas and therefore is not detectable by human senses alone. At standard temperature and pressure, radon forms a monatomic gas with a density of 9.73 kg/m3, about 8 times the density of the Earth's atmosphere at sea level, 1.217 kg/m3. Radon is the densest of the noble gases. Although colorless at standard temperature and pressure, when cooled below its freezing point of 202 K, radon emits a brilliant radioluminescence that turns from yellow to orange-red as the temperature lowers. Upon condensation, radon glows. Radon is sparingly more soluble than lighter noble gases. Radon is appreciably more soluble in organic liquids than in water. Being a noble gas, radon is chemically not reactive.
However, the 3.8-day half-life of radon-222 makes it useful in physical sciences as a natural tracer. Because radon is a gas at standard conditions, unlike its parents, it can be extracted from them for research. Radon is a member of the zero-valence elements, it is inert to most common chemical reactions, such as combustion, because the outer valence shell contains eight electrons. This produces a stable, minimum energy configuration in which the outer electrons are bound. 1037 kJ/mol is required to extract one electron from its shells. In accordance with periodic trends, radon has a lower electronegativity than the element one period before it, is therefore more reactive. Early studies concluded that the stability of radon hydrate should be of the same order as that of the hydrates of chlorine or sulfur dioxide, higher than the stability of the hydrate of hydrogen sulfide; because of its cost and radioactivity, experimental chemical research is performed with radon, as a result there are few reported compounds of radon, all either fluorides or oxides.
Radon can be oxidized by powerful oxidizing agents such as fluorine. It decomposes back to its elements at a temperature of above 250 °C, is reduced by water to radon gas and hydrogen fluoride: it may be reduced back to its elements by hydrogen gas, it has a low volatility and was thought to be RnF2. Because of the short half-life of radon and the radioactivity of its compounds, it has not been possible to study the compound in any detail. Theoretical studies on this molecule predict that it should have a Rn–F bond distance of 2.08 Å, that the compound is thermodynamically more stable and less volatile than its lighter counterpart XeF2. The octahedral molecule RnF6 was predicted to have an lower enthalpy of formation than the difluoride; the higher fluorides RnF4 and RnF6 have been claimed, are calculated to be stable, but it is doubtful whether they have yet been synthesized. The + ion is believed to form by the following reaction: Rn + 2 +− → +− + 2 O2 For this reason, antimony pentafluoride together with chlorine trifluoride and N2F2Sb2F11 have been considered for radon gas removal in uranium mines due to the formation of radon–fluorine compounds.
The existence of RnF2 allows
Sodium hydride is the chemical compound with the empirical formula NaH. This alkali metal hydride is used as a strong, yet combustible base in organic synthesis. NaH is representative of the saline hydrides, meaning it is a salt-like hydride, composed of Na+ and H− ions, in contrast to the more molecular hydrides such as borane, methane and water, it is an ionic material, insoluble in organic solvents, consistent with the fact that H− remains an unknown anion in solution. Because of the insolubility of NaH, all reactions involving NaH occur at the surface of the solid. NaH is produced by the direct reaction of liquid sodium. Pure NaH is colorless, although samples appear grey. NaH is ca. 40% denser than Na. NaH, like LiH, KH, RbH, CsH, adopts the NaCl crystal structure. In this motif, each Na+ ion is surrounded by six H− centers in an octahedral geometry; the ionic radii of H − and F − are comparable, as judged by the Na − Na − F distances. A unusual situation occurs in a compound dubbed "inverse sodium hydride", which contains Na− and H+ ions.
Na− is an alkalide, this compound differs from ordinary sodium hydride in having a much higher energy content due to the net displacement of two electrons from hydrogen to sodium. A derivative of this "inverse sodium hydride" arises in the presence of the base adamanzane; this molecule irreversibly encapsulates the H+ and shields it from interaction with the alkalide Na−. Theoretical work has suggested that an unprotected protonated tertiary amine complexed with the sodium alkalide might be metastable under certain solvent conditions, though the barrier to reaction would be small and finding a suitable solvent might be difficult. NaH is a base of wide utility in organic chemistry; as a superbase, it is capable of deprotonating a range of weak Brønsted acids to give the corresponding sodium derivatives. Typical "easy" substrates contain O-H, N-H, S-H bonds, including alcohols, phenols and thiols. NaH most notably is employed to deprotonate carbon acids such as 1,3-dicarbonyls and analogues such as malonic esters.
The resulting sodium derivatives can be alkylated. NaH is used to promote condensation reactions of carbonyl compounds via the Dieckmann condensation, Stobbe condensation, Darzens condensation, Claisen condensation. Other carbon acids susceptible to deprotonation by NaH include sulfonium salts and DMSO. NaH is used to make sulfur ylides, which in turn are used to convert ketones into epoxides, as in the Johnson–Corey–Chaykovsky reaction. NaH reduces certain main group compounds, but analogous reactivity is rare in organic chemistry. Notably boron trifluoride reacts to give diborane and sodium fluoride: 6 NaH + 2 BF3 → B2H6 + 6 NaFSi-Si and S-S bonds in disilanes and disulfides are reduced. A series of reduction reactions, including the hydrodecyanation of tertiary nitriles, reduction of imines to amines, amides to aldehydes, can be effected by a composite reagent composed of sodium hydride and an alkali metal iodide; the use of sodium hydride has been proposed for hydrogen storage for use in fuel cell vehicles, the hydride being encased in plastic pellets which are crushed in the presence of water to release the hydrogen.
Sodium hydride is sold by many chemical suppliers as a mixture of 60% sodium hydride in mineral oil. Such a dispersion is safer to handle and weigh than pure NaH; the compound is used in this form but the pure grey solid can be prepared by rinsing the oil with pentane or THF, with care being taken because the washings will contain traces of NaH that can ignite in air. Reactions involving NaH require an inert atmosphere, such as argon gas. NaH is used as a suspension in THF, a solvent that resists deprotonation but solvates many organosodium compounds. NaH can ignite in air upon contact with water to release hydrogen, flammable. Hydrolysis converts NaH into a caustic base. In practice, most sodium hydride is dispensed as a dispersion in oil, which can be safely handled in air
Osmium is a chemical element with symbol Os and atomic number 76. It is a hard, bluish-white transition metal in the platinum group, found as a trace element in alloys in platinum ores. Osmium is the densest occurring element, with an experimentally measured density of 22.59 g/cm3. Manufacturers use its alloys with platinum and other platinum-group metals to make fountain pen nib tipping, electrical contacts, in other applications that require extreme durability and hardness; the element's abundance in the Earth's crust is among the rarest. Osmium is the densest stable element. Calculations of density from the X-ray diffraction data may produce the most reliable data for these elements, giving a value of 22.587±0.009 g/cm3 for osmium denser than the 22.562±0.009 g/cm3 of iridium. Osmium is a hard but brittle metal that remains lustrous at high temperatures, it has a low compressibility. Correspondingly, its bulk modulus is high, reported between 395 and 462 GPa, which rivals that of diamond; the hardness of osmium is moderately high at 4 GPa.
Because of its hardness, low vapor pressure, high melting point, solid osmium is difficult to machine, form, or work. Osmium forms compounds with oxidation states ranging from −2 to +8; the most common oxidation states are +2, +3, +4, +8. The +8 oxidation state is notable for being the highest attained by any chemical element aside from iridium's +9 and is encountered only in xenon, ruthenium and iridium; the oxidation states −1 and −2 represented by the two reactive compounds Na2 and Na2 are used in the synthesis of osmium cluster compounds. The most common compound exhibiting the +8 oxidation state is osmium tetroxide; this toxic compound is formed. It is a volatile, water-soluble, pale yellow, crystalline solid with a strong smell. Osmium powder has the characteristic smell of osmium tetroxide. Osmium tetroxide forms. With ammonia, it forms the nitrido-osmates OsO3N−. Osmium tetroxide is a powerful oxidizing agent. By contrast, osmium dioxide is black, non-volatile, much less reactive and toxic.
Only two osmium compounds have major applications: osmium tetroxide for staining tissue in electron microscopy and for the oxidation of alkenes in organic synthesis, the non-volatile osmates for organic oxidation reactions. Osmium pentafluoride is known; the lower oxidation states are stabilized by the larger halogens, so that the trichloride, tribromide and diiodide are known. The oxidation state +1 is known only for osmium iodide, whereas several carbonyl complexes of osmium, such as triosmium dodecacarbonyl, represent oxidation state 0. In general, the lower oxidation states of osmium are stabilized by ligands that are good σ-donors and π-acceptors; the higher oxidation states are stabilized by strong σ- and π-donors, such as O2− and N3−. Despite its broad range of compounds in numerous oxidation states, osmium in bulk form at ordinary temperatures and pressures resists attack by all acids and alkalis, including aqua regia. Osmium has seven occurring isotopes, six of which are stable: 184Os, 187Os, 188Os, 189Os, 190Os, 192Os.
186Os undergoes alpha decay with such a long half-life ×1015 years 140000 times the age of the universe, that for practical purposes it can be considered stable. Alpha decay is predicted for all seven occurring isotopes, but it has been observed only for 186Os due to long half-lives, it is predicted that 184Os and 192Os can undergo double beta decay but this radioactivity has not been observed yet.187Os is the descendant of 187Re and is used extensively in dating terrestrial as well as meteoric rocks. It has been used to measure the intensity of continental weathering over geologic time and to fix minimum ages for stabilization of the mantle roots of continental cratons; this decay is a reason. However, the most notable application of osmium isotopes in geology has been in conjunction with the abundance of iridium, to characterise the layer of shocked quartz along the Cretaceous–Paleogene boundary that marks the extinction of the non-avian dinosaurs 65 million years ago. Osmium was discovered in 1803 by William Hyde Wollaston in London, England.
The discovery of osmium is intertwined with that of platinum and the other metals of the platinum group. Platinum reached Europe as platina, first encountered in the late 17th century in silver mines around the Chocó Department, in Colombia; the discovery that this metal was not an alloy, but a distinct new element, was published in 1748. Chemists who studied platinum dissolved it in aqua regia to create soluble salts, they always observed a small amount of a insoluble residue. Joseph Louis Proust thought. Victor Collet-Descotils, Antoine François, comte de Fourcroy, Louis Nicolas Vauquelin observed iridium in the black platinum residue in 1803, but did not obtain enough material for further experiments; the two French ch