Vaska's complex is the trivial name for the chemical compound trans-carbonylchlorobisiridium, which has the formula IrCl2. This square planar diamagnetic organometallic complex consists of a central iridium atom bound to two mutually trans triphenylphosphine ligands, carbon monoxide, a chloride ion; the complex was first reported by J. W. DiLuzio and Lauri Vaska in 1961. Vaska's complex can undergo oxidative addition and is notable for its ability to bind to O2 reversibly, it is a bright yellow crystalline solid. The synthesis involves heating any iridium chloride salt with triphenylphosphine and a carbon monoxide source; the most popular method uses dimethylformamide as a solvent, sometimes aniline is added to accelerate the reaction. Another popular solvent is 2-methoxyethanol; the reaction is conducted under nitrogen. In the synthesis, triphenylphosphine serves as both a ligand and a reductant, the carbonyl ligand is derived by decomposition of dimethylformamide via a deinsertion of an intermediate Ir-CH species.
The following is a possible balanced equation for this complicated reaction. IrCl33 + 3 P3 + HCON2 + C6H5NH2 → IrCl2 + Cl + OP3 + Cl + 2 H2OTypical sources of iridium used in this preparation are IrCl3·xH2O and H2IrCl6. Studies on Vaska's complex helped provide the conceptual framework for homogeneous catalysis. Vaska's complex, with 16 valence electrons, is considered "coordinatively unsaturated" and can thus bind to one two-electron or two one-electron ligands to become electronically saturated with 18 valence electrons; the addition of two one-electron ligands is called oxidative addition. Upon oxidative addition, the oxidation state of the iridium increases from Ir to Ir; the four-coordinated square planar arrangement in the starting complex converts to an octahedral, six-coordinate product. Vaska's complex undergoes oxidative addition with conventional oxidants such as halogens, strong acids such as HCl, other molecules known to react as electrophiles, such as iodomethane. Vaska's complex binds O2 reversibly: IrCl2 + O2 ⇌ IrCl2O2The dioxygen ligand is bonded to Ir by both oxygen atoms, so-called side-on bonding.
In myoglobin and hemoglobin, by contrast, O2 binds end-on, attaching to the metal via only one of the two oxygen atoms. The resulting dioxygen adduct reverts to the parent complex upon heating or purging the solution with an inert gas, signaled by a colour change from orange back to yellow. Infrared spectroscopy can be used to analyse the products of oxidative addition to Vaska's complex because the reactions induce characteristic shifts of the stretching frequency of the coordinated carbon monoxide; these shifts are dependent on the amount of π-back bonding allowed by the newly associated ligands. The CO stretching frequencies for Vaska's complex and oxidatively added ligands have been documented in the literature. Vaska's complex: 1967 cm−1 Vaska's complex + O2: 2015 cm−1 Vaska's complex + MeI: 2047 cm−1 Vaska's complex + I2: 2067 cm−1Oxidative addition to give Ir products reduces the π-bonding from Ir to C, which causes the increase in the frequency of the carbonyl stretching band; the stretching frequency change depends upon the ligands that have been added, but the frequency is always greater than 2000 cm−1 for an Ir complex
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, an alkyne is an unsaturated hydrocarbon containing at least one carbon—carbon triple bond. The simplest acyclic alkynes with only one triple bond and no other functional groups form a homologous series with the general chemical formula CnH2n−2. Alkynes are traditionally known as acetylenes, although the name acetylene refers to C2H2, known formally as ethyne using IUPAC nomenclature. Like other hydrocarbons, alkynes are hydrophobic but tend to be more reactive. Alkynes are characteristically more unsaturated than alkenes, thus they add two equivalents of bromine. Other reactions are listed below. In some reactions, alkynes are less reactive than alkenes. For example, in a molecule with an -ene and an -yne group, addition occurs preferentially at the -ene. Possible explanations involve the two π-bonds in the alkyne delocalising, which would reduce the energy of the π-system or the stability of the intermediates during the reaction, they show greater tendency to oligomerize than alkenes do.
The resulting polymers, called polyacetylenes are conjugated and can exhibit semiconducting properties. In acetylene, the H–C≡C bond angles are 180°. By virtue of this bond angle, alkynes are rod-like. Correspondingly, cyclic alkynes are rare. Benzyne is unstable; the C≡C bond distance of 121 picometers is much shorter than the C=C distance in alkenes or the C–C bond in alkanes. The triple bond is strong with a bond strength of 839 kJ/mol; the sigma bond contributes 369 kJ/mol, the first pi bond contributes 268 kJ/mol and the second pi-bond of 202 kJ/mol bond strength. Bonding discussed in the context of molecular orbital theory, which recognizes the triple bond as arising from overlap of s and p orbitals. In the language of valence bond theory, the carbon atoms in an alkyne bond are sp hybridized: they each have two unhybridized p orbitals and two sp hybrid orbitals. Overlap of an sp orbital from each atom forms one sp–sp sigma bond; each p orbital on one atom overlaps one on the other atom, forming two pi bonds, giving a total of three bonds.
The remaining sp orbital on each atom can form a sigma bond to another atom, for example to hydrogen atoms in the parent acetylene. The two sp orbitals project on opposite sides of the carbon atom. Internal alkynes feature carbon substituents on each acetylenic carbon. Symmetrical examples include 3-hexyne. Terminal alkynes have the formula RC2H. An example is methylacetylene. Terminal alkynes, like acetylene itself, are mildly acidic, with pKa values of around 25, they are far more acidic than alkenes and alkanes, which have pKa values of around 40 and 50, respectively. The acidic hydrogen on terminal alkynes can be replaced by a variety of groups resulting in halo-, silyl-, alkoxoalkynes; the carbanions generated by deprotonation of terminal alkynes are called acetylides. In systematic chemical nomenclature, alkynes are named with the Greek prefix system without any additional letters. Examples include octyne. In parent chains with four or more carbons, it is necessary to say. For octyne, one can either write oct-3-yne when the bond starts at the third carbon.
The lowest number possible is given to the triple bond. When no superior functional groups are present, the parent chain must include the triple bond if it is not the longest possible carbon chain in the molecule. Ethyne is called by its trivial name acetylene. In chemistry, the suffix -yne is used to denote the presence of a triple bond. In organic chemistry, the suffix follows IUPAC nomenclature. However, inorganic compounds featuring unsaturation in the form of triple bonds may be denoted by substitutive nomenclature with the same methods used with alkynes. "-diyne" is used when there are two triple bonds, so on. The position of unsaturation is indicated by a numerical locant preceding the "-yne" suffix, or'locants' in the case of multiple triple bonds. Locants are chosen. "-yne" is used as an infix to name substituent groups that are triply bound to the parent compound. Sometimes a number between hyphens is inserted before it to state which atoms the triple bond is between; this suffix arose as a collapsed form of the end of the word "acetylene".
The final" - e" disappears. Commercially, the dominant alkyne is acetylene itself, used as a fuel and a precursor to other compounds, e.g. acrylates. Hundreds of millions of kilograms are produced annually by partial oxidation of natural gas: 2 CH4 + 3/2 O2 → HC≡CH + 3 H2OPropyne industrially useful, is prepared by thermal cracking of hydrocarbons. Most other industrially useful alkyne derivatives are prepared from acetylene, e.g. via condensation with formaldehyde. Specialty alkynes are prepared by dehydrohalogenation of vicinal alkyl dihalides or vinyl halides. Metal acetylides can be coupled with primary alkyl halides. Via the Fritsch–Buttenberg–Wiechell rearrangement, alkynes are prepared from vinyl bromides. Alkynes can be prepared from aldehydes using the Corey–Fuchs reaction and from aldehydes or ketones by the Seyferth–Gilbert homologation. In the alkyne zipper reaction, alkynes are generated from other alkynes by treatment with a strong base. Featuring a reactive functional group, alkynes participate in many organic reactions.
Such use was pioneered by Ralph Raphael, who in 1955 wrote the first book describing their versatility as intermediates in synthesis. Alkynes character
Oxygen is the chemical element with the symbol O and atomic number 8. It is a member of the chalcogen group on the periodic table, a reactive nonmetal, an oxidizing agent that forms oxides with most elements as well as with other compounds. By mass, oxygen is the third-most abundant element in the universe, after helium. At standard temperature and pressure, two atoms of the element bind to form dioxygen, a colorless and odorless diatomic gas with the formula O2. Diatomic oxygen gas constitutes 20.8% of the Earth's atmosphere. As compounds including oxides, the element makes up half of the Earth's crust. Dioxygen is used in cellular respiration and many major classes of organic molecules in living organisms contain oxygen, such as proteins, nucleic acids and fats, as do the major constituent inorganic compounds of animal shells and bone. Most of the mass of living organisms is oxygen as a component of water, the major constituent of lifeforms. Oxygen is continuously replenished in Earth's atmosphere by photosynthesis, which uses the energy of sunlight to produce oxygen from water and carbon dioxide.
Oxygen is too chemically reactive to remain a free element in air without being continuously replenished by the photosynthetic action of living organisms. Another form of oxygen, ozone absorbs ultraviolet UVB radiation and the high-altitude ozone layer helps protect the biosphere from ultraviolet radiation. However, ozone present at the surface is a byproduct of thus a pollutant. Oxygen was isolated by Michael Sendivogius before 1604, but it is believed that the element was discovered independently by Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, Joseph Priestley in Wiltshire, in 1774. Priority is given for Priestley because his work was published first. Priestley, called oxygen "dephlogisticated air", did not recognize it as a chemical element; the name oxygen was coined in 1777 by Antoine Lavoisier, who first recognized oxygen as a chemical element and characterized the role it plays in combustion. Common uses of oxygen include production of steel and textiles, brazing and cutting of steels and other metals, rocket propellant, oxygen therapy, life support systems in aircraft, submarines and diving.
One of the first known experiments on the relationship between combustion and air was conducted by the 2nd century BCE Greek writer on mechanics, Philo of Byzantium. In his work Pneumatica, Philo observed that inverting a vessel over a burning candle and surrounding the vessel's neck with water resulted in some water rising into the neck. Philo incorrectly surmised that parts of the air in the vessel were converted into the classical element fire and thus were able to escape through pores in the glass. Many centuries Leonardo da Vinci built on Philo's work by observing that a portion of air is consumed during combustion and respiration. In the late 17th century, Robert Boyle proved. English chemist John Mayow refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus. In one experiment, he found that placing either a mouse or a lit candle in a closed container over water caused the water to rise and replace one-fourteenth of the air's volume before extinguishing the subjects.
From this he surmised that nitroaereus is consumed in both combustion. Mayow observed that antimony increased in weight when heated, inferred that the nitroaereus must have combined with it, he thought that the lungs separate nitroaereus from air and pass it into the blood and that animal heat and muscle movement result from the reaction of nitroaereus with certain substances in the body. Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "De respiratione". Robert Hooke, Ole Borch, Mikhail Lomonosov, Pierre Bayen all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as a chemical element; this may have been in part due to the prevalence of the philosophy of combustion and corrosion called the phlogiston theory, the favored explanation of those processes. Established in 1667 by the German alchemist J. J. Becher, modified by the chemist Georg Ernst Stahl by 1731, phlogiston theory stated that all combustible materials were made of two parts.
One part, called phlogiston, was given off when the substance containing it was burned, while the dephlogisticated part was thought to be its true form, or calx. Combustible materials that leave little residue, such as wood or coal, were thought to be made of phlogiston. Air did not play a role in phlogiston theory, nor were any initial quantitative experiments conducted to test the idea. Polish alchemist and physician Michael Sendivogius in his work De Lapide Philosophorum Tractatus duodecim e naturae fonte et manuali experientia depromti described a substance contained in air, referring to it as'cibus vitae', this substance is identical with oxygen. Sendivogius, during his experiments performed between 1598 and 1604, properly recognized that the substance is equivalent to the gaseous byproduct released by the thermal decomposition of potassium nitrate. In Bugaj’s view, the isolation of oxygen and the proper association of the substance to that part of air, required for life, lends sufficient weight to the discovery of oxygen by Sendivogius.
Ruthenium is a chemical element with symbol Ru and atomic number 44. It is a rare transition metal belonging to the platinum group of the periodic table. Like the other metals of the platinum group, ruthenium is inert to most other chemicals. Russian-born scientist of Baltic-German ancestry Karl Ernst Claus discovered the element in 1844 at Kazan State University and named it after the Latin name of his homeland, Ruthenia. Ruthenium is found as a minor component of platinum ores. Most ruthenium produced is used in thick-film resistors. A minor application for ruthenium is as a chemistry catalyst. A new application of ruthenium is as the capping layer for extreme ultraviolet photomasks. Ruthenium is found in ores with the other platinum group metals in the Ural Mountains and in North and South America. Small but commercially important quantities are found in pentlandite extracted from Sudbury, Ontario and in pyroxenite deposits in South Africa. Ruthenium, a polyvalent hard white metal, is a member of the platinum group and is in group 8 of the periodic table: Whereas all other group 8 elements have 2 electrons in the outermost shell, in ruthenium, the outermost shell has only one electron.
This anomaly is observed in the neighboring metals niobium and rhodium. Ruthenium does not tarnish unless subject to high temperatures. Ruthenium dissolves in fused alkalis to give ruthenates, is not attacked by acids but is attacked by halogens at high temperatures. Indeed, ruthenium is most attacked by oxidizing agents. Small amounts of ruthenium can increase the hardness of palladium; the corrosion resistance of titanium is increased markedly by the addition of a small amount of ruthenium. The metal can be plated by thermal decomposition. A ruthenium-molybdenum alloy is known to be superconductive at temperatures below 10.6 K. Ruthenium is the last of the 4d transition metals that can assume the group oxidation state +8, then it is less stable there than the heavier congener osmium: this is the first group from the left of the table where the second and third-row transition metals display notable differences in chemical behavior. Like iron but unlike osmium, ruthenium can form aqueous cations in its lower oxidation states of +2 and +3.
Ruthenium is the first in a downward trend in the melting and boiling points and atomization enthalpy in the 4d transition metals after the maximum seen at molybdenum, because the 4d subshell is more than half full and the electrons are contributing less to metallic bonding. Unlike the lighter congener iron, ruthenium is paramagnetic at room temperature, as iron is above its Curie point; the reduction potentials in acidic aqueous solution for some common ruthenium ions are shown below: Naturally occurring ruthenium is composed of seven stable isotopes. Additionally, 34 radioactive isotopes have been discovered. Of these radioisotopes, the most stable are 106Ru with a half-life of 373.59 days, 103Ru with a half-life of 39.26 days and 97Ru with a half-life of 2.9 days. Fifteen other radioisotopes have been characterized with atomic weights ranging from 89.93 u to 114.928 u. Most of these have half-lives that are less than five minutes except 105Ru; the primary decay mode before the most abundant isotope, 102Ru, is electron capture and the primary mode after is beta emission.
The primary decay product before 102Ru is the primary decay product after is rhodium. As the 74th most abundant element in Earth's crust, ruthenium is rare, found in about 100 parts per trillion; this element is found in ores with the other platinum group metals in the Ural Mountains and in North and South America. Small but commercially important quantities are found in pentlandite extracted from Sudbury, Canada, in pyroxenite deposits in South Africa; the native form of ruthenium is a rare mineral. 12 tonnes of ruthenium are mined each year with world reserves estimated at 5,000 tonnes. The composition of the mined platinum group metal mixtures varies depending on the geochemical formation. For example, the PGMs mined in South Africa contain on average 11% ruthenium while the PGMs mined in the former USSR contain only 2%. Ruthenium and iridium are considered the minor platinum group metals. Ruthenium, like the other platinum group metals, is obtained commercially as a by-product from nickel, copper, platinum metals ore processing.
During electrorefining of copper and nickel, noble metals such as silver and the platinum group metals precipitate as anode mud, the feedstock for the extraction. The metals are converted to ionized solutes by any of several methods, depending on the composition of the feedstock. One representative method is fusion with sodium peroxide followed by dissolution in aqua regia, solution in a mixture of chlorine with hydrochloric acid. Osmium, ruthenium and iridium are insoluble in aqua regia and precipitate, leaving the other metals in solution. Rhodium is separated from the residue by treatment with molten sodium bisulfate; the insoluble residue, containing
Triruthenium dodecacarbonyl is the chemical compound with the formula Ru312. Classified as metal carbonyl cluster, it is a dark orange-colored solid, soluble in nonpolar organic solvents; the compound serves as a precursor to other organoruthenium compounds. The cluster has D3h symmetry, consisting of an equilateral triangle of Ru atoms, each of which bears two axial and two equatorial CO ligands. Os312 has the same structure, whereas Fe312 is different, with two bridging CO ligands, resulting in C2v symmetry. Ru312 is prepared by treating solutions of ruthenium trichloride with carbon monoxide under high pressure; the stoichiometry of the reaction is uncertain, one possibility being the following: 6 RuCl3 + 33 CO + 18 CH3OH → 2 Ru312 + 9 CO2 + 18 HCl The chemical properties of Ru312 have been studied, the cluster has been converted to hundreds of derivatives. High pressures of CO convert the cluster to the monomeric ruthenium pentacarbonyl, which reverts to the parent cluster upon standing. Ru312 + 3 CO ⇌ 3 Ru5 Keq = 3.3 x 10−7 mol dm−3 at room temperatureThe instability of Ru5 contrasts with the robustness of the corresponding Fe5.
The condensation of Ru5 into Ru312 proceeds via initial, rate-limiting loss of CO to give the unstable, coordinatively unsaturated species Ru4. This tetracarbonyl binds Ru5. Upon warming under a pressure of hydrogen, Ru312 converts to the tetrahedral cluster H4Ru412. Ru312 undergoes substitution reactions with Lewis bases: Ru312 + n L → Ru312-nLn + n CO where L is a tertiary phosphine or an isocyanide. At high temperatures, Ru312 converts to a series of clusters that contain interstitial carbido ligands; these include Ru6C17 and Ru5C15. Anionic carbido clusters are known, including 2− and the bioctahedral cluster 2−. Ru312 -derived carbido compounds have been used to synthesize nanoparticles for catalysis; these particles consist of 6-7 atoms and thus are all surface, resulting in extraordinary activity
Ruthenium borides are compounds of ruthenium and boron. Their most remarkable property is high hardness. Vickers hardness HV = 50 GPa was reported for thin films composed of Ru2B3 phases; this value is higher than those of bulk RuB2 or Ru2B3, but it has to be confirmed independently, as measurements on superhard materials are intrinsically difficult. For example, note that the initial report on extreme hardness of related material rhenium diboride was too optimistic. Ruthenium diboride was first thought to have a hexagonal structure, as in rhenium diboride, but it was tentatively determined to possess an orthorhombic structure