1.
Electron
–
The electron is a subatomic particle, symbol e− or β−, with a negative elementary electric charge. Electrons belong to the first generation of the lepton particle family, the electron has a mass that is approximately 1/1836 that of the proton. Quantum mechanical properties of the include a intrinsic angular momentum of a half-integer value, expressed in units of the reduced Planck constant. As it is a fermion, no two electrons can occupy the same state, in accordance with the Pauli exclusion principle. Like all elementary particles, electrons exhibit properties of particles and waves, they can collide with other particles and can be diffracted like light. Since an electron has charge, it has an electric field. Electromagnetic fields produced from other sources will affect the motion of an electron according to the Lorentz force law, electrons radiate or absorb energy in the form of photons when they are accelerated. Laboratory instruments are capable of trapping individual electrons as well as electron plasma by the use of electromagnetic fields, special telescopes can detect electron plasma in outer space. Electrons are involved in applications such as electronics, welding, cathode ray tubes, electron microscopes, radiation therapy, lasers, gaseous ionization detectors. Interactions involving electrons with other particles are of interest in fields such as chemistry. The Coulomb force interaction between the positive protons within atomic nuclei and the negative electrons without, allows the composition of the two known as atoms, ionization or differences in the proportions of negative electrons versus positive nuclei changes the binding energy of an atomic system. The exchange or sharing of the electrons between two or more atoms is the cause of chemical bonding. In 1838, British natural philosopher Richard Laming first hypothesized the concept of a quantity of electric charge to explain the chemical properties of atoms. Irish physicist George Johnstone Stoney named this charge electron in 1891, electrons can also participate in nuclear reactions, such as nucleosynthesis in stars, where they are known as beta particles. Electrons can be created through beta decay of isotopes and in high-energy collisions. The antiparticle of the electron is called the positron, it is identical to the electron except that it carries electrical, when an electron collides with a positron, both particles can be totally annihilated, producing gamma ray photons. The ancient Greeks noticed that amber attracted small objects when rubbed with fur, along with lightning, this phenomenon is one of humanitys earliest recorded experiences with electricity. In his 1600 treatise De Magnete, the English scientist William Gilbert coined the New Latin term electricus, both electric and electricity are derived from the Latin ēlectrum, which came from the Greek word for amber, ἤλεκτρον
2.
Noble gas
–
The noble gases make a group of chemical elements with similar properties, under standard conditions, they are all odorless, colorless, monatomic gases with very low chemical reactivity. The six noble gases that occur naturally are helium, neon, argon, krypton, xenon, oganesson is predicted to be a noble gas as well, but its chemistry has not yet been investigated. For the first six periods of the table, the noble gases are exactly the members of group 18 of the periodic table. Noble gases are typically highly unreactive except when under particular extreme conditions, the inertness of noble gases makes them very suitable in applications where reactions are not wanted. The melting and boiling points for a noble gas are close together, differing by less than 10 °C. Neon, argon, krypton, and xenon are obtained from air in an air separation unit using the methods of liquefaction of gases, Noble gases have several important applications in industries such as lighting, welding, and space exploration. After the risks caused by the flammability of hydrogen became apparent, it was replaced with helium in blimps, Noble gas is translated from the German noun Edelgas, first used in 1898 by Hugo Erdmann to indicate their extremely low level of reactivity. The name makes an analogy to the noble metals, which also have low reactivity. The noble gases have also referred to as inert gases. Rare gases is another term that was used, but this is inaccurate because argon forms a fairly considerable part of the Earths atmosphere due to decay of radioactive potassium-40. Pierre Janssen and Joseph Norman Lockyer discovered a new element on August 18,1868 while looking at the chromosphere of the Sun, no chemical analysis was possible at the time, but helium was later found to be a noble gas. Before them, in 1784, the English chemist and physicist Henry Cavendish had discovered that air contains a proportion of a substance less reactive than nitrogen. A century later, in 1895, Lord Rayleigh discovered that samples of nitrogen from the air were of a different density than nitrogen resulting from chemical reactions, with this discovery, they realized an entire class of gases was missing from the periodic table. During his search for argon, Ramsay also managed to isolate helium for the first time while heating cleveite, Ramsay continued to search for these gases using the method of fractional distillation to separate liquid air into several components. In 1898, he discovered the elements krypton, neon, and xenon, and named them after the Greek words κρυπτός, νέος, and ξένος, respectively. Rayleigh and Ramsay received the 1904 Nobel Prizes in Physics and in Chemistry, respectively, for their discovery of the noble gases, the discovery of the noble gases aided in the development of a general understanding of atomic structure. In 1895, French chemist Henri Moissan attempted to form a reaction between fluorine, the most electronegative element, and argon, one of the noble gases, but failed. Scientists were unable to prepare compounds of argon until the end of the 20th century, in 1962, Neil Bartlett discovered the first chemical compound of a noble gas, xenon hexafluoroplatinate
3.
Irving Langmuir
–
Irving Langmuir /ˈlæŋmjʊr/ was an American chemist and physicist. The Langmuir Laboratory for Atmospheric Research near Socorro, New Mexico, was named in his honor as was the American Chemical Society journal for Surface Science, Irving Langmuir was born in Brooklyn, New York, on January 31,1881. He was the third of the four children of Charles Langmuir and Sadie, during his childhood, Langmuirs parents encouraged him to carefully observe nature and to keep a detailed record of his various observations. When Irving was eleven, it was discovered that he had poor eyesight, when this problem was corrected, details that had previously eluded him were revealed, and his interest in the complications of nature was heightened. During his childhood, Langmuir was influenced by his older brother, Arthur was a research chemist who encouraged Irving to be curious about nature and how things work. Arthur helped Irving set up his first chemistry lab in the corner of his bedroom, Langmuirs hobbies included mountaineering, skiing, piloting his own plane, and classical music. In addition to his professional interest in the politics of atomic energy, Langmuir attended his early education at various schools and institutes in America and Paris. Langmuir graduated high school from Chestnut Hill Academy, a private school located in the affluent Chestnut Hill area in Philadelphia. He graduated with a Bachelor of Science degree in engineering from the Columbia University School of Mines in 1903. He earned his Ph. D. degree in 1906 under Nobel laureate Walther Nernst in Göttingen, for research using the Nernst glower. His doctoral thesis was entitled “On the Partial Recombination of Dissolved Gases During Cooling. ”He later did work in chemistry. Langmuir then taught at Stevens Institute of Technology in Hoboken, New Jersey, until 1909 and his initial contributions to science came from his study of light bulbs. His first major development was the improvement of the diffusion pump and he also discovered that twisting the filament into a tight coil improved its efficiency. These were important developments in the history of the incandescent light bulb and his assistant in vacuum tube research was his cousin William Comings White. As he continued to study filaments in vacuum and different gas environments and he was one of the first scientists to work with plasmas and was the first to call these ionized gases by that name, because they reminded him of blood plasma. Langmuir and Tonks discovered electron density waves in plasmas that are now known as Langmuir waves, the current of a biased probe tip is measured as a function of bias voltage to determine the local plasma temperature and density. He also discovered atomic hydrogen, which he put to use by inventing the atomic hydrogen welding process, plasma welding has since been developed into gas tungsten arc welding. In 1917, he published a paper on the chemistry of oil films that became the basis for the award of the 1932 Nobel Prize in chemistry
4.
Oxygen
–
Oxygen is a chemical element with symbol O and atomic number 8. It is a member of the group on the periodic table and is a highly reactive nonmetal. By mass, oxygen is the third-most abundant element in the universe, after hydrogen, at standard temperature and pressure, two atoms of the element bind to form dioxygen, a colorless and odorless diatomic gas with the formula O2. This is an important part of the atmosphere and diatomic oxygen gas constitutes 20. 8% of the Earths atmosphere, additionally, as oxides the element makes up almost half of the Earths crust. Most of the mass of living organisms is oxygen as a component of water, conversely, oxygen is continuously replenished by photosynthesis, which uses the energy of sunlight to produce oxygen from water and carbon dioxide. Oxygen is too reactive to remain a free element in air without being continuously replenished by the photosynthetic action of living organisms. Another form of oxygen, ozone, strongly absorbs ultraviolet UVB radiation, but ozone is a pollutant near the surface where it is a by-product of smog. At low earth orbit altitudes, sufficient atomic oxygen is present to cause corrosion of spacecraft, the name oxygen was coined in 1777 by Antoine Lavoisier, whose experiments with oxygen helped to discredit the then-popular phlogiston theory of combustion and corrosion. 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, 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 later Leonardo da Vinci built on Philos work by observing that a portion of air is consumed during combustion and respiration, Oxygen was discovered by the Polish alchemist Sendivogius, who considered it the philosophers stone. In the late 17th century, Robert Boyle proved that air is necessary for combustion, English chemist John Mayow refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus. From this he surmised that nitroaereus is consumed in both respiration and combustion, Mayow observed that antimony increased in weight when heated, and inferred that the nitroaereus must have combined with it. 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, and 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, which was then the favored explanation of those processes. Established in 1667 by the German alchemist J. J. Becher, 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. The fact that a substance like wood gains overall weight in burning was hidden by the buoyancy of the combustion products
5.
Ferrocene
–
Ferrocene is an organometallic compound with the formula Fe2. It is the prototypical metallocene, a type of chemical compound consisting of two cyclopentadienyl rings bound on opposite sides of a central metal atom. Such organometallic compounds are known as sandwich compounds. The rapid growth of organometallic chemistry is often attributed to the excitement arising from the discovery of ferrocene, in 1951, Pauson and Kealy at Duquesne University reported the reaction of cyclopentadienyl magnesium bromide and ferric chloride with the goal of oxidatively coupling the diene to prepare fulvalene. Instead, they obtained a light orange powder of remarkable stability, a second group at British Oxygen also unknowingly discovered ferrocene. Miller, Tebboth and Tremaine were trying to synthesise amines from hydrocarbons such as cyclopentadiene and they published this result in 1952 although the actual work was done three years earlier. The stability of the new compound was accorded to the aromatic character of the negatively charged cyclopentadienyls. Robert Burns Woodward and Geoffrey Wilkinson deduced the structure based on its reactivity, independently Ernst Otto Fischer also came to the conclusion of the sandwich structure and started to synthesize other metallocenes such as nickelocene and cobaltocene. The structure of ferrocene was confirmed by NMR spectroscopy and X-ray crystallography, in 1973 Fischer of the Technische Universität München and Wilkinson of Imperial College London shared a Nobel Prize for their work on metallocenes and other aspects of organometallic chemistry. The carbon–carbon bond distances are 1.40 Å within the rings. The staggered conformation is believed to be most stable in the condensed phase due to crystal packing, the point group of the staggered conformation is D5d and the point group of the eclipsed conformation is D5h. The Cp rings rotate with a low barrier about the Cp–Fe–Cp axis, as observed by measurements on substituted derivatives of ferrocene using 1H, for example, methylferrocene exhibits a singlet for the C5H5 ring. In terms of bonding, the center in ferrocene is usually assigned to the +2 oxidation state. Each cyclopentadienyl ring is then allocated a single charge, bringing the number of π-electrons on each ring to six. These twelve electrons are shared with the metal via covalent bonding. When combined with the six d-electrons on Fe2+, the complex attains an 18-electron configuration, the first reported syntheses of ferrocene were nearly simultaneous. Pauson and Kealy synthesised ferrocene using iron chloride and a Grignard reagent, iron chloride is suspended in anhydrous diethyl ether and added to the Grignard reagent, which is prepared by reacting cyclopentadiene with magnesium and bromoethane in anhydrous benzene. The other early synthesis of ferrocene was by Miller et al. who reacted metallic iron directly with gas-phase cyclopentadiene at elevated temperature, an approach using iron pentacarbonyl was also reported
6.
Atomic orbital
–
In quantum mechanics, an atomic orbital is a mathematical function that describes the wave-like behavior of either one electron or a pair of electrons in an atom. This function can be used to calculate the probability of finding any electron of an atom in any specific region around the atoms nucleus. The term, atomic orbital, may refer to the physical region or space where the electron can be calculated to be present. Each such orbital can be occupied by a maximum of two electrons, each with its own quantum number s. The simple names s orbital, p orbital, d orbital and these names, together with the value of n, are used to describe the electron configurations of atoms. They are derived from the description by early spectroscopists of certain series of alkali metal spectroscopic lines as sharp, principal, diffuse, Orbitals for ℓ >3 continue alphabetically, omitting j because some languages do not distinguish between the letters i and j. Atomic orbitals are the building blocks of the atomic orbital model. In this model the electron cloud of an atom may be seen as being built up in an electron configuration that is a product of simpler hydrogen-like atomic orbitals. The lowest possible energy an electron can take is therefore analogous to the frequency of a wave on a string. Higher energy states are similar to harmonics of the fundamental frequency. The electrons are never in a point location, although the probability of interacting with the electron at a single point can be found from the wave function of the electron. Particle-like properties, There is always a number of electrons orbiting the nucleus. Electrons jump between orbitals in a particle-like fashion, for example, if a single photon strikes the electrons, only a single electron changes states in response to the photon. The electrons retain particle-like properties such as, each state has the same electrical charge as the electron particle. Each wave state has a single discrete spin and this can depend upon its superposition. Thus, despite the popular analogy to planets revolving around the Sun, in addition, atomic orbitals do not closely resemble a planets elliptical path in ordinary atoms. A more accurate analogy might be that of a large and often oddly shaped atmosphere, atomic orbitals exactly describe the shape of this atmosphere only when a single electron is present in an atom. This is due to the uncertainty principle, atomic orbitals may be defined more precisely in formal quantum mechanical language
7.
Catalysis
–
Catalysis is the increase in the rate of a chemical reaction due to the participation of an additional substance called a catalyst. In most cases, reactions occur faster with a catalyst because they require less activation energy, furthermore, since they are not consumed in the catalyzed reaction, catalysts can continue to act repeatedly. Often only tiny amounts are required in principle, in the presence of a catalyst, less free energy is required to reach the transition state, but the total free energy from reactants to products does not change. A catalyst may participate in multiple chemical transformations, the effect of a catalyst may vary due to the presence of other substances known as inhibitors or poisons or promoters. Catalyzed reactions have an activation energy than the corresponding uncatalyzed reaction, resulting in a higher reaction rate at the same temperature. However, the mechanics of catalysis is complex. Usually, the catalyst participates in this slowest step, and rates are limited by amount of catalyst, in heterogeneous catalysis, the diffusion of reagents to the surface and diffusion of products from the surface can be rate determining. A nanomaterial-based catalyst is an example of a heterogeneous catalyst, analogous events associated with substrate binding and product dissociation apply to homogeneous catalysts. Although catalysts are not consumed by the reaction itself, they may be inhibited, deactivated, in heterogeneous catalysis, typical secondary processes include coking where the catalyst becomes covered by polymeric side products. Additionally, heterogeneous catalysts can dissolve into the solution in a system or sublimate in a solid–gas system. The production of most industrially important chemicals involves catalysis, similarly, most biochemically significant processes are catalysed. Research into catalysis is a field in applied science and involves many areas of chemistry, notably organometallic chemistry. Catalysis is relevant to aspects of environmental science, e. g. the catalytic converter in automobiles. Many transition metals and transition metal complexes are used in catalysis as well, Catalysts called enzymes are important in biology. A catalyst works by providing a reaction pathway to the reaction product. The rate of the reaction is increased as this route has a lower activation energy than the reaction route not mediated by the catalyst. The disproportionation of hydrogen peroxide creates water and oxygen, as shown below,2 H2O2 →2 H2O + O2 This reaction is preferable in the sense that the reaction products are more stable than the starting material, though the uncatalysed reaction is slow. In fact, the decomposition of hydrogen peroxide is so slow that hydrogen peroxide solutions are commercially available and this reaction is strongly affected by catalysts such as manganese dioxide, or the enzyme peroxidase in organisms
8.
Octahedral molecular geometry
–
The octahedron has eight faces, hence the prefix octa. The octahedron is one of the Platonic solids, although octahedral molecules typically have an atom in their centre, a perfect octahedron belongs to the point group Oh. Examples of octahedral compounds are sulfur hexafluoride SF6 and molybdenum hexacarbonyl Mo6, the term octahedral is used somewhat loosely by chemists, focusing on the geometry of the bonds to the central atom and not considering differences among the ligands themselves. For example, 3+, which is not octahedral in the mathematical sense due to the orientation of the N-H bonds, is referred to as octahedral, the concept of octahedral coordination geometry was developed by Alfred Werner to explain the stoichiometries and isomerism in coordination compounds. His insight allowed chemists to rationalize the number of isomers of coordination compounds, octahedral transition-metal complexes containing amines and simple anions are often referred to Werner-type complexes. When two or more types of ligands are coordinated to a metal centre, the complex can exist as isomers. The naming system for these isomers depends upon the number and arrangement of different ligands, for MLa 4Lb 2, two isomers exist. These isomers of MLa 4Lb 2 are cis, if the Lb ligands are mutually adjacent and it was the analysis of such complexes that led Alfred Werner to the 1913 Nobel Prize–winning postulation of octahedral complexes. For MLa 2Lb 2Lc 2, a total of six isomers are possible, one isomer in which all three pairs of identical ligands are trans Three distinct isomers in which one pair of identical ligands is trans while the other two are cis. Two enantiomeric chiral isomers in which all three pairs of identical ligands are cis, the number of possible isomers can reach 30 for an octahedral complex with six different ligands. One can see that octahedral coordination allows much greater complexity than the tetrahedron that dominates organic chemistry, the tetrahedron MLaLbLcLd exists as a single enatiomeric pair. To generate two diastereomers in a compound, at least two carbon centers are required. For compounds with the formula MX6, the alternative to octahedral geometry is a trigonal prismatic geometry. In this geometry, the six ligands are also equivalent, there are also distorted trigonal prisms, with C3v symmetry, a prominent example is W6. The interconversion of Δ- and Λ-complexes, which is slow, is proposed to proceed via a trigonal prismatic intermediate. An alternative pathway for the racemization of these complexes is the Ray–Dutt twist. Pairs of octahedra can be fused in a way that preserves the octahedral coordination geometry by replacing terminal ligands with bridging ligands, two motifs for fusing octahedra are common, edge-sharing and face-sharing. Edge- and face-shared bioctahedra have the formulas 2 and M2L63, respectively, polymeric versions of the same linking pattern give the stoichiometries ∞ and ∞, respectively
9.
Nickel tetracarbonyl
–
Nickel carbonyl is the organonickel compound with the formula Ni4. This pale-yellow liquid is the carbonyl of nickel. It is an intermediate in the Mond process for the purification of nickel, Nickel carbonyl is one of the most toxic substances encountered in industrial processes. In nickel tetracarbonyl, the state for nickel is assigned as zero. The formula conforms to 18-electron rule, the molecule is tetrahedral, with four carbonyl ligands attached to nickel. The CO ligands, in which the C and the O are connected by bonds, are covalently bonded to the nickel atom via the carbon ends. Electron diffraction studies have been performed on this molecule, and the Ni–C, Ni4 was first synthesised in 1890 by Ludwig Mond by the direct reaction of nickel metal with CO. This pioneering work foreshadowed the existence of other metal carbonyl compounds, including those of V, Cr, Mn, Fe. It was also applied industrially to the purification of nickel by the end of the 19th century, at 323 K, carbon monoxide is passed over impure nickel. The optimal rate occurs at 130 °C, Ni4 is not readily available commercially. It is conveniently generated in the laboratory by carbonylation of commercially available bisnickel, on moderate heating, Ni4 decomposes to carbon monoxide and nickel metal. Combined with the formation from CO and even impure nickel. Thermal decomposition commences near 180 °C and increases at higher temperature, like other low-valent metal carbonyls, Ni4 is susceptible to attack by nucleophiles. Attack can occur at nickel center, resulting in displacement of CO ligands, or at CO. Thus, donor ligands such as triphenylphosphine react to give Ni3, bipyridine and related ligands behave similarly. The monosubstitution of nickel tetracarbonyl with other ligands can be used to determine the Tolman electronic parameter, treatment with hydroxides gives clusters such as 2− and 2−. These compounds can also be obtained by reduction of nickel carbonyl, Thus, treatment of Ni4 with carbon nucleophiles results in acyl derivatives such as −. Chlorine oxidizes nickel carbonyl into NiCl2, releasing CO gas and this reaction provides a convenient method for destroying unwanted portions of the toxic compound. Reactions of Ni4 with alkyl and aryl halides often result in carbonylated organic products, vinylic halides, such as PhCH=CHBr, are converted to the unsaturated esters upon treatment with Ni4 followed by sodium methoxide
10.
Square planar molecular geometry
–
The square planar molecular geometry in chemistry describes the stereochemistry that is adopted by certain chemical compounds. As the name suggests, molecules of this geometry have their atoms positioned at the corners of a square on the plane about a central atom. The addition of two ligands to linear compounds, ML2, can afford square planar complexes, for example, XeF2 adds fluorine to give square planar XeF4. In principle, square planar geometry can be achieved by flattening a tetrahedron, as such, the interconversion of tetrahedral and square planar geometries provides an intramolecular pathway for the isomerization of tetrahedral compounds. This pathway does not operate readily for hydrocarbons, but tetrahedral nickel complexes, e. g. NiBr22, square planar geometry can also be achieved by the removal of a pair of ligands from the z-axis of an octahedron, leaving four ligands in the xy plane. For transition metal compounds, the crystal field splitting diagram for square planar geometry can thus be derived from the octahedral diagram, the removal of the two ligands stabilizes the dz2 level, leaving the dx2−y2 level as the most destabilized. Consequently, the dx2−y2 remains unoccupied in complexes of metals with the d8 configuration and these compounds typically have 16 valence electrons. Numerous compounds adopt this geometry, examples being especially numerous for transition metal complexes, the noble gas compound XeF4 adopts this structure as predicted by VSEPR theory. The geometry is prevalent for transition metal complexes with d8 configuration, which includes Rh, Ir, Pd, Pt, notable examples include the anticancer drugs cisplatin and carboplatin. Many homogeneous catalysts are square planar in their state, such as Wilkinsons catalyst. Other examples include Vaskas complex and Zeises salt, when the two axial ligands are removed to generate a square planar geometry, the dz2 orbital is driven lower in energy as electron-electron repulsion with ligands on the z-axis is no longer present. However, for purely σ-donating ligands the dz2 orbital is still higher in energy than the dxy, dxz and it bears electron density on the x- and y-axes and therefore interacts with the filled ligand orbitals. The dxy, dxz and dyz orbitals are generally presented as degenerate and their relative ordering depends on the nature of the particular complex. Furthermore, the splitting of d-orbitals is perturbed by π-donating ligands in contrast to octahedral complexes
11.
Vaska's complex
–
Vaskas 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, and a chloride ion, the complex was first reported by J. W. DiLuzio and Lauri Vaska in 1961. Vaskas complex can undergo oxidative addition and is notable for its ability to bind to O2 reversibly and it is a bright yellow crystalline solid. The synthesis involves heating virtually any iridium chloride salt with triphenylphosphine, the most popular method uses dimethylformamide as a solvent, and sometimes aniline is added to accelerate the reaction. The reaction is conducted under nitrogen. The following is a possible balanced equation for this complicated reaction, irCl33 +3 P3 + HCON2 + C6H5NH2 → IrCl2 + Cl + OP3 + Cl +2 H2O Typical sources of iridium used in this preparation are IrCl3·xH2O and H2IrCl6. Studies on Vaskas complex helped provide the framework for homogeneous catalysis. 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, Vaskas complex undergoes oxidative addition with conventional oxidants such as halogens, strong acids such as HCl, and other molecules known to react as electrophiles, such as iodomethane. Vaskas complex binds O2 reversibly, IrCl2 + O2 ⇌ IrCl2O2 The 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 and these shifts are dependent on the amount of π-back bonding allowed by the newly associated ligands. The CO stretching frequencies for Vaskas complex and oxidatively added ligands have been documented in the literature, 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
12.
Iron pentacarbonyl
–
Iron pentacarbonyl, also known as iron carbonyl, is the compound with formula Fe5. Under standard conditions Fe5 is a free-flowing, straw-colored liquid with a pungent odour and this compound is a common precursor to diverse iron compounds, including many that are useful in small scale organic synthesis. Iron pentacarbonyl is a metal carbonyl, where carbon monoxide is the only ligand complexed with a metal. Other examples include octahedral Cr6 and tetrahedral Ni4, most metal carbonyls have 18 valence electrons, and Fe5 fits this pattern with 8 valence electrons on Fe and five pairs of electrons provided by the CO ligands. Reflecting its symmetrical structure and charge neutrality, Fe5 is volatile, Fe5 adopts a trigonal bipyramidal structure with the Fe atom surrounded by five CO ligands, three in equatorial positions and two axially bound. The Fe–C–O linkages are each linear, Fe5 exhibits a relatively low rate of interchange between the axial and equatorial CO groups via the Berry mechanism. Iron carbonyl is sometimes confused with carbonyl iron, a high-purity metal prepared by decomposition of iron pentacarbonyl, Fe5 is produced by the reaction of fine iron particles with carbon monoxide. The compound was described in a journal by Mond and Langer in 1891 as a viscous liquid of a pale-yellow colour. Samples were prepared by treatment of finely divided, oxide-free iron powder with carbon monoxide at room temperature, industrial synthesis of the compound requires relatively high temperatures and pressures as well as special, chemically resistant equipment. Preparation of the compound at the laboratory scale avoids these complications by using an iodide intermediate, photodissociation of Fe5 produces Fe29, a yellow-orange solid, also described by Mond. When heated, Fe5 converts to small amounts of the metal cluster Fe312, simple thermolysis, however, is not a useful synthesis, and each iron carbonyl complex exhibits distinct reactivity. The industrial production of compound is somewhat similar to the Mond process in that the metal is treated with carbon monoxide to give a volatile gas. In the case of iron pentacarbonyl, the reaction is more sluggish and it is necessary to use iron sponge as the starting material, and harsher reaction conditions of 5–30 MPa of carbon monoxide and 150–200 °C. Similar to the Mond process, sulfur acts as a catalyst, the crude iron pentacarbonyl is purified by distillation. Most iron pentacarbonyl produced is decomposed on site to give pure iron in analogy to carbonyl nickel. Some iron pentacarbonyl is burned to give pure iron oxide, other uses of pentacarbonyliron are small in comparison. Many compounds are derived from Fe5 by substitution of CO by Lewis bases, L, common Lewis bases include isocyanides, tertiary phosphines and arsines, and alkenes. Usually these ligands displace only one or two CO ligands, but certain acceptor ligands such as PF3 and isocyanides can proceed to tetra- and these reactions are often induced with a catalyst or light
