1.
Organometallic chemistry
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Moreover, some related compounds such as transition metal hydrides and metal phosphine complexes are often included in discussions of organometallic compounds. The field of organometallic chemistry combines aspects of inorganic and organic chemistry. Organometallic compounds are distinguished by the prefix organo- e. g. organopalladium compounds, examples of such organometallic compounds include all Gilman reagents, which contain lithium and copper. Tetracarbonyl nickel, and ferrocene are examples of compounds containing transition metals. The term metalorganics usually refers to metal-containing compounds lacking direct metal-carbon bonds, metal beta-diketonates, alkoxides, and dialkylamides are representative members of this class. Representative Organometallic Compounds Many complexes feature coordination bonds between a metal and organic ligands, the organic ligands often bind the metal through a heteroatom such as oxygen or nitrogen, in which case such compounds are considered coordination compounds. However, if any of the form a direct M-C bond, then complex is usually considered to be organometallic. Furthermore, many compounds such as metal acetylacetonates and metal alkoxides are called metalorganics. Many organic coordination compounds occur naturally, for example, hemoglobin and myoglobin contain an iron center coordinated to the nitrogen atoms of a porphyrin ring, magnesium is the center of a chlorin ring in chlorophyll. The field of inorganic compounds is known as bioinorganic chemistry. In contrast to these compounds, methylcobalamin, with a cobalt-methyl bond, is a true organometallic complex. This subset of complexes are often discussed within the subfield of bioorganometallic chemistry, the metal-carbon bond in organometallic compounds are generally highly covalent. For highly electropositive elements, such as lithium and sodium, the carbon ligand exhibits carbanionic character, but free carbon-based anions are extremely rare, as in other areas of chemistry, electron counting is useful for organizing organometallic chemistry. The 18-electron rule is helpful in predicting the stabilities of metal carbonyls, most organometallic compounds do not however follow the 18e rule. Chemical bonding and reactivity in organometallic compounds is often discussed from the perspective of the isolobal principle, as well as X-ray diffraction, NMR and infrared spectroscopy are common techniques used to determine structure. The dynamic properties of compounds is often probed with variable-temperature NMR. The abundant and diverse products from coal and petroleum led to Ziegler-Natta, Fischer-Tropsch, hydroformylation catalysis which employ CO, H2, recognition of organometallic chemistry as a distinct subfield culminated in the Nobel Prizes to Ernst Fischer and Geoffrey Wilkinson for work on metallocenes. In 2005, Yves Chauvin, Robert H. Grubbs and Richard R. Schrock shared the Nobel Prize for metal-catalyzed olefin metathesis,1900 Paul Sabatier works on hydrogenation organic compounds with metal catalysts
2.
Electron shell
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In chemistry and atomic physics, an electron shell, or a principal energy level, may be thought of as an orbit followed by electrons around an atoms nucleus. The closest shell to the nucleus is called the 1 shell, followed by the 2 shell, then the 3 shell, the shells correspond with the principal quantum numbers or are labeled alphabetically with letters used in the X-ray notation. Each shell can contain only a number of electrons, The first shell can hold up to two electrons, the second shell can hold up to eight electrons, the third shell can hold up to 18. The general formula is that the nth shell can in principle hold up to 2 electrons, since electrons are electrically attracted to the nucleus, an atoms electrons will generally occupy outer shells only if the more inner shells have already been completely filled by other electrons. However, this is not a requirement, atoms may have two or even three incomplete outer shells. For an explanation of why electrons exist in these shells see electron configuration, the electrons in the outermost occupied shell determine the chemical properties of the atom, it is called the valence shell. Each shell consists of one or more subshells, and each consists of one or more atomic orbitals. The shell terminology comes from Arnold Sommerfelds modification of the Bohr model, sommerfeld retained Bohrs planetary model, but added mildly elliptical orbits to explain the fine spectroscopic structure of some elements. The multiple electrons with the principal quantum number had close orbits that formed a shell of positive thickness instead of the infinitely thin circular orbit of Bohrs model. The existence of electron shells was first observed experimentally in Charles Barklas, barkla labeled them with the letters K, L, M, N, O, P, and Q. The origin of this terminology was alphabetic, a J series was also suspected, though later experiments indicated that the K absorption lines are produced by the innermost electrons. These letters were found to correspond to the n values 1,2,3. They are used in the spectroscopic Siegbahn notation, the physical chemist Gilbert Lewis was responsible for much of the early development of the theory of the participation of valence shell electrons in chemical bonding. Linus Pauling later generalized and extended the theory while applying insights from quantum mechanics. The electron shells are labeled K, L, M, N, O, P, and Q, or 1,2,3,4,5,6, and 7, going from innermost shell outwards. Electrons in outer shells have higher energy and travel farther from the nucleus than those in inner shells. This makes them important in determining how the atom reacts chemically and behaves as a conductor, because the pull of the atoms nucleus upon them is weaker. In this way, a given elements reactivity is highly dependent upon its electronic configuration, each shell is composed of one or more subshells, which are themselves composed of atomic orbitals
3.
Transition metal
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Many scientists describe a transition metal as any element in the d-block of the periodic table, which includes groups 3 to 12 on the periodic table. In actual practice, the lanthanide and actinide series are also considered transition metals and are called inner transition metals. Cotton and Wilkinson expand the brief IUPAC definition by specifying which elements are included, as well as the elements of groups 4 to 11, they add scandium and yttrium in group 3 which have a partially filled d subshell in the metallic state. These last two elements are included even though they do not seem to possess the properties which are so characteristic of the transition metals in general. Lanthanum and actinium in Group 3 are however classified as lanthanides and actinides respectively and these elements are now known as the d-block. In the d-block the atoms of the elements have between 1 and 10 d electrons, Sc and Y in Group 3 are also generally recognized as transition metals. However the elements La–Lu and Ac–Lr and Group 12 attract different definitions from different authors, many chemistry textbooks and printed periodic tables classify La and Ac as Group 3 elements and transition metals, since their atomic ground-state configurations are s2d1 like Sc and Y. The elements Ce-Lu are considered as the series and Th-Lr as the actinide series. The two series together are classified as elements, or as inner transition elements. Some inorganic chemistry textbooks include La with the lanthanides and Ac with the actinides, a third classification defines the f-block elements as La-Yb and Ac-No, while placing Lu and Lr in Group 3. This is based on the Aufbau principle for filling electron subshells, in which 4f is filled before 5d, zinc, cadmium, and mercury are generally excluded from the transition metals as they have the electronic configuration d10s2, with no incomplete d shell. In the oxidation state +2 the ions have the electronic configuration d10, however, these elements can exist in other oxidation states, including the +1 oxidation state, as in the diatomic ion Hg2+2. The group 12 elements Zn, Cd and Hg may therefore, under certain rules, however, it is often convenient to include these elements in a discussion of the transition elements. Another example occurs in the Irving-Williams series of stability constants of complexes, although meitnerium, darmstadtium, and roentgenium are within the d-block and are expected to behave as transition metals, this has not yet been experimentally confirmed. The general electronic configuration of the elements is d1, 10n s1,2. The period 6 and 7 transition metals also add f1,14 electrons and this rule is however only approximate – it only holds for some of the transition elements, and only then in their neutral ground state. The d-sub-shell is the next-to-last sub-shell and is denoted as d -sub-shell, the number of s electrons in the outermost s sub-shell is generally one or two except palladium, with no electron in that s-sub shell in its ground state. The s-sub-shell in the shell is represented as the ns sub-shell
4.
Atomic orbital
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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
5.
Electron configuration
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In atomic physics and quantum chemistry, the electron configuration is the distribution of electrons of an atom or molecule in atomic or molecular orbitals. For example, the configuration of the neon atom is 1s2 2s2 2p6. Electronic configurations describe electrons as each moving independently in an orbital, mathematically, configurations are described by Slater determinants or configuration state functions. Knowledge of the configuration of different atoms is useful in understanding the structure of the periodic table of elements. This is also useful for describing the chemical bonds that hold atoms together, in bulk materials, this same idea helps explain the peculiar properties of lasers and semiconductors. An electron shell is the set of allowed states that share the principal quantum number, n. An atoms nth electron shell can accommodate 2n2 electrons, e. g. the first shell can accommodate 2 electrons, the second shell 8 electrons, a subshell is the set of states defined by a common azimuthal quantum number, ℓ, within a shell. The values ℓ =0,1,2,3 correspond to the s, p, d, for example the 3d subshell has n =3 and ℓ =2. The maximum number of electrons that can be placed in a subshell is given by 2 and this gives two electrons in an s subshell, six electrons in a p subshell, ten electrons in a d subshell and fourteen electrons in an f subshell. Physicists and chemists use a notation to indicate the electron configurations of atoms. For atoms, the notation consists of a sequence of atomic subshell labels with the number of electrons assigned to each subshell placed as a superscript, for example, hydrogen has one electron in the s-orbital of the first shell, so its configuration is written 1s1. Lithium has two electrons in the 1s-subshell and one in the 2s-subshell, so its configuration is written 1s2 2s1, phosphorus is as follows, 1s2 2s2 2p6 3s2 3p3. For atoms with many electrons, this notation can become lengthy, phosphorus, for instance, is in the third period. It differs from the neon, whose configuration is 1s2 2s2 2p6. This convention is useful as it is the electrons in the outermost shell that most determine the chemistry of the element, for a given configuration, the order of writing the orbitals is not completely fixed since only the orbital occupancies have physical significance. For example, the configuration of the titanium ground state can be written as either 4s2 3d2 or 3d2 4s2. The first notation follows the order based on the Madelung rule for the configurations of atoms, 4s is filled before 3d in the sequence Ar, K, Ca, Sc. The superscript 1 for a singly occupied subshell is not compulsory and it is quite common to see the letters of the orbital labels written in an italic or slanting typeface, although the International Union of Pure and Applied Chemistry recommends a normal typeface
6.
D orbitals
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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.
Electron
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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
8.
Ligand
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In coordination chemistry, a ligand is an ion or molecule that binds to a central metal atom to form a coordination complex. The bonding with the metal generally involves formal donation of one or more of the electron pairs. The nature of bonding can range from covalent to ionic. Furthermore, the bond order can range from one to three. Ligands are viewed as Lewis bases, although rare cases are known to involve Lewis acidic ligand, metals and metalloids are bound to ligands in virtually all circumstances, although gaseous naked metal ions can be generated in high vacuum. Ligands in a complex dictate the reactivity of the atom, including ligand substitution rates, the reactivity of the ligands themselves. Ligand selection is a consideration in many practical areas, including bioinorganic and medicinal chemistry, homogeneous catalysis. Ligands are classified in many ways, including, charge, size, the identity of the atom. The size of a ligand is indicated by its cone angle, the composition of coordination complexes have been known since the early 1800s, such as Prussian blue and copper vitriol. The key breakthrough occurred when Alfred Werner reconciled formulas and isomers and he showed, among other things, that the formulas of many cobalt and chromium compounds can be understood if the metal has six ligands in an octahedral geometry. The first to use the term ligand were Alfred Stock and Carl Somiesky, the theory allows one to understand the difference between coordinated and ionic chloride in the cobalt ammine chlorides and to explain many of the previously inexplicable isomers. He resolved the first coordination complex called hexol into optical isomers, in general, ligands are viewed as electron donors and the metals as electron acceptors. This is because the ligand and central metal are bonded to one another, bonding is often described using the formalisms of molecular orbital theory. The HOMO can be mainly of ligands or metal character, ligands and metal ions can be ordered in many ways, one ranking system focuses on ligand hardness. Metal ions preferentially bind certain ligands, in general, soft metal ions prefer weak field ligands, whereas hard metal ions prefer strong field ligands. According to the orbital theory, the HOMO of the ligand should have an energy that overlaps with the LUMO of the metal preferential. Metal ions bound to strong-field ligands follow the Aufbau principle, whereas complexes bound to weak-field ligands follow Hunds rule. Binding of the metal with the results in a set of molecular orbitals, where the metal can be identified with a new HOMO and LUMO
9.
Molecular orbital
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In chemistry, a molecular orbital is a mathematical function describing the wave-like behavior of an electron in a molecule. This function can be used to calculate chemical and physical properties such as the probability of finding an electron in any specific region, the term orbital was introduced by Robert S. Mulliken in 1932 as an abbreviation for one-electron orbital wave function. At an elementary level, it is used to describe the region of space in which the function has a significant amplitude, Molecular orbitals are usually constructed by combining atomic orbitals or hybrid orbitals from each atom of the molecule, or other molecular orbitals from groups of atoms. They can be calculated using the Hartree–Fock or self-consistent field methods. A molecular orbital can be used to represent the regions in a molecule where an electron occupying that orbital is likely to be found, Molecular orbitals are obtained from the combination of atomic orbitals, which predict the location of an electron in an atom. A molecular orbital can specify the configuration of a molecule. Most commonly a MO is represented as a combination of atomic orbitals. They are invaluable in providing a model of bonding in molecules. Most present-day methods in computational chemistry begin by calculating the MOs of the system, a molecular orbital describes the behavior of one electron in the electric field generated by the nuclei and some average distribution of the other electrons. In the case of two electrons occupying the orbital, the Pauli principle demands that they have opposite spin. Necessarily this is an approximation, and highly accurate descriptions of the electronic wave function do not have orbitals. Molecular orbitals arise from allowed interactions between orbitals, which are allowed if the symmetries of the atomic orbitals are compatible with each other. Efficiency of atomic orbital interactions is determined from the overlap between two atomic orbitals, which is significant if the atomic orbitals are close in energy. Finally, the number of molecular orbitals that form must equal the number of orbitals in the atoms being combined to form the molecule. Here, the orbitals are expressed as linear combinations of atomic orbitals. Molecular orbitals were first introduced by Friedrich Hund and Robert S. Mulliken in 1927 and 1928, the linear combination of atomic orbitals or LCAO approximation for molecular orbitals was introduced in 1929 by Sir John Lennard-Jones. His ground-breaking paper showed how to derive the electronic structure of the fluorine and this qualitative approach to molecular orbital theory is part of the start of modern quantum chemistry. Linear combinations of atomic orbitals can be used to estimate the molecular orbitals that are formed upon bonding between the constituent atoms
10.
Noble gas
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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
11.
Octet rule
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The rule is especially applicable to carbon, nitrogen, oxygen, and the halogens, but also to metals such as sodium or magnesium. The valence electrons can be counted using a Lewis electron dot diagram as shown at the right for carbon dioxide, the electrons shared by the two atoms in a covalent bond are counted twice, once for each atom. In carbon dioxide each oxygen shares four electrons with the central carbon, all four of these electrons are counted in both the carbon octet and the oxygen octet. Ionic bonding is common between pairs of atoms, where one of the pair is a metal of low electronegativity, a chlorine atom has seven electrons in its outer electron shell, the first and second shells being filled with two and eight electrons respectively. The first electron affinity of chlorine is +328.8 kJ per mole of chlorine atoms, adding a second electron to chlorine requires energy, energy that cannot be recovered by formation of a chemical bond. The result is that chlorine will very often form a compound in which it has eight electrons in its outer shell, a sodium atom has a single electron in its outermost electron shell, the first and second shells again being full with two and eight electrons respectively. To remove this outer electron requires only the first ionization energy, which is +495.8 kJ per mole of sodium atoms, a small amount of energy. By contrast, the electron resides in the deeper second electron shell. Thus sodium will, in most cases, form a compound in which it has lost an electron and have a full outer shell of eight electrons. The energy required to transfer an electron from an atom to a chlorine atom is small. This energy is offset by the lattice energy of sodium chloride. This completes the explanation of the rule in this case. In 1893, Alfred Werner showed that the number of atoms or groups associated with an atom is often 4 or 6, other coordination numbers up to a maximum of 8 were known. Abegg noted that the difference between the positive and negative valences of an element under his model is frequently eight. In 1919 Irving Langmuir refined these concepts further and renamed them the cubical octet atom, the octet theory evolved into what is now known as the octet rule. The quantum theory of the atom explains the eight electrons as a shell with an s2p6 electron configuration. A closed-shell configuration is one in which low-lying energy levels are full, for example, the neon atom ground state has a full n =2 shell and an empty n =3 shell. According to the rule, the atoms immediately before and after neon in the periodic table, tend to attain a similar configuration by gaining, losing
12.
Irving Langmuir
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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
13.
Spin states (d electrons)
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Spin states when describing transition metal coordination complexes refers to the potential spin configurations of the metal centers d electrons. In many these spin states vary between high-spin and low-spin configurations, the Δ splitting of the d orbitals plays an important role in the electron spin state of a coordination complex. There are three factors that affect the Δ, the period of the ion, the charge of the metal ion. If the energy required to two electrons is greater than the energy cost of placing an electron in an eg, Δ. If the separation between the orbitals is large, then the lower energy orbitals are filled before population of the higher orbitals according to the Aufbau principle. Complexes such as this are called low-spin since filling an orbital matches electrons, so, one electron is put into each of the five d orbitals before any pairing occurs in accord with Hunds rule resulting in what is known as a high-spin complex. Complexes such as this are called high-spin since populating the upper orbital avoids matches between electrons with opposite spin, within a transition metal group moving down the series corresponds with an increase in Δ. The observed result is larger Δ splitting for complexes in octahedral geometries based around transition metal centers of the second or third row and this Δ splitting is generally large enough that these complexes do not exist as high-spin state. This is true even when the center is coordinated to weak field ligands. It is only octahedral coordination complexes which are centered on first row transition metals that fluctuate between high and low-spin states, the charge of the metal center plays a role in the ligand field and the Δ splitting. For example, Fe2+ and Co3+ are both d6, however, the charge of Co3+ creates a stronger ligand field than Fe2+. All other things being equal, Fe2+ is more likely to be high spin than Co3+, ligands also affect the magnitude of Δ splitting of the d orbitals according to their field strength as described by the spectrochemical series. Strong-field ligands, such as CN− and CO, increase the Δ splitting and are likely to be low-spin. Weak-field ligands, such as I− and Br− cause a smaller Δ splitting and are likely to be high-spin. The Δ splitting energy for tetrahedral metal complexes, Δtet is smaller than that for an octahedral complex and it is unknown to have a Δtet sufficient to overcome the spin pairing energy. Tetrahedral complexes are always high spin, there are no known ligands powerful enough to produce the strong-field case in a tetrahedral complex Most spin-state transitions are between the same geometry, namely octahedral. However, in the case of d8 complexes is a shift in geometry between spin states, there is no possible difference between the high and low-spin states in the d8 octahedral complexes. However, d8 complexes are able to shift from paramagnetic tetrahedral geometry to a diamagnetic low-spin square planar geometry, the rationale for why the spin states exist according to ligand field theory is essentially the same as the crystal field theory explanation
14.
Ferrocene
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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
15.
Iron pentacarbonyl
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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
16.
Chromium hexacarbonyl
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Chromium carbonyl, also known as chromium hexacarbonyl, is the chemical compound with the formula Cr6. At room temperature the solid is stable to air, although it does have a vapor pressure. Cr6 is zerovalent, meaning that Cr has a state of zero, and it is a homoleptic complex. The complex is octahedral with Cr–C and C–O distances of 1.91 and 1.14 Å, when heated or photolyzed in tetrahydrofuran solution, Cr6 converts to Cr5 with loss of one CO ligand. The products adopt a piano-stool structure and these species are typically yellow solids, which dissolve well in common organic solvents. The arene can be liberated from the chromium with iodine or by photolysis in air, in general, substituted derivatives of Cr6 decompose upon exposure to air. Alkyl and aryl organolithium reagents RLi add to a ligand to give anionic acyl complexes. These species react with alkylating agents such as Me3O+ to form 5Cr=CR, if the R group is a vinyl or an aryl group, then the resulting carbene complex can react with an acetylene to form a new benzene ring to which is bonded the chromium tricarbonyl fragment. The two acetylene carbon atoms become part of the new ring, as does a carbon from one of the carbonyl ligands, also the three carbons from the vinyl carbene become part of the new benzene ring. In common with many of the other metal carbonyls, chromium hexacarbonyl is toxic. Its vapor pressure is high for a metal complex,1 mmHg at 36 °C). National Pollutant Inventory - Chromium and compounds fact sheet
17.
Nickel tetracarbonyl
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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
18.
Ligand field theory
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Ligand field theory describes the bonding, orbital arrangement, and other characteristics of coordination complexes. It represents an application of molecular orbital theory to transition metal complexes, a transition metal ion has nine valence atomic orbitals - consisting of five nd, three p, and one s orbitals. These orbitals are of appropriate energy to form bonding interaction with ligands, the LFT analysis is highly dependent on the geometry of the complex, but most explanations begin by describing octahedral complexes, where six ligands coordinate to the metal. Other complexes can be described by reference to field theory. Griffith and Orgel championed ligand field theory as an accurate description of such complexes. In their paper, they propose that the cause of color differences in transition metal complexes in solution is the incomplete d orbital subshells. That is, the d orbitals of transition metals participate in bonding. In an octahedral complex, the orbitals created by coordination can be seen as resulting from the donation of two electrons by each of six σ-donor ligands to the d-orbitals on the metal. In octahedral complexes, ligands approach along the x-, y- and z-axes, so their σ-symmetry orbitals form bonding and anti-bonding combinations with the dz2, the dxy, dxz and dyz orbitals remain non-bonding orbitals. The irreducible representations that these span are a1g, t1u and eg, the six σ-bonding molecular orbitals result from the combinations of ligand SALCs with metal orbitals of the same symmetry. π bonding in octahedral complexes occurs in two ways, via any ligand p-orbitals that are not being used in σ bonding, and via any π or π* molecular orbitals present on the ligand. In the usual analysis, the p-orbitals of the metal are used for σ bonding, so the π interactions take place with the appropriate metal d-orbitals, i. e. dxy, dxz and these are the orbitals that are non-bonding when only σ bonding takes place. One important π bonding in coordination complexes is metal-to-ligand π bonding and it occurs when the LUMOs of the ligand are anti-bonding π* orbitals. These orbitals are close in energy to the dxy, dxz and dyz orbitals, the ligands end up with electrons in their π* molecular orbital, so the corresponding π bond within the ligand weakens. The other form of coordination π bonding is ligand-to-metal bonding and this situation arises when the π-symmetry p or π orbitals on the ligands are filled. They combine with the dxy, dxz and dyz orbitals on the metal and it is filled with electrons from the metal d-orbitals, however, becoming the HOMO of the complex. For that reason, ΔO decreases when ligand-to-metal bonding occurs, the greater stabilization that results from metal-to-ligand bonding is caused by the donation of negative charge away from the metal ion, towards the ligands. This allows the metal to accept the σ bonds more easily, the combination of ligand-to-metal σ-bonding and metal-to-ligand π-bonding is a synergic effect, as each enhances the other
19.
Alkene
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In organic chemistry, an alkene is an unsaturated hydrocarbon that contains at least one carbon–carbon double bond. The words alkene, olefin, and olefine are used often 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 less than the corresponding alkane, the simplest alkene, ethylene, 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. This double bond is stronger than a single covalent bond and also shorter, 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 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, 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 two isomers, 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 groups of -but-2-ene appear on the same side of the double bond, and in -but-2-ene the methyl groups appear on opposite sides. These two isomers of butene are slightly different in their chemical and physical properties, a 90° twist of the C=C bond requires less energy than the strength of a pi bond, and the bond still holds. This contradicts a common assertion that the p orbitals would be unable sustain such a bond. In truth, the misalignment of the p orbitals is less than expected because pyramidalization takes place, trans-Cyclooctene is a stable strained alkene and the orbital misalignment is only 19° with a dihedral angle of 137° and a degree of pyramidalization of 18°. The trans isomer of cycloheptene is stable only 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 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, Bredts rule states that a double bond cannot occur at the bridgehead of a bridged ring system unless the rings are large enough. The physical properties of alkenes and alkanes are similar and they are colourless, nonpolar, combustable, and almost odorless. The physical state depends on mass, like the corresponding saturated hydrocarbons, the simplest alkenes, ethene, propene. Linear alkenes of approximately five to sixteen carbons are liquids, Alkenes are relatively 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
20.
Phosphine
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Phosphine is the compound with the chemical formula PH3. It is a colorless, flammable, toxic gas and pnictogen hydride, pure phosphine is odorless, but technical grade samples have a highly unpleasant odor like garlic or rotting fish, due to the presence of substituted phosphine and diphosphane. With traces of P2H4 present, PH3 is spontaneously flammable in air, phosphines are also a group of organophosphorus compounds with the formula R3P. Organophosphines are important in catalysts where they complex to various ions, complexes derived from a chiral phosphine can catalyze reactions to give chiral. Philippe Gengembre, a student of Lavoisier, first obtained phosphine in 1783 by heating phosphorus in a solution of potash. He considered diphosphine’s formula to be PH2, and thus an intermediate between elemental phosphorus, the polymers, and phosphine. Calcium phosphide produces more P2H4 than other phosphides because of the preponderance of P-P bonds in the starting material, the name phosphine first appeared in combined form in 1857. The gas PH3 was named phosphine by 1865, PH3 is a trigonal pyramidal molecule with C3v molecular symmetry. The length of the P-H bond is 1.42 Å, the H-P-H bond angles are 93. 5°. The dipole moment is 0.58 D, which increases with substitution of methyl groups in the series, CH3PH2,1.10 D, 2PH,1.23 D, 3P,1.19 D. In contrast, the moments of amines decrease with substitution, starting with ammonia. For this reason, the pair on phosphorus may be regarded as predominantly formed by the 3s orbital of phosphorus. The upfield chemical shift of the atom in the 31P NMR spectrum accords with the conclusion that the lone pair electrons occupy the 3s orbital. This electronic structure leads to a lack of nucleophilicity and an ability to form only weak hydrogen bonds, the aqueous solubility of PH3 is slight,0.22 mL of gas dissolve in 1 mL of water. Phosphine dissolves more readily in non-polar solvents than in water because of the non-polar P-H bonds and it is technically amphoteric in water, but acid and base activity is poor. Proton exchange proceeds via a phosphonium ion in solutions and via PH2− at high pH, with equilibrium constants Kb =4 × 10−28. Phosphine burns producing a white cloud of phosphorus pentoxide,2 PH3 +4 O2 → P2O5 +3 H2O Phosphine may be prepared in a variety of ways. Industrially it can be made by the reaction of phosphorus with sodium or potassium hydroxide
21.
Carbonyl group
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In organic chemistry, a carbonyl group is a functional group composed of a carbon atom double-bonded to an oxygen atom, C=O. It is common to several classes of compounds, as part of many larger functional groups. A compound containing a group is often referred to as a carbonyl compound. The term carbonyl can also refer to carbon monoxide as a ligand in an inorganic or organometallic complex, the remainder of this article concerns itself with the organic chemistry definition of carbonyl, where carbon and oxygen share a double bond. A carbonyl group characterizes the types of compounds, Note that the most specific labels are usually employed. For example, ROR structures are known as acid anhydride rather than the more generic ester, other organic carbonyls are urea and the carbamates, the derivatives of acyl chlorides chloroformates and phosgene, carbonate esters, thioesters, lactones, lactams, hydroxamates, and isocyanates. Examples of inorganic compounds are carbon dioxide and carbonyl sulfide. A special group of compounds are 1, 3-dicarbonyl compounds that have acidic protons in the central methylene unit. Examples are Meldrums acid, diethyl malonate and acetylacetone, because oxygen is more electronegative than carbon, carbonyl compounds often have resonance structures which affect their reactivity. This relative electronegativity draws electron density away from carbon, increasing the bonds polarity, carbon can then be attacked by nucleophiles or a negatively charged part of another molecule. During the reaction, the double bond is broken. This reaction is known as addition-elimination or condensation, the electronegative oxygen also can react with an electrophile, for example a proton in an acidic solution or with Lewis acids to form an oxocarbenium ion. The polarity of oxygen also makes the alpha hydrogens of carbonyl compounds much more acidic than typical sp3 C-H bonds, for example, the pKa values of acetaldehyde and acetone are 16.7 and 19 respectively, while the pKa value of methane is extrapolated to be approximately 50. This is because a carbonyl is in resonance with an enol. The deprotonation of the enol with a base produces an enolate. Amides are the most stable of the carbonyl couplings due to their high resonance stabilization between the nitrogen-carbon and carbon-oxygen bonds, carbonyl groups can be reduced by reaction with hydride reagents such as NaBH4 and LiAlH4, with bakers yeast, or by catalytic hydrogenation. Ketones give secondary alcohols while aldehydes, esters and carboxylic acids give primary alcohols, carbonyls can be alkylated in nucleophilic addition reactions using organometallic compounds such as organolithium reagents, Grignard reagents, or acetylides. Carbonyls also may be alkylated by enolates as in aldol reactions, carbonyls are also the prototypical groups with vinylogous reactivity
22.
Pi backbonding
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π backbonding, also called π backdonation, is a concept from chemistry in which electrons move from an atomic orbital on one atom to a π* antibonding orbital on a π-acceptor ligand. It is especially common in the chemistry of transition metals with multi-atomic ligands such as carbon monoxide. Electrons from the metal are used to bond to the ligand, compounds where π backbonding occurs include Ni4 and Zeises salt. The electrons are transferred from a d-orbital of the metal to anti-bonding molecular orbitals of CO. This electron-transfer strengthens the bond and weakens the C–O bond. The strengthening of the M–CO bond is reflected in increases of the frequencies for the M–C bond. Furthermore, the M–CO bond length is shortened, for this reason, IR spectroscopy is an important diagnostic technique in metal–carbonyl chemistry. The article infrared spectroscopy of metal carbonyls discusses this in detail, many ligands other than CO are strong backbonders. Nitric oxide is an even stronger π-acceptor than is CO and νNO is a tool in metal–nitrosyl chemistry. Isocyanides, RNC, are class of ligands that are able of π-backbonding. In contrast with CO, the lone pair on the C atom of isocyanides is antibonding in nature and upon complexation the CN bond is strengthened. At the same time, π-backbonding lowers the νCN, depending on the balance of σ-bonding versus π-backbonding, the νCN can either be raised or lowered. For the isocyanides, a parameter is the MC=N–C angle. Other ligands have weak π-backbonding abilities, which creates an effect of CO. As in metal–carbonyls, electrons are transferred from a d-orbital of the metal to antibonding molecular orbitals of the alkenes and alkynes. This electron transfer strengthens the bond and weakens the C–C bonds within the ligand. In the case of metal-alkenes and alkynes, the strengthening of the M–C2R4 and M–C2R2 bond is reflected in bending of the C–C–R angles which assume greater sp3 and sp2 character, thus strong π backbonding causes a metal-alkene complex to assume the character of a metallacyclopropane. Electronegative substituents exhibit greater π backbonding, thus, strong π backbonding ligands are tetrafluoroethylene, tetracyanoethylene, and hexafluoro-2-butyne
23.
Hexamminecobalt(III) chloride
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Hexaamminecobalt chloride is the chemical compound with the formula Cl3. This coordination compound is considered an archetypal Werner complex, named after the pioneer of coordination chemistry and this salt consists of 3+ trications with three Cl− anions. The term ammine refers to ammonia in its complexes. Originally this compound was described as a complex, but this name has been discarded as modern chemistry considers color less important than molecular structure. Other similar complexes also had names, such as purpureo for a pentammine complex. 3+ is diamagnetic, with a low-spin octahedral Co center, the cation obeys the 18-electron rule and is considered to be a classic example of an exchange inert metal complex. In contrast, labile metal complexes, such as Cl2. Upon heating, hexamminecobalt begins to lose some of its ammine ligands, the chloride ions in Cl3 can be exchanged with a variety of other anions such as nitrate, bromide, and iodide to afford the corresponding X3 derivative. Such salts are yellow and display varying degrees of water solubility. Since CoCl3 is not available, Cl3 is prepared from cobalt chloride, the latter is treated with ammonia and ammonium chloride followed by oxidation. Oxidants include hydrogen peroxide or oxygen in the presence of charcoal catalyst and this salt appears to have been first reported by Fremy. The acetate salt can be prepared by oxidation of cobalt acetate, ammonium acetate. The acetate salt is highly water-soluble to the level of 1.9 M, 3+ is a component of some structural biology methods, to help solve their structures by X-ray crystallography or by nuclear magnetic resonance
24.
Dissociative substitution
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Dissociative substitution describes a pathway by which compounds interchange ligands. The term is applied to coordination and organometallic complexes. This pathway can be described by the cis effect, or the labilization of CO ligands in the cis position. The opposite pathway is associative substitution, being analogous to Sn2 pathway, intermediate pathways exist between the pure dissociative and pure associative pathways, these are called interchange mechanisms. Complexes that undergo dissociative substitution are often coordinatively saturated and often have octahedral molecular geometry, the entropy of activation is characteristically positive for these reactions, which indicates that the disorder of the reacting system increases in the rate determining step. Dissociative pathways are characterized by a rate determining step that involves release of a ligand from the sphere of the metal undergoing substitution. The concentration of the nucleophile has no influence on this rate. Interchange pathways apply to substitution reactions where intermediates are not observed, if the reaction rate is insensitive to the nature of the attacking nucleophile, the process is called dissociative interchange, abbreviated Id. This rate is relevant to toxicity, catalysis, magnetic resonance imaging, for octahedral mono- and dicationic aquo complexes, these exchange processes occur via an interchange pathway that has more or less dissociative character. Rates vary by a factor of 1018, 3+ being the slowest, charge has a significant influence on these rates but non-electrostatic effects are also important. The rate for the hydrolysis of cobalt ammine halide complexes are deceptive, appearing to be associative, the hydrolysis of 2+ follows second order kinetics, the rate increases linearly with concentration of hydroxide as well as the starting complex. Studies show, however, that in the hydroxide deprotonates one NH3 ligand to give the base of the starting complex. In this monocation, the chloride spontaneously dissociates from this base of the starting complex. This pathway is called the Sn1CB mechanism
25.
Stoichiometry
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Stoichiometry /ˌstɔɪkiˈɒmᵻtri/ is the calculation of relative quantities of reactants and products in chemical reactions. This means that if the amounts of the reactants are known. Conversely, if one reactant has a quantity and the quantity of product can be empirically determined. This is illustrated in the image here, where the equation is. Here, one molecule of methane reacts with two molecules of gas to yield one molecule of carbon dioxide and two molecules of water. Stoichiometry measures these quantitative relationships, and is used to determine the amount of products/reactants that are produced/needed in a given reaction, describing the quantitative relationships among substances as they participate in chemical reactions is known as reaction stoichiometry. In the example above, reaction stoichiometry measures the relationship between the methane and oxygen as they react to form carbon dioxide and water. Gas stoichiometry deals with reactions involving gases, where the gases are at a temperature, pressure. For gases, the ratio is ideally the same by the ideal gas law. In practice, due to the existence of isotopes, molar masses are used instead when calculating the mass ratio, the term stoichiometry was first used by Jeremias Benjamin Richter in 1792 when the first volume of Richters Stoichiometry or the Art of Measuring the Chemical Elements was published. The term is derived from the Greek words στοιχεῖον stoicheion element, in patristic Greek, the word Stoichiometria was used by Nicephorus to refer to the number of line counts of the canonical New Testament and some of the Apocrypha. In general, chemical reactions combine in definite ratios of chemicals, since chemical reactions can neither create nor destroy matter, nor transmute one element into another, the amount of each element must be the same throughout the overall reaction. Chemical reactions, as unit operations, consist of simply a very large number of elementary reactions. As the reacting molecules consist of a set of atoms in an integer ratio. A reaction may consume more than one molecule, and the number counts this number, defined as positive for products. Different elements have a different atomic mass, and as collections of atoms, molecules have a definite molar mass. By definition, carbon-12 has a mass of 12 g/mol. The mass ratios can be calculated by dividing each by the total in the whole reaction, Elements in their natural state are mixtures of isotopes of differing mass, thus atomic masses and thus molar masses are not exactly integers
26.
Catalysis
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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
27.
Basis set (chemistry)
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The use of basis sets is equivalent to the use of an approximate resolution of the identity. The single-particle states are expressed as linear combinations of the basis functions. The basis set can either be composed of atomic orbitals, which is the choice within the quantum chemistry community. Several types of atomic orbitals can be used, Gaussian-type orbitals, Slater-type orbitals, out of the three, Gaussian-type orbitals are by far the most often used, as they allow efficient implementations of Post-Hartree–Fock methods. In modern computational chemistry, quantum chemical calculations are performed using a set of basis functions. When the finite basis is expanded towards a set of functions, calculations using such a basis set are said to approach the complete basis set limit. Within the basis set, the wavefunction is represented as a vector, one-electron Operators correspond to matrices, in this basis, whereas two-electron operators are rank four tensors. When molecular calculations are performed, it is common to use a basis composed of atomic orbitals, the physically best motivated basis set are Slater-type orbitals, which are solutions to the Schrödinger equation of hydrogen-like atoms, and decay exponentially far away from the nucleus. While hydrogen-like atoms lack many-electron interactions, it can be shown that the orbitals of Hartree-Fock. Furthermore, S-type STOs also satisfy Katos cusp condition at the nucleus, however, calculating integrals with STOs is computationally difficult and it was later realized by Frank Boys that STOs could be approximated as linear combinations of Gaussian-type orbitals instead. Because the product of two GTOs can be written as a combination of GTOs, integrals with Gaussian basis functions can be written in closed form. Dozens of Gaussian-type orbital basis sets have been published in the literature, basis sets typically come in hierarchies of increasing size, giving a controlled way to obtain a more accurate solutions, however at a higher cost. The smallest basis sets are called minimal basis sets, a minimal basis set is one in which, on each atom in the molecule, a single basis function is used for each orbital in a Hartree–Fock calculation on the free atom. For atoms such as lithium, basis functions of p type are also added to the functions that correspond to the 1s and 2s orbitals of the free atom. For example, each atom in the period of the periodic system would have a basis set of five functions. The minimal basis set is close to exact for the gas-phase atom, in the next level, additional functions are added to describe polarization of the electron density of the atom in molecules. For example, while the minimal basis set for hydrogen is one function approximating the 1s atomic orbital and this adds flexibility to the basis set, effectively allowing molecular orbitals involving the hydrogen atom to be more asymmetric about the hydrogen nucleus. This is very important for modeling chemical bonding, because the bonds are often polarized, similarly, d-type functions can be added to a basis set with valence p orbitals, and f-functions to a basis set with d-type orbitals, and so on
28.
Oxidation state
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The oxidation state, often called the oxidation number, is an indicator of the degree of oxidation of an atom in a chemical compound. This is never exactly true for real bonds, the term oxidation was first used by Lavoisier to signify reaction of a substance with oxygen. Much later, it was realized that the substance, upon being oxidized, loses electrons, oxidation states are typically represented by integers. In some cases, the oxidation state of an element is a fraction. It is predicted that even a +10 oxidation state may be achieveable by platinum in the platinum tetroxide dication, the increase in oxidation state of an atom, through a chemical reaction, is known as an oxidation, a decrease in oxidation state is known as a reduction. Such reactions involve the transfer of electrons, a net gain in electrons being a reduction. For pure elements, the state is zero. There are various methods for determining oxidation states, in inorganic nomenclature, the oxidation state is determined and expressed as an oxidation number, and is represented by a Roman numeral placed after the element name. In coordination chemistry, oxidation number is defined differently from oxidation state, a “Comprehensive definition of oxidation state ” has been published with free access. This article is a distillation of an IUPAC technical report Toward a comprehensive definition of oxidation state. Exceptions to this are that hydrogen has a state of −1 in hydrides of active metals, e. g. LiH. There are two different methods for determining the state of elements in chemical compounds. First a method based on the rules in the IUPAC definition, second, a method based on the relative electronegativity of the elements in the compound, where the more electronegative element is assumed to take the negative charge. Any pure element—even if it forms diatomic molecules like chlorine —has an oxidation state of zero, examples of this are Cu or O2. For monatomic ions, the state is the same as the charge of the ion. For example, the anion has an oxidation state of −2. The sum of oxidation states for all atoms in a molecule or polyatomic ion is equal to the charge of the molecule or ion, thus, the oxidation state of one element can be calculated from the oxidation states of the other elements. An application of this rule is that the sum of the states of all atoms in a neutral molecule must be zero
29.
Octahedral molecular geometry
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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
30.
Tetrahedral molecular geometry
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In a tetrahedral molecular geometry, a central atom is located at the center with four substituents that are located at the corners of a tetrahedron. The bond angles are cos−1 =109.4712206. ° ≈109. 5° when all four substituents are the same, the perfectly symmetrical tetrahedron belongs to point group Td, but most tetrahedral molecules have lower symmetry. Aside from virtually all saturated organic compounds, most compounds of Si, Ge, often tetrahedral molecules feature multiple bonding to the outer ligands, as in xenon tetroxide, the perchlorate ion, the sulfate ion, the phosphate ion. Thiazyl trifluoride is tetrahedral, featuring a triple bond. Other molecules have an arrangement of electron pairs around a central atom. However the usual classification considers only the atoms and not the lone pair. The H–N–H angles are 107°, contracted from 109.5 and this difference is attributed to the influence of the lone pair which exerts a greater repulsive influence than a bonded atom. Again the geometry is widespread, particularly so for complexes where the metal has d0 or d10 configuration, illustrative examples include tetrakispalladium, nickel carbonyl, and titanium tetrachloride. Many complexes with incompletely filled d-shells are often tetrahedral, e. g. the tetrahalides of iron, cobalt, however in liquid water or in ice, the lone pairs form hydrogen bonds with neighboring water molecules. The most common arrangement of hydrogen atoms around an oxygen is tetrahedral with two hydrogen atoms bonded to oxygen and two attached by hydrogen bonds. Since the hydrogen bonds vary in many of these water molecules are not symmetrical. Many compounds and complexes adopt bitetrahedral structures, in this motif, the two tetrahedra share a common edge. The inorganic polymer silicon disulfide features a chain of edge-shared tetrahedra. Inversion of tetrahedral occurs widely in organic and main group chemistry, the so-called Walden inversion illustrates the stereochemical consequences of inversion at carbon. Nitrogen inversion in ammonia also entails transient formation of planar NH3, geometrical constraints in a molecule can cause a severe distortion of idealized tetrahedral geometry. In compounds featuring inverted tetrahedral geometry at a carbon atom, all four groups attached to carbon are on one side of a plane. The carbon atom lies at or near the apex of a pyramid with the other four groups at the corners. The simplest examples of organic molecules displaying inverted tetrahedral geometry are the smallest propellanes, such as propellane, or more generally the paddlanes, such molecules are typically strained, resulting in increased reactivity
31.
Square planar molecular geometry
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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
32.
Vaska's complex
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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
33.
Zeise's salt
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Zeises salt, potassium trichloroplatinate, is the chemical compound with the formula K·H2O. The anion of this air-stable, yellow, coordination complex contains an η2-ethylene ligand, the anion features a platinum atom with a square planar geometry. The salt is of importance in the area of organometallic chemistry as one of the first examples of a transition metal alkene complex. This compound is available as a hydrate. The hydrate is commonly prepared from K2 and ethylene in the presence of an amount of SnCl2. The water of hydration can be removed in vacuo, the alkene C=C bond is approximately perpendicular to the PtCl3 plane. In Zeises salt and related compounds, the alkene rotates about the bond with a modest activation energy. Analysis of the barrier heights indicates that the π-bonding between most metals and the alkene is weaker than the σ-bonding, in Zeises anion, this rotational barrier has not been assessed. Zeises salt was one of the first organometallic compounds to be reported and it was discovered by William Christopher Zeise, a professor at the University of Copenhagen, who prepared this compound in 1830 while investigating the reaction of PtCl4 with boiling ethanol. Following careful analysis he proposed that the compound contained ethylene. Justus von Liebig, an influential chemist of that era, often criticised Zeises proposal. Zeises salt received a deal of attention during the second half of the 19th century because chemists could not explain its molecular structure. This question remained unanswered until the determination of its X-ray crystal structure in the 20th century, Zeises salt stimulated much scientific research in the field of organometallic chemistry and would be key in defining new concepts in chemistry. The Dewar–Chatt–Duncanson model explains how the metal is coordinated to the C=C double bond, Zeises dimer,2, derived from Zeises salt by elimination of KCl followed by dimerisation. COD-platinum dichloride, PtCl2, derived from platinum chloride and 1, many other ethylene complexes have been prepared. For example, ethylenebisplatinum, 2Pt, wherein the platinum is three-coordinate and in oxidation state 0
34.
Catalytic cycle
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A catalytic cycle in chemistry is a term for a multistep reaction mechanism that involves a catalyst. The catalytic cycle is the method for describing the role of catalysts in biochemistry, organometallic chemistry, bioinorganic chemistry, materials science. Often such cycles show the conversion of a precatalyst to the catalyst, since catalysts are regenerated, catalytic cycles are usually written as a sequence of chemical reactions in the form of a loop. In such loops, the initial step entails binding of one or more reactants by the catalyst, articles on the Monsanto process, the Wacker process, and the Heck reaction show catalytic cycles. A catalytic cycle is not necessarily a full reaction mechanism, for example, it may be that the intermediates have been detected, but it is not known by which mechanisms the actual elementary reactions occur. Catalytic cycles also have an important role in the atmosphere, for example in the reactions that lead to ozone depletion, the complicated interplay between different schemes, including null cycles that have no overall effect, determines the rate of production and destruction of many atmospheric components. Often a so-called sacrificial catalyst is also part of the system with the purpose of regenerating the true catalyst in each cycle. As the name implies, the sacrificial catalyst is not regenerated and irreversibly consumed and this sacrificial compound is also known as a stoichiometric catalyst when added in stoichiometric quantities compared to the main reactant. Usually the true catalyst is an expensive and complex molecule and added in quantities as small as possible, the stoichiometric catalyst on the other hand should be cheap and abundant. Sacrificial catalysts are more accurately referred to by their role in the catalytic cycle, for example as a reductant
35.
Monsanto process
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The Monsanto process is an industrial method for the manufacture of acetic acid by catalytic carbonylation of methanol. The Monsanto process has largely supplanted by the Cativa process. This process operates at a pressure of 30–60 atm and a temperature of 150–200 °C and it was developed in 1960 by the German chemical company, BASF, and improved by the Monsanto Company in 1966, which introduced a new catalyst system. The catalytically active species is the anion cis-−, the first organometallic step is the oxidative addition of methyl iodide to cis-− to form the hexacoordinate species −. This anion rapidly transforms, via the migration of a group to an adjacent carbonyl ligand. This five-coordinate complex then reacts with carbon monoxide to form the six-coordinate dicarbonyl complex, the catalytic cycle involves two non-organometallic steps, conversion of methanol to methyl iodide and the hydrolysis of the acetyl iodide to acetic acid and hydrogen iodide. The reaction has shown to be first-order with respect to methyl iodide. Hence the oxidative addition of iodide is proposed as rate-determining step. Acetic anhydride is produced by carbonylation of methyl acetate in a process that is related to the Monsanto acetic acid synthesis. CH3CO2CH3 + CO → 2O In this process lithium iodide converts methyl acetate to lithium acetate and methyl iodide, acetyl iodide reacts with acetate salts or acetic acid to give the anhydride. Rhodium iodides and lithium salts are employed as catalysts, because acetic anhydride is not stable in water, the conversion is conducted under anhydrous conditions in contrast to the Monsanto acetic acid synthesis
36.
Hydrogenation
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Hydrogenation – to treat with hydrogen – is a chemical reaction between molecular hydrogen and another compound or element, usually in the presence of a catalyst such as nickel, palladium or platinum. The process is employed to reduce or saturate organic compounds. Hydrogenation typically constitutes the addition of pairs of atoms to a molecule. Catalysts are required for the reaction to be usable, non-catalytic hydrogenation takes place only at high temperatures. Hydrogenation reduces double and triple bonds in hydrocarbons and it has three components, the unsaturated substrate, the hydrogen and, invariably, a catalyst. The reduction reaction is carried out at different temperatures and pressures depending upon the substrate, the same catalysts and conditions that are used for hydrogenation reactions can also lead to isomerization of the alkenes from cis to trans. This process is of great interest because hydrogenation technology generates most of the fat in foods. A reaction where bonds are broken while hydrogen is added is called hydrogenolysis, some hydrogenations of polar bonds are accompanied by hydrogenolysis. For hydrogenation, the source of hydrogen is H2 gas itself. The hydrogenation process often uses greater than 1 atmosphere of H2, usually conveyed from the cylinders, gaseous hydrogen is produced industrially from hydrocarbons by the process known as steam reforming. For many applications, hydrogen is transferred from donor molecules such as acid, isopropanol. These hydrogen donors undergo dehydrogenation to, respectively, carbon dioxide, acetone and these processes are called transfer hydrogenations. Typical substrates are listed in the table With rare exceptions, H2 is unreactive toward organic compounds in the absence of metal catalysts, the unsaturated substrate is chemisorbed onto the catalyst, with most sites covered by the substrate. In heterogeneous catalysts, hydrogen forms surface hydrides from which hydrogens can be transferred to the chemisorbed substrate, platinum, palladium, rhodium, and ruthenium form highly active catalysts, which operate at lower temperatures and lower pressures of H2. Non-precious metal catalysts, especially based on nickel have also been developed as economical alternatives. The trade-off is activity vs. cost of the catalyst and cost of the apparatus required for use of high pressures, notice that the Raney-nickel catalysed hydrogenations require high pressures, Catalysts are usually classified into two broad classes, homogeneous catalysts and heterogeneous catalysts. Homogeneous catalysts dissolve in the solvent that contains the unsaturated substrate, heterogeneous catalysts are solids that are suspended in the same solvent with the substrate or are treated with gaseous substrate. Some well known homogeneous catalysts are indicated below and these are coordination complexes that activate both the unsaturated substrate and the H2
37.
Hydroformylation
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Hydroformylation, also known as oxo synthesis or oxo process, is an industrial process for the production of aldehydes from alkenes. The process was developed by the German chemist Otto Roelen in 1938 and this chemical reaction entails the net addition of a formyl group and a hydrogen atom to a carbon-carbon double bond. This process has undergone continuous growth since its invention, Production capacity reached 6. 6×106 tons in 1995 and it is important because aldehydes are easily converted into many secondary products. For example, the resulting aldehydes are hydrogenated to alcohols that are converted to plasticizers or detergents, Hydroformylation is also used in speciality chemicals, relevant to the organic synthesis of fragrances and drugs. The development of hydroformylation, which originated within the German coal-based industry, is considered one of the achievements of 20th-century industrial chemistry. The process typically entails treatment of an alkene with high pressures of carbon monoxide, invariably, the catalyst dissolves in the reaction medium, i. e. hydroformylation is an example of homogeneous catalysis. The discovery of this reaction is attributed to Otto Roelen, who was investigating the Fischer-Tropsch reaction, aldehydes and diethylketone were obtained when ethylene was added to an F-T reactor. Through these studies, Roelen discovered the involvement of cobalt catalysts, hCo4, which had been isolated prior to Roelen, was shown to be an excellent catalyst. The term oxo synthesis was created by the Ruhrchemie patent department, subsequent work demonstrated that the ligand tributylphosphine improved the selectivity of the cobalt-catalysed process. In the 1960s, highly active rhodium catalysts were discovered, since the 1970s, most hydroformylation relies on catalysts based on rhodium. Subsequent research led to the development of catalysts that facilitate the separation of the products from the catalyst. The overall mechanism resembles that for homogeneous hydrogenation with additional steps, the reaction begins with the generation of coordinatively unsaturated metal hydrido carbonyl complex such as HCo3 and HRh3. Such species bind alkenes, and the complex undergoes a migratory insertion reaction to form an alkyl complex. A key consideration of hydroformylation is the normal vs. iso selectivity, of course, both products are not equally desirable. Much research was dedicated to the quest for catalyst that favored the normal isomer, when the hydrogen is transferred to the carbon bearing the most hydrogen atoms the resulting alkyl group has a larger steric bulk close to the ligands on the cobalt. If the ligands on the cobalt are bulky, then this steric effect is greater, hence, the mixed carbonyl/phosphine complexes offer a greater selectivity toward the straight chain products. In addition, the more electron-rich the hydride complex is the less proton-like the hydride is, thus, as a result, the electronic effects that favour the Markovnikov addition to an alkene are less able to direct the hydride to the carbon atom bearing the most hydrogens already. Thus, as a result, as the centre becomes more electron-rich
38.
Pentamethylcyclopentadiene
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1,2,3,4, 5-Pentamethylcyclopentadiene is a cyclic dialkene with the formula C5Me5H. 1,2,3,4, 5-Pentamethylcyclopentadiene is the precursor to the ligand 1,2,3,4, 5-pentamethylcyclopentadienyl, in contrast to less-substituted cyclopentadiene derivatives, Cp*H is not prone to dimerization. It was first prepared from tiglaldehyde via 2,3,4, an instructive but obsolete route to Cp* complexes involves the use of hexamethyl Dewar benzene. This method was used for preparation of the chloro-bridged dimers 2 and 2. Complexes of pentamethylcyclopentadienyl differ in ways from the more common cyclopentadienyl derivatives. Being more electron-rich, Cp* is a donor and is less easily displaced from the metal. Its steric bulk stabilizes complexes with fragile ligands and its bulk also attenuates intermolecular interactions, decreasing the tendency to form polymeric structures. Its complexes also tend to be soluble in non-polar solvents. The methyl group in Cp* complexes can undergo C–H activation leading to tuck-in complexes
39.
Vanadium hexacarbonyl
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Vanadium hexacarbonyl is the inorganic compound with the formula V6. It is a volatile solid. This highly reactive species is noteworthy from theoretical perspectives as a rare isolable homoleptic metal carbonyl that is paramagnetic, most species with the formula Mxy follow the 18-electron rule, whereas V6 has 17 valence electrons. Traditionally V6 is prepared in two-steps via the intermediacy of V−6, in the first step, VCl3 is reduced with metallic sodium under 200 atm CO at 160 °C. The solvent for this reduction is typically diglyme, CH3OCH2CH2OCH2CH2OCH3 and its primary reaction is reduction to the monoanion V−6, salts of which are well studied. It is also susceptible to substitution by tertiary phosphine ligands, often leading to disproportionation, V6 reacts with sources of the cyclopentadienyl anion to give the orange four-legged piano stool complex V4. Like many charge-neutral organometallic compounds, this species is volatile. In the original preparation of this species, C5H5HgCl was employed as the source of C 5H−5, V6 adopts an octahedral coordination geometry and is isostructural with chromium hexacarbonyl, even though they have differing valence electron counts. High resolution X-ray crystallography indicates that the molecule is distorted with two shorter V–C distances of 1.993 Å vs. four 2.005 Å. Even though V is a larger ion than V, the V–C distances in V−6 are 0.07 Å shorter than in the neutral precursor, original synthesis, Calderazzo, F. Ercoli, R. Synthesis of V6 and Hexacarbonyl Vanadates
40.
Butyl group
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In organic chemistry, butyl is a four-carbon alkyl radical or substituent group with general chemical formula −C4H9, derived from either of the two isomers of butane. In the convention of skeletal formulas, every line ending and line intersection specifies a carbon atom saturated with single-linked hydrogen atoms, the R symbol indicates any radical or other non-specific functional group. Butyl is the largest substituent for which names are commonly used for all isomers. The butyl groups carbon that is connected to the rest of the molecule is called the RI or R-prime carbon, the prefixes sec and tert refer to the number of additional side chains connected to the first butyl carbon. The prefix iso means equal while the prefix n- stands for normal, the four isomers of butyl acetate demonstrate these four isomeric configurations. In that progression, Butyl is the fourth, and the last to be named for its history, the word butyl is derived from butyric acid, a four-carbon carboxylic acid found in rancid butter. The name butyric acid comes from Latin butyrum, butter, subsequent alkyl radicals in the series are simply named from the Greek number that indicates the number of carbon atoms in the group, pentyl, hexyl, heptyl, etc. The tert-butyl substituent is very bulky and is used in chemistry for kinetic stabilization, the effect of the tert-butyl group on the progress of a chemical reaction is called the tert-butyl effect, illustrated in the Diels-Alder reaction below. Compared to a substituent, the tert-butyl substituent accelerates the reaction rate by a factor of 240
41.
Norbornyl
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Norbornane is an organic compound and a saturated hydrocarbon with chemical formula C7H12. It is a compound with melting point 88 °C. The carbon skeleton is derived from cyclohexane ring with a bridge in the 1, 4- position. The compound is a prototype of a class of strained bicyclic hydrocarbons, the compound was originally synthesized by reduction of norcamphor. The name norbornane is derived from bornane, which is 1,7, 7-trimethylnorbornane, the prefix nor refers to the stripping of the methyl groups from the parent molecule bornane. 2-Norbornyl cation Norbornene Norbornadiene Bornane endo-Norborneol exo-Norborneol Norcamphor, the derivative of norbornane Norbornane in 3D Datasheet at Sigma-Aldrich
42.
Agostic interaction
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Many catalytic transformations, e. g. oxidative addition and reductive elimination, are proposed to proceed via intermediates featuring agostic interactions. Agostic interactions are observed throughout organometallic chemistry in alkyl, alkylidene, often such agostic interactions involve alkyl or aryl groups that are held close to the metal center through an additional σ-bond. Short interactions between hydrocarbon substituents and coordinatively unsaturated metal complexes have been noted since the 1960s, for example, in tris ruthenium dichloride, a short interaction is observed between the ruthenium center and a hydrogen atom on the ortho position of one of the nine phenyl rings. Complexes of borohydride are described as using the three-center two-electron bonding model, the nature of the interaction was foreshadowed in main group chemistry in the structural chemistry of trimethylaluminium. Agostic interactions are best demonstrated by crystallography, neutron diffraction data has shown that C−H and M┄H bond distances are 5-20% longer than expected for isolated metal hydride and hydrocarbons. The distance between the metal and the hydrogen is typically 1. 8–2.3 Å, and the M┄H−C angle falls in the range 90°–140°. The presence of a 1H NMR signal that is shifted upfield from that of an aryl or alkane. The coupling constant 1JCH is typically lowered to 70–100 Hz versus the 125 Hz expected for a normal sp3 carbon–hydrogen bond, on the basis of experimental and computational studies, the stabilization arising from an agostic interaction is estimated to be 10–15 kcal/mol. Recent calculations using compliance constants point to a weaker stabilisation, thus, agostic interactions are stronger than most hydrogen bonds. Agostic bonds sometimes play a role in catalysis by increasing rigidity in transition states, for instance, in Ziegler–Natta catalysis the highly electrophilic metal center has agostic interactions with the growing polymer chain. This increased rigidity influences the stereoselectivity of the polymerization process, the term agostic is reserved to describe two-electron, three-center bonding interactions between carbon, hydrogen, and a metal. Two-electron three-center bonding is clearly implicated in the complexation of H2, e. g. in W32H2, silane binds to metal centers often via agostic-like, three-centered Si┄H−M interactions. Because these interactions do not include carbon, however, they are not classified as agostic, certain M┄H−C interactions are not classified as agostic but are described by the term anagostic. Anagostic interactions are more electrostatic in character, in terms of structures of anagostic interactions, the M┄H distances and M┄H−C angles fall into the ranges 2. 3–2.9 Å and 110°–170°, respectively. Agostic interactions serve a key function in alkene polymerization and stereochemistry, as well as migratory insertion
43.
Oxygen
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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
44.
Tetrahydrofuran
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Tetrahydrofuran, whose preferred IUPAC name was changed in 2013 to oxolane, is an organic compound with the formula 4O. The compound is classified as heterocyclic compound, specifically a cyclic ether and it is a colorless, water-miscible organic liquid with low viscosity. It is mainly used as a precursor to polymers, being polar and having a wide liquid range, THF is a versatile solvent. About 200,000 tonnes of tetrahydrofuran are produced annually, the most widely used industrial process involves the acid-catalyzed dehydration of 1, 4-butanediol. Ashland/ISP is one the biggest producers of this chemical route, the method is similar to the production of diethyl ether from ethanol. The butanediol is derived from condensation of acetylene with formaldehyde followed by hydrogenation, duPont developed a process for producing THF by oxidizing n-butane to crude maleic anhydride, followed by catalytic hydrogenation. A third major industrial route entails hydroformylation of allyl alcohol followed by hydrogenation to the butanediol, THF can also be synthesized by catalytic hydrogenation of furan. Certain sugars can be converted to THF, although this method is not widely practiced, furan is thus derivable from renewable resources. The other main application of THF is as a solvent for polyvinyl chloride. It is a solvent with a dielectric constant of 7.6. It is a polar solvent and can dissolve a wide range of nonpolar and polar chemical compounds. THF is water-miscible and can form solid clathrate hydrate structures with water at low temperatures, in the laboratory, THF is a popular solvent when its water miscibility is not an issue. It is more basic than diethyl ether and forms complexes with Li+, Mg2+. It is a solvent for hydroboration reactions and for organometallic compounds such as organolithium. Although similar to diethyl ether, THF is a stronger base, commercial THF contains substantial water that must be removed for sensitive operations, e. g. those involving organometallic compounds. Although THF is traditionally dried by distillation from an aggressive desiccant, aqueous THF augments the hydrolysis of glycans from biomass and dissolves the majority of biomass lignin making it a suitable solvent for biomass pretreatment. THF is often used in polymer science, for example, it can be used to dissolve polymers prior to determining their molecular mass using gel permeation chromatography. THF dissolves PVC as well, and thus it is the ingredient in PVC adhesives
