1.
Organometallic chemistry
–
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
–
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
–
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
–
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
–
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
–
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
–
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
–
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
–
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
–
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
