Central atom concepts metals

The original concepts of metal-ligand bonding were essentially related to the dative covalent bond the development of organometallic chemistry has revealed a further way in which ligands can supply more than one electron pair to a central atom. This is exemplified by the classical cases of bis(benzene)chromium and bis(cyclopentadienyl)iron, trivial name ferrocene. These molecules are characterised by the bonding of a formally unsaturated system (in the organic chemistry sense, but expanded to include aromatic systems) to a central atom, usually a metal atom. 

A metal sandwich compound does not strictly fit our previous concept of an ABfl-type molecule, since the ligand atoms interact strongly with each other as well as with the central atom. It is desirable to extend the discussion to these molecules, however, since they provide clear and important examples of how to treat the situation in which the ligands in a complex are themselves polyatomic entities with an internal set of MOs perturbed by interaction with the AOs of the central atom. 

Some of the chemical concepts with little or no quantum-mechanical meaning outside the Bohmian formulation but, well explained in terms of the new interpretation, include electronegativity, the valence state, chemical potential, metallization, chemical bonding, isomerism, chemical equilibrium, orbital angular momentum, bond strength, molecular shape, phase transformation, chirality and barriers to rotation. In addition, atomic stability is explained in terms of a simple physical model. The central new concepts in Bohmian mechanics are quantum potential and quantum torque. 

Let us turn now to the nature of the bond that is normally found in M—CO groups. In most cases, it is adequate for practical, everyday purposes to regard a ligand simply as an electron pair donor and to think of the bond to the central atom simply as L—>M. However, there are important classes of compounds for which this simple concept is seriously inadequate. The most prominent examples are the metal carbonyls, metal nitrosyls, and compounds of low-valent metals containing phosphines or isonitriles as ligands. 

The idea of lone pairs was originated by W. J. Pope of Cambridge in 1900 who extended the concept of the three-dimensionality of carbon and nitrogen compounds to those of sulfur. His resolution of sulfonium cations RR R"S+ with three different substituents into optically active enantiomers suggested that these species were tetrahedral with an invisible substituent. The influence of these lone pairs can hardly be detected in transition metal compounds, but the situation is different for post-transition group central atoms such as Ge(II) As(III), Se(IV), and Br(V) with 30 electrons, In(I), Sn(II), Sb(III), Te(IV), I(V), and Xe(VI) with 48 electrons, and Au( —I), T1(I), Pb(II), and Bi(III) with 80 electrons (90). 

While these concepts have usually been applied to metal compounds, a wide range of other species can be considered to consist of a central atom or central atoms to which a number of other groups are bound. The application of additive nomenclature to such species is briefly described and exemplified in Chapter IR-7, and abundantly exemplified for inorganic acids in Chapter IR-8. 

Electron counting in transition metal complexes (18 e rule) is an useful tool to understand their stability and structure, although it does not apply to all transition metal complexes—only for a majority of compounds containing 7i-acceptor ligands. The 18 electron rule is an extension of the idea of the octet rule, which applies to atoms having only s and p orbitals. The idea is that the molecule will be stable when the central atom has the same electronic structure as noble gases of the same row. A similar concept can be applied to transition metal complexes having d electrons. The compound is considered most stable when the total number of electrons around the atom becomes the... 

In most coordination compounds it is possible to identify a central or core atom or ion that is bonded not simply to one other atom, ion or group through a coordinate bond, but to several of these entities at once. The central atom is an acceptor, with the surrounding species each bringing (at least) one lone pair of electrons to donate to an empty orbital on the central atom, and each of these electron-pair donors is called a ligand when attached. The central atom is a metal or metalloid, and the compound that results from bond formation is called a coordination compound, coordination complex or often simply a complex. We shall explore these concepts further below. 

In a generalized sense, acids are electron pair acceptors. They include both protic (Bronsted) acids and Lewis acids such as AlCb and BF3 that have an electron-deficient central metal atom. Consequently, there is a priori no difference between Bronsted (protic) and Lewis acids. In extending the concept of superacidity to Lewis acid halides, those stronger than anhydrous aluminum chloride (the most commonly used Friedel-Crafts acid) are considered super Lewis acids. These superacidic Lewis acids include such higher-valence fluorides as antimony, arsenic, tantalum, niobium, and bismuth pentafluorides. Superacidity encompasses both very strong Bronsted and Lewis acids and their conjugate acid systems. 

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