Key Concepts in Coordination Chemistry

The simple yet distinctive concept of the coordinate bond (also sometimes called a dative bond) lies at the core of coordination chemistry. Molecular structure, in its simplest sense, is interpreted in terms of covalent bonds formed through shared pairs of electrons. The coordinate bond, however, arises not through the sharing of electrons, one from each of two partner atoms, as occurs in a standard covalent bond, but from the donation of a pair of electrons from an orbital on one atom (a lone pair) to occupy an empty orbital on what will become its partner atom. First introduced by G.N. Lewis almost a century ago, the concept of a covalent bond formed when two atoms share an electron pair remains as a firm basis of chemistry, giving us a basic understanding of single, double and triple bonds, as well as of a lone pair of electrons on an atom. Evolving from these simple concepts came valence bond theory, an early quantum mechanical theory which expressed the concepts of Lewis in terms of wave functions. These concepts still find traditional roles in coordination chemistry. However, coordination chemistry is marked by a need to employ the additional concept of coordinate bond formation, where the bond pair of electrons originates on one of the two partner atoms alone. In coordinate bond formation, the bonding arrangement between electron-pair acceptor (designated as A) and electron-pair donor (designated as :D, where the pair of dots represent the lone pair of electrons) can be represented simply as Equation (1.1):

A+:D→A:D

The product alternatively may be written as A←:D or A←D, where the arrow denotes the direction of electron donation, or, where the nature of the bonding is understood, simply as A D.Thislatter standard representation is entirely appropriate since the covalent bond, once formed, is indistinguishable from a standard covalent bond. The process should be considered reversible in the sense that, if the A D bond is broken, the lone pair of electrons originally donated by :D remains entirely with that entity. 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 ofelectrons todonate toanemptyorbitalonthe 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.

Figure 1.1

A schematic view of ammonia acting as a donor ligand to a metalloid acceptor and to a metal ion acceptor to form coordinate bonds.

The species providing the electron pair (the electron-pair donor) is thought of as being coordinated to the species receiving that lone pair of electrons (the electron-pair acceptor). The coordinating entity, the ligand, can be as small as a monatomic ion (e.g. F−)oraslarge as a polymer– the key characteristic is the presence of one or more lone pairs of electrons on an electronegative donor atom. Donor atoms often met are heteroatoms like N, O, S and P as well as halide ions, but this is by no means the full range. Moreover, the vast majority of existing organic molecules can act as ligands, or else can be converted into molecules capable of acting as ligands. A classical and successful ligand is ammonia, NH3, which has one lone pair (Figure 1.1). Isoelectronic with ammonia is the carbanion −CH3, which can also be considered a ligand under the simple definition applied; even hydrogen as its hydride, H−, has a pair of electrons and can act as a ligand. It is not the type of donor atom that is the key, but rather its capacity to supply an electron pair. The acceptor with which a coordinate covalent bond is formed is conventionally either a metal or metalloid. With a metalloid, covalent bond formation is invariably associated with an increase in the number of groups or atoms attached to the central atom, and simple electron counting based on the donor–acceptor concept can account for the number of coordinate covalent bonds formed. With a metal ion, the simple model is less applicable, since the number of new bonds able to be generated through complexation doesn’t necessarily match the number of apparent vacancies in the valence shell of the metal; a more sophisticated model needs to be applied, and will be developed herein. What is apparent with metal ions in particular is the strong drive towards complexation– ‘naked’ ions are extremely rare, and even in the gaseous state complexation will occur. It is a case of the whole being better than the sum of its parts, or, put more appropriately, coordinate bond formation is energetically favorable. Amor elaborate example than those shown above is the anionic compound SiF6^2- (Fig ure 1.2), which adopts a classical octahedral shape that we will meet also in many metal complexes. Silicon lies below carbon in the Periodic Table, and there are some limited similarities in their chemistry. However the simple valence bond theory and octet rule that

Figure 1.2 The octahedral [SiF6]^2- molecular ion, and a simple valence bond approach to explaining its forma tion. Overlap of a p orbital containing two electrons on each of the six fluoride anions with one of six empty hybrid orbitals on the Si (IV) cation, arranged in an octahedral array, generates the octahedral shape with six equivalent covalent bonds.

works so well for carbon cannot deal with a silicon compound with six bonds, particularly one where all six bonds are equivalent. One way of viewing this molecular species is as being composed of a Si4+ or Si(IV) centre with six F− anions bound to it through each f luoride anionusinganelectronpair (: F−) to donate to an empty orbital on the central Si (IV) ion, which has lost all of its original four valence electrons in forming the Si4+ ion. Using traditional valence bond theory concepts, a process of hybridization is necessary to accom modate the outcome (Figure 1.2). The generation of the shape arises through asserting that the silicon arranges a combination of one 3s, three 3p and two of five available 3d valence orbitals into six equivalent sp3d2 hybrid orbitals that are directed as far apart as possible and towards the six corners of an octahedron. Each empty hybrid orbital then accommodates an electron pair from a fluoride ion, each leading in effect to a coordinate covalent bond that is a bondbecause electron density in the bond lies along the line joining the two atomic centres. The shape depends on the type and number of orbitals that are involved in the hy bridization process. Above, a combination resulting in an octahedral shape (sp3d2 hybrids) is developed; however, different combinations of orbitals yield different shapes, perhaps the most familiar being the combination of one s and three p orbitals to yield tetrahedral sp3; others examples are linear (sp hybrids) and trigonal planar (sp2 hybrids) shapes. A central atom or ion with vacant or empty orbitals and ionic or neutral atoms or molecules joining it, with each bringing lone pairs of electrons ,is the classic requirement for formation of what we have termed coordinate bonds, leading to a coordination compound. The very basic valence bonding model described above can be extended to metal ions, as we will see, but with some adjustments due to the presence of electrons in the d orbitals; more sophisticated models are required. Of developed approaches, molecular orbital theory is the most sophisticated, and is focused on the overlap of atomic orbitals of comparable energy on different atoms to form molecular orbitals to which electrons are allocated. While providing accurate descriptions of molecules and their properties, it is relatively complicated and time-consuming, and somewhat difficult to comprehend for large complexes; consequently, simpler models still tend to be used. In the simple theory based on Lewis’ concepts exemplified above, the key aspects are an empty orbital on one atom and a filled or bital (with a pair of electrons present, the lone pair) on the other. Many of the ligand species providing the lone pair are considered bases in the classical Brønsted–Lowry concept of acids and bases (which has as its focus the transfer  of a proton), since these species are able to accept a proton. However, in the description we have developed here, no proton is involved, but the concept of accepting an electron deficient species does apply. The broader and more general concept of an electron-pair donor as a base and an electron-pair acceptor as the acid evolved, and these are called a Lewis base(electron-pairdonor)and a Lewis acid (electron-pair acceptor).Consequently, an H3B

NH3 compound is traditionally considered a coordination compound, arising through coordination of the electron deficient  (or Lewis acid) H3B and the electron lone-pair containing (or Lewis base) compound :NH3 (Figure 1.1). It is harder, in part as a result of entrenched views of covalent bonding in carbon-based compounds, to accept [H3C NH3] + in similar terms purely as a H3C+ and: NH₃ assembly. This need to consider and debate the nature of the assembly limits the value of the model for non-metals and metalloids. With metal ions, however, you tend to know where you stand– almost invariably, you may start by considering them as forming coordination compounds; perhaps it is not surprising that coordination chemistry is focused mainly on compounds of metals and their ions.

Coordination has a range of consequences for the new assembly. It leads to structural change, seen in terms of change in the number of bonds and/or bond angles and distances. This is inevitably tied to a change in the physical properties of the assembly, which differ from those of its separate components. With metal atoms or ions at the centre of a coordination complex, even changing one of a set of ligands will be reflected in readily observable change in physical properties, such as colour. With growing sophistication in both synthesis and our understanding of physical methods, properties can often be ‘tuned’ through varying ligands to produce a particular result, such as a desired reduction potential. It should also be noted that a coordination compound adopts one of a limited number of basic shapes, with the shape determined by the nature of the central atom and its attached ligands. Moreover, the physical properties of the coordination compound depend on and reflect the nature of the central atom, ligand set and molecular shape. Whereas only one central atom occurs in many coordination compounds (a compound wemaythusdefineasa monomer), it should also be noted that there exists a large and growing range of compounds where there are two or more ‘central atoms’, either of the same or different types. These ‘central atoms’ are linked together through direct atom-to-atom bonding, or else are linked by ligands that as a result are joined to at least two ‘central atoms’ at the same time. This latter arrangement, where one or even several ligands are said to ‘bridge’ between central atoms, is the more common of these two options. The resulting species can usually be thought of as a set of monomer units linked together, leading to what is formally a polymer or, more correctly when only a small number of units are linked, an oligomer. We shall concentrate largely on simple monomeric species herein, but will introduce examples of larger linked compounds where appropriate. Although as we have seen, the metalloid elements can form molecular species that we call coordination compounds, the decision on what constitutes a coordination compound is perhaps more subtle with these than is the case with metals. Consequently, in this tale of complexes and ligands, it is with metals and particularly their cations as the central atom that we will almost exclusively meet examples.

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مواضيع ذات صلة

Meeting More Metals - Bridging Ligands

More Teeth Stronger Bite - Polydentates

Greed is Good– Polydentate Ligands The Simple Chelate

Monodentate Ligands– The Simple Type Basic Binders

Ligands with More Bite– Denticity

Different Tribes– Donor Group Variation

Do I Look Big on That?– Chelate Ring Size

Putting the Bite on Metals– Chelation

Making Attachments– Coordination

The Road Ahead

Natural or Made-to-Measure Complexes

Metals in Contrived Environments