Complex Shape - Not Just Any Which Way

It is now over two hundred years since the realization that chemical compounds have a distinctive three-dimensional shape was expressed by Wollaston in a paper published in 1808. As early as 1860 Pasteur suggested a tetrahedral grouping around carbon and both Kekulé (in 1867) and Paterno in (1869) developed three-dimensional models involving tetrahedral carbon. From 1864 Crum Brown was developing graphical representations of inorganic compounds. By the second half of the nineteenth century, the tetrahedral carbon had been well defined experimentally by van't Hoff and LeBel (in 1874) through optical resolution of compounds. Following this approach, the tetrahedral nitrogen was established (by LeBel in 1891), then other tetrahedral centres, and subsequently other stereochemistries. The use of single crystal X-ray diffraction commenced with the work of William and Lawrence Bragg in the early twentieth century and from the 1920s was established as the supreme method for determination of the absolute three-dimensional shape of molecules in the solid state.

Shapes of transition metal compounds received limited attention until late in the nineteenth century. From 1875 for several decades, Sophus Mads Jørgensen championed graphical chain formulae for metal complexes, building on an original proposal of C.W. Blomstrand of 1871. For cobalt(III) three direct bonds to the cobalt was assumed by defining a relationship between metal oxidation state and number of bonds, with other groups linked in chains reminiscent of the chains in organic compounds that had been developed a little earlier. Three representations are shown in Figure 3.2, along with modern formulations; for these chain formulae, a key aspect was that Co-Cl bonds were considered unable to ionize in solution, whereas NH₃-Cl bonds could do so. This fitted with the ionization behaviour of the first three compounds presented - a case of a model interpreting (albeit very limited) experimental evidence satisfactorily. However, the nonelectrolyte CoCl₃ (NH₃)₃ could not be fitted at all with this model with its three-bond limit to cobalt(III, with at best a 1:1 electrolyte proposed. Alfred Werner rejected this model, and in a series of studies between 1891 and 1893 introduced two important concepts to coordination chemistry: stereochemistry and 'affinities'. His formulation for CoCl₃ (NH₃)₃ satisfies the nonelectrolytic nature of the compound (see Figure 3.2), and takes the 'modern' shape, also now proven absolutely by structural studies.

Werner's theory is generally seen as the birth of modern coordination chemistry. He pro- posed two valencies for metal ions: primary (ionizable; corresponds to oxidation state) and secondary (nonionizable; corresponds to coordination number), with the latter (usually four or six) able to be satisfied by neutral or anionic molecules, and with every element usually satisfying both its primary and secondary valence. Further, he recognized and proposed that these 'secondary valencies' would be distributed around the metal in fixed positions in space in such a way as to lead to the least repulsion between the groups or atoms attached. Subsequently, he devoted himself to finding experimental proof for his theory, using the limited methods available at the time, which included conductivity and resolution of com- plexes into optical isomers. His suite of studies of six-coordinate compounds, which he argued should exist in many cases as geometric and/or optical isomers, convincingly sup- ported his predictions. By isolating optical isomers of a complex without any carbon atoms present, he also put to rest the view that optical activity was a property of carbon compounds

Figure 3.2

Jørgensen's chain theory for some cobalt(III) complexes (top), compared with the Werner represen- tations (bottom). Werner's representation for [CoCl (NH3)3] fits the nonelectrolyte behaviour seen experimentally, unlike the Jørgensen model.

alone. Werner also proposed in 1912 the concept of a second (outer) coordination sphere through directed but weaker interactions of groups in the first (inner) coordination sphere with species arranged in an outer shell around the main coordination sphere. These effects, involving what we would now call specific hydrogen bonding interactions, were later shown by Pfeiffer (in 1931) to exist through his work on optically active substances interacting with an optically inactive central metal complex, where an effect was induced in the metal complex despite its not being bonded coordinatively to the optically active species.

Werner's early representations of coordination complexes go a long way towards the representations we use today. This is illustrated in Figure 3.2 for what were initially formu- lated as CoCl₃.6NH₃ and CoCl₃ (NH₃)₃. The former is formulated as [Co(NH₃)₆]Cl₃, and was assigned a primary valence (oxidation state) of three and a secondary valence (coordi- nation number) of six. The three chloride ions that neutralize the charge on the cobalt(III) ion fully satisfy the primary valence, and the six ammonia molecules use the secondary valence fully, through being coordinated to the metal centre. Because the ammonias fully satisfy the coordination number, the chlorides do not participate in the same way and are disposed further away, as represented in Figure 3.2. This exposes them more to reaction, such as combining readily with silver ion to precipitate AgCl. For CoCl₃-4NH₃ Werner applied his postulate that both primary and secondary valence seek to be satisfied. Thus, with only four ammonia molecules present, he included two of the three chloride ions as additional components of the coordination sphere to satisfy the coordination number of six. These two then serve a dual function, and satisfy both valencies; Werner represented them by a combined solid and dashed line, as exemplified in Figure 3.2. Unlike the remaining chloride, the two dual-purpose chlorides were considered firmly held and not available for reaction with silver ion to precipitate AgCl. The complex CoCl₃.3NH₃, also represented in Figure 3.2 is not only a nonelectrolyte in Werner's theory, but unable to release chloride ion to form AgCl. Overall, his model interpreted all experimental observations of the time correctly, including being able to explain how CoCl₃-4NH₃ can exist in two forms of what Werner showed were geometric isomers, species with the same formula but different spatial dispositions of groups around the cobalt. Only for the octahedral shape are two isomers predicted for CoCl3-4NH3, matching the experimental observations. Using either a flat hexagonal shape or else a trigonal prismatic shape, three isomers are predicted, supporting (but not proving) the octahedral shape for six-coordination. The proof came from Werner's success in resolving some six-coordinate complexes into optical isomers - something that could occur if they exist in the octahedral shape but not some other proposed shapes.

We now know conclusively that Werner was correct, or at least nearly so. Interestingly, although his octahedral shape is supremely dominant in six-coordination, we now know that there are trigonal prismatic six-coordinate shapes also - albeit with ligand systems to which Werner had no access. What sophisticated modern physical methods tell us conclusively is that coordination complexes don't just happen upon a shape. Each different coordination number known, and these vary from 2 to 14 supports a limited number of basic shapes. The shape, or stereochemistry, depends both on the metal and its oxidation state and on the form of the ligand - each contributes in different ways.

In 1940, Sidgwick and Powell proposed a simple approach to inorganic stereochemistry based on the concept that the broad features of molecular shape could be related to the distribution achieved by all bonding and nonbonding electron pairs (bond pairs and lone pairs) around a central atom as a result of minimized repulsion. A basic but very useful model for shape evolving out of this is the valence shell electron pair repulsion (VSEPR) model. This, an extension of Lewis' ideas, is based on a simple concept - it considers the electron pairs as point charges placed on a spherical surface with the central atom in the centre of the sphere. These charges of identical type will distribute themselves on the surface so as to minimize repulsive interactions. In other words, this is a simple electrostatic model. For any set of like charges, there is usually either only one lowest energy arrangement or a very limited set of arrangements of the same or at least very similar energy. This very simple model has proven to be of great value as a predictor of polyatomic molecular shape for p-block compounds, since in these any lone pairs present play an important structural and directional role in defining molecular shape. The concept remains valid decades after its inception. Although developed for main group (p-block) chemistry, the concepts can be adapted to some extent for metal complexes, and was first applied by Nyholm and Gillespie in the late 1950s to coordination complexes, although the strong role of lone pairs found in p-block chemistry is absent in the d block.

For d-block metal ions in particular, it was known that a series of structurally common [M(OH₂)6¹²⁺ complexes form for metal ions with different numbers of d electrons, although from a VSEPR viewpoint one would expect structures to vary since electronic configurations differ- the nonbonding electrons are clearly not playing a major role. Kepert amended the VSEPR model for use with transition metals by ignoring nonbonding electrons, and considering only the set of ligand donor groups, treating them as point charges on a spherical surface as in the VSEPR model and considering repulsions between them. For two to six point charges, the outcomes are represented in Figure 3.3; only for five-coordination are two geometries of essential equal energy predicted. These dispositions of point charge on the surface then can be taken to represent the location of the donor atoms in the attached ligands  joined by coordinate bonds to the metal that is placed at the body centre of the sphere. Once bonds are inserted, this then represents the basic shape of the complex (Figure 3.3). The solid lines shown in the right half of Figure 3.3 represent the actual coordinate covalent bonds; the dashed lines in the left half of the figure only define the shape of the framework and do not relate to bonding at all.

Figure 3.3

Distributions predicted for from two to six point charges on a spherical surface (left), and shapes of complexes evolving from this point-charge model, when a central atom is placed at the core of the sphere and considered to bond to each donor atom located where a point charge occurs.

As with any simple model, it is not universally correct, and for each coordination number, there may be alternate shapes. For example, the predicted shape for four-coordination is tetrahedral, but we know that one well-known shape found for four-coordination in metal complexes is square planar, a flat geometry where all four ligands lie in the same plane as the metal, disposed in a square arrangement with the metal ion at the centre. Obviously, there are other influences on the shape or stereochemistry of complexes. We shall look at real outcomes for each coordination number in turn in Chapter 4, but at this stage it is sufficient to recognize that there are several basic shapes, most of those being defined perfectly well by the simple model depicted in Figure 3.3 above.

It is also appropriate at this stage to recognize that the way we represent or draw molecules for display and discussion is important for communication of concepts. Technology has provided a number of ways in which molecules can be illustrated (Figure 3.4). For everyday use and discussion between chemists, it is most likely that the simple basic drawing will be used, as it can be hand or computer sketched rapidly. Views with perspective, or else ball and stick models, tend to be met in formal presentations, as will be the case in this book.

Figure 3.4

Various ways in which metal complexes can be represented, illustrated for the simple octahedral complex ion [IrCl₆]₂.

Space-filling models, while giving a view closer to the actual situation for the molecular assembly are difficult to visualize even for very simple molecules because atoms at the front tend to obscure those behind. As a consequence, the ball and stick models are met more often in formal presentations.

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مواضيع ذات صلة

A Covalent Bonding Model - Embracing Molecular Orbital Theory

An Ionic Bonding Model - Introducing Crystal Field Theory

Holding On - The Nature of Bonding in Metal Complexes

The Foundation of Coordination Chemistry

Photo-Oxidative Degradations of Polymers

Oxidation of Step-Growth Polymers

Hydrolytic Degradation of Polymers at Elevated Temperatures

The Termination Reaction

صناعة حمض الكبريتيك

تركيب حامض الكبريتوز

كلوريد الثيونيل Thionyl Chloride

حامض الكبريتوز: Sulphureous Acid