Holding On - The Nature of Bonding in Metal Complexes

Having established the basic concepts of coordination complexes it is now time to attempt to understand how these complexes hold together or bond. To pursue this aspect, we need to develop models for bonding that not only provide a satisfactory basis for dealing with the array of shapes that exist but also can provide interpretation of the spectroscopic and other physical properties of this class of compounds. It is useful to introduce the core concepts and models that we use to interpret observations immediately as they pervade discussion throughout the field.

The metals of the d block characteristically exist in stable oxidation states for which the nd subshell has only partial occupancy by electrons. This differs from the situation with main group elements, for which it is common or at least more usual to have stable oxidation states associated with subshells that are either completely empty or filled. Importantly the chemical and physical properties characteristic of the transition elements are determined by the partly filled nd subshells (likewise in the f block, but to a lesser extent, by partly- filled nf subshells). In the simple atomic model of the first-row d-block elements, however, the set of closely-spaced levels involving 4s, 4p and 3d orbitals can be considered as the valence orbitals. This provides the luxury of nine orbitals (one s three p and five d) giving rise to what is called the nine-orbital (or 18-electron) rule that attempts to explain metal-donor coordination numbers of up to nine. The valence bond model approach to describing complexes is limited, but worth a short review. Valence bond theory, which describes bonding in terms of hybrid orbitals and electron pairs, evolved from the 1927 Heitler-London model for the covalent bond, evolving from the original 1902 concept of covalency proposed by Lewis, and augmented at a later date to include the hybridization concept by Pauling; it is obviously an early model.

Figure 3.5

A simple valence bond description of bonding for the [Co(NH3)6]+ complex ion.

Consider cobalt as an example, forming the cobalt(III) cation and subsequently an octahedral [Co(NH3)6]3+ complex ion where all Co-N bonds are identical. We commence with our basic model of the coordinate covalent bond, which requires the ligand donor group to supply a lone pair of electrons to an empty orbital on the metal. We can just about deal with this for the cobalt(III) cation, through a somewhat complicated process of ion formation, electron rearrangement, orbital hybridization and filling of empty hybrid orbitals, as depicted in Figure 3.5.

The model relies on energy expended in hybridization and electron relocation being recovered through the set of bond formation processes that occur. It is an adequate model, but rather limited. It can't accommodate, for example, six-coordination in complexes with more d electrons even though there are many examples known experimentally without recourse to employing the empty but higher energy 4d levels. The two systems, 3d-4s-4p and 4s-4p-4d, were introduced in the early 1950s by Nobel laureate Henry Taube and are termed inner shell and outer shell respectively to distinguish them and reflect reactivity differences for species fitted to the different models, but the need to move to the latter higher energy levels in the model is somewhat unsatisfactory. The concept is illustrated in Figure 3.6 for Co (III) in two different d-electron arrangements, as met in [Co(NH3)6]3+ and [COF6]. The model allows an understanding of magnetic properties relating to the number of unpaired electrons and aspects of reactivity, but deals very modestly with stereochemistry and spectroscopy. When a model meets with difficulty in explaining experimental observa- tions it becomes time to set it aside and look for alternative, perhaps more sophisticated, models. After all, the only value of a model is that it explains observations - otherwise it's about as useful as a nose on a pumpkin.

The higher stability of inner shell complexes that employ just the nine 3d-4s-4p orbital set suggests that there may be some special stability associated with the systems that employ nine orbitals that can accommodate no more than 18 electrons. There arose the 18-electron rule, which suggests that coordination complexes whose total number of valence

Figure 3.6

The extended valence bond description of inner shell and outer shell bonding for octahedral cobalt(III) complexes, required as a result of different d-electron arrangements on the metal. The six empty hybridized orbitals can in each case accommodate six bonding lone pairs from six ligand donor atoms.

electrons approaches or equals 18 (but does not exceed it) are more stable assemblies, with those actually achieving 18-electron sets being most stable. This rule works best for organometallic compounds, which are more covalent in character, whereas Werner-type compounds, with more ionic character, tend to break the rule more often, and consequently it is much less used for these. It may be helpful to illustrate it in operation, nevertheless. We shall do this with the simple organometallic compound, octahedral [Mn(CO), I]. There are two methods of electron counting: the closed shell (or oxidation state) method and neutral ligand (or formal charge method). Both exist because accounting for electrons is not always straightforward. Employing the former method as an illustration, [Mn(CO);I] may be considered as composed of Mn+, five CO and one I ligands. Each ligand contributes two electrons from a lone pair to the complex, and the metal contributes its valence electrons. This amounts to 6 electrons (for Mn+), 10 electrons (for 5 x (CO)) and 2 electrons (for 1 x I), a total of 18 electrons. This suggests that [Mn(CO)51] should be stable, which is the case. This 'rule' more correctly acts as a guideline for stability, but we shall not dwell on it here, given our focus on Werner-type rather than organometallic-type complexes.

Here, it is perhaps valuable to remind ourselves of the nature of the d orbitals that we are employing. The five d orbitals (dy. dyz. dz. d2-2 and. d) occupy different spatial orientations and involve two different basic shapes. These orbitals are shown in Figure 3.7. For elements of the f block, there are seven f orbitals (fayz. fz-y2) fy(22x2) fx(22-12) fy f and f) which offer three different basic shapes and also, like the d orbitals, occupy quite different spatial positions. While we will not be discussing f orbitals in any detail it is useful to be reminded that they do many of the things that d orbitals do except that the 4f orbitals of the lanthanoids for example, are less 'exposed' due to being screened by larger 5s, 5p and 6s shells which leads to weaker interactions with their surrounding ligands and very similar chemistry across the series. The nd orbitals are more exposed and hence more

Figure 3.7

The five d orbitals of d-block elements, each represented in two different views.

involved in their element's chemistry. It is the set of five d orbitals that will be of particular interest in this book. To start to understand how they may differ, notice that two of the five d orbitals have lobes lying along axes of the defined coordinate system, whereas for the other three the lobes lie between axes. It is these two classes of spatial arrangement that are the key to much of the discussion that will follow in coming sections suffice to say that this spatial difference will be significant.

Another way of viewing complex formation is to invoke a purely ionic model. In this, we recognize that the metal at the centre of coordination complexes is usually a cation, and that ligands are either anionic or else have regions of high electron density on the donors that make them attractive to the cation. Thus, we can conceive of a situation where a set of donors is arranged around a metal ion centre in an array defined mostly by their electrostatic attraction to the metal ion and electrostatic repulsion towards other ligands. While very limited in what it says about the metal-donor interaction, this ionic model does offer a useful and surprisingly successful way of predicting and understanding geometric shape in complexes, as we have seen in Figure 3.3 and we shall return to this in detail in Chapter 4. It is at first sight an eccentricity of metal coordination compounds that we can invoke either a covalent bonding model or an ionic bonding model. However, if we think of ionic and covalent bonds as the two extremes of a bonding continuum, the apparent contradiction becomes more acceptable. It's almost a state of mind; we would rarely consider CH as anything but covalently bonded but at the same time recognize that representation as CH3 and H at least provides a way of understanding some nucleophilic and electrophilic organic reactions. We are likely to be reasonably comfortable in considering a neutral compound such as [Mo(CO)6] composed formally of uncharged Mo(0) bonded to the carbon atoms of neutral CO molecules, as covalently bonded. When presented with a highly ionic assembly, such an [Mn(OH)6] (stable only in very strongly basic solution), we can at least accept there is some wisdom in thinking of this as a Mn+ ion with six hydroxide anions attached by strong ionic bonds. What is common to both approaches to bonding in coordination chemistry is that the bonding models tend to be holistic in nature rather than focused on  individual atom-to-atom bonding; this is probably a better description of systems where a set of entities are assembled to form a much larger whole. Two theories have developed to explain the properties of metal complexes. For the former example, the molecular orbital theory represents an appropriate treatment of holistic covalent bonding, whereas for the latter example, the crystal field theory (CFT) is a useful holistic ionic bonding model. Both models have obvious limitations - not least of all an assumption of 'purity' in the bonding character that is unlikely to be so. But then, these are models, after all; complexes don't need either theory to go about their business. Our models are developed not only to allow some level of understanding of physical and chemical properties of complexes, but also to provide a means of prediction of properties where changes or new species are involved. A model is only useful if it performs the tasks we ask of it.

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المزيد

مواضيع ذات صلة

A Covalent Bonding Model - Embracing Molecular Orbital Theory

An Ionic Bonding Model - Introducing Crystal Field Theory

Complex Shape - Not Just Any Which Way

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