![]() That means that we would have the six valence orbitals that we would need to explain the six bonds. In addition, it would also be justified to consider the three 4p orbitals as valence orbitals because the 4p orbitals are energetically only slightly higher than the 4s orbital. How many empty valence orbitals remain? These would be two 3d and the 4s orbitals. The three remaining 3d electrons are expected to be spin up in three different d orbitals according to Hund’s rule. For chromium this means that we must remove the one 4s electron, and two of the five 3d-electrons. When a transition metal loses electrons to form a cation, it always loses its two valence electrons first, and then its d electrons. A neutral Cr atom has the electron configuration 4s 13d 5. Next, we need to know the electron configuration of the Cr 3 +. ![]() Therefore, the chromium is a Cr 3 + cation. It is +3 because the ligands are all neutral when the bonds are cleaved heteroleptically, and the complex cation has a 3+ charge. Does chromium have six empty valence orbitals? In order to assess this, we first need to know the oxidation state of the chromium. We can see that the six ammine ligands have one electron lone pair each that can serve as the valence orbitals. Can valence bond theory explain the six bonds and the octahedral shape satisfactorily? In order to explain the six dative Cr-N bonds we would need to overlap six empty chromium valence orbitals with six filled valence orbitals of N. From experiment we know that it has an octahedral shape, with six dative Cr-N bonds. The first example is the hexaammine chromium (3+) cation (Fig. Figure 7.1.2 Valence bond theory applied to the hexaammine chromium (3+) cation Let us have a closer look at the valence bond theory, and assess valence bond theory for complexes by a number of examples. Overall, valence bond theory is far more suitable for main group element molecules compared to transition metal complexes. By its nature, valence bond theory cannot explain optical properties. We will also see that valence bond theory can explain magnetism in some cases, but also here the valence bond theory has significant deficits. We will see that this concept can explain the shapes of coordination compounds in some cases, but overall it does not work very well. To adapt valence bond theory to suit coordination compounds, Pauling suggested that a dative bond is formed via the overlap of a full valence orbital of the donor and an empty valence orbital of the acceptor. This is not consistent with the dative bonding in coordination compounds where it is assumed that one partner donates an electron pair and the other partner accepts it. The valence bond concept in its original form assumes that each bonding partner contributes one electron to the covalent bond. ![]() The concept works very well to explain the shapes of molecules of main group elements. These orbitals can either be atomic orbitals, or hybridized atomic orbitals. The basic idea is to overlap half-filled valence orbitals to form covalent bonds in which the two electrons are shared between the bonding partners (Fig. Figure 7.1.1 Electron sharing and valence bond theory The valence bond concept was introduced by Linus Pauling in 1931 to explain covalent bonding in molecules of main group elements. The first one is the valence bond theory. There are essentially three bonding concepts that are used to describe the bonding in coordination compounds. Optical properties of compounds are linked to bonding because they are related to electronic states. It should further be able to explain the stability and reactivity of complexes, as well as the optical properties of complexes. A paramagnetic molecule is attracted by an external magnetic field. A diamagnetic molecule is repelled by an external magnetic field. It is paramagnetic when there are unpaired electrons. Remember, a molecule is diamagnetic when it has no unpaired electrons. In addition, it should be able to explain the magnetism of molecules, in particular dia- and paramagnetism. What does this mean for a bonding theory? What would a good bonding theory for coordination compounds be able to do? It should certainly be able to explain and predict the number of bonds and the shape of a molecule. The more the theory can explain and predict, and the fewer the necessary assumptions, the better the theory. In addition, it should be able to predict experimental observations. ![]() The answer is, that it should be able to make many correct explanations for experimental observations based on a few, sensible, assumptions. Let us first think about, what a good theory should be able to do in general. This chapter is devoted to bonding theories for coordination compounds.
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