Orbital coordination complexes

Figure 6.21 Relation between C2B9Hn and C sHs". In this formalism the c/o.ro-carboranes C2B10H12 are considered as a coordination complex between the pentahapto 6-electron donor C2B9H] - and the acceptor BH (which has 3 vacant orbitals). The o/oso-structure can be regained by capping the open pentagonal face with an equivalent metal acceptor that has 3 vacant orbitals.

The effect of inner orbital splitting on the thermodynamic properties of transition metal compounds and coordination complexes. P. George and D. S. McClure, Prog. Inorg. Chem., 1959,1, 381-463 (36). 

This procedure is strictly invalid, of course, since the symmetry of a six-coordinate complex with dissimilar ligands cannot be exactly octahedral. In this case, further splitting of the d orbitals takes place which is not representable by a single splitting parameter like 4oct-However, if the departure from Oh symmetry is slight, so that spectral bands are broadened rather than split, the law of average environments retains utility. 

The crystal field energy level diagram for octahedral coordination complexes. The energies of the d orbitals differ because of differing amounts of electron-electron repulsion. The... 

Color is a spectacular property of coordination complexes. For example, the hexaaqua cations of 3 transition metals display colors ranging from orange through violet (see photo at right). The origin of these colors lies in the d orbital energy differences and can be understood using crystal field theory. 

C20-0096. Use orbital sketches to explain why the d, and j2 y2 orbitals have different stabilities in octahedral coordination complexes. 

Four-coordinate complexes exhibit lower isomer shifts than six-coordinate compounds. Metal-ligand bonds are shorter and more covalent if the coordination number is smaller because of less steric hindrance and less overlap with antibonding 2g orbitals in the case of four as compared to six bonds. 

In their pursuit of modeling Type I copper proteins, Kitajima et al. reported112 a rare, tetrahedrally coordinated complex (105), which displayed an EPR spectrum consistent with the presence of the unpaired electron in the dz2 orbital.1 They also isolated a square-pyramidal DMF adduct (complex (106)). They were successful in providing structural proof of a copper(II) complex (trigonal pyramidal) with C6F5S -coordinated complex (107), with CuN3S chromo-phore.113 The X-ray analysis (poor data set) of a closely similar complex with Ph3CS as the... 

Pauling further extended the sp"dm hybridization approach to the d-block compounds.3 By varying the relative importance of p and d orbitals, Pauling was able to construct hybrid orbitals that rationalized the geometries and magnetic properties of many transition-metal coordination complexes. For example, the square-planar... 

The specified spin multiplicity in (4.76) is in each case that expected from Hund s rule, with electrons maximally unpaired in remaining orbitals of the tungsten valence shell. Further discussion of daughter radical species is beyond the scope of this work, but certain aspects of open-shell coordination complexes will be considered in Section 4.6.4. below. 

From the above considerations, we can recognize that an ideal hypersaturated coordination complex would arise from what may be denoted as a 3cu/3a/3n metal configuration, with three orthogonal cu bonds (3cu one each in the x,y, and z directions) built from three parent sigma bonds (3 a from sd2 hybrids at 90° angles), and with three lone pairs (3n pure d orbitals) in the duodectet of the parent... 

Figure 4.45 A metal-ligand m,—orbital splitting diagram depicting interaction of the metal-atom d NAO and ligand nL NBO to form semi-localized NLMOs of the coordination complex, with splitting energy Aed. = < d/NLMO — fd> (NAO).

The four-coordinate model complex RuHCl(PH3)2 is not planar but has a saw-horse geometry withtrans phosphines and H-Ru-Cl = 101.3°. This angle illustrates that a d6 tetra coordinated complex prefers to be a piece of an octahedron with two empty coordination sites in order to keep the six electrons of the metal in nonbonding orbitals (essentially similar to the t2g set of an octahedron). 

Formation of coordination complexes is typical of transition metals, but other metals also form complexes. The tendency to form complexes is a function of the metal s electron configuration and the nature of its outer electron orbitals. Metal cations can be classified into types A and B based on their coordination characteristics, as shown in Table 3.5. A-type cations, which tend to be from the left side of the Periodic Table, have the inert-gas type electron configuration with largely empty d-orbitals. They can be imagined as having electron sheaths not easily deformed under the influence of the electric fields around neighbouring ions. B-type cations have a more readily deformable electron sheath. 

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