Complexes heptacoordinated

In addition, a chlorine ion most probably replaces the fluoride ion located on the rotation diad axis of the heptacoordinated complex with C2v symmetry otherwise splitting of the band would be observed. Replacement of the ligand in any other position would reduce the symmetry from C2v to Cj. 

The most common property of molten systems containing tantalum or niobium is the ionic equilibrium between hexa- and heptacoordinated complexes, with the general compositions MeFg and MeF6X(n+I) ... 

The third factor that strongly affects the equilibrium between hexa- and heptacoordinated complexes (85) is the nature of the second, outer-sphere cations. Increasing the ionic radii of the cations causes the equilibrium in Equation (85) to shift to the left, forming mostly hexacoordinated complexes MeF6 The mechanism of interionic equilibrium in fluoride melts can be presented schematically as follows ... 

The anions MeF6 and X approach each other closely to form the heptacoordinated complex MeF6X(n+1)", or separate from one another, according to the polarization potential of the outer-sphere cation (alkali metal cation -M+). This process is unique in that the mode frequencies of the complexes remain practically unchanged despite varying conditions. This particular stability of the complexes is due to the high charge density of Ta5+ and Nbs+. 

The scheme of the interaction mechanism (Equation 88) testifies to an electro-affinity of MeFe" ions. In addition, MeFe" ions have a lower negative charge, smaller size and higher mobility compared to MeF6X(n+1> ions. The above arguments lead to the assumption that the reduction to metal form of niobium or tantalum from melts, both by electrolysis [368] and by alkali metals, most probably occurs due to interaction with MeF6 ions. The kinetics of the reduction processes are defined by flowing equilibriums between hexa-and heptacoordinated complexes. 

It should be noted that in addition to changes in K2NbF7 or K2TaF7 concentrations that afford control over the complex structure and electrolysis parameters, the cation type also affects the equilibrium between the complex ions. The heptacoordinated complexes become increasingly dominant when progressing along the cation series from Li to Cs. 

Higher temperatures, increased K2TaF7 concentrations and other factors mentioned above shift the equilibrium in Equation (174) to the left and lead to the formation of coarser tantalum particles. Form this point of view, it can be concluded that smaller hexacoordinated complexes, TaF6 lead to the formation of coarser tantalum powder, whereas the predominant presence of larger heptacoordinated complexes ions initiates the formation of finer particles. 

Going back to mechanistic studies it is not clear if the reactions of nucleophiles with hexacoordinated silicon compounds are pure nucleophilic substitutions or if they take a more complex route. However there is another challenge to find whether the silicon atom can accept being in heptacoordination. Such a possible situation has been observed with a tricapped tetrahedron structure of a silane which has been proved to be isosteric with the corresponding germane of which the X-ray structure determination has been carried out. 

The quite reactive Re(III) heptacoordinated complex, [Re (terpy)2(H20)]3+, denoted by ReCR20)TER, is also included. 

The same happens for the last haemerythrin model complex [Fe2(Htpp-do)(C6H5C02)]3 + (Htppdo = V,V,A A+tetrakis(6-pivalamido-2-pyridyl-methyl)-l,3-diaminopropa-2-ol), in which one Fe(II) atom is heptacoordinate, and the other Fe(II) atom is hexacoordinate. At low temperature (—50°C), in acetone solution this complex reversibly reacts with dioxygen.37 The only electrochemical information is that, in MeCN solution, the original trication exhibits two one-electron oxidations ( Fe(II)Fe(II)/Fe(II)Fe(III) = +0.29 V VS. NHE Fe(II)Fe(III)/Fe(III)Fe(III) =... 

Fig. 12. Bond distances and angles in the heptacoordinated dimeric form of Pr DPM) 3 complex...

Fig. 14. The heptacoordinated polyhedron of Yb(acac)3 acimin) complex. The coordinated acimin oxygen (0(7)) is shaded...

In the above structures, the uncoordinated chloride ion is surrounded by a polyhedron of six water molecules forming a distorted octahedron with average Cl—H2O distances of 3.18 and 3.21 A for Eu(III) and Gd(III) complexes respectively. The coordinated chloride ions are, however, heptacoordinated being bonded to six water molecules and to a lanthanide ion. The average Cl—H2O distances for this chlorine is somewhat larger than the uncoordinated case and are 3.23 and 3.16 A for Eu(III) and Gd(III) complexes respectively. 

It is interesting to see that aU three complexes have eight essentially equivalent M—0 distances. The bite in the case of [Y(HFA)4] is 2.772 A and compares well with that for square antiprismatic [Y(acac)3 (0112)2] (2.80 A) and the heptacoordinated [Y(bzac)s (OH2)] (2.76 A) complexes (64, 116). The chelate rings in the HFA-complexes also show familiar folding of 8° along the 0—0 line of the chelate rings. 

The first reported synthesis of a heptacoordinate complex of silicon (197) was accomplished by the reaction of IlSiCh with 2-lithio-(dimethylaminomethyl)-benzene227. The crystal structure was determined later and confirmed the tricapped tetrahedron geometry81,82. Replacement of the hydrido ligand by chloro (198) led to the formation of a tetracoordinate silane without any dative bonds, apparently due to the steric bulk associated with the chloro ligand81. 

FIGURE 28. Crystallographic structure of the heptacoordinate complex 199. Reprinted with permission from Reference 228. Copyright 1994 American Chemical Society. 

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