Hexaaqua complexes

Table 4. Water Exchange Rates and Activation Parameters of Hexaaqua Complexes at 25°C, ...

Iron(II) formate dihydrate, 14 537 Iron(II) fumarate, 14 537 Iron gelbs, 19 399, 400 Irondl) gluconate dihydrate, 14 541 Iron group carbides, 4 690-692 Iron halides, 14 537-540 Iron hydroxide, water exchange rates and activation parameters of hexaaqua complexes, 7 589t Iron(II) hydroxide, 14 542 Iron(III) hydroxide, 14 542 Iron hydroxides, 14 541—542 Iron(II) iodide, 14 540 Iron(III) iodide, 14 540 Iron/iron alloy plating, 9 813—814. See also Fe entries... 

Deviations from the Point-Dipole in Hexaaqua Complexes of Divalent 3d Metal... 

In aqueous solutions, in which the most probable ligand is the water molecule, most of the lower oxid ation states (i.e. + 2, + 3 and some of the + 4 states) of transition metal ions are best regarded as hexaaqua complex ions, e.g. [Feu(H20)6]2 +. In these ions the six coordinated water molecules are those that constitute the first hydration sphere, and it is normally accepted that such ions would have a secondary hydration sphere of water molecules that would be electrostatically attracted to the positive central ion. The following discussion includes only the aqua cations that do not, at pH = 0, undergo hydrolysis. For example, the iron(III) ion is considered quite correctly as [Fe(H20)6]3 +, but at pH values higher than 1.8 the ion participates in several hydrolysis reactions, which lead to the formation of polymers and the eventual precipitation of the iron(III) as an insoluble compound as the pH value increases, e.g. ... 

As noted earlier, metal ions in polar solvents will form complexes with the solvent molecules. X-Ray diffraction, EXAFS, and visible absorption spectroscopy show that nickel(II) ion in dilute aqueous solution is present as the green hexaaqua complex Ni(H20)62+, just as in solids such as NiS04-7H20, which is actually [Ni(H20)e]S04-H20. In the crystal, the extra water molecule is loosely associated with the sulfate ion independently of the nickel-aqua complex it is sometimes referred to as lattice water, as distinct from complexed water. 

The pATa of aqueous iron(III) ion in very dilute solution at 25 °C is 2.17. On the basis of this information alone, estimate (a) the minimal H+ concentration required to ensure that 99% of the iron(III) is present as the hexaaqua complex, and (b) the [H+] at which you would expect Fe(H20)63+ and Fe(H20)s0H2+ to be present in equal concentrations, (c) What factors (polymerization would be one) could interfere with the accuracy of the predictions, and how would you attempt to minimize their effects ... 

Somewhat better data are available for the enthalpies of hydration of transition metal ions. Although this enthalpy is measured at (or more property, extrapolated to) infinite dilution, only six water molecules enter the coordination sphere of the metal ion lo form an octahedral aqua complex. The enthalpy of hydration is thus closely related to the enthalpy of formation of the hexaaqua complex. If the values of for the +2 and +3 ions of the first transition elements (except Sc2, which is unstable) are plotted as a function of atomic number, curves much like those in Fig. 11.14 are obtained. If one subtracts the predicted CFSE from the experimental enthalpies, the resulting points lie very nearly on a straight line from Ca2 lo Zn2 and from Sc to Fe3 (the +3 oxidation state is unstable in water for Ihe remainder of the first transition series). Many thermodynamic data for coordination compounds follow this pattern of a douUe-hunped curve, which can be accounted for by variations in CFSE with d orbital configuration. 

As anticipated in Sections 2.2.2 and 3.1, the unpaired electrons should not be considered as point-dipoles centered on the metal ion. They are at the least delocalized over the atomic orbitals of the metal ion itself. The effect of the deviation from the point-dipole approximation under these conditions is estimated to be negligible for nuclei already 3-4 A away [31]. Electron delocalization onto the ligands, however, may heavily affect the overall relaxation phenomena. In this case the experimental Rm may be higher than expected, and the ratios between the Rim values of different nuclei does not follow the sixth power of the ratios between metal to nucleus distances. In the case of hexaaqua metal complexes the point-dipole approximation provides shorter distances than observed in the solid state (Table 3.2) for both H and 170. This implies spin density delocalization on the oxygen atom. Ab initio calculations of R m have been performed for both H and 170 nuclei in a series of hexaaqua complexes (Table 3.2). The calculated metal nucleus distances in the assumption of a purely metal-centered dipolar relaxation mechanism are sizably smaller than the crystallographic values for 170, and the difference dramatically increases from 3d5 to 3d9 metal ions [32]. The differences for protons are quite smaller [32]. 

Crystallographic values of metal-hydrogen and metal-oxygen distances in hexaaqua complexes of divalent 3d metal ions compared with calculated effective distances... 

The hexaaqua complex can be converted into the tetrachloro complex by addition of chloride ion through the reaction... 

FIGURE 14.2 Sketch of the change with time of the concentrations of products and reactants in the spontaneous reactions illustrated in Figure 14.1. For ease of display, concentrations are expressed as percent of the total Co(ll) present in each species, (a) Partial conversion of pink hexaaqua complex A into blue tetrachloro complex C. (b) Partial conversion of blue tetrachloro complex C into pink hexaaqua complex A. After changes in the slope of each species concentration become imperceptibly small, we say the reaction has arrived at chemical equilibrium. 

The questions raised in the first paragraph require quantitative investigations of the reaction mixture, which we carry out as follows. In the preceding reaction let A represent the pink hexaaqua complex, B the chloride ion, and C the blue tetrachloro complex. In the first experiment, we start the reaction by mixing initial concentrations of A and B, denoted as [A]o and [B]q. As the reaction proceeds, we periodically sample the reaction mixture. For each sample, we measure the concentration of A, B, and C and plot concentration of each species versus time. The results of the first experiment are represented schematically in Figure 14.2a, which shows the consumption of A and the production of C. Similarly, we start the second experiment with the initial concentration [C]o, and add water. The results are represented schematically in Figure 14.2b, which shows the consumption of C and the production of A. 

It may appear surprising to find such imbalance in the concentration of the Co(II) species when the color intensities of the pink and lavender solutions in Figure 14.1a and Figure 14.1c appear quite similar. This difference is explained by the fact that the blue tetrachloro complex absorbs light much more efficiently than the pink hexaaqua complex. 

Volumes of activation for water exchongo in hexaaqua complexes of transition metal ions of the first row"... 

If one adds water, the green solution changes back into the known blue solution, the monochloropentaaqua complex disappears, the hexaaqua complex is formed again a chemical equilibrium exists between both complexes, different stabilities of complexes must be considered and can be calculated using the stability constants. 

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