Hydrated hydroxo complexes

Chromium(III) hydroxide dissolves in acids to form hydrated chromium(lll) salts in concentrated alkali, hydroxo-complexes [Cr(OH)g] are formed. 

We also found that iridium hydrido(hydroxo) complexes like [ lrH(diphos-phine) 2( x-OH)2( x-Cl)]Cl (43) and the precursor diphosphine complexes 42 can also catalyze the hydration of nitriles. In the presence of catalyhc amounts of these complexes, heating acetonitrile and benzonitrile with excess water at 120°C gave the corresponding amides [47, 50]. 

The dependence-of redox potentials in redox systems that involve metal ions is due to the fact that the complex stability differs at different -pH. values in alkaline media deprotonation of aquo ions is enhanced in accordance with the pK. values of the hydrated ions, and the resulting hydroxo complexes are stronger than the aquo complexes. The differences in stability can be measured by means of the differences in redox potential. Thus the redox potentials of alkali metals (M)... 

Woolley230 has found that a number of hydroxo metal complexes of the macrocycle (64) are active in the hvdration of C02 and acetaldehyde. Electronic spectroscopy231 and X-ray crystallography232 suggest coordination numbers in aqueous solution of five for [CoL]2+ and [ZnL]z+, and six for [NiL]2+ and [CuL]2+. For the [ZnL]2+ system, the p/Ca of the coordinated water molecule is 8.69 at 25 °C and the hydroxo complex is a reasonable catalyst for C02 hydration. The complex [Cu(glycylglycinate)OH] is also active in C02 hydration.233... 

We have Investigated the kinetics of base hydrolysis reactions of the cobalt acldate complexes In aqueous solution and In several mlcroemulslon solutions In which detergent concentrations are at least twice the respective cmc. The results are complicated by the onset of a slow secondary reaction which Is presumably formation of Insoluble, polymeric hydroxo, or hydrated hydroxo compounds. 

If the pressure of H2 is maintained at 1 bar, apphcation of the Nernst equation (equation 7.21) allows us to calculate E over a range of values of [H ]. For neutral water (pH 7),, = -0.41V, and at pH 14, °[qh = i = -0.83V. Whether or not the water (pH 7) or molar aqueous alkali (pH 14) is reduced by a species present in solution depends upon the reduction potential of that species relative to that of the 2H /H2 couple. Bear in mind that we might be considering the reduction of H2O to H2 as a competitive process which could occur in preference to the desired reduction. The potential of —0.83 V for the 2H /H2 electrode in molar alkali is of hmited importance in isolation. Many M /M systems that should reduce water under these conditions are prevented from doing so by the formation of a coating of hydroxide or hydrated oxide. Others, which are less powerfully reducing, bring about reduction because they are modified by complex formation. An example is the formation of [Zn(OH)4] in alkahne solution (equation 7.26). The value of E° = —0.76 V for the Zn jZn half-cell (Table 7.1) applies only to hydrated Zn ions. When they are in the form of the stable hydroxo complex [Zn(OH)4] , °[oh 1 = 1 = -1.20 V (equation 7.27). 

Soluble forms of aluminium present in waters are the simple hydrated Al " " cation, cationic and anionic hydroxo complexes, sulphate complexes of [AISO4] and A1[(S04)2], and following fluoridation it also occurs as fluo-roaJuminates with 1-6 coordinated atoms of fluorine, for example [AlFg]. ... 

Evaluating the distribution diagram of different forms of the aluminium occurrence, at pH < 4 simple hydrated Al cation predominates in the solution, and at pH > 7 anionic hydroxo complexes prevail. The lowest solubility of hydrated aluminium oxide is achieved at about pH 5.5. Solubility rapidly increases in the alkaline region. Low concentrations of soluble aluminium can be maintained within a rather narrow pH scale. Precipitation of aluminium salts appears at pH < 4 and at the initial A1 concentration below 0.01 mol 1 . ... 

We now consider Fe hydrolysis. The hexaaquaflFerric cation[Fe(H20)e] is more acid than hexaaquaferrous cation [Fe(H20)g]. The equilibrium constant of hydrolysis is approximately one order lower than that in phosphoric acid, whereas the equilibrium constant of the hydrolysis of Fe " is approximately one order higher than that in boric add. During the hydrolysis the following essentially mononuclear complexes are produced [FeOH] ", [Fe(OH)2]" , [Fe(OH)3(aq)]° and [Fe(OH)4]. By other reactions a series of polynuclear complexes is formed, for example, [Fe2(OH)2], [Fe3(OH)4] , [Fe4(OH)g] , etc. (for simplicity, the coordinated water molecules are omitted). First, colloid hydroxo complexes are formed and finally there is a precipitate of hydrated ferric oxide which is in fact a mixture of different polynuclear complexes. The distribution of polynudear complexes depends not only on pH, but also on the initial concentration of iron. In diluted solutions of ferric salts a precipitate of hydrated Fe203 is separated only at a higher pH. The equilibrium between particular polynuclear complexes is established only very slowly. 

Fluorescence kinetic measurements were applied to the determination of hydration numbers in lanthanide-polyaminocarboxylate complexes the results are given in table 5 (Brittain et al. 1991, Brittain and Jasinski 1988). Measurements of hydration numbers as a function of the solution pH for Eu " and Tb " complexes showed the presence of three buffer regions the first one at low pH in which the hydration number of the cations is equivalent to that of the free ions the second one in the pH range of 4-8 in which complexation decreases the hydration numbers and the third one at high pH (ca. 9-12) being associated with the formation of ternary hydroxo complexes. 

At pH > 6, the D4 Fj J = 3 ) transitions of the Tb(NTA) spectrum were enhanced in intensity by formation of the ternary hydroxo complex which caused selfassociation and formation of oligomers (Brittain and Jasinski 1988). In all these systems, the total coordination number, i.e. the sum of the number of ligand donor groups and the number of primary water molecules, was 8.8 0.5 for Eu and 8.5 + 0.5 for Tb complexes. Such nonintegral hydration numbers were interpreted as the equilibrium weighted average of mixtures of species with different hydration numbers (Brittain and Jasinsky 1988). 

Hydrated cations can act as Brbnsted acids by releasing protons from the bound water molecules to form hydroxo complexes ... 

Two mechanisms were proposed for the hydrolysis reaction. Since the hydroxide ions were generated elect roly tically, the lanthanide ion may have reacted at the cathode surface with replacement of coordinated water molecules to form Ln(OH) or Ln(OH)2. Polynuclear complex formation could also have been formed by this mechanism. Alternatively, the aquo complexes could have been electrolyzed at the cathode surface to form the hydroxo complexes. The authors (Suzuki et al. 1986) suggested, without explanation, that hydrolysis of the elements with large ionic radii and higher hydration number occurs more likely by the first mechanism. However, the second mechanism would seem more likely for the formation of all hydrolyzed species. 

Stopped-flow methods have also been used to study the rapid reactions of CO2 with [Co(NH8)8(OH)] + ion in H2O and D2O. The solvent isotope effect is 1.0 as expected for a reaction not involving a rate-determining proton transfer. However, an isotope effect of ca. 3 has been reported for the bovine carbonic anhydrase-catalysed hydration of CO2. It is argued that if the zinc-hydroxo-complex in the enzyme is acting as a nucleophile in the hydration of CO2, then an isotope effect of unity would be expected. On the other hand, if the zinc-hydroxo-complex, or any other amino-... 

Prominent among the complexes coordinating via oxygen are the hydrates and the hydroxo complexes. On account of the fundamental importance of these... 

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