Water loss from metal cations

Rate of water loss from metal cations as a function of the ratio of the charge to the radius of the metal ion. Rate constants from Margerum et al. (1978). 2Jx ratios calculated from ionic radii tabulated in the CRC Handbook of Chemistry and Physics. 

In the case of trace metals, adsorption is typically much faster than the time intervals for which it is practically possible to separate the cells. Therefore, in practice, values of kf and kr are most often estimated by assuming that water loss from the hydrated cation is rate-limiting (Eigen-Wilkins mechanism, see Section 4.3.1 above). In some cases, uptake transients can be observed at the start of a short-term uptake experiment or by using pulse-chase experiments for which a metal solution containing a radioactive tracer is replaced by a solution... 

The protons come from the water molecules that hydrate these metal cations in solution (Fig. 10.19). The water molecules act as Lewis bases and share electrons with the metal cations. This partial loss of electrons weakens the O -H bonds and allows one or more hydrogen ions to be lost from the water molecules. Small, highly charged cations exert the greatest pull on the electrons and so form the most acidic solutions. 

The equilibrium constant of this reaction has been measured for a variety of metal cations in aqueous solutions and is known to correlate positively with both ionic potential and Lewis acid softness. As the ionic potential increases, the intensity of the positive coulomb field of the cation increases and repulsion of a solvating water proton becomes more likely. As the Lewis acid softness increases, the covalency of the M—O bond in a solvation complex increases and electron density is withdrawn from the O—H bond, thereby promoting the loss of the proton. 

Finally, the route to metal oxide cations from metal nitrate complexes can involve multistep reactions. For example, under harsh electrospray conditions (raised cone voltages) of aqueous iron nitrate solutions, FeO can be formed [98]. CID reactions of various precursor ions suggest that the likely reactions that lead to this ion involve NO2 loss to form the dihydroxide (Eq. (6.115)), which then undergoes water loss (Eq. (6.116)). 

Experiments and calculations both indicate that electron transfer from potassium to water is spontaneous and rapid, whereas electron transfer from silver to water does not occur. In redox terms, potassium oxidizes easily, but silver resists oxidation. Because oxidation involves the loss of electrons, these differences in reactivity of silver and potassium can be traced to how easily each metal loses electrons to become an aqueous cation. One obvious factor is their first ionization energies, which show that it takes much more energy to remove an electron from silver than from potassium 731 kJ/mol for Ag and 419 kJ/mol for K. The other alkali metals with low first ionization energies, Na, Rb, Cs, and Fr, all react violently with water. 

In proposed mechanism I, the loss of water from the complex is the rate-determining step, but removal of water from the coordination sphere of the metal ion should be independent of the nature of the anion that is not part of the coordination sphere of the metal ion. On the other hand, if mechanism II is correct, the entry of X into the coordination sphere of the metal would be dependent on the nature of the anion, because different anions would be expected to enter the coordination sphere at different rates. Because there is an observed anion effect, it was concluded that the anation reaction must be an Sn2 process. However, it is not clear how a process can be "second-order" when both the complex cation and the anion are parts of the same formula. As discussed in Chapter 8, it is not always appropriate to try to model reactions in solids by the same kinetic schemes that apply to reactions in solutions. 

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