Copper , free transfer energy

If we were to place a piece of zinc metal into an aqueous copper(II) sulfate solution, we would see a layer of metallic copper begin to deposit on the surface of the zinc (see Fig. K.5). If we could watch the reaction at the atomic level, we would see that, as the reaction takes place, electrons are transferred from the Zn atoms to adjacent Cu2 r ions in the solution. These electrons reduce the Cu2+ ions to Cu atoms, which stick to the surface of the zinc or form a finely divided solid deposit in the beaker. The piece of zinc slowly disappears as its atoms give up electrons and form colorless Zn2+ ions that drift off into the solution. The Gibbs free energy of the system decreases as electrons are transferred and the reaction approaches equilibrium. However, although energy is released as heat, no electrical work is done. 

The reactivities of potassium and silver with water represent extremes in the spontaneity of electron-transfer reactions. The redox reaction between two other metals illustrates less drastic differences in reactivity. Figure 19-5 shows the reaction that occurs between zinc metal and an aqueous solution of copper(II) sulfate zinc slowly dissolves, and copper metal precipitates. This spontaneous reaction has a negative standard free energy change, as does the reaction of potassium with water ... 

Electron transfer rate constants, kt, free energy changes, - AG°, and stability constants, K1 and Kz, for the reactions of Cr(III)(phen)3 and Ru(II)(bpy)3 with reduced blue copper proteins at 295 K [96]... 

The cuprous-cupric electron transfer reaction is believed to be the rate-limiting step in the process of stress corrosion cracking in some engineering environments [60], Experimental studies of the temperature dependence of this rate at a copper electrode were carried out at Argonne. Two remarkable conclusions arise from the study reviewed here [69] (1) Unlike our previous study of the ferrous-ferric reaction [44], we find the cuprous-cupric electron transfer reaction to be adiabatic, and (2) the free energy barrier to the cuprous cupric reaction is dominated in our interpretation by the energy required to approach the electrode and not, as in the ferrous-ferric case, by solvent rearrangement. 

The rate constants and the associated free-energy snrfaces available to the peroxide and native intermediates deserve comment since they differ overall by nearly 10 (or ca. Vkcalmol in absolnte valne). Given the relatively electronentral nature of electron transfers between the copper sites (the E° values for the three sites differ overall by only 60 mV), the differences in rate in the first instance reflect the difference in the E° value for le versus 2e reduction of dioxygen (leading to the peroxy intermediate) and peroxide (leading to the native intermediate). Second, the differences reflect the work available from the favorable 4e reduction that drives the turnover from native intermediate to fully reduced enzyme primed, now, to react with O2. This latter process, k 100 s (compare to k = 0.34s for decay of the native intermediate to fully oxidized enzyme), is functionally equivalent to the reductive release of Fe + from Fe +-transferrin catalyzed by the membrane metalloreductase, Dcytb in both cases, the lower valent metal species is more loosely coordinating. Whereas Fe + dissociates in the latter case, in MCO turnover the bound water(s) dissociate. 

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