Reduction potential diagrams

We can summarize the half-reactions of copper in a reduction potential diagram of the form... 

Recall from Example 11.10 that disproportionation is the process in which a single substance is both oxidized and reduced. Reduction potential diagrams enable us to determine which species are stable with respect to disproportionation. 

Use reduction potential diagrams to determine strengths of oxidizing and reducing agents and stability toward disproportionation (Section 17.2, problems 25-26). 

A voltaic cell containing a standard Fe /Fe electrode and a standard Ga /Ga electrode is constructed, and the circuit is closed. Without consulting the table of standard reduction potentials, diagram and completely describe the cell from the following experimental observations, (i) The mass of the gallium electrode decreases, and the gallium ion concentration increases around that elec-... 

These potentials are compiled in reduction potential diagrams, often referred to as Latimer diagrams, in table 5. These are formal potentials measured in or calculated for 1.00 mol dm HC104(aq) or for another acid if no value in HCIO4 has been measured. In a few cases the formal potentials have been converted to standard potentials for incorporation into tables of standard thermodynamic properties using the specific ion interaction theory (Grenthe et al. 1992). Values in basic solution have been calculated using solubility equilibria (Morss 1985, 1986). 

In water, many of the oxyacids and their anions are unstable with respect to disproportionation (self-oxidation and reduction). The tendency for this kind of instability may be conveniently determined from the reduction potential diagrams shown in Fig. 1. A large, positive standard reduction potential (the number over the arrows) indicates a strong tendency for the particular reaction indicated by the arrow (a reduction). A large, negative standard reduction potential indicates a strong tendency for change in the opposite direction (an oxidation). A particular oxyacid or its anion will not be stable with respect to self oxidation and reduction if a reduction... 

Fig. 17.2 Standard (or formal) reduction potential diagrams for the actinide ions in (a) acidic (pH 0) and (b) basic (pH 14) solutions. (Values in volts versus standard hydrogen electro. ) Note that the solubility of PuOz(OH)i increases from 1 m KOH to 10 m KOH solution [57]. Thus there is evidence for amphoteric behavior of PuOi (OH) by forming Pu02(OH)1 in strong base.

The oxoacids of P are dearly very different structurally from those of N (p. 459) and this difference is accentuated when the standard reduction potentials (p. 434) and oxidation-stale diagrams (p. 437) for the two sets of compounds are compared. Some reduction potentials ( /V) in acid solution are in Table 12.8 (p. 513) and these are shown schematically below, together with the corresponding data for alkaline solutions. 

Note that, because the right side of the cell diagram corresponds to reduction, E° = °(for reduction) — E°(for oxidation) where both values of E° are the standard reduction potentials. 

Fig. 7. Schematic diagram showing the oxidation and reduction potentials of conducting polymers relative to oxygen reduction and water oxidation.

Fig. 14. Schematic representation of a mixed potential diagram for a generic electroless deposition reaction. The dashed line represents the current for metal ion reduction in the presence of a stronger complexing agent.

Figure 14 shows a schematic representation of a mixed potential diagram for the electroless deposition reaction. Oxidation of the reductant, in this case hypophos-phite, is considered to be under 100% kinetic control. A mixed kinetic-diffusion curve is shown for the reduction of the metal ion, in our case Co2+, in the region close to the mixed potential, Em. Thus, since Co deposition occurs under a condition of mixed kinetic and diffusion control, features small relative to the diffusion layer thickness for Co2+ will experience a higher concentration of the metal ion, and hence... 

Figure L Modified Latimer diagram illustrating the relative reduction potentials of a metal complex (M) and its excited state (M ).

Figure 3, Standard reduction potentials associated with the chemistry of oxygen values in upper and lower halves of diagram refer to pH 0,0 and pH 7,0 conditions,...

The redox potential diagram in eq. 1 illustrates that the effect of optical excitation is to create an excited state which has enhanced properties both as an oxidant and reductant, compared to the ground state. The results of a number of experiments have illustrated that it is possible for the excited state to undergo either oxidative or reductive electron transfer quenching (2). An example of oxidative electron transfer quenching is shown in eq. 2 where the oxidant is the alkyl pyridinium ion, paraquat (3). 

In order to add CIO2 to the Latimer diagram drawn above, we must calculate the voltages denoted by ( ) and ( ). The equation associated with the reduction potential ( ) is... 

Table D-4 contains the following data Cr3+/Cr2+ reduction potential = -0.424 V, Cr2072 /Cr3+ reduction potential = 1.33 V and Cr2+/Cr reduction potential = -0.90 V. By using the additive nature of free energies and the fact that AG° = -nFE°, we can determine the two unknown potentials and complete the diagram. 

Figure 4. Band diagram showing the relative positions of the conduction and valence bands in WSe2 with respect to the reduction potentials in aqueous solutions for the redox couples shown in Figure 3.

A "potential ladder" diagram models the potential difference. The rungs on the ladder correspond to the values of the reduction potentials. For a galvanic cell, the half-reaction at the cathode is always on the upper rung, and the subtraction... 

In accordance with the general features of the chemistry of first-, second-, and third-row transition elements, the reduction potentials for the tetraoxometallates follow the sequences [MnOJ 3> [Tc04] > [ReOJ. The comparatively facile reduction of Tc , as evident from the Latimer diagram, limits the range of chemistry accessible for this oxidation state. 

The reader should recall that the concentrations of the electron, pure solids, and the solvent (water), are defined as 1. The calculated value of —2.77 v matches the value of —2.8 v which was estimated from the diagram. It is interesting to note from the Nernst equation that the reduction potential for the half-reaction is dependent only upon the concentration of the sodium ion, Na" ". Neither the concentration of the hydrogen ion nor the hydroxide ion influences the potential at which the half-reaction occurs since they do not appear in the above equation. Similar calculations may be made for other concentrations of Na" ". It will be found that the horizontal line separating Na" " and Na moves from —2.71 v at 1.00 M Na+ to —2.89 v at 10 M, to —3.06 V at 10 M, to —3.24 v at 10 M, and so on. 

Draw a cell diagram for each of the following reactions, and calculate the cell potentials by using standard reduction potentials. 

Latimer diagrams were invented by W. M. Latimer and consist of lines of text of the various oxidation states of an element arranged in descending order from left to right, with the appropriate standard reduction potentials (in volts) placed between each pair of states. The diagram for chromium in acid solution is written as ... 

For conciseness in the remainder of the chapter, the Latimer diagrams are presented with the relevant reduction potentials between the two oxidation states of the couple, as shown below, or in tabular form. 

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