Electrodes, oxidation-reduction calculation

To compensate partially for activity effects and errors resulting from side reactions, such as those described in the previous section. Swift proposed substituting a quantity called the formal potential in place of the standard electrode potential in oxidation-reduction calculations. The formal potential, sometimes referred to as the conditional potential, of a system is the potential of the half-cell with respect to the SHE when the concentrations of reactants and products are 1 M and the concentrations of any other constituents of the solution are carefully specified. Thus, for example, the formal potential tor the reduction of iron(III) is +0.732 V in 1 M perchloric acid and +0.700 V in 1 M hydrochloric acid. Using these values in place of the standard electrode potential in the Nernst equation will yield better agreement between calculated and experi-... 

The first report of the SERS spectrum of a species adsorbed at the electrode/ electrolyte interface was by Fleischman et al (1974) and concerned pyridine on silver. The Raman spectrum of the adsorbed pyridine was only observed after repeated oxidation/reduction cycles of the silver electrode, which resulted in a roughened surface. Initially, it was thought that the 106-fold enhancement in emission intensity arose as a result of the substantially increased surface area of the Ag and thus depended simply on the amount of adsorbate. However, Jeanmarie and Van Duync (1977) and Albrecht and Creighton (1977), independently reported that only a single oxidation/reduction cycle was required to produce an intense Raman spectrum and calculations showed that the increase in surface area could not possibly be sufficient to give the observed enhancement. 

STRATEGY The standard potential of one of the electrodes and the overall potential are known the value for the other electrode can be calculated from them. The cell diagram reveals which electrode is the anode (the site of oxidation, the one on the left) and which is the cathode (the site of reduction, the one on the right). The difference of the standard potentials, E°(cathode) — °(anode), is equal to the overall potential of the cell. 

Holm, T.R. and Curtiss, C.D. (1989) A comparison of oxidation-reduction potentials calculated from the As(V)/As(III) and Fe(III)/Fe(II) couples with measured platinum-electrode potentials in groundwater. Journal of Contaminant Hydrology, 5(1), 67-81. 

When a biochemical half-reaction involves the production or consumption of hydrogen ions, the electrode potential depends on the pH. When reactants are weak acids or bases, the pH dependence may be complicated, but this dependence can be calculated if the pKs of both the oxidized and reduced reactants are known. Standard apparent reduction potentials E ° have been determined for a number of oxidation-reduction reactions of biochemical interest at various pH values, but the E ° values for many more biochemical reactions can be calculated from ArG ° values of reactants from the measured apparent equilibrium constants K. Some biochemical redox reactions can be studied potentiometrically, but often reversibility cannot be obtained. Therefore a great deal of the information on reduction potentials in this chapter has come from measurements of apparent equilibrium constants. 

Because any two oxidation-reduction reactions can be combined to make a cell, the tabulation of standard electrode potentials becomes a very efficient way of calculating cell potentials under standard conditions. As indicated by Eq. (54), if the electrode reactions involve the metals of the cell terminals, the metal-metal potential due to the cell terminals is automatically included in the result. A short table of standard electrode potentials is given in Table 2. 

The cell potential of an electrochemical cell is calculated from the electrode potentials (reduction potentials) of the respective halfreactions1. Given that, by convention, the half-reaction on the left is considered to be an oxidation and that on the right a reduction we have... 

When the electromotive force of the system under examination equals E and at the reference electrode the reduction process takes place, the sought for oxidation potential e of the half cell investigated can be calculated from equation (VT-Ga) ... 

While the redox titration method is potentiometric, the spectroelectrochemistry method is potentiostatic [99]. In this method, the protein solution is introduced into an optically transparent thin layer electrochemical cell. The potential of the transparent electrode is held constant until the ratio of the oxidized to reduced forms of the protein attains equilibrium, according to the Nemst equation. The oxidation-reduction state of the protein is determined by directly measuring the spectra through the tranparent electrode. In this method, as in the redox titration method, the spectral characterization of redox species is required. A series of potentials are sequentially potentiostated so that different oxidized/reduced ratios are obtained. The data is then adjusted to the Nemst equation in order to calculate the standard redox potential of the proteic species. Errors in redox potentials estimated with this method may be in the order of 3 mV. 

Partition coefficients in the octanol-pH 7.4-phosphate-buffer system. c Nitrothiazole oxidation-reduction potentials (volts) as calculated from their half-wave potentials, as determined using a Polarecord E 261 polarograph (Metrohm AG, Herisau, Switzerland) and a saturated Ag/AgCl reference electrode. Measurements were performed at 20°C and a drop time of 1 drop/2.8 sec. The compounds were dissolved in 1 ml dimethyl formamide and added to 24 ml of a borax-potassium biphosphate buffer of pH 7.3 [prepared according to J. M. Kolthoff, J. Biol. Chem. (1925) 68, 135]. A pH of 7.4 resulted. Standard error of determination 3 mv. 

Determination of Standard Oxidation-Reduction Potentials.—In principle, the determination of the standard potential of an oxidation-reduction system involves setting up electrodes containing the oxidized and reduced states at known activities and measuring the potential B by combination with a suitable reference electrode insertion of the value of B in the appropriate form of equation (3) then permits B to be calculated. The inert metal employed in the oxidation-reduction electrode is frequently of smooth platinum, clthough platinized platinum, mercury and particularly gold are often used. 

How does calculation of the electrode potential of the system at the equivalence point differ from that for any other point of an oxidation/reduction titration ... 

Special information about the electrode reactions is often needed in order to calculate the equivalent mass for electrolysis, just as in ordinary oxidation-reduction reactions. If a solution containing Fe " is electrolyzed at low voltages, the electrode reaction for the iron might be... 

A nonspontaneous oxidation-reduction reaction, for which the calculated cell potential is negative, may be induced by electrolysis, i.e., by using an external electrical potential to force electrons into the couple undergoing reduction and to extract electrons from the couple undergoing oxidation. The minimum external potential required for electrolysis has the magnitude of the computed cell potential of the reaction. (The actual electrolyzing potential exceeds this minimum because of the irreversibility of electrode processes occurring at nonzero rates.)... 

Applying models of equilibrium oxidation-reduction, such as Figs. 4.2,4.4, and 4.6, quantitatively to soils requires that the electrode potential be known. From the electrode potential one could then calculate the soil solution concentrations of Fe2+, Mn2+, and NO and the sulfate/sulfide ratio from Eq. 4.20. Ideally, the potential of an inert electrode in the system should equal the electrode potential, because the electrode should take on a potential corresponding to the electron availability. This measurement is called the redox potential. 

In view of the numerous oxidation states, an extensive oxidation-reduction chemistry of technetium is expected. Polarographic reductions of pcrtechnetate in aqueous and in non-aqueous solutions, supplemented by coulometric and cyclic voltammetric measurements, were conducted to study the electrochemical behavior of technetium, to identify some oxidation stales and to synthesize new technetium compounds. Electrode reactions frequently proved to be irreversible and therefore not adequate for calculating thermodynamic data. The electrochemistry of technetium is reported in detail in several review articles [11-13]. 

Advantage of t ie oxidation potential is that it provides an opportunity to compare between themselves potentials of various atoms-electrodes participating in oxidation-reduction reactions. Their values for specific reactions may be found in reference literature or calculated. 

The natural cause is probably the absence in natural waters of the total equilibrium of oxidation-reduction reactions. That is why Eh values from the electrode method, as a rule, cannot be referred to any specific redox-couples and used for the interpretation of interaction between their components. An exception may be a case when the used electrode is sensitive only to one redox-couple. And the other way aroxmd. Eh calculated from concentrations of an individual redox-couple cannot be applied for characterization of the oxidation-reduction potential of a solution as a whole. 

The magnitude of the net cell potential AV° will signify the spontaneity of the oxidation-reduction reaction. However, it does not indicate the rate at which corrosion will occur. As noted before, we apply the superscript 0 to denote that we are considering the Standard Electrode Potentials. Engineers may be required to calculate the potential of a particular half-cell at concentrations and temperatures other than the standard conditions. For this purpose, we shall use the Nernst equation, which allows us to account for non-standard temperatures and solution concentrations. 

It has already been seen that whole oxidation-reduction reactions can be constructed from half-reactions. The direction in which a reaction goes is a function of the relative tendencies of its constituent half-reactions to go to the right or left These tendencies, in turn, depend upon the concentrations of the half-reaction reactants and products and their relative tendencies to gain or lose electrons. The latter is expressed by a standard electrode potential The tendency of the whole reaction to proceed to the right as written is calculated from the Nernst equation, which contains both EP and the concentrations of the reaction participants. These concepts are explained further in this section and the following section. 

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