Oxidation-reduction electrodes determination

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. 

Reactions that take place consecutive to the electrode process can be studied polarographioally only in those cases in which the electrode process is reversible. In these cases the wave-heights and the wave-shape remain unaffected by the chemical processes. However, the half-wave potentials are shifted relative to the equilibrium oxidation-reduction potential, determined e.g. potentiometrically. Hence, whereas in all above examples, limiting currents were measured to determine the rate constant, it is the shifts of half-wave potentials which are measured here. First- and second-order chemical reactions will be discussed in the following. 

The Standard Potential of the Mercurous-Mercuric Electrode. A method for obtaining standard potentials of oxidation-reduction electrodes which utilizes the best procedure so far developed in this field is the one that was used by Fopoff and associates. The method may be illustrated by the determination of the standard potential of the mercurous-mercuric electrode. The type of cell used by Popoff, Riddick, Worth and Ough1 was... 

Although the applied potential at the working electrode determines if a faradaic current flows, the magnitude of the current is determined by the rate of the resulting oxidation or reduction reaction at the electrode surface. Two factors contribute to the rate of the electrochemical reaction the rate at which the reactants and products are transported to and from the surface of the electrode, and the rate at which electrons pass between the electrode and the reactants and products in solution. 

Many dehydrogenase enzymes catalyze oxidation/reduction reactions with the aid of nicotinamide cofactors. The electrochemical oxidation of nicotinamide adeniiw dinucleotide, NADH, has been studied in depthThe direct oxidation of NADH has been used to determine concentration of ethanol i s-isv, i62) lactate 157,160,162,163) pyTuvate 1 ), glucose-6-phosphate lactate dehydrogenase 159,161) alanine The direct oxidation often entails such complications as electrode surface pretreatment, interferences due to electrode operation at very positive potentials, and electrode fouling due to adsorption. Subsequent reaction of the NADH with peroxidase allows quantitation via the well established Clark electrode. 

Polarisation titrations are often referred to as amper-ometric or biamperometric titrations. It is necessary that one of the substances involved in the titration reaction be oxidisable or reducible at the working electrode surface. In general, the polarisation titration method is applicable to oxidation-reduction, precipitation and complex-ation titrations. Relatively few applications involving acid/base titration are found. Amperometric titrations can be applied in the determination of analyte solutions as low as ICE5 M to 10-6 M in concentration. 

As described in Sec. 11.3, the spontaneous corrosion potential of a corroding metal is represented by the intersection of the anodic polarization curve of metal dissolution with the cathodic polarization curve of oxidant reduction (Figs. 11—5 and 11-6). Then, whether a metal electrode is in the active or in the passive state is determined by the intersection of the anodic and cathodic polarization curves. 

In a fluid such as milk, which contains several oxidation-reduction systems, the effect of each system on the potential depends on several factors. These include the reversibility of the system, its E0 value or position on the scale of potential, the ratio of oxidant to reductant, and the concentration of active components of the system. Only a reversible system gives a potential at a noble metal electrode, and this measured potential is an intensity factor analogous to the potential measured on a hydrogen electrode in determining hydrogen ion concentrations. 

When comparing Equations f. 37 and 1.38, it is clear that it is the relation of two parameters, namely the transport coefficient, m, and the formal reaction rate constant k° (incorporated in k(E) in Equation 1.37 see also Equation 1.20), that decides whether the current is determined mainly by the BV relation or by transport. In a voltammogram, the potential area in which the BV relation applies decreases when k° increases relatively in relation to m. In electrochemical terms, one speaks of an increase in the reversible character of the voltammetric wave. When having sufficient positive (negative) potentials for an oxidation (reduction), transport mostly prevails. Providing that the appropriate cell and/or electrode configuration is present, transport-determined currents are very reproducible and suitable for analytical purposes. 

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. 

These laws (determined by Michael Faraday over a half century before the discovery of the electron) can now be shown to be simple consequences of the electrical nature of matter. In any electrolysis, an oxidation must occur at the anode to supply the electrons that leave this electrode. Also, a reduction must occur at the cathode removing electrons coming into the system from an outside source (battery or other DC source). By the principle of continuity of current, electrons must be discharged at the cathode at exactly the same rate at which they are supplied to the anode. By definition of the equivalent mass for oxidation-reduction reactions, the number of equivalents of electrode reaction must be proportional to the amount of charge transported into or out of the electrolytic cell. Further, the number of equivalents is equal to the number of moles of electrons transported in the circuit. The Faraday constant (F) is equal to the charge of one mole of electrons, as shown in this equation ... 

Redox titrations involve determination of equilibrium between the enzyme and a redox agent of known redox potential. The method requires a redox agent with redox potential close to the protein of interest, to ensure reversibility. The protein is exposed to different concentrations of the redox agent, and once equilibrium is attained, the half cell potential is measured with electrodes and the oxidation-reduction state of the proteins is measured by some physical technique, usually UV-Vis spectrophotometry. The concentration of the oxidized and reduced forms is determined at isosbestic points, and thus spectral characterization of redox species (ferric enzyme,... 

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. 

Electrometric Titration Precipitation Reactions.—One of the most important practical applications of electrode potentials is to the determination of the end-points of various typos of titration the subject will be treated here from the standpoint of precipitation reactions, while neutralization and oxidation-reduction processes are described more conveniently in later chapters. 

Approximate Determination of Standard Potentials.—Many studies have been made of oxidation-reduction systems with which, for one reason or another, it is not possible to obtain accurate results this may be due to the difficulty of applying activity corrections, uncertainty as to the exact concentrations of the substances involved, or to the slowness of the establishment of equilibrium with the inert metal of the electrode. It is probable that whenever the difference in the number of electrons between the oxidized and reduced states, i.e., the value of n for the oxidation-reduction system, is relatively large the processes of oxidation and reduction occur in stages, one or more of which may be slow. In that event equilibrium between the system in the solution and the electrode will be established slowly, and the measured potential may be in error. To expedite the attainment of the equilibrium a potential mediator may be emploj cd this is a substance that undergoes reversible oxidation-reduction and rapidly reaches equilibrium with the electrode. 

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