Oxidation-reduction titration equilibrium

In the case of two flavoenzyme oxidase systems (glucose oxidase (18) and thiamine oxidase s where both oxidation-reduction potential and semiquinone quantitation values are available, semiquinone formation is viewed to be kinetically rather than thermodynamically stabilized. The respective one-electron redox couples (PFl/PFl- and PFI7PFIH2) are similar in value (from essential equality to a 50 mV differential) which would predict only very low levels of semiquinone (32% when both couples are identical) at equilibrium. However, near quantitative yields (90%) of semiquinone are observed either by photochemical reduction or by titration with dithionite which demonstrates a kinetic barrier for the reduction of the semiquinone to the hydroquinone form. The addition of a low potential one-electron oxidoreductant such as methyl viologen generally acts to circumvent this kinetic barrier and facilitate the rapid reduction of the semiquinone to the hydroquinone form. 

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. 

In an oversimplified way one can say that acids of the volcanoes have reacted with the bases of the rocks the composition of the ocean (which is at the first endpoint (pH = 8) of the titration of a strong acid with a carbonate) and the atmosphere (which with its pco2 = 10 3 5 atm is nearly in equilibrium with the ocean) reflect the proton balance of reaction (5.25). Oxidation and reduction are accompanied by proton release and proton consumption, respectively. (In order to maintain charge balance, the production of e will eventually be balanced by the production of H+.) Furthermore, the dissolution of rocks and the precipitation of minerals are accompanied by H+-consumption and H+-release, respectively. Thus, as shown by Broecker (1971), the pe and pH of the surface of our global environment reflect the levels where the oxidation states and the H+ ion reservoirs of the weathering sources equal those of the sedimentary products. 

The exchange current density, which governs the rate of attainment of electrode equilibrium, varies enormously from one potential-determining redox couple to another. It varies not only with the initial concentration but also with the ratio of oxidant to reductant [Equation (14-14)]. In titrations performed at great dilution, the equilibrium near the end point may be reached slowly. Therefore, it may be advantageous to select a method of end-point detection that is not dependent on equilibrium near the end point. 

In the foregoing discussion the indicator has tacitly been assumed to come rapidly to equilibrium at each point of the titration curve. That this is an over-simplihcation is evident from a number of experimental observations. Kolthoflf and Sarver found that the oxidation of diphenylamine with dichromate is induced by the Fe(II)-dichromate reaction. The direct oxidation is so slow that the indicator blank is best determined by comparison of the visual with the potentiometric end point. With ferroin. Smith and Brandt and Stockdale foimd that the reverse titration, dichromate with iron, gave satisfactory results at sufficiently high acidities, whereas the direct titration failed because the indicator could not be oxidized. Here the oxidation seems to be slow and the reduction rapid because of the irreversible nature of the oxidant and the reversible nature of the reductant. 

Fig. 6. (A) Redox titration of the primary electron acceptor in Rb. sphaeroides chromatophores at pH 11. The amplitude of absorption changes Induced by short flashes is plotted as a function of the redox potential solid dots and empty circles represent reductive and oxidative titrations, respectively. The solid line is the theoretical Nernst curve. (B) Equilibrium midpoint potentials of the primary acceptor as determined in (A) plotted as a function of pH. Figure source Prince and Dutton (1978) Protonation and the reducing potential of the primary electron acceptor. In RK Clayton and WR Sistrom (eds) The Photosynthetic Bacteria, p 443, 444. Plenum.

Fig. 9 shows the titration results for the following samples chloroplast lamellae and TSF-1 particles, both measured at 820 nm, and the CPI complex measured at 820 as well as 703 nm. Each sample was titrated oxidatively (starting with 100 pM ferrocyanide and adding ferricyanide to a maximum concentra tion of 10 mM) and reductively (starting with 1-5 mM ferricyanide and adding ferrocyanide to a maximum concentration of 10 mM). The titration is a plot of the light-induced AA V5. the actual redox-potential of the medium or the ferri-/ferrocyanide ratio as shown in Fig. 9. The plot of the data points clearly show that the titration was completely reversible and that P700 was in redox equilibrium with the ferri-/ferro-cya-nide couple. The solid line is the theoretical Nernst curve for a one-electron transition and the data points agree well with the theoretical course. The titration curve for both the chloroplast lamellae and the TSF-1, as well as D144 (data not shown here), yielded an value of+492 mV.

Because of the close analogy between acid-base and redox behavior, it will come as no surprise that one can use redox titrations, and also simulate them on a spreadsheet. In fact, the expressions for redox progress curves are often even simpler than those for acid-base titrations, because they do not take the solvent into account. (Oxidation and reduction of the solvent are almost always kinetically controlled, and therefore do not fit the equilibrium description given here. In the examples given below, they need not be taken... 

Before we discuss redox titration curves based on reduction-oxidation potentials, we need to learn how to calculate equilibrium constants for redox reactions from the half-reaction potentials. The reaction equilibrium constant is used in calculating equilibrium concentrations at the equivalence point, in order to calculate the equivalence point potential. Recall from Chapter 12 that since a cell voltage is zero at reaction equilibrium, the difference between the two half-reaction potentials is zero (or the two potentials are equal), and the Nemst equations for the halfreactions can be equated. When the equations are combined, the log term is that of the equilibrium constant expression for the reaction (see Equation 12.20), and a numerical value can be calculated for the equilibrium constant. This is a consequence of the relationship between the free energy and the equilibrium constant of a reaction. Recall from Equation 6.10 that AG° = —RT In K. Since AG° = —nFE° for the reaction, then... 

Iodide colors a solution of PtCU red to brown (sensitive to 0.3 mmol Pt) (Fe , Cu , and other oxidants interfere) and may precipitate black PU4. Excess of KI forms K2[Ptl6], brown, shghtly soluble, and unstable enough that platinum may be determined volumetrically by treating [PtCle] with excess r and titrating the liberated iodine with thiosulfate, which shifts the Pt -Pl equilibrium completely toward reduction (as the O2 and acidic CO2 three paragraphs above shift it toward oxidation) ... 

C60 with strong acids, a C q solution was generated by controlled potential electrolysis and titrated with concentrated triflic acid. UV/visible-near IR spectra as well as steady state voltammetry with ultramicroelectrodes were used to monitor 50 and C60H. Addition of the triflic acid led to a decrease of absorbance and steady state oxidation current of Cgo". At the same time, a reduction current appeared at potentials more negative than the half wave potential of Cgg oxidation. This current was attributed to CeoH reduction. With some simple assumptions, the current changes were used to calculate the equilibrium constant of the reaction ... 

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