Electron Transfer Processes Redox Potentials

Before we continue with the main theme of this chapter, we may consider briefly an entirely different type of process that may lead to equilibrium, namely electron transfer. Processes of this kind are very familiar in inorganic chemistry in the form of the typical oxidation-reductions undergone by metal ions e.g., 

In organic chemistry, similar electron transfer reactions can be observed, interconverting species that differ only by one or two electrons e.g.. 

When a metal electrode is placed in a solution of one of its salts, an equilibrium is set up between the metal and its ion e.g., 

The rates of the forward and backward reactions in equation (4.71) will depend on the potential difference between the metal and the solution. The more positive the metal, the more easily will its atoms be oxidized to ions, the electron remaining behind in the electrode. The more negative the metal, the more easily will it give up electrons to ions, reducing them to atoms. Since each of these processes leads to production on the electrode of charges which oppose them, equilibrium will be reached when the difference in potential between the metal and the solution has some definite value (the electrode potential). This potential will of course depend on the concentration of the ion since the reverse reaction involves the ion. 

In order to measure the potential difference, we have to connect the electrode and the solution to a voltmeter or potentiometer. We cannot of course use a metal connection to the solution since the metal would itself behave as an electrode. What we can do is to take a second metal in a solution of one of its salts and connect the two solutions together electrically. The potential difference between the two metal electrodes will then be equal to the difference between the corresponding electrode potentials. Such measurements do not of course tell us the absolute values of electrode potentials, but these in fact are of little importance. The relative values are sufficient. Since the electrode potentials vary with concentration, it is also necessary to specify this. The potential when the electrolyte is present in molar concentration is called the standard electrode potential 

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Electron Transport Between Photosystem I and Photosystem II Inhibitors. The interaction between PSI and PSII reaction centers (Fig. 1) depends on the thermodynamically favored transfer of electrons from low redox potential carriers to carriers of higher redox potential. This process serves to communicate reducing equivalents between the two photosystem complexes. Photosynthetic and respiratory membranes of both eukaryotes and prokaryotes contain stmctures that serve to oxidize low potential quinols while reducing high potential metaHoproteins (40). In plant thylakoid membranes, this complex is usually referred to as the cytochrome b /f complex, or plastoquinolplastocyanin oxidoreductase, which oxidizes plastoquinol reduced in PSII and reduces plastocyanin oxidized in PSI (25,41). Some diphenyl ethers, eg, 2,4-dinitrophenyl 2 -iodo-3 -methyl-4 -nitro-6 -isopropylphenyl ether [69311-70-2] (DNP-INT), and the quinone analogues,... 

As described in the introduction, certain cosurfactants appear able to drive percolation transitions. Variations in the cosurfactant chemical potential, RT n (where is cosurfactant concentration or activity), holding other compositional features constant, provide the driving force for these percolation transitions. A water, toluene, and AOT microemulsion system using acrylamide as cosurfactant exhibited percolation type behavior for a variety of redox electron-transfer processes. The corresponding low-frequency electrical conductivity data for such a system is illustrated in Fig. 8, where the water, toluene, and AOT mole ratio (11.2 19.2 1.00) is held approximately constant, and the acrylamide concentration, is varied from 0 to 6% (w/w). At about = 1.2%, the arrow labeled in Fig. 8 indicates the onset of percolation in electrical conductivity. 

Upon further contact with a redox reagent or at higher redox potentials, additional electrons can be transferred. After a two-electron transfer, each redox unit can accept one charge with formation of a singlet or triplet dianion, or (less favourably from an electrostatic point of view) both charges can enter one redox unit. Here again, an intramolecular electron-exchange process is possible. 

It should be mentioned that Spasojevic et al. [57] recently determined the two-electron reduction potential of lucigenin in water as —0.14 V. As this value is close to the one-electron reduction potential of dioxygen °[02 702] = — 0.16 V, these authors regarded their finding as a support for lucigenin redox cycling. However, it has been demonstrated long ago that two-electron reduction potentials cannot be used for the calculation of equilibrium for one-electron transfer processes [58]. 

Diacyl peroxides are, however, also electron transfer oxidants, which according to a theoretical analysis should possess standard potentials, °[(ArCOO)2/RCOO RCOO ) of around 0.6 V in water, provided that the electron transfer process is of the dissociative type (50) (Eberson, 1982c). Such a value brings thermal ET steps involving DBPO within reach for redox-active organic molecules, as for example suggested by the so-called CIEEL mechanism of chemiluminescence (Schuster, 1982). 

A preliminary electrochemical overview of the redox aptitude of a species can easily be obtained by varying with time the potential applied to an electrode immersed in a solution of the species under study and recording the relevant current-potential curves. These curves first reveal the potential at which redox processes occur. In addition, the size of the currents generated by the relative faradaic processes is normally proportional to the concentration of the active species. Finally, the shape of the response as a function of the potential scan rate allows one to determine whether there are chemical complications (adsorption or homogeneous reactions) which accompany the electron transfer processes. 

The redox potential of an electron transfer process involving a metal complex is influenced by various factors typically, in addition to the many times cited inductive effects of the ligands (together with the eventual substituents of the ligands themselves) and the stereochemistry of the redox couples, the degree of solvation of the complex and the temperature also play an important role. In an attempt to rationalize the effects of these factors on the redox potential, as well as the relationship between the redox potential and the spectroscopic properties of the complex, linear correlations between redox potential and widely differing chemical and physico-chemical properties have been investigated. 

In general, the study of the variation of the formal electrode potential of a redox process with temperature has thermodynamic implications. Hence, one is interested in the measurement of AG°, AS° and AH° for the electron transfer process. It is recalled from thermodynamics that, under standard conditions, AE° is directly proportional to the free energy of the redox reaction according to the equation ... 

An electroreductive Barbier-type allyla-tion of imines (434) with allyl bromide (429) also occurs inaTHF-PbBr2/Bu4NBr-(Al/Pt) system to give homoallyl amine (436) (Scheme 151) [533]. The combination of Pb(II)/Pb(0) redox and a sacrificial metal anode in the electrolysis system plays a role as a mediator for both cathodic and anodic electron-transfer processes. The metals used in the anode must have a less positive anodic dissolution potential than the oxidation potentials of the organic materials in order to be present or to be formed in situ. In addition, the metal ion plays the role of a Lewis acid to form the iminium ion (437) by associating with imine (435) (Scheme 151). 

Emission quenching is also observed with mononucleotides. In that case the quenching efficiency decreases from GMP (guanosine 5 monophosphate) to AMP (adenosine 5 monophosphate) i.e. it also follows the redox potentials of the bases, as G is more easily oxidisable than A, although the oxidation potential valura reported in the literature are rather different from one author to the other [101-104], Moreover the quenching rate constant by GMP in a Kries of different TAP and HAT complexes plotted versus the reduction potential of the excited state (Fig. 12) [95] is consistent with an electron transfer process. Indeed, as will be demonstrated in Sect. 4.3.1, these quenchings (by the mono-and polynucleotides) originate from such processes. 

When interaction between the metal-based components is weak, polynuclear transition metal complexes belong to the field of supramolecular chemistry. At the roots of supramolecular chemistry is the concept that supramolecular species have the potential to achieve much more elaborated tasks than simple molecular components while a molecular component can be involved in simple acts, supramolecular species can performIn other words, supramolecular species have the potentiality to behave as molecular devices. Particularly interesting molecular devices are those which use light to achieve their functions. Molecular devices which perform light-induced functions are called photochemical molecular devices (PMD). Luminescent and redox-active polynuclear complexes as those described in this chapter can play a role as PMDs operating by photoinduced energy and electron transfer processes. ... 

These potentials theoretically allow water photolysis. However, multi-electron processes have to occur at the catalyst in order to photolyze water with this complex. The lifetime of the excited state is 650 ns, and the excited state is quenched efficiently through electron transfer with redox reagents. The conversion model with this complex is described in Chapter 4. 

When the above factors are put under control, the possibility of changing the ligand L in the pentacyano(L)ferrate complexes adds a further dimension for studying systematic reactivity changes, brought out by the controlled modification of the redox potentials of the Fe(II)-Fe(III) redox couples. In this way, the rates of electron transfer reactions between a series of [Fen(CN)5L]re complexes toward a common oxidant like [Coin(NH3)5(dmso)]3+ showed a variation in agreement with Marcus predictions for outer-sphere electron transfer processes, as demonstrated by linear plots of the rate constants versus the redox potentials (123). 

An interesting finding in the CB7-MV2+ system is that, in clear contrast to host-guest systems involving CD hosts, the voltammetric data do not contain any indication that complex dissociation must precede any of the electron transfer processes. Furthermore, the electrochemistry of the inclusion complex is as fast—in the timescale accessible in these cyclic voltammetric experiments—as that of the free guest. This is clearly illustrated by the voltammograms depicted in Fig. 3.2, which show the comparative results of a scan-rate study on the MV2+/MV+ and CB7 MV2 + / 7 MV + redox couples. In both cases, the observed anodic and cathodic peak potentials are basically invariant as the scan rate is increased up to... 

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