Reduction-oxidation potentials redox defined

The formal potential of a reduction-oxidation electrode is defined as the equilibrium potential at the unit concentration ratio of the oxidized and reduced forms of the given redox system (the actual concentrations of these two forms should not be too low). If, in addition to the concentrations of the reduced and oxidized forms, the Nernst equation also contains the concentration of some other species, then this concentration must equal unity. This is mostly the concentration of hydrogen ions. If the concentration of some species appearing in the Nernst equation is not equal to unity, then it must be precisely specified and the term apparent formal potential is then employed to designate the potential of this electrode. 

The midpoint potential of a half-reaction E, is the value when the concentrations of oxidized and reduced species are equal, [Aox] = [Aredl- In biological systems the standard redox potential of a compound is the reduction/oxidation potential measured under standard conditions, defined at pH = 7.0 versus the hydrogen electrode. On this scale, the potential of 02/water is +815 mV, and the potential of water/H2 is 414 mV. A characteristic of redox reactions involving hydrogen transfer is that the redox potential changes with pH. The oxidation of hydrogen H2 = 2H + 2e is an m = 2 reaction, for which the potential is —414 mV at pH 7, changing by 59.2 mV per pH unit at 30°C. 

The reduction-oxidation potential (typically expressed in volts) of a compound or molecular entity measured with an inert metallic electrode under standard conditions against a standard reference half-cell. Any oxidation-reduction reaction, or redox reaction, can be divided into two half-reactions, one in which a chemical species undergoes oxidation and one in which another chemical species undergoes reduction. In biological systems the standard redox potential is defined at pH 7.0 versus the hydrogen electrode and partial pressure of dihydrogen of 1 bar. 

Reduction-oxidation reactions are mediated by micro-organisms and involve the transfer of electrons between reactants and products. Free electrons do not exist in solution, so an oxidation reaction (loss of electrons) must be balanced by a reduction reaction (gain of electrons). Redox potential is defined by the Nemst equation and is the energy gained in the transfer of 1 mol of electrons from an oxidant to H2. 

The property of chemotropicity testifies to the balance of the redox layer system with respect to the vertical fluxes of the oxidants and reductants supplied. This should be the well-defined sequence of changes with depth of the favorability of the potential redox reactions [ 17,75] that can be realized by the bacterial community. The development of bacteria in this case should affect the distributions of nutrients. By modern estimation [79] the chemosynthetic production is comparable with photosynthetic production, and that should in the same manner affect the consumption of inorganic nutrients and production of their organic forms. Besides this the possible abiotic chemical reactions and the sedimentation of particulate matter of different densities should also play their roles in this mechanism. 

A measure of the oxidation/reduction capability of a solution (liquid or solid) measured with an -> inert electrode. For -> electrochemically reversible systems it is defined by the - Nernst equation. For -> electrochemically irreversible systems it is a conditional measuring quantity, i.e., depending on the experimental conditions. See also -> potential, - redox potential. 

Redox potential is defined by the half cell reduction potential that is created by redox couples that are primarily due to GSH, NAD+ and nicotinamide dinucleotide phosphate. These couples are in ratios of the oxidized to reduced form of the molecules (NAD /NAD, NADP /NADPH, and GSSG/2GSH). The redox couples can be independent, as well linked to each other to form related couples. The redox environment is a reflection of these couples. These ratios can be measured by the Nemst equation, similar to a voltaic cell. 

Redox parameters analogous to those for acid-base chemistry can be defined for all aqueous systems. The redox intensity factor pE is an energy parameter in non-dimensional form that describes the ratio of electron acceptors (oxidants) and donors (reductants) in a redox couple. The redox potential (Ej ) of the system is an alternative and equivalent intensity factor. Table I summarizes the complete thermodynamic analogy between pH and pE. An analogy between acid-base and redox systems can also be made for capacity factors. 

The excess potential that causes the faradic current flow is defined as the overpotential, 7, which is a measure of the degree of polarization or the departure of the electrode potential from the Nemstian (equilibrium) value (Bard and Faulkner, 1980). The Nemstian equilibrium applies to reversible (nonpolarizable) electrode surfaces, where dynamic equilibrium is established rapidly. When a large electrical field is applied, a large current fiows, and the equilibrium at the electrodes cannot adjust quickly to overcome the electrode reactions, thus giving rise to irreversible electrodes. To maintain the current flow, the electrodes adapt a new potential, differing from the equilibrium by the overpotential, rj. As electrons migrate and reduction-oxidation (redox) reactions take place, the equilibrium potential is restored. 

In natural waters occur not one but several oxidation-reduction reactions. These reactions are associated with the presence of several elements, which are capable of changing their charge, and run in parallel. For this reason, total oxidation potential of the solution is defined by the nature and concentration of all redox-couples. Components which noticeably affect the solution s oxidation-reduction potential are called electroactive. Elements whose concentration and form of existence actually control solution s oxidation are culled potential-setting. In natural waters these are usually O, S, C, N and Fe. The medium whose oxidation potential value almost does not change with the addition of oxidizers or reducers is called redox-buffers. The redox-buffer may be associated with composition of the water itself, of its host rocks or with the effect of atmosphere. In the subsurface redox-buffers are associated, as a rule, with the content of iron, sulphur or manganese. Stably high Eh value in the surface and ground waters is caused by the inexhaustible source of in the atmosphere. 

To be able to understand how computational approaches can and should be used for electrochemical prediction we first of all need to have a correct description of the precise aims. We start from the very basic lithium-ion cell operation that ideally involves two well-defined and reversible reduction and oxidation redox) reactions - one at each electrode/electrolyte interface - coordinated with the outer transport of electrons and internal transport of lithium ions between the positive and negative electrodes. However, in practice many other chemical and physical phenomena take place simultaneously, such as anion diffusion in the electrolyte and additional redox processes at the interfaces due to reduction and/or oxidation of electrolyte components (Fig. 9.1). Control of these additional phenomena is crucial to ensure safe and stable ceU operation and to optimize the overall cell performance. In general, computations can thus be used (1) to predict wanted redox reactions, for example the reduction potential E ) of a film-forming additive intended for a protective solid electrolyte interface (SEI) and (2) to predict unwanted redox reactions, for example the oxidation potential (Eox) limit of electrolyte solvents or anions. As outlined above, the additional redox reactions involve components of the electrolyte, which thus is a prime aim of the modelling. The working agenda of different electrolyte materials in the cell -and often the unwanted reactions - are addressed to be able to mitigate the limitations posed in a rational way. 

Both harmonic and electrochemical frequency modulation (EFM) methods take advantage of nonlinearity in the E-I response of electrochemiced interfaces to determine corrosion rate [47-50]. A special application of harmonic methods involves harmonic impedance spectroscopy [5i]. The EFM method uses one or more a-c voltage perturbations in order to extract corrosion rate. The electrochemical frequency modulation method has been described in the literature [47-50] and has recently been reviewed [52]. In the most often used EFM method, a potential perturbation by two sine waves of different frequencies is applied across a corroding metal interface. The E-I behavior of corroding interfaces is typically nonlinear, so that such a potential perturbation in the form of a sine wave at one or more frequencies can result in a current response at the same and at other frequencies. The result of such a potential perturbation is various AC current responses at various frequencies such as zero, harmonic, and intermodulation. The magnitude of these current responses can be used to extract information on the corrosion rate of the electrochemical interface or conversely the reduction-oxidation rate of an interface dominated by redox reactions as well as the Tafel parameters. This is an advantage over LPR and EIS methods, which can provide the Z( ) and, at = 0, the polarization resistance of the corroding interface, but do not uniquely determine Tafel parameters in the same set of data. Separate erqreriments must be used to define Tafel parameters. A special extension of the method involves... 

Note that AG° q has a liquid-phase standard state of 1 moI/L, and we can use a gas-phase standard state of either 1 atm or 1 mol/L, as long as we use the same convention for the oxidized and reduced forms. Often the 1 mol/L standard state is used. There are several sources of uncertainties in the calculations of reduction potentials, and we will comment on them by scanning the published literature on actinide elements, for which the most studied redox systems are actinyl aqua ions, with the exception of one study on Pu(VII)/Pu(VIII) [172], The first comment is that redox potentials are defined with respect to the standard hydrogen electrode corresponding to the following half-equation... 

Electronic influences on reduction-oxidation (redox) potentials. The electrons released to the rest of a molecule by a C-methyl substituent lower the redox potential (Eq), As a result, the affected substance becomes a more active reducing agent and a less active oxidizing agent (i.e. it is more resistant to becoming reduced) than is the unmethylated homologue. Redox potentials are defined in Section 11.4 and refer to the equilibrium between oxidized and reduced forms (all values in this book are versus the normal hydrogen electrode). 

Anaerobic metabolism occnrs nnder conditions in which the diffusion rate is insufficient to meet the microbial demand, and alternative electron acceptors are needed. The type of anaerobic microbial reaction controls the redox potential (Eh), the denitrification process, reduction of Mu and SO , and the transformation of selenium and arsenate. Keeney (1983) emphasized that denitrification is the most significant anaerobic reaction occurring in the subsurface. Denitrification may be defined as the process in which N-oxides serve as terminal electron acceptors for respiratory electron transport (Firestone 1982), because nitrification and NOj" reduction to produce gaseous N-oxides. hi this case, a reduced electron-donating substrate enhances the formation of more N-oxides through numerous elechocarriers. Anaerobic conditions also lead to the transformation of organic toxic compounds (e.g., DDT) in many cases, these transformations are more rapid than under aerobic conditions. 

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