Electrochemistry standard electrode potentials

Standard electrode potential is an important concept in electrochemistry. Standard potentials for many half-reactions have been measured or calculated. It is designated as EY and expressed in volts (V). From the values of E° one can... 

First of all, the important role of platinum as the metal part of the standard hydrogen electrode (SHE), which is the primary standard in electrochemistry should be mentioned. The standard potential of an electrode reaction (standard electrode potential) is defined as the value of the standard potential of a cell reaction when that involves the oxidation of molecular hydrogen to solvated (hydrated) protons (hydrogen ions) ... 

A significant quantity of the standard electrode potentials, if, has been experimentally estimated at ambient conditions (25 °C and 1 bar) [8] however, only a limited number of A(0 values can be found at high temperatures and pressures [9], Therefore, experimental studies to derive a comprehensive list of values for a wide range of temperatures are a serious challenge in high-temperature experimental electrochemistry. 

Whereas the standard electrode potentials of many half-cell reactions have been known at ambient conditions and can be easily found in a number of reference books, almost none of them are documented for a region of high-temperature subcritical and supercritical conditions. Therefore, the creation of well-established approaches for developing a comprehensive list of the standard potentials measured over a wide range of temperatures remains a challenge for high-temperature experimental electrochemistry. The recently developed instruments for poten-tiometric studies at temperatures above 300 °C can be useful for developing such a database. 

Numerous applications of standard electrode potentials have been made in various aspects of electrochemistry and analytical chemistry, as well as in thermodynamics. Some of these applications will be considered here, and others will be mentioned later. Just as standard potentials which cannot be determined directly can be calculated from equilibrium constant and free energy data, so the procedure can be reversed and electrode potentials used for the evaluation, for example, of equilibrium constants which do not permit of direct experimental study. Some of the results are of analjrtical interest, as may be shown by the following illustration. Stannous salts have been employed for the reduction of ferric ions to ferrous ions in acid solution, and it is of interest to know how far this process goes toward completion. Although the solutions undoubtedly contain complex ions, particularly those involving tin, the reaction may be represented, approximately, by... 

Chapter 18 Introduction to Electrochemistry 490 Chapter 19 Applications of Standard Electrode Potentials 523 Chapter 20 Applications of Oxidation/Reduction Titrations 560 Chapter 21 Potentiometry 588... 

Redox Titrations Electrochemistry Chapter 19 Standard Electrode Potentials Chapter 20 Oxidation/ReductionTitrations Chapter 21 Potentiometry Chapter 17 Using Electrode Potentials Chapter 18 Oxidation/Reduction Titrations Chapter 19 Potentiometry... 

As already discussed, the standard hydrogen electrode (SHE) is the chosen reference half-cell upon which tables of standard electrode potentials are based. The potential of this system is zero by definition at all temperatures. Although this reference electrode was often used in early work in electrochemistry, it is almost never seen in chemical laboratories at the present time. It is simply too awkward to use because of the requirement for H2 gas at 1 bar pressure and safety considerations. 

In electrochemistry, the electrode potential is defined by the electronic energy level in a solid electrode referred to the energy level of the standard gaseous electron just outside the surface of an electrolyte (aqueous solution) in which the electrode is immersed [6] ... 

Basic equations for almost every subfield of electrochemistry from first principles, referring at all times to the soundest and most recent theories and results unusually useful as text or as reference. Covers coulometers and Faraday s Law, electrolytic conductance, the Debye-Hueckel method for the theoretical calculation of activity coefficients, concentration cells, standard electrode potentials, thermodynamic ionization constants, pH, potentiometric titrations, irreversible phenomena. Planck s equation, and much more, a indices. Appendix. 585-item bibliography. 197 figures. 94 tables, ii 4. 478pp. 5-% x 8. ... 

The data in this table are mainly taken from A. J. Bard, J. Jordan, and R. Parsons, Eds., Standard Potentials in Aqueous Solutions, Marcel Dekker, New York, 1985 (prepared under the auspices of the Electrochemistry and Electroanalytical Chemistry Commissions of lUPAC). Other sources of standard potentials and thermodynamic data include (1) A. J, Bard and H. Lund, Eds., The Encyclopedia of the Electrochemistry of the Elements, Marcel Dekker, New York, 1973-1986. (2) G. Milazzo and S. Caroli, Tables of Standard Electrode Potentials, Wiley-Interscience, New York, 1977. The data here are referred to the NHE based on a 1-atm standard state for H2. See the footnote in Section 2.1.5 concerning the recent change in standard state. 

In the above equation, is the standard electrode potential [d(ox] = [ redl = 1, / refers to the gas constant 1.987 cal (g-mol) K, T to the absolute temperature in degrees Kelvin, and F to Faraday s constant F = 23.06kcal (g-equivalent) V ). Equation (26) is called the Nemst equation relating the electrode potential to the concentrations, and it is one of the most important relationships in electrochemistry. values for various reactions are presented in Table 4.1.2. 

Theoretically, by poising the potential of an electrochemical cell at a value which is sufficient to reduce chromium(III) but not aluminium(III), chromium could be removed preferentially from solution. As chromium is a common contaminant of bauxitic alloys (the main feedstock for aluminium industry) electrochemistry may provide a means of selectively removing chromium from aluminium products. However, this process may be impractically slow. Much depends on the relative concentrations of aluminium and chromium, temperature, pH and cell design. Nevertheless, standard electrode potentials can be used as a preliminary evaluation of the feasibility of electrochemical methods for clean-up. 

For the study of electrochemistry and corrosion one must be able to compare the equilibrium potentials of different electrode reactions. To these ends, by convention, a scale of standard electrode potentials is defined by arbitrarily assigning the value of zero to the equilibrium potential of the electrode (2.42), under standard conditions (Pfj2 = 1 bar =1.013 atm, T = 298 K, au+ = 1) ... 

Since the measured cell potential difference is actually the potential difference between two electrodes, it immediately comes to mind to assimilate each of the bracketed terms into the potential of each of the electrodes. They are called electrode potentials. E° and °2, which are in the two subgroups, exhibit characteristic values of both couples Oxi/Redi and Oxa/Reda. These constants are called standard potentials of both couples and are symbolized (Oxi/Redi) and °(Ox2/Red2). Assigning numerical values to and E°2 has been a problem since the experimental determination of absolute electrode potentials hence, assigning those to standard electrode potentials is impossible (see the electrochemistry part). It was solved by assigning relative values to them. The strategy was based on the fact that if absolute electrode potentials are not measurable, the difference between them can be. Thus, an electrode standard potential has been chosen conventionally for the couple H+w/H2(g) (hydrogen electrode). Its standard electrode has been set definitively to the value 0.0000 V at every temperature ... 

Another troublesome aspect of the reactivity ratios is the fact that they must be determined and reported as a pair. It would clearly simplify things if it were possible to specify one or two general parameters for each monomer which would correctly represent its contribution to all reactivity ratios. Combined with the analogous parameters for its comonomer, the values rj and t2 could then be evaluated. This situation parallels the standard potential of electrochemical cells which we are able to describe as the sum of potential contributions from each of the electrodes that comprise the cell. With x possible electrodes, there are x(x - l)/2 possible electrode combinations. If x = 50, there are 1225 possible cells, but these can be described by only 50 electrode potentials. A dramatic data reduction is accomplished by this device. Precisely the same proliferation of combinations exists for monomer combinations. It would simplify things if a method were available for data reduction such as that used in electrochemistry. 

Electrochemistry is in many aspects directly comparable to the concepts known in heterogeneous catalysis. In electrochemistry, the main driving force for the electrochemical reaction is the difference between the electrode potential and the standard potential (E — E°), also called the overpotential. Large overpotentials, however, reduce the efficiency of the electrochemical process. Electrode optimization, therefore, aims to maximize the rate constant k, which is determined by the catalytic properties of the electrode surface, to maximize the surface area A, and, by minimization of transport losses, to result in maximum concentration of the reactants. 

In electrochemistry, the electron level of the normal hydrogen electrode is important, because it is used as the reference zero level of the electrode potential in aqueous solutions. The reaction of normal hydrogen electrode in the standard state (temperature 25°C, hydrogen pressure 1 atm, and unit activity of hydrated protons) is written in Eqn. 2-54 ... 

The relative electrode potential nhe referred to the normal (or standard) hydrogen electrode (NHE) is used in general as a conventional scale of the electrode potential in electrochemistry. Since the electrode potential of the normal hydrogen electrode is 4.5 or 4.44 V, we obtain the relationship between the relative electrode potentiEd, Ema, and the absolute electrode potential, E, as shown in Eqn. 4-36 ... 

TABLE 6-L Die standard equilibrium potentials for redox electrode reactions of h rdrated redox particles at 25 C nhe = relative electrode potential referred to the normal hydrogen electrode. [Handbooks of electrochemistry.]... 

In electrochemistry, the chemical potential of hydrated ions has been determined from the equilibrium potential of ion transfer reactions referred to the normal hydrogen electrode. For the reaction of metal ion transfer (metal dissolution-deposition reaction) of Eqns. 6-16 and 6-17, the standard equilibriiun potential Sive in terms of the standard chemical potential, li, by Eqn. 

Although the entire discussion of electrochemistry thus far has been in terms of aqueous solutions, the same principles apply equaly well to nonaqueous solvents. As a result of differences in solvation energies, electrode potentials may vary considerably from those found in aqueous solution. In addition the oxidation and reduction potentials characteristic of the solvent vary with the chemical behavior of the solvent. as a result of these two effects, it is often possible to carry out reactions in a nonaqueous solvent that would be impossible in water. For example, both sodium and beryllium are too reactive to be electroplated from aqueous solution, but beryllium can be electroplated from liquid ammonia and sodium from solutions in pyridine. 0 Unfortunately, the thermodynamic data necessary to construct complete tables of standard potential values are lacking for most solvents other than water. Jolly 1 has compiled such a table for liquid ammonia. The hydrogen electrode is used as the reference point to establish the scale as in water ... 

ACTIVITY SERIES- Also referred to as the electromotive series or the displacement series, this is an arrangement of the metals (other elements can be included) in the order of their tendency to react with water and acids, so that each metal displaces from solution those below itiu the series and is displaced by those above it. See Table 1. Since the electrode potential of a metal in equilibrium with a solution of its ions cannot be measured directly, the values in the activity series are, in each case, the difference between the electrode potential of the given metal tor element) in equilibrium with a solution of its ions, and that of hydrogen in equilibrium with a solution of its ions. Thus in the table, it will be noted that hydrogen lias a value of 0.000. In experimental procedure, the hydrogen electrode is used as the standard with which the electrode potentials of other substances are compared. The theory of displacement plays a major role in electrochemistry and corrosion engineering. See also Corrosion and Electrochemistry. 

There is an inherent similarity between the spectrum and an electrochemical current-voltage curve that is important from the point of view of chemical selectivity. In both cases, the x-axis (voltage or wavelength) is directly related to the energy. In electrochemistry, this energy corresponds to the transfer of electrons between the analyte and the electrode. It is related to the standard electrochemical potential. In optical interactions, molar absorptivity is probabilistically related to the excitation energy of the molecule. 

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