Electrode potentials, standard tabulating

The cell potential is the difference between two electrode potentials, one associated with the cathode and the other associated with the anode. By convention, the potential associated with each electrode is chosen to be the potential for reduction to occur af fhaf elecfrode. Thus, standard electrode potentials are tabulated for reduction reactions they are standard reduction potentials, denoted E°ed The cell potential, E°eii, is given by the standard reduction potential of fhe cafhode... 

Electrode potentials are customarily tabulated on the standard hydrogen electrode (SHE) scale (although the SHE is never actually used experimentally because it is inconvenient in many respects). Therefore, conversion of potentials into the UHV scale requires the determination of E°(H+/H2) vs. UHV. According to the concepts developed above, such a potential would measure the energy of electrons in the Pt wire of the hydrogen electrode, modified by the contact with the solution. 

The relationships of the type (3.1.54) and (3.1.57) imply that the standard electrode potentials can be derived directly from the thermodynamic data (and vice versa). The values of the standard chemical potentials are identified with the values of the standard Gibbs energies of formation, tabulated, for example, by the US National Bureau of Standards. On the other hand, the experimental approach to the determination of standard electrode potentials is based on the cells of the type (3.1.41) whose EMFs are extrapolated to zero ionic strength. 

In the discussion of the Daniell cell we indicated that this cell produces 1.10 volts. This voltage is really the difference in potential between the two half-cells. There are half-cell potentials associated with all half-cells A list of all possible combinations of half-cells would be tremendously long. Therefore, a way of combining desired half-cells has been developed. The cell potential (really the half-cell potentials) depends on concentration and temperature, but initially we ll simply look at the half-cell potentials at the standard temperature of 298 K (25°C) and all components in their standard states (1 M concentration of all solutions, 1 atmosphere pressure for any gases, and pure solid electrodes). All the half-cell potentials are tabulated as the reduction potentials, that is, the potentials associated with the reduction reaction. The hydrogen half-reaction has been defined as the standard and has been given a value of exactly 0.00 V. All the other half-reactions have been measured relative to it, some positive and some negative. The table of standard reduction potentials provided on the AP exam is shown in Table 16.1 and in the back of this book. 

The SHE. The H" " H2 couple is the basis of the primary standard around which the whole edifice of electrode potentials rests. We call the H H2 couple, under standard conditions, the standard hydrogen electrode (SHE). More precisely, we say that hydrogen gas at standard pressure, in equilibrium with an aqueous solution of the proton at unity activity at 298 K has a defined value of of 0 at all temperatures. Note that all other standard electrode potentials are temperature-dependent. The SHE is shown schematically in Figure 3.3, while values of Eq r are tabulated in Appendix 3. 

As it has been shown that the Gibbs function for formation of an individual ion has no operational meanings [12], no way exists to determine such a quantity experimentally. However, for the purposes of tabulation and calculation, it is possible to separate AfGm of an electrolyte arbitrarily into two or more parts, which correspond to the number of ions formed, in a way analogous to that used in tables of standard electrode potentials. In both cases, the standard Gibbs function for formation of aqueous H" " is defined to be zero at every temperature ... 

Solubility products can be derived indirectly from standard electrode potentials and other thermochemical data, and directly from tabulated standard Gibbs energies of formation, AfG°, of the ions in aqueous solution [12]. Thus, the use of... 

Because electrode potentials are defined with reference to the H+/H2 electrode under standard conditions, E° values apply implicitly to (hypothetically ideal) acidic solutions in which the hydrogen ion concentration is 1 mol kg-1. Such E° values are therefore tabulated in Appendix D under the heading Acidic Solutions. Appendix D also lists electrode potentials for basic solutions, meaning solutions in which the hydroxide ion concentration is 1.0 mol kg-1. The conversion of E° values to those appropriate for basic solutions is effected with the Nernst equation (Eq. 15.15), in which the hydrogen ion concentration (if it appears) is set to 1.0 x 10-14 mol kg-1 and the identity and concentrations of other solute species are adjusted for pH 14. For example, for the Fc3+/2+ couple in a basic medium, the relevant forms of iron(III) and iron(II) are the solid hydroxides, and the concentrations of Fe3+ (aq) and Fe2+ (aq) to be inserted into the Nernst equation are those determined for pH 14 by the solubility products of Fe(OH)3(s) and Fe(OH)2(s), respectively. Examples of calculations of electrode potentials for nonstandard pH values are given in Sections 15.2 and 15.3. 

Because any two oxidation-reduction reactions can be combined to make a cell, the tabulation of standard electrode potentials becomes a very efficient way of calculating cell potentials under standard conditions. As indicated by Eq. (54), if the electrode reactions involve the metals of the cell terminals, the metal-metal potential due to the cell terminals is automatically included in the result. A short table of standard electrode potentials is given in Table 2. 

Standard half-cell potentials are tabulated as reduction potentials. These are sometimes referred to as standard electrode potentials E°. Therefore,... 

Here, AH(A-B) is the partial molar net adsorption enthalpy associated with the transformation of 1 mol of the pure metal A in its standard state into the state of zero coverage on the surface of the electrode material B, ASVjbr is the difference in the vibrational entropies in the above states, n is the number of electrons involved in the electrode process, F the Faraday constant, and Am the surface of 1 mol of A as a mono layer on the electrode metal B [70]. For the calculation of the thermodynamic functions in (12), a number of models were used in [70] and calculations were performed for Ni-, Cu-, Pd-, Ag-, Pt-, and Au-electrodes and the micro components Hg, Tl, Pb, Bi, and Po, confirming the decisive influence of the choice of the electrode material on the deposition potential. For Pd and Pt, particularly large, positive values of E5o% were calculated, larger than the standard electrode potentials tabulated for these elements. This makes these electrode materials the prime choice for practical applications. An application of the same model to the superheavy elements still needs to be done, but one can anticipate that the preference for Pd and Pt will persist. The latter are metals in which, due to the formation of the metallic bond, almost or completely filled d orbitals are broken up, such that these metals tend in an extreme way towards the formation of intermetallic compounds with sp-metals. The perspective is to make use of the Pd or Pt in form of a tape on which the tracer activities are electrodeposited and the deposition zone is subsequently stepped between pairs of Si detectors for a-spectroscopy and SF measurements. 

Standard electrode potential series - potential, and subentries -> standard potential, and - tabulated standard potentials... 

The electrochemical series tabulates standard electrode potentials. Some sources call the electrochemical series oxi-dation/reduction potentials, electromotive series, and so on. The reference state of electrochemical series is the hydrogen evolution reaction, or H+/H2 reaction. Its standard electrode potential has been universally assigned as 0 V. This electrode is the standard hydrogen electrode (SHE) against which all others are compared. For example, the standard electrode potential of the Fe/Fe2+ reaction is —0.440 V and that of Cu/Cu2+ reaction is +0.337 V. The standard electrode potentials are calculated from Gibbs free energy values by Eq. (8) that is applicable only in the above-mentioned standard state. 

Comprehensive sources for standard electrode potentials include Standard Potentials in Aqueous Solution, A. J. Bard, R. Parsons, and J. Jordan, Eds. New York Marcel Dekker, 1985 G. Milazzo and S. Caroli, Tables of Standard Electrode Potentials. New York Wiley-Interscience, 1977 M. S. Antelman with F. J. Harris. Jr., The Encyclopedia of Chemical Electrode Potentials. New York Plenum Press, 1982. Some compilations are arranged alphabetically by element others are tabulated according to the numerical value oiEf. 

Reference works, particularly those published before 1953, often contain tabulations of electrode potentials that are not in accord with the lUPAC recommendations. For example, in a classic source of standard-potential data complied by Latimer," one finds... 

In the above equation, if a = 1, then E = eP. The standard potential of an electrode eP is the potential of an electrode in contact with a solution of its ions of unit activity. The standard potentials are always expressed against the standard hydrogen electrode (SHE), the potential of which is zero by definition. The standard potentials are a function of temperature they are usually tabulated for 25° C. Standard electrode potential is also called normal electrode potential. 

F is the Faraday constant, K is the equilibrium constant of the reaction, R is the gas constant, and T is the thermodynamic temperature. However, E jj is not the standard potential of the electrode reaction (or sometimes called half-cell reaction), which is tabulated in the tables mentioned. It is the standard potential of the reaction in a chemical cell which is equal to the standard potential of an electrode reaction (abbreviated as standard electrode potential), E when the reaction involves the oxidation of molecular hydrogen to solvated protons... 

The standard potential of an oxidation half-reaction is equal in magnitude but has the opposite sign to the potential of the reverse reduction reaction. Standard half-cell potentials are tabulated as reduction potentials. These are sometimes referred to as standard electrode potentials Therefore,... 

The standard electrode potentials (electromotoric forces, EMK, for different reduction reactions are tabulated in most textbooks on physical chemistry. - ... 

An electrode potential is a measure of the thermodynamics of a redox reaction. It may be expressed as the difference between two half-cell potentials, which by convention are measured against a hydrogen electrode. Tabulated values refer to standard conditions (ions at unit activity). 

Both direct potentiometry and the potentiometric titration method (see next Section) require the measurement of emf between an indicator electrode system and a reference electrode system, the two comprising a cell system. Despite the fact that potentials are usually referred to the standard hydrogen electrode (SHE) for tabulation purposes, they can be expressed relative to any acceptable reference system. In practice, the common reference electrodes are saturated KCl calomel (SCE), 3M KCl calomel or 3M KCl Ag/AgCl. In many present-day applications, the indicator and reference electrodes are combined in a single electrode system called a "combination electrode". 

The numerical value of an electrode potential depends on the nature of the particular chemicals, on the temperature, and on the concentrations of the various members of the couple. For purposes of reference, half-cell potentials are tabulated for standard states of all the chemicals, defined as 1 atm... 

By measuring the cell potential at various concentrations of Cu ", we can determine < cu2+/cu = 0cu2+/cu- This standard potential is tabulated along with the standard potentials of other half-cells in Table 17.1. Such a table of half-cell potentials, or electrode potentials, is equivalent to a table of standard Gibbs energies f rom which we can calculate values of equilibrium constants for chemical reactions in solution. Note that the standard potential is the potential of the electrode when all of the reactive species are present with unit activity, a = 1. 

Since many disciplines now use pe as much as Eh to express electron activity in a system, it is worthwhile to discuss the relationships between these two variables (Lindsay, 1979). Eive decades ago the Swedish chemist Lars Gunnar Sillen suggested that the electrons (e ) can be considered as any other reactant or product in chemical reactions. Sillen and Martell (1964) tabulated equilibrium constants for redox reactions in terms of both E° (standard electrode potentials) and log K (equilibrium activity constants), and encouraged the use of log K to calculate pe values for redox systems. Like pH, the electron activity in a reaction can be defined as... 

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