Potential tabulated standard potentials

The absolute potential for a half-reaction cannot be determined, but the potential for a half-reaction relative to another half-reaction can be determined the tabulated standard potentials for half-reactions are relative to the H2/H30+ half-reaction, which is arbitrarily assigned a standard potential of zero. 

Tabulated standard potentials — The - standard potential of various redox couples have been tabulated by using either the values obtained in electrochemical experiments or calculated from thermochemical data. The electrochemical determination of these values based on -> emf measurements is possible provided that -> diffusion potentials and thermo emf s are elimi-... 

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. 

The tabulated standard potentials are 0.799 V for the Ag Ag" " electrode and 0.222 V for the silver silver chloride electrode. It follows that is equal to 1.76 xlO. Since a saturated aqueous solution of AgCl is very dilute, this result may be used on either the concentration or the molality scale, and the activity coefficients set equal to one. Thus,... 

See Skill 11.3 for information on cell potentials. The standard potential of an oxidation half-reaction E 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 E°. Therefore,... 

Standard potentials are tabulated except when a solution composition is stated the latter are formal potentials and the concentrations are in mol/liter. 

The quantity RT/F is readily evaluated at 25°C, the temperature at which standard potentials are tabulated. 

E° values have been measured for many reactions and tabulated as standard half-cell potentials. Table 9.3 summarizes half-cell potentials as standard reduction potentials for a select set of reactions.aa In the tabulations, E° for... 

In addition to defined standard conditions and a reference potential, tabulated half-reactions have a defined reference direction. As the double arrow in the previous equation indicates, E ° values for half-reactions refer to electrode equilibria. Just as the value of an equilibrium constant depends on the direction in which the equilibrium reaction is written, the values of S ° depend on whether electrons are reactants or products. For half-reactions, the conventional reference direction is reduction, with electrons always appearing as reactants. Thus, each tabulated E ° value for a half-reaction is a standard reduction potential. 

Tabulated standard reduction potentials allow us to determine the potential of any cell under standard conditions. This net standard cell potential is obtained by subtracting the more negative standard reduction potential from the more positive standard reduction potential, giving a positive overall potential. 

Equation expresses an important link between two standard quantities. The equation lets us calculate standard electrical potentials from tabulated values for standard free energies. Equally important, accurate potential measurements on galvanic cells yield experimental values for standard potentials that can be used to calculate standard free energy changes for reactions. 

Use tabulated standard reduction potentials to determine for the following redox reaction ... 

Here, R is the gas constant, 7k is absolute temperature, and XB is the mole fraction of B in the solution phase. Using this equation, we can calculate the equilibrium point of reactions in ideal systems directly from tabulated values of standard potentials p°. 

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... 

Table 3 provides entropies for those species that are needed to determine the temperature dependence of standard potentials for alkali metal redox couples. AU of these entropies were obtained from values published by NIST [11]. The resulting temperature dependences agree well with values tabulated by Bratsch [17]. 

For dissociative electron transfer, an analogous thermochemical cycle can be derived (Scheme 2). In this case the standard potential includes a contribution from the bond fragmentation. Using equations (40) and (41) one can derive another useful expression for BDFEab-, equation (42). While direct electrochemical measurements on solutions may provide b. b, for example, of phenoxides and thiophenoxides (Section 4), the corresponding values for alkoxyl radicals are not as easily determined. Consequently, these values must be determined from a more circuitous thermochemical cycle (Scheme 3), using equation (43). The values of E°h+/h io a number of common solvents are tabulated elsewhere. Values of pKa in organic solvents are available from different sources. " A comparison of some estimated E° values with those determined by convolution voltammetry can be found in Section 3. 

The oxidation [Zn (s) — Zn2+ (aq) + 2e ] and reduction [Cu2+ (aq) + 2e —> Cu (s)] half-reactions are confined to the left (anode) and right (cathode) chambers, respectively, separated by a salt bridge (allowing passage of S04- counterions to preserve electroneutrality). Use the Nemst equation and tabulated standard oxidation potentials... 

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. 

In electrochemistry we make it a rule that the standard chemical potential ju. of hydrogen ions is set zero as the level of reference for the chemical potentials of all other hydrated ions. The standard chemical potentials of various hydrated ions tabulated in electrochemical handbooks are thus relative to the standard chemical potential of hydrogen ions at unit activity in aqueous solutions. Table 9.3 shows the numerical values of the standard chemical potential, the standard partial molar enthalpy h°, and the standard partial molar entropy. 5 ,° for a few of hydrated ions. 

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. 

As stated in the previous chapter, to determine the reversible potential of any electrode in an arbitrary state, it is first of all necessary to know its standard potential. The required values of these potentials, stated in terms of the hydrogen scale and valid for a temperature of 25 °C, arc tabulated. Such data do not express the absolute potentials but the electromotive force of the combination of the given half coll and the standard hydrogen electrode. This fact must be remembered, when making calculations based on these potentials. 

Electrochemical electromotive 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. 

Half-cell standard potentials at 25 °C with respect to the SHE are tabulated [17-19], The procedure explained above to determine the open-circuit potential of a cell can be used to calculate the half-cell open-circuit potential with respect to a SHE. Once again, the first step is to represent the schematic of the cells. 

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... 

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