Standard electrode potentials complex constant

SOURCES The most authoritative source is S. G. Bratsch, J. Phys. Chem. Ref. Data 1989,18, 1. Additional data come from L. G. Sillen and A. E. Martell, Stability Constants of Metal-Ion Complexes (London The Chemical Society, Special Publications Nos. 17 and 25. 1964 and 1971) G. Milazzo and S. Caroti, Tables of Standard Electrode Potentials (New York Wiley, 1978) T. Muss ini,... 

Table 4 Standard Electrode Potentials at 25 °C and Stability Constants for Gold(III) Complexes in Aqueous Solution22,44...

Table 8 Absorption maxima, extinction coefficients, stability constants, and standard electrode potentials for selected Tris-(l,4-diimine)iron(II) complexes in water...

One of the first questions one might ask about forming a metal complex is how strong is the metal ion to ligand binding In other words, what is the equilibrium constant for complex formation A consideration of thermodynamics allows us to quantify this aspect of complex formation and relate it to the electrode potential at which the complex reduces or oxidizes. This will not be the same as the electrode potential of the simple solvated metal ion and will depend on the relative values of the equilibrium constants for forming the oxidized and reduced forms of the complex. The basic thermodynamic equations which are needed here show the relationships between the standard free energy (AG ) of the reaction and the equilibrium constant (K), the heat of reaction, or standard enthalpy (A// ), the standard entropy (AS ) and the standard electrode potential (E for standard reduction of the complex (equations 5.1-5.3). 

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... 

We will use standard electrode potentials throughout the rest of this text to calculate cell potentials and equilibrium constants for redox reactions as well as to calculate data for redox titration curves. You should be aware that such calculations sometimes lead to results that are significantly different from those you would obtain in the laboratory. There are two main sources of these differences (1) the necessity of using concentrations in place of activities in the Nernst equation and (2) failure to take into account other equilibria such as dissociation, association, complex formation, and solvolysis. Measurement of electrode potentials can allow us to investigate these equilibria and determine their equilibrium constants, however. 

The application of standard electrode potential data to many systems of interest in analytical chemistry is further complicated by association, dissociation, complex formation, and solvolysis equilibria involving the species that appear in the Nemst equation. These phenomena can be taken into account only if their existence is known and appropriate equilibrium constants are available. More often than not, neither of these requirements is met and significant discrepancies arise as a consequence. For example, the presence of 1 M hydrochloric acid in the iron(Il)/iron(llI) mixture we have just discussed leads to a measured potential of + 0.70 V in 1 M sulfuric acid, a potential of -I- 0.68 V is observed and in 2 M phosphoric acid, the potential is + 0.46 V. In each of these cases, the iron(II)/iron(III) activity ratio is larger because the complexes of iron(III) with chloride, sulfate, and phosphate ions are more stable than those of iron(II) thus, the ratio of the species concentrations, [Fe ]/[Fe ], in the Nemst equation is greater than unity and the measured potential is less than the standard potential. If fomnation constants for these complexes were available, it would be possible to make appropriate corrections. Unfortunately, such data are often not available, or, if they are, they are not very reliable. 

Stability constant (of a complex) see Formation constant. Standard electrode potential E°) the electrode potential when the concentrations of solutes are 1 M, the gas pressures are 1 atm, and the temperature has a specified value—usually 25°C. (20.5)... 

Like solubility products (page 344), these constants are calculated from standard electrode potentials (Table 22.3), Table 22.4 gives the dissociation constants of several complex ions. 

Voltammetric methods also provide a convenient approach to establish the thermodynamic reversibility of an electrode reaction and for the evaluation of the electron stoichiometry for the electrode reaction. As outlined in earlier sections, the standard electrode potential, the dissociation constants of weak acids and bases, solubility products, and the formation constants of complex ions can be evaluated from polarographic half-wave potentials, if the electrode process is reversible. Furthermore, studies of half-wave potentials as a function of ligand concentration provide the means to determine the formula of a metal complex. 

Formal potential — Symbol Efr (SI Unit V), has been introduced in order to replace the standard potential of -> cell reaction when the values of - activity coefficients of the species involved in the cell reaction are unknown, and therefore concentrations used in the equation expressing the composition dependence of ceii instead of activities. It also involves the activity effect regarding the -+ standard hydrogen electrode, consequently in this way the formal electrode potential is also defined. Formal potentials are similar to conditional (apparent) equilibrium constants (-> equilibrium constant), in that, beside the effect of the activity coefficients, side reaction equilibria are also considered if those are not known or too complex to be taken into account. It follows that when the logarithmic term which contains the ratio of concentrations in the -> Nernst... 

Complex constants and standard potentials of some metal ion complex electrodes with various complex agents are given in Table 3.4. ... 

Starting with the standard potential of the electrode reaction M" + n e = M, it is possible, using the Nernst equation, to calculate the equilibrium constant for complex formation between the ion and a dissolved complexing agent. With the following calculation, the complexation constant for equation (2.93) is determined. 

By using a copper amalgam electrode, which in effect is a copper(I)- and copper(II)-aqua ion electrode, and a pure mercury electrode, which is a redox electrode, IB remeasured and confirmed the gross complexity constant Pa and the fifth consecutive constant for the copper(II)-ammine system and provided, in addition to a number of standard redox potentials, the two step constants of the copper(I)... 

The preceding equations are valid for any ion /. To describe the phase-boundary potential of an electrode for different ions, it is useful to apply one single constant term, Fj, which is defined for the so-called primary ion, I. For another ion, /, the E value changes for two reasons (1) the standard chemical potential, s , is not equal to ej and (2) the free ion activity in the membrane, flj(m), changes owing to the differences in the complex formation constants and the strength of ion pairs. 

Similarly, there are no reliable data on the standard reduction potential for the gold sulfite complex, for which values in the range 0-0.4 were reported [25], However, assuming for the stability constant the previously quoted value of 26.8 [23], the standard reduction potential of the sulfite complex can be calculated as 0.11 V, which is in excellent agreement with the formal potential of the sulfite complex at a mercury amalgam electrode, namely, 0.116 V, as measured by Baltrunas et al. [26]. 

Other Coordination Complexes. Because carbonate and bicarbonate are commonly found under environmental conditions in water, and because carbonate complexes Pu readily in most oxidation states, Pu carbonato complexes have been studied extensively. The reduction potentials vs the standard hydrogen electrode of Pu(VI)/(V) shifts from 0.916 to 0.33 V and the Pu(IV)/(III) potential shifts from 1.48 to -0.50 V in 1 Tf carbonate. These shifts indicate strong carbonate complexation. Electrochemistry, reaction kinetics, and spectroscopy of plutonium carbonates in solution have been reviewed (113). The solubiUty of Pu(IV) in aqueous carbonate solutions has been measured, and the stabiUty constants of hydroxycarbonato complexes have been calculated (Fig. 6b) (90). 

The use of ISEs in non-aqueous media(for a survey see [125,128]) is limited to electrodes with solid or glassy membranes. Even here there are further limitations connected with membrane material dissolution as a result of complexation by the solvent and damage to the membrane matrix or to the cement between the membrane and the electrode body. Silver halide electrodes have been used in methanol, ethanol, n-propanol, /so-propanol and other aliphatic alcohols, dimethylformamide, acetic acid and mixtures with water [40, 81, 121, 128]. The slope of the ISE potential dependence on the logarithm of the activity decreases with decreasing dielectric constant of the medium. With the fluoride ISE, the theoretical slope was found in ethanol-water mixtures [95] and in dimethylsulphoxide [23], and with PbS ISE in alcohols, their mixtures with water, dioxan and dimethylsulphoxide [134]. The standard Gibbs energies for the transfer of ions from water into these media were also determined [27, 30] using ISEs in non-aqueous media. 

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