Applications of Standard Electrode Potentials

This composite satellite image displays areas on the surface of the Earth where chlorophyll-bearing plants are located. Chlorophyll, which is one of nature s most important biomolecules, is a member of a class of compounds called porphyrins. This Glass also includes hemoglobin and cytochrome c, which is discussed in Feature 19-1. Many analytical techniques have been used to measure the chemical and physical properties of chlorophyll to explore Its role in photosynthesis. The redox titration of chlorophyll with other standard redox couples reveals the oxidation/ reduction properties of the molecule that help explain the photophysics of the complex process that green plants use to oxidize water to molecular oxygen. 

We can use standard electrode potentials and the Nernst equation to calculate the potential obtainable from a galvanic cell or the potential required to operate an electrolytic cell. The calculated potentials (sometimes called thermodynamic potentials) are theoretical in the sense that they refer to cells in which there is no current. As we show in Chapter 22, additional factors must be taken into account if a current is involved. 

The thermodynamic potential of an electrochemical cell is the difference between the electrode potential of the right-hand electrode and the electrode potential of the left-hand electrode. That is. 

Gustav Robert Kirchhoff (1824—1877) was a German physicist who made many important contributions to physics and chemistry. In addition to his work in spectroscopy, he is known for Kirchhoff s laws of cuirent and voltage in electrical circuits. These laws can be summarized by the following equations 2/ = 0 and 2 = 0. These equations state that the sum of the currents into any circuit point (node) is zero and the sum of the potential differences around any circuit loop is zero. 

Calculate the thermodynamic potential of the following cell and the free energy change associated with the cell reaction. 

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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... 

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. 

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... 

The application of standard electrode potentials is further complicated by solvation, dissociation, association. and complex-formation reactions involving the species of interest. An example of this problem is the behavior of the potential of the iron(lll)-iron(ll) couple. As noted earlier, an equimoiar mixture of these two ions in 1 M perchloric acid has an electrode potential of -f O.liZ V, Substituting hydroehloric acid of the same concentration for perchioric acid alters the observed potential to +0.700 V, and we observe a value of + 0,600 V in I M phosphoric acid, These changes in potential occur because iron(III) forms more stable complexes with chloride and phosphate ions than does iron(ll), As a result, the actual concentration of iincomplexed iron(III) in such solutions is less than that of uncomplexeci iron(U), and the net effect is a shift in the observed potential. 

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