Oxidation-reduction equations standard cell potential

Since oxidation occurs at the anode and reduction at the cathode, the standard cell potential can be calculated from the standard reduction potentials of the two half-reactions involved in the overall reaction by using the equation ... 

In voltaic cells, it is possible to carry out the oxidation and reduction halfreactions in different places when suitable provision is made for transporting the electrons over a wire from one half-reaction to the other and to transport ions from each half-reaction to the other in order to preserve electrical neutrality. The chemical reaction produces an electric current in the process. Voltaic cells, also called galvanic cells, are introduced in Section 17.1. The tendency for oxidizing agents and reducing agents to react with each other is measured by their standard cell potentials, presented in Section 17.2. In Section 17.3, the Nernst equation is introduced to allow calculation of potentials of cells that are not in their standard states. 

To compensate partially for activity effects and errors resulting from side reactions, such as those described in the previous section. Swift proposed substituting a quantity called the formal potential E" in place of the standard electrode potential in oxidation-reduction caleulations. The formal potential, sometimes referred to as the conditional potential, of a system is the potential of the half-cell with respect to tite SI IE when the concentrations of reactants and products are 1. M and the concentrations of any other constituents of the solution are earefully specilted. Thus, for example, the formal potential for the reductionof iron(lll) is +0.732 V in 1 M perchloric acid and +0.7(K) V in 1 M hydrochloric acid. Using these values in place of the standard electrode potential in the Nernst equation will yield heller agreement between calculated and experi-... 

You are given the half-cell equations and can find standard reduction potentials in Table 20.1. The half-reaction with the lower reduction potential will be an oxidation. With this information, you can write the overall cell reaction, calculate the standard cell potential, and describe the cell in cell notation. 

Plan When an aqueous solution of an ionic compound is electrolyzed, the possible reactants are H2O and the ions of the solute (in this case Ag" " and F ). Because the products are Ag and O2, the reactants must be Ag" " and H2O. Writing the half-reactions for these processes reveals which is oxidation and which is reduction and therefore which occurs at the anode and which at the cathode. The minimum emf is found by using Equation 20.8 to calculate the standard cell potential. 

While the redox titration method is potentiometric, the spectroelectrochemistry method is potentiostatic [99]. In this method, the protein solution is introduced into an optically transparent thin layer electrochemical cell. The potential of the transparent electrode is held constant until the ratio of the oxidized to reduced forms of the protein attains equilibrium, according to the Nemst equation. The oxidation-reduction state of the protein is determined by directly measuring the spectra through the tranparent electrode. In this method, as in the redox titration method, the spectral characterization of redox species is required. A series of potentials are sequentially potentiostated so that different oxidized/reduced ratios are obtained. The data is then adjusted to the Nemst equation in order to calculate the standard redox potential of the proteic species. Errors in redox potentials estimated with this method may be in the order of 3 mV. 

The standard reduction potential for the Sn +/Sn + couple is +0.15 volt that for the Cr +/Cr couple is —0.74 volt. The equation for the reaction shows Cr being oxidized to Cr +, so the sign of the value for the Cr +/Cr couple is reversed. The overall reaction, the sum of the two half-reactions, has a cell potential equal to the sum of the two half-reaction potentials. 

Equation 4.6a is valid only when standard electrode potentials of the redox reactions in both half cells are written as reduction reactions (as shown in Table 4.1). If the reaction in the oxidation cell L is written as an oxidation reaction, the positive E° changes to negative and the negative E° changes to positive. Under these conditions, the following eqnation applies to obtain the electrode potential for the overall reaction ... 

The emf of a cell can be calculated from the standard electrode potentials of the half-reactions. In order to find the emf, we have to look at the two halfreactions involved in the reaction. Then, set up the two half-reactions so that when they are added we will get the net reaction. Once we have set the equations properly and assigned the prpper potentials to those half-reactions, we can add the standard electrode potentials. A common mistake that students make is that they forget the fact that the standard electrode potentials are given in terms of reduction reactions. Redox reactions involve both oxidation and reduction. If one half-reaction is reduction, the other should be oxidation. So we must be careful about the signs of the half-reaction potentials, before we add the two half-reaction potentials to get the emf value. Do the next example. 

The magnitude of the net cell potential AV° will signify the spontaneity of the oxidation-reduction reaction. However, it does not indicate the rate at which corrosion will occur. As noted before, we apply the superscript 0 to denote that we are considering the Standard Electrode Potentials. Engineers may be required to calculate the potential of a particular half-cell at concentrations and temperatures other than the standard conditions. For this purpose, we shall use the Nernst equation, which allows us to account for non-standard temperatures and solution concentrations. 

From the values of the standard reduction potentials, we can see that copper is oxidized, and silver is reduced. We then use equation 13.2 to find the cell potential ... 

In tables of standard potentials, all of the half-cell reactions are expressed as reductions. The sign is reversed if the reaction is reversed to become an oxidation. In a spontaneous reaction, when both half-cells are written as reductions, the half-cell with the more negative potential will be the one that oxidizes. The negative sign in Equation 15.9 reverses one of the reduction processes to an oxidation. Some standard reduction potentials for common half-cells are given in Appendix 15.A. These can be used to calculate for other electrochemical cell combinations as we have done for the Zn/Cu cell. More complete lists of half-cell potentials can be found in references such as Bard et al. listed in the bibliography. 

We first identify the reactant that is reduced and the reactant that is oxidized. We then use the data in Table 19.1 and Appendix D to obtain the standard reduction potentials and hence the cell potential. Finally, we use equation (19.17) to obtain K from E°eH. 

Describe the role of non-fVwork in electrochemical systems. Define the roles of the anode, cathode, and electrolyte in an electrochemical cell. Given shorthand notation for an electrochemical cell, identify the oxidation and reduction reactions. Use data for the standard half-cell potential for reduction reactions, E°, to calculate the standard potential of reaction E. Apply the Nernst equation to determine the potential in an electrochemical cell given a reaction and reactant concentrations. 

The reduction potential of a half-cell depends not only on the chemical species present but also on their activities, approximated by their concentrations. About a century ago, Walther Nemst derived an equation that relates standard reduction potential ( ") to the reduction potential (E) at any concentration of oxidized and reduced species in the cell ... 

The cell reaction for cells without liquid junction can be written as the sum of an oxidation reaction and a reduction reaction, the so-called half-cell reactions. If there are C oxidation reactions, and therefore C reduction reactions, there are C C — 1) possible cells. Not all such cells could be studied because of irreversible phenomena that would take place within the cell. Still, a large number of cells are possible. It is therefore convenient to consider half-cell reactions and to associate a potential with each such reaction or electrode. Because of Equation (12.88), there would be (C - 1) independent potentials. We can thus assign an arbitrary value to the potential associated with one half-cell reaction or electrode. By convention, and for aqueous solutions, the value of zero has been assigned to the hydrogen half-cell when the hydrogen gas and the hydrogen ion are in their standard states, independent both of the temperature and of the pressure on the solution. 

For biological systems this tendency is expressed by the standard reduction potential, Eq, defined as the electromotive force (emf) in volts given by a half-cell in which the reductant and oxidant species are both present at 1.0 M concentration unit activity, at 25°C and pH 7.0 in equilibrium with an electrode which can reversibly accept electrons from the reductant species, according to the equation ... 

In writing the equation this way, we have dropped the subscript cell to indicate that the calculated emf does not necessarily refer to a voltaic cell. Also, we have generalized the standard reduction potentials by using the general terms reduction and oxidation rather than the terms specific to voltaic cells, cathode and anode. We can now make a general statement about the spontaneity of a reaction and its associated emf, E A positive value of E indicates a spontaneous process a negative value of E indicates a nonspontaneous process. We use E to represent the emf under nonstandard conditions and E° to indicate the standard emf. 

We are now in a position to discuss Equation 20.8 more fully. For each of the half-cells in a voltaic cell, the standard reduction potential provides a measure of the driving force for reduction to occur The more positive the value ofE° the greater the driving force for reduction. In any voltaic cell the reaction at the cathode has a more positive value of E°ed than does the reaction at the anode. In essence, the greater driving force of the cathode half-reaction is used to force the anode reaction to occur "in reverse," as an oxidation. 

The Ag half-reaction, with the more positive standard reduction potential, will occur as a reduction, and the Zn half-reaction will occur as an oxidation. Balancing the equation for the overall cell reaction requires multiplying the reduction (the Ag half-reaction) by 2 ... 

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