Oxidation-reduction reactions Nernst equation

Most analytical oxidation/reduction reactions are carried out in solutions that have such high ionic strengths that activity coefficients cannot be obtained via the Debye-Hiickel equation (see Equation 10-1, Section lOB-2). Significant errors may result, however, if concentrations are used in the Nernst equation rather than activities. For example, the standard potential for the half-reaction... 

The Nemst equation was given before (Eq. 10.115), and in this chapter the effect of pH on the reduction potential of the hydrogen ion has been mentioned, but the effect in general should be emphasized. There are several types of reactions in which concentrations of the reactants and products affect the stability of various oxidation stales. This can be understood through application of the Nernst equation. The reduction potential of hydrogen will vary with the concentration of the hydrogen ion hence the commonly known fact that many reasonably active metals dissolve in acid but not in base. 

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. 

It has already been seen that whole oxidation-reduction reactions can be constructed from half-reactions. The direction in which a reaction goes is a function of the relative tendencies of its constituent half-reactions to go to the right or left These tendencies, in turn, depend upon the concentrations of the half-reaction reactants and products and their relative tendencies to gain or lose electrons. The latter is expressed by a standard electrode potential The tendency of the whole reaction to proceed to the right as written is calculated from the Nernst equation, which contains both EP and the concentrations of the reaction participants. These concepts are explained further in this section and the following section. 

The difference between the potential actually required to initiate an oxidation or reduction reaction, and the potential predicted by the Nernst equation. 

The reducing equivalents transferred can be considered either as hydrogen atoms or electrons. The driving force for the reaction, E, is the reduction/oxidation (redox) potential, and can be measured by electrochemistry it is often expressed in millivolts. The number of reducing equivalents transferred is n. The redox potential of a compound A depends on the concentrations of the oxidized and reduced species [Aqx] and [Area] according to the Nernst equation ... 

The equilibrium constant is used for the first of these reactions since no electrons are involved in the equation, that is, no change of oxidation state is occurring. But since there are electrons in the second reaction, the Nernst equation which handles reduction reactions must be used. 

You then use the Nernst equation to calculate the EMF, so you have to determine the values of Q and n. You can find n by examining the oxidation and reduction half-reactions presented earlier in this chapter, which indicate that two electrons are exchanged in the process. Q is expressed as follows (if you can t recall how to calculate reaction quotients, see Chapter 14) ... 

The standard reversible potential is that listed in the EMF series of Table 1 and represents a special case of the Nernst equation in which the second term is zero. The influence of the solution composition manifests itself through the logarithmic term. The ratio of activities of the products and reactants influences the potential above which the reaction is thermodynamically favorted toward oxidation (and conversely, below which reduction is favored). By convention, all solids are considered to be at unit activity. Activities of gases are equal to their fugacity (or less strictly, their partial pressure). 

Similarly, from the Tables 1.17 and 1.18 we can see, for example, that permanganate ions (in acid medium) can oxidize chloride, bromide, iodide, iron(II), and hexacyanoferrate(II) ions, also that iron(III) ions may oxidize arsenite or iodide ions but never chromium(III) or chloride ions etc. It must be emphasized that the standard potentials are to be used only as a rough guide the direction of a reaction will depend on the actual values of oxidation-reduction potentials. These, if the concentrations of the species are known, can be calculated easily by means of the Nernst 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. 

As the potential is increased, there is a point at which no equilibrium state is reached, but instead, an appreciable steady current flows which will obey Ohm s law over a reasonable range of applied potential. The potential at which this steady current is observed is called the decomposition potential because it is accompanied by chemical reaction (electrolysis) at the electrode surfaces. These electrode reactions are quite generally the oxidation (anode) and reduction (cathode) of ionic or molecular species present in the solution. If the reactions at the electrodes are reversible, then the decomposition potential Ed is related by the Nernst equation to the free energy changes of the electrode reactions... 

A very important point in electrochemistry is that the exchange of electrons between an electrode and an reactant, that is, an oxidant or reductant, can only take place at the electrode surface. In the Nernst equation, this fact might seem to be partly obscured by the use of equilibrium concentrations, but as no net reaction occurs at equilibrium, the surface concentrations must be equal to the equilibrium concentrations. In most applications of electrochemistry, the electrode potential is varied, to achieve surface concentrations Co(0, t) that are different from the concentrations found far from the electrode surface, that is, in the bulk (C ). In other words, the electrode potential is used for creating a nonequilibrium situation where the dynamic response of the chemical system can be examined. 

The reduction potential of a half-reaction in which the oxidized and reduced forms of the substance are present at nonstandard concentrations may be calculated from the following expression, called the Nernst equation. 

Recall that the Nernst equation is the mathematical model describing the relationship between cell potential and concentrations and is readily derived from the fact that cell potential shows a concentration dependence due to its relationship to free energy, equations (A.2.2) and (A.2.3), where Q is the concentration ratio of oxidized (e.g., [Fe3+L ]) to reduced (e.g., [Fe2+L ]) species. In our system, E° is the reduction potential for the one electron transfer half reaction [equation (5.4.4)]. 

Ox refers to the oxidized species and Red to the reduced species x and y are their coefficients, respectively, in the balanced equation. The Nernst equation for any cathode half-cell reduction half-reaction) is... 

Fig. 6 (C) shows the redox titration curve consisting of data points from triplet signals produced in samples poised at various potentials at pH 11. Again, the solid curve is a fit for the Nernst equation based on a one-electron change. The empty-circled data points are taken from the reductive titrations, and several data points (solid-dots) are shown for the reverse oxidative titration. All points coincide reasonably well with the theoretical curve, confirming that the redox reaction is reversible.The redox potential of estimated from the titration curve is -604 mV, very close to the value derived from the attenuated absorbance-change measurements by Klimov etal.. ...

Fig. 6. (A) Plot of amplitude of light-induced pheophytin-reduction signal vs. ambient potential of the medium. (B) Effed of ambient redox potential on the extent of light-induced PS-II reaction-center triplet signal in pea chloroplast particles. (C) Plot of the extent of the light-induced triplet EPR signal in (B) vs. redox potential. Open and closed circles are for reductive and oxidative titrations, respedively. The solid curve is a computer fit of the Nernst equation with n=1 and E , was estimated to be -604 mV. Figure source (A) Klimov. Allakhverdiev. Demeter and Krasnovsky (1979) Photoreduction ofpheophytin in photosystem 2 ofchloroplasts with respect to the redox potential of the medium. DokI Akad NaukSSSR 249 229 (B and C) Rutherford, Mullet and Crofts (1981) Measurement of the midpoint potential of the pheophytin acceptor of photosystem II. FEBS Lett 123 236,237...

Plan The standard half-cell reactions are identical, so °eii is zero, and we calculate Eceii from the Nernst equation. Because half-cell A has a higher [Ag ], Ag ions will be reduced and plate out on electrode A. In half-cell B, Ag will be oxidized and Ag ions will enter the solution. As in all voltaic cells, reduction occurs at the cathode, which is positive. 

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 ve heard electrochemistry of corrosion as a lecture I shouldn t spend much time on it but I d like to describe some electrochemical effects for film formers. First the general principles. If you put a good electronic conductor (a metal) in an aqueous solution, you will typically find that an electrical potential is developed between the piece of conductor and the solution. When ions of the metal enter the solution and leave extra electrons behind a negative potential is developed. All oxidation reactions occurring on the surface are expected to produce this result. Similarly, reduction reactions that use electrons from the metal are expected to produce a more positive potential in the metal. The solution potential of the metal influences the rate of an electrochemical half-cell reaction in accordance with Le Chatelier s Principle, so it is possible to predict through the use of the Nernst Equation the potential that will exist when the only significantly rapid reactions are the oxidation and reduction parts of a reversible reaction. When more than one potentially reversible process occurs, the rate of oxidation will be expected to exceed the rate of reduction for at least one and the converse for at least one. At... 

Reactants and products are not typically in their standard states. The Nernst equation accounts for the effects of activity on reversible potential. Equation (4) presents the free energy change for a generalized reaction in which the reactants and products are not in their standard states. If the reaction is a reduction reaction (the reactants and products are the oxidized and reduced form of a species, respectively), we can consider that AG = —nFE and AG = —nFE. Substituting into Eq. (4) ... 

The Nernst equation gives us the very important relationship between E, the emf of the halfcell, and the concentration of the oxidized and reduced forms of the components of the solution. Measurements using pH electrodes and ISEs are based on this relationship. If the two electrode reactions are written as reductions, the potentials can be calculated for the cathode and anode (or the electrode inserted into the positive terminal of the potentiometer, E+ and the electrode inserted into the negative terminal of the potentiometer, EJ. The cell voltage is the difference... 

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