Nernst equation applications

In potentiometry the potential of an electrochemical cell is measured under static conditions. Because no current, or only a negligible current, flows while measuring a solution s potential, its composition remains unchanged. For this reason, potentiometry is a useful quantitative method. The first quantitative potentiometric applications appeared soon after the formulation, in 1889, of the Nernst equation relating an electrochemical cell s potential to the concentration of electroactive species in the cell. ... 

It must be emphasised that standard electrode potential values relate to an equilibrium condition between the metal electrode and the solution. Potentials determined under, or calculated for, such conditions are often referred to as reversible electrode potentials , and it must be remembered that the Nernst equation is only strictly applicable under such conditions. 

An important application of the Nernst equation is the measurement of concentration. In a concentration cell, the two electrodes are identical except for their concentrations. For such a cell, E° = 0 and at 25°C the potential corresponding to the cell reaction is related to Q by E = —(0.025693 V//z) In Q. For example, a concentration cell having two Ag+/Ag electrodes is... 

An important application of the Nernst equation is the measurement of pH (and, through pFI, acidity constants). The pH of a solution can be measured electro-... 

The Nernst equation is of limited use at low absolute concentrations of the ions. At concentrations of 10 to 10 mol/L and the customary ratios between electrode surface area and electrolyte volume (SIV 10 cm ), the number of ions present in the electric double layer is comparable with that in the bulk electrolyte. Hence, EDL formation is associated with a change in bulk concentration, and the potential will no longer be the equilibrium potential with respect to the original concentration. Moreover, at these concentrations the exchange current densities are greatly reduced, and the potential is readily altered under the influence of extraneous effects. An absolute concentration of the potential-determining substances of 10 to 10 mol/L can be regarded as the limit of application of the Nernst equation. Such a limitation does not exist for low-equilibrium concentrations. 

R is the ideal gas constant, T is the Kelvin temperature, n is the number of electrons transferred, F is Faraday s constant, and Q is the activity quotient. The second form, involving the log Q, is the more useful form. If you know the cell reaction, the concentrations of ions, and the E°ell, then you can calculate the actual cell potential. Another useful application of the Nernst equation is in the calculation of the concentration of one of the reactants from cell potential measurements. Knowing the actual cell potential and the E°ell, allows you to calculate Q, the activity quotient. Knowing Q and all but one of the concentrations, allows you to calculate the unknown concentration. Another application of the Nernst equation is concentration cells. A concentration cell is an electrochemical cell in which the same chemical species are used in both cell compartments, but differing in concentration. Because the half reactions are the same, the E°ell = 0.00 V. Then simply substituting the appropriate concentrations into the activity quotient allows calculation of the actual cell potential. 

Values of E° cannot be used for the last two reactions as given because the H" " concentration (pH) in both cases is not the standard value of 1.00 M. Hence the Nernst equation must be used to ascertain the applicable values of E when [H+] is 10 ° M inthe A1(0H)3 to A1 reaction and 10 ° Minthe Al(OH) reaction. 

The pH values at which the transformations occur may be calculated by use of the equilibrium constant K or the Nernst equation, the latter being applicable to reactions in which the oxidation state changes, that is, reaction equations... 

Application of the equilibrium constant equation along with the value for [Fe+ ] and then the Nernst equation containing values for E and E° and the concentration changes (0.10 M for all other soluble species) will give accurate pH values for comparison with the estimates from the diagram. The estimated values from the diagram are again presented in brackets. 

Tower, Stephen. All About Electrochemistry. Available online. URL http //www.cheml.com/acad/webtext/elchem/. Accessed May 28, 2009. Part of a virtual chemistry textbook, this excellent resource explains the basics of electrochemistry, which is important in understanding how fuel cells work. Discussions include galvanic cells and electrodes, cell potentials and thermodynamics, the Nernst equation and its applications, batteries and fuel cells, electrochemical corrosion, and electrolytic cells and electrolysis. 

In the case of gases, properties may be tabulated til terms of their existence at 0°C and 760 mm pressure, To determine the volume of a gas at some different temperature and pressure, corrections derived from known relationships (Charles , Amonton s. Gay-Lussac s, and other laws) must be applied as appropriate. In tile case of pH values given at some measured value (standard for comparison), the same situation applies. Commonly, lists of pH values are based upon measurements taken at 25°C. The pH of pure water at 22°C is 7.00 at 25,JC, 6.998 and at 100°C. 6.13. Modern pH instruments compensate for temperature differences through application of the Nernst equation. 

The physical situation in the solution adjacent to the electrode during the potential scan is illustrated by the concentration-distance profiles included in Figure 3.18 for selected potentials [38]. The C-x profiles in Figure 3.18A are for O and R when Ej is imposed. Note that the application of Ej does not measurably alter the concentration of O at the electrode surface as compared to the solution bulk. As the potential is scanned positively, the concentration of O at the electrode surface decreases in order to establish a Cr/Cq ratio that satisfies the Nernst equation for the applied potential at any particular instant. This is illustrated by profiles in B-D. Note that the profile in B (for which the concentration of O at the electrode surface equals the concentration of R) corresponds to an Eappljed that is the formal electrode potential (vs. SCE) of the couple. 

Figure 3.37C shows the profiles that result when the applied potential is sufficiently negative that the concentration of O at the electrode surface is effectively zero. In this case, essentially all of O at the electrode surface must be electrolyzed to R in order to satisfy the Nernst equation. Consequently, O is converted to R as rapidly as it can diffuse to the electrode surface. Since this is the limiting condition, application of even more negative potentials causes no measurable change in the profiles. 

The shape of the potential-time response is determined by the concentration changes of O and R at the electrode surface during electrolysis. The potential is related to Cq/Cr via the Nernst equation for a reversible system. The initial potential before current application is simply the rest potential or open-circuit potential (E ) of the solution, which reflects the initial Cq/C in solution. At the instant of current application, this ratio becomes finite and the potential changes to a value consistent with the Nernst equation. Early in the... 

One important application of the Nernst equation is the measurement of pH (and, through pH, acidity constants). The pH of a solution can be measured electrochemically with a device called a pH meter. The technique makes use of a cell in which one electrode is sensitive to the H30+ concentration and the second electrode serves as a reference. An electrode sensitive to the concentration of a particular ion is called an ion-selective electrode. One combination is a hydrogen electrode connected through a salt bridge to a calomel electrode. The reduction half-reaction for the calomel electrode is... 

The electrochemical determination of pH using a pH meter is a particularly important application of the Nernst equation. Consider, for example, a cell with a hydrogen electrode as the anode and a second reference electrode as the cathode ... 

In the context of RTILs the criterion (3) raises considerable problems since the concept of activity and activity coefficients of ions is largely unexplored in such media. Accordingly, validation of the applicability of the Nernst equation in such media is a non-simple exercise, given that RTILs are likely to exhibit gross non-ideality. Rather, electrochemical measurements based on otherwise validated reference electrodes, may likely in the future provide a methodology for the study of RTIL non-ideality. 

This type of reasoning is applicable to any cell and always leads to the corresponding Nernst equation. 

The foregoing discussion was applicable to redox potentials, and the concentrations of reactants and products were assumed to be 1 M each. If it is not so, a correction must be made as was the case with AGq. Such a correction may be made via the Nernst equation, which for the reaction A + B —> C + D can be written as... 

The knowledge of the surface potential for the dispersed systems, such as metal oxide-aqueous electrolyte solution, is based on the model calculations or approximations derived from zeta potential measurements. The direct measurement of this potential with application of field-effect transistor (MOSFET) was proposed by Schenk [108]. These measurements showed that potential is changing far less, then the potential calculated from the Nernst equation with changes of the pH by unit. On the other hand, the pHpzc value obtained for this system, happened to be unexpectedly high for Si02. These experiments ought to be treated cautiously, as the flat structure of the transistor surface differs much from the structure of the surface of dispersed particle. The next problem may be caused by possible contaminants and the surface property changes made by their presence. 

The correct answer is (B). This problem requires the use of the Nernst equation (or at least its application). However, before you can solve the equation, you need to know which electrode is the cathode and which is the anode. Remember, the substance with the most negative reduction potential is the most easily oxidized, making it the anode. In this problem, lead is the anode and cobalt the cathode. The emf of the second cell can be determined by ... 

The correct answer is (A). This question is a conceptual application of the Nernst Equation. The Nernst equation is used to determine the impact of differences in concentration on cell emf. The equation takes the form ... 

The Nernst equation is applicable also to interfacial ionic equilibria where the activity of an ion is equal to a (1) at one side of the interface and a (2) at the other side, as follows ... 

Nernstian electrode reaction — An electrode reaction proceeding at electrochemical equilibrium. Therefore the -> Nernst equation is applicable. See - reversibility. 

Tracer methods — [i] The application of radiotracer methods in electrochemistry dates back to the pioneering works by Hevesy in 1914. The aim of these studies was to demonstrate that isotopic elements can replace each other in both -> electrodeposition and equilibrium processes (Nernst law -> Nernst equation). Nevertheless, Joliot s fundamental work in 1930 is considered by electrochemists as a landmark in the application of -> radiochemical (nuclear) methods in electrochemistry. 

Of course, the solution pH is a major factor that can be tuned to favor hydrogen reduction or water oxidation. Application of the Nernst equation for the two manifolds yields expressions 18-21 ... 

Because the potential of an electrochemical cell depends on the concentrations of the participating ions, the observed potential can be used as a sensitive method for measuring ion concentrations in solution. We have already mentioned the ion-selective electrodes that work by this principle. Another application of the relationship between cell potential and concentration is the determination of equilibrium constants for reactions that are not redox reactions. For example, consider a modified version of the silver concentration cell shown in Fig. 11.11. If the 0.10 M AgN03 solution in the left-hand compartment is replaced by 1.0 M NaCl and an excess of solid AgCl is added to the cell, the observed cell potential can be used to determine the concentration of Ag+ in equilibrium with the AgCl(s). In other words, at 25°C we can write the Nernst equation as... 

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