From Measurements of Cell Potentials

It is possible to select a cell that contains a weak acid in solution whose potential depends on the ion concentrations in the solution and hence on the dissociation constant of the acid. As an example, we will consider acetic acid in a cell that contains a hydrogen electrode and a silver-silver chloride electrode  

As the molality m[j+ depends on the acetic acid equilibrium, which we can indicate in a simplified notation by the equation 

All terms on the left in Equation (20.10) are known from previous experiments (see Section 19.2 for the determination of °) or can be calculated from the composition of the solution in the cell. Thus [see Equation (20.6)], 

Generally, mn+ C m2 or m3, so it can be estimated from Equation (20.8) by inserting an approximate value of K and neglecting the activity coefficients. Thus, it is possible to obtain tentative values of — (RT/S ) In fC and hence K at various concentrations of acetic acid, sodium acetate, and sodium chloride, respectively. The ionic strength /  

It can be observed from the limiting behavior of activity coefficients [Equation (19.11)] that 

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The method of determination from measurements of cell potentials depends on the possibility of carrying out a transformation reversibly in an electrical cell. (See Fig. 7.2.) In this case, the spontaneous tendency of the transformation wUl be opposed by an opposing potential just sufficient to balance the potential in the electrical ceU produced by that spontaneous tendency. The potential observed under such circumstances is related to the change of the Gibbs function for the reaction by Equation (7.84)... 

Dickson et al. [5], calculated the Gibbs function for the ionization of the bisulfate ion by measurement of cell potentials in the temperature range from 50° to 250°C. They found that the Gibbs function could be represented by the equation... 

Ozkaya (76) studied conceptual difficulties experienced by prospective teachers in a number of electrochemical concepts, namely half-cell potential, cell potential, and chemical and electrochemical equilibrium in galvanic cells. The study identified common misconceptions among student teachers from different countries and different levels of electrochemistry. Misconceptions were also identified in relation to chemical equilibrium, electrochemical equilibrium, and the instrumental requirements for die measurement of cell potentials. Learning difficulties were attributed mainly to failure of students to acquire adequate conceptual understanding, and the insufficient explanation of the relevant... 

In addition, Chattopadhyay and coworkers [75CHA/KAR] determined Gibbs energies of reaction from measurements of galvanic potentials for cells of the type... 

When dealing with potentials of electrodes and cells we can define a cell potential as the amount of work needed to transfer a unit charge from one electrode to the other. Let us now deal with the ways by which these cell potentials can be measured. Consider first the measurement of cell potential using Fig. 3. Here there are two electrodes and M2 dipping into the solutions Si and S2, respectively. The electrodes are connected through copper wires Cu to a resistance R, By a suitable calorimetric device the amount of heat generated at i in a unit time can be measured. 

This equation is a reasonable model of electrokinetic behavior, although for theoretical studies many possible corrections must be considered. Correction must always be made for electrokinetic effects at the wall of the cell, since this wall also carries a double layer. There are corrections for the motion of solvated ions through the medium, surface and bulk conductivity of the particles, nonspherical shape of the particles, etc. The parameter zeta, determined by measuring the particle velocity and substituting in the above equation, is a measure of the potential at the so-called surface of shear, ie, the surface dividing the moving particle and its adherent layer of solution from the stationary bulk of the solution. This surface of shear ties at an indeterrninate distance from the tme particle surface. Thus, the measured zeta potential can be related only semiquantitatively to the curves of Figure 3. 

The determination of the real energies of solvation from measurements of the voltaic cells (Section VI) makes it possible to find the absolute electrode potentials in nonaqueous solvents owing to the relation... 

Consequently, a wealth of information on the energetics of electron transfer for individual redox couples ("half-reactions") can be extracted from measurements of reversible cell potentials and electrochemical rate constant-overpotential relationships, both studied as a function of temperature. Such electrochemical measurements can, therefore, provide information on the contributions of each redox couple to the energetics of the bimolecular homogeneous reactions which is unobtainable from ordinary chemical thermodynamic and kinetic measurements. 

The de Broglie wavelength of thermal He atoms is comparable with the interatomic distances of surfaces and adsorbed layers. Thus, from measurements of the angular positions of the diffraction peaks the size and orientation of the 2D unit cell, i.e. the structure of the outermost layer, can be straightforwardly determined. Analysis of the peak intensities yields the potential corrugation, which usually reflects the geometrical arrangement of the atoms within the 2D unit celP. 

Ion-selective microelectrodes [18, 70,71, 164] are chiefly used for measurement of ion activities in individual cells and in intracellular liquid. They were developed from micropipettes, which are miniature liquid bridges used for measurement of cell membrane potentials [94]. Micropipettes and ion-selective microelectrodes are made using commercial drawing devices. Ion-selective... 

Figure 6-3 Device for measurement of electrode potentials. The electrode reactions are indicated below each half-cell. The maximum electrical work that can be done by such a cell on its surroundings is - AG = nEF, where E = V2-V1as measured by a potentiometer. If A is reduced to AH2 by H2/ electrons will flow through an external circuit as indicated. A will be reduced in the right-hand cell. H2 will be oxidized to H+ in the left-hand cell. Protons will flow through the gel bridge from left to right as one of the current carriers in the internal circuit.

With an understanding of the meaning and measurement of the difference of electrical potential, we can develop the thermodynamics of a galvanic cell. We choose a specific cell, but one in which many of the principles related to the obtaining of thermodynamic data from measurement of the electromotive forces (emf) of the cell are illustrated. The specific cell is depicted as... 

Hematologic Studies Hematologic study includes estimations of hemoglobin content, packed-cell volume, total erythrocytes, total leukocytes, platelets, or other measures of clotting potential. These should be performed on blood samples collected from all nonrodents, from 10 rats of both genders, from all groups at 3 months, 6 months, thereafter at approximately 6-month intervals, and at termination. If possible, these collections should be from the same rats of each interval. In additions, a pretest sample should be collected from nonrodents. 

Electrophysiology — Electrophysiology is the study of the electrochemical phenomena associated with biological cells and tissues in animals, plants, bacteria, and insects. It involves measurements of electrical -> potentials or - current on a wide variety of scales from single ion channel, to whole tissues like the heart, muscles, phloem,... 

V, then measure the potential of another half-cell, such as the copper halfcell, with that half-cell. The entire potential of this cell can then be assigned to the other (copper) half-cell, because the potential of the hydrogen half-cell is zero. Of course, we can do the same thing for every other half-cell, but we need not do so, since the hydrogen/hydrogen ion half-cell is hard to work with because it involves a gas. We can get an unknown half-cell potential from its cell potential with any half-cell of known potential. For example, once we get the copper half-cell potential, we can use it to calculate the unknown zinc halfcell potential from the Daniell cell potential. A collection of half-cell potentials, all written as reductions, is presented as Table 17.2. 

The standard potential of the silver-silver bromide electrode has been determined from emf measurements of cells with hydrogen electrodes and silver-silver bromide electrodes in solutions of hydrogen bromide in mixtures of water and N-methylacetamide (NMA). The mole fractions of NMA in the mixed solvents were 0.06, 0.15, 0.25, and 0.50, and the dielectric constants varied from 87 to 110 at 25°C. The molality of HBr covered the range 0.01-0.1 mol kg 1. Data for the mixed solvents were obtained at nine temperatures from 5° to 45°C. The results were used to derive the standard emf of the cell as well as the mean ionic activity coefficients and standard thermodynamic constants for HBr. The information obtained sheds some light on the nature of ion-ion and ion-solvent interactions in this system of high dielectric constant. 

The measured potential is a sum of several potential differences. When a current is made to flow through the cell, these potential differences are affected to different degrees, and the change in cell potential reflects the sum of all these changes. If we consider one of the examples shown in Fig. 3B we can express the change of cell potential resulting from an applied current i as follows ... 

An operational definition endorsed by the International Union of Pure and Applied Chemistry (lUPAC) and based on the work of Bates determines pH relative to that of a standard buffer (where pH has been estimated in terms of p"H) from measurements on cells with liquid junctions the NBS (National Bureau of Standards) pH scale. This operational pH is not rigorously identical to p H defined in equation 30 because liquid junction potentials and single ion activities cannot be evaluated without nonthermodynamic assumptions. In dilute solutions of simple electrolytes (ionic strength, I < 0.1) the measured pH corresponds to within 0.02 to p H. Measurement of pH by emf methods is discussed in Chapter 8. 

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