Cell potentials, their measurement

This review is arranged as follows. After a short review of the basic definitions and significance of the various potentials which are assumed to exist at free surfaces and interfaces, the nature and most important features of voltaic cells, including their measurement techniques, are... 

Compare this equation with Eqs. (15.7) and (15.15). By convention, the reference electrode is connected to the negative terminal of the potentiometer (the readout device). The common reference electrodes used in potentiometry are the SCE and the silver/silver chloride electrode, which have been described. Their potentials are fixed and known over a wide temperature range. Some values for these electrode potentials are given in Table 15.3. The total cell potential is measured experimentally, the reference potential is known, and therefore the variable indicator electrode potential can be calculated and related to the concentration of the analyte through the Nemst equation. In practice, the concentration of the unknown analyte is determined after calibration of the potentiometer with suitable standard solutions. The choice of reference electrode depends on the application. For example, the Ag/AgCl electrode cannot be used in solutions containing species such as halides or sulfides that will precipitate or otherwise react with silver. 

Compare this equation with Equations 15.7 and 15.15. By convention, the reference electrode is connected to the negative terminal of the potentiometer (the readout device). The common reference electrodes used in potentiometry are the SCE and the silver/silver chloride electrode, which have been described. Their potentials are fixed and known over a wide temperature range. Some values for these electrode potentials are given in Table 15.3. The total cell potential is measured... 

Each half-cell reaction has a specific standard potential reported as the potential of the reduction reaction vs. the normal hydrogen electrode (NHE). In an elecdochemical cell, there is a half-cell corresponding to the working electrode (WE), where the reactions under study take place, and a reference half-cell. Experimentally the cell potential is measured as the difference between the potentials of the WE half-cell and the reference electrode/ref-erence half-cell (see Chapter 4). The archetypal reference electrode is the NHE, also known as the standard hydrogen electrode (SHE) and is defined, by convention, as 0.000 V for any temperature. Although the NHE is not typically encountered due to difficulty of operation, all conventional electrodes are in turn referenced to this standard to define their absolute potential (i.e., the Ag/AgCl, 3 M KCl reference has a potential of 203 mV vs. the NHE). In practice, experimental results are either stated as being obtained vs. a specific reference electrode, or converted to potentials vs. NHE. 

Accurate control of potential, stability, frequency response and uniform current distribution required the following low resistance of the cell and reference electrode small stray capacitances small working electrode area small solution resistance between specimen and point at which potential is measured and a symmetrical electrode arrangement. Their design appears to have eliminated the need for the usual Luggin capillary probe. 

We would like to measure the contribution each half-reaction makes to the voltage of a cell. Yet every cell involves two half-reactions and every cell voltage measures a difference between their half-cell potentials. We can never isolate one half-reaction to measure its E°. An easy escape is to assign an arbitrary value to the potential of some selected half-reaction. Then we can combine all other half-reactions in turn with this reference half-reaction and find values for them relative to our reference. The handiest arbitrary value to assign is zero and chemists have decided to give it to the half-reaction... 

Since concentration variations have measurable effects on the cell voltage, a measured voltage cannot be interpreted unless the cell concentrations are specified. Because of this, chemists introduce the idea of standard-state. The standard state for gases is taken as a pressure of one atmosphere at 25°C the standard state for ions is taken as a concentration of 1 M and the standard state of pure substances is taken as the pure substances themselves as they exist at 25°C. The half-cell potential associated with a halfreaction taking place between substances in their standard states is called ° (the superscript zero means standard state). We can rewrite equation (37) to include the specifications of the standard states ... 

Reliable markers are needed to monitor the activity of antiangiogenic dtugs. Circulating endothelial cells and their progenitor subset are a potential candidate, as is MRI dynamic measurement of vascular permeability/ flow in response to angiogenesis inhibitors, but neither has been clinically validated. 

In the discussion of the Daniell cell, we indicated that this cell produces a voltage of 1.10 V. This voltage is really the difference in potential between the two half-cells. The cell potential (really the half-cell potentials) is dependent upon concentration and temperature, but initially we ll simply look at the half-cell potentials at the standard state of 298 K (25°C) and all components in their standard states (1M concentration of all solutions, 1 atm pressure for any gases and pure solid electrodes). Half-cell potentials appear in tables as the reduction potentials, that is, the potentials associated with the reduction reaction. We define the hydrogen half-reaction (2H+(aq) + 2e - H2(g)) as the standard and has been given a value of exactly 0.00 V. We measure all the other half-reactions relative to it some are positive and some are negative. Find the table of standard reduction potentials in your textbook. 

In this investigation, you will build some galvanic cells and measure their cell potentials. 

In section 11.1, you learned that a cell potential is the difference between the potential energies at the anode and the cathode of a cell. In other words, a cell potential is the difference between the potentials of two half-cells. You cannot measure the potential of one half-cell, because a single half-reaction cannot occur alone. However, you can use measured cell potentials to construct tables of half-cell potentials. A table of standard half-cell potentials allows you to calculate cell potentials, rather than building the cells and measuring their potentials. Table 11.1 includes a few standard half-cell potentials. A larger table of standard half-cell potentials is given in Appendix E. 

The galvanic cell pictured in Figure 7.1 is not at equilibrium. If switch S is closed, electrons will spontaneously flow from the zinc (anode) to the copper (cathode) electrode. This flow will continue imtil the reactants and products attain their equilibrium concentrations. If switch S is opened before the cell reaches equilibrium, the electron flow will be interrupted. The voltmeter would register a positive voltage, which is a measure of the degree to which the redox reaction drives electrons from the anode to the cathode. Since this voltage is a type of energy that has the potential to do work, it is referred to as a redox potential or cell potential, denoted as... 

Coated specimens were placed in an electrochemical cell. After 4 hours of temperature, open-circuit potentials were measurements were made on duplicate samples, in a salt spray test cabinet (ASTM B117-73) for 1, 17 and 96 hours respectively and their surfaces photographed in order to calculate the percentage of surface covered by corroded spots and blisters (ASTM D610-68). 

In this expression, E° is the standard cell potential, the cell potential measured when all the species taking part are in their standard states. In practice, that means all gases are at 1 bar and all ions are at 1 mol-L 1. For example, to measure the standard potential of the Daniell cell, we should use 1 M CuS04(aq) in one electrode compartment and 1 M ZnS04(aq) in the other. 

A problem with this procedure is that we know only the overall cell potential, not the contribution of each individual electrode. A voltmeter placed between the two electrodes of a galvanic cell measures the difference of their potentials, not the individual values. To get around this difficulty, we arbitrarily set the standard potential of one particular electrode, the hydrogen electrode, equal to 0 at all temperatures ... 

The cell potential E (also called the cell voltage or electromotive force) is an electrical measure of the driving force of the cell reaction. Cell potentials depend on temperature, ion concentrations, and gas pressures. The standard cell potential E° is the cell potential when reactants and products are in their standard states. Cell potentials are related to free-energy changes by the equations AG = —nFE and AG° = —mFE°, where F = 96,500 C/mol e is the faraday, the charge on 1 mol of electrons. 

The similarities between the cytoplasmic processes of bones cells and the stereocilia of hair cells are that they both (1) measure mechanical deformations (vibrations of a fluid domain), (2) communicate their measurement to a network, (3) do this with dendritic structures, (4) the dendrites of both cells are constructed of similar materials (e.g., actin and fimbrin) and (5) the initial signaling in both cases consists of opening ion channels. While the hair cells communicate their information to a network that feeds to the brain, the bones cells connect to a lower level network (CCN) with (potentially) local decision-making software. 

It is not possible to measure the potential difference between the solution and the electrode, because in order to do this the solution must be connected to a conductor, i.e. a piece of another metal must be dipped into it. On the phase boundary another electrical double layer will be formed and in fact another, unknown electrode potential is developed. It is impossible therefore to measure absolute electrode potentials, only their differences. As seen before, the e.m.f. of a cell can be measured relatively easily, and this e.m.f. is the algebraic difference of the two electrode potentials. Building up cells from two electrodes,... 

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. 

Galvanic cell reactions, snch as those involving the everyday use of batteries, follow the same eqnations as electrolysis cells do. When we mea-snre cell potentials, however, we do not allow the reactions to proceed, becanse if they did, their potentials wonld change as the concentrations changed. The potentials are measured without a complete circuit. 

The differences in the hydration of a solnte in H2O and D2O have been extensively stndied by measnring their thermodynamic properties, the change of free energy (AG°t), enthalpy (A//°t), and entropy (AY°t) at the transfer of 1 mol of solnte from a highly dilute solution in H2O to the same concentration in D2O under reversible conditions (mostly 25 °C and atmospheric pressure). Greyson measured the electromotive force (emf) of electrochemical cells of several alkali halides containing heavy and normal water solutions. The cell potentials had been combined with available heat of solution data to determine the entropy of transfer of the salts between the isotopic solvents. The thermodynamic properties for the transfer from H2O to D2O and the solubilities of alkali halides at 25° in H2O and D2O are shown in Table 4. 

To resolve the apparent inconsistency of the summary statement of Section 5.2, let us consider yet another example. Assume that a gold electrode and a platinum electrode are dipped into the same solution of Fe VFe VH SO and both electrodes behave reversibly with respect to this redox couple. We have already established that the metal-solution potential difference at the two interphases is different, since each includes the term (J., as seen in Eq. 3IB. Yet the measured potential of this cell must be zero, since both electrodes are assumed to behave reversibly in the same solution. We could tentatively state that since electrons are neither produced nor consumed in the overall cell reaction, their chemical potential in the different phases cannot affect the measured potentials. 

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