Cell potential defined

A combination of any two dissimilar metallic conductors can be used to construct a galvanic cell. The cell potential defines the measure of the energy available in a cell. A high cell potential signifies a vigorous spontaneous redox reaction. The unit of potential is the volt, V. A Daniell cell, for example, has a potential of 1.1 V. 

The thermodynamic convention of cell potential defined in Eq. 5 is not practical for most electrochemical work. For a battery switching from discharge to charge, cathode and anode interchange, but the electrode polarities remain the same. It is not practical to switch the leads of the voltmeter when the battery switches to the charge mode. Rather, the voltmeter stays connected during the switch with the positive terminal attached to the positive terminal of the battery and likewise for the negative terminal. The practical convention for ceU potential is... 

Equations 11.19-11.21 are defined for a potentiometric electrochemical cell in which the pH electrode is the cathode. In this case an increase in pH decreases the cell potential. Many pH meters are designed with the pH electrode as the anode so that an increase in pH increases the cell potential. The operational definition of pH then becomes... 

The reduction potentials for the actinide elements ate shown in Figure 5 (12—14,17,20). These ate formal potentials, defined as the measured potentials corrected to unit concentration of the substances entering into the reactions they ate based on the hydrogen-ion-hydrogen couple taken as zero volts no corrections ate made for activity coefficients. The measured potentials were estabhshed by cell, equihbrium, and heat of reaction determinations. The potentials for acid solution were generally measured in 1 Af perchloric acid and for alkaline solution in 1 Af sodium hydroxide. Estimated values ate given in parentheses. 

The net electrochemical driving force is determined by two factors, the electrical potential difference across the cell membrane and the concentration gradient of the permeant ion across the membrane. Changing either one can change the net driving force. The membrane potential of a cell is defined as the inside potential minus the outside, i.e. the potential difference across the cell membrane. It results from the separation of charge across the cell membrane. 

Since the reaction in the working electrode is an oxidation when the overall reaction is (7.14), the cell potential in (7.15) is defined as = ,ork - ref = anode Scaihode and the sign in this equation is opposite to that obtained with the more common convention that defines the cell potential as = caUiode - T anode ) From the temperature variation of the cell potential, the following equation can be written for the entropy of the overall reaction ... 

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. 

Thus far, all of our calculations have been based on the standard cell potential or standard halfcell potentials—that is, the standard state conditions that were defined previously. However, many times the cell is not at standard conditions—commonly the concentrations are not 1 M. The actual cell potential, E, can be calculated by the use of the Nemst equation ... 

The design and construction of an electrochemical cell derives from consideration of the system being examined. Potential sources of contamination must be carefully evaluated. Cell components are typically made of inert materials such as Teflon or Kel-F. Alternatively, electrolyte contact with confining materials may be avoided altogether by letting the cell be defined by the geometry of a hanging meniscus. The latter method has been incor-... 

The ideal performance of a fuel cell is defined by its Nemst potential represented as cell voltage. The overall cell reactions corresponding to the individual electrode reactions listed in Table 2-1 are given in Table 2-2, along with the corresponding form of the Nemst equation. The Nemst... 

While this potential cannot he determined for a single electrode, a potential can be derived if the potential of the other electrode in a cell is defined, i.e. the potential of the standard hydrogen electrode (SHE) is arbitrarily taken as 0.(XXX)V. In this way. a potential scale can then be devised for single electrode potentials - see Section 3.2. t The abbreviation emf , in upright script, is often used in other lextNmks as a direct , i.e. non-variable, acronym for the electromotive force. Note, however, that in this present text it is used to represent a variable (cell potential) and is therefore. shown in italic script. 

A reference electrode is defined as a constant-potential device . In other words, if we make a cell in which one half cell is the SHE and then measure the emf. we then simply employ the equation emf = Eright-hand side — ieft-hand side (equation (3.3)) to determine the other half-cell potential. 

Standard hydrogen electrode (SHE) The standard against which redox potentials are measured. The SHE consists of a platinum electrode electroplated with Pt black (to catalyse the electrode reaction), over which hydrogen at a pressure of 1 atm is passed. The electrode is immersed in a solution containing hydrogen ions at unit activity (e.g. 1.228 mol dm of aqueous HCl at 20°C). The potential of the SHE half cell is defined as 0.000 V at all temperatures. 

By this convention, the potential of the cell is defined as the potential of the electrode on the right, at which reduction occurs, minus the potential of the electrode on the left, at which oxidation occurs. 

Finally, although the free energies of transfers found in Table 7 are quite useful in comparing cell potentials measured in various solvents, these values are insufiicient to adequately interpret potentials measured under many nonaqueous conditions. The primary reason for this difficulty stems from the manner in which AG°(M", w s) is defined. For example, AGj (Na, w CH3CN) is associated with the transfer of sodium ions... 

This is an important relationship, and it defines the cell potential, V, as the difference in the inner potentials of phases M and M. ... 

The nonpolarizable interface has been defined above (Section 6.3.3) as one which, at constant solution composition, resists any change in potential due to a change in cell potential. This implies that (3s Ma< )/3V)jl = 0. However, the inner potential difference at such an interface can change with solution composition hence, Eq. (6.89) can be rewritten in the form of dM7ds< > = (RT/ZjF) d In a, which is the Nemst equation [see Eq. (7.51)] in differential form for a single interface. 

The example here stresses simplicity. To determine the pathway of more complex reactions (glucose oxidation, for example), one needs to do more, just because there are several alternative pathways. However, if the reaction is complex, it may be enough to find the pathway up to the rds. That, after all, is what determines the overpotential in the steady state and therefore the cell potential and how much electricity will be needed to produce a given amount of product. If it occurs in the first three steps of, e.g., a six-step sequence, the work needed to define the path up to the third step is much less (not just 1/2) than one that involves an investigation of the steps after the rds. 

The potential difference of the cell is defined as the difference between the Galvani potentials at the right (R) and left (L) terminals named here as the copper wires,... 

Evaluation of VOC and SVOC emission potential of individual products and materials under indoor-related conditions and over defined timescales requires the use of climate-controlled emission testing systems, so-called emission test chambers and cells, the size of which can vary between a few cm3 and several m3, depending on the application. In Figure 5.1 the dots ( ) represent volumes of test devices reported in the literature. From this size distribution they can be classified as large scale chambers, small scale chambers, micro scale chambers and cells. The selection of the systems, the sampling preparation and the test performance all depend on the task to be performed. According to ISO, chambers and cells are defined as follows ... 

Throughout this discussion we have considered cells in which the electrolytic solution is an aqueous solution. The same methods can be used to define standard half-cell potentials in any solvent system. However, it is important to remember that when the reference state is defined as the infinitely dilute solution of a solute in a particular solvent, the standard state depends upon that solvent. The values so obtained are not interchangeable between the different solvent systems. Only if the standard states could all be defined independently of the solvent would the values be applicable to all solvent systems. 

The numerical value of an electrode potential depends on the nature of the particular chemicals, the temperature, and on the concentrations of the various members of the couple. For the purposes of reference, half-cell potentials are taken at the standard states of all chemicals. Standard state is defined as 1 atm pressure of each gas (the difference between 1 bar and 1 atm is insignificant for the purposes of this chapter), the pure substance of each liquid or solid, and 1 molar concentrations for every nongaseous solute appearing in the balanced half-cell reaction. Reference potentials determined with these parameters are called standard electrode potentials and, since they are represented as reduction reactions (Table 19-1), they are more often than not referred to as standard reduction potentials (E°). E° is also used to represent the standard potential, calculated from the standard reduction potentials, for the whole cell. Some values in Table 19-1 may not be in complete agreement with some sources, but are used for the calculations in this book. 

In the cell Zn(s) Zn2+(ag) I I Cu2+(aq) I Cu(.s) the zinc appears on the left side, indicating that it is being oxidized, not reduced. For this reason, the potential difference contributed by the left half-cell has the opposite sign to its conventional half-cell potential. More generally, we can define the cell potential or cell EMF as... 

Reference half-cells The fact that individual half-cell potentials are not directly measurable does not prevent us from defining and working with them. Although we cannot determine the absolute value of a half-cell potential, we can still measure its value in relation to the potentials of other half cells. In particular, if we adopt a reference half-cell whose potential is arbitrarily defined as zero, and measure the potentials of various other electrode systems against this reference cell, we are in effect measuring the half-cell potentials on a scale that is relative to the potential of the reference cell. 

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