Half-cells redox potential

Consider the difference between two related half-cell redox potentials. 

Ostwald had what appeared to be a very elegant concept. It involved the measurement of a single electrode potential. The method of measurement was in good accordance with his philosophical views and with the chemistry of the times, and it would, in his opinion, yield an absolute potential. An absolute potential was a sharp contrast to the relative potential obtained by referring a measured half-cell to another single electrode reaction arbitrarily set at zero. Ostwald s measurements of half-cell potentials could be directly related to heats of ionization (1 ). In his opinion, an absolute half-cell redox potential would allow the establishment of an electromotive series which would be analogous to the absolute temperature scale. 

Some typical half-cell reactions and their respective standard reduction potentials are listed in Table 21.1. Whenever reactions of this type are tabulated, they are uniformly written as reduction reactions, regardless of what occurs in the given half-cell. The sign of the standard reduction potential indicates which reaction really occurs when the given half-cell is combined with the reference hydrogen half-cell. Redox couples that have large positive reduction potentials... 

Tabulated E values can be used to calculate the for any reaction, as illustrated in Table 7.2 for the Zn/Cu galvanic cell. The redox reaction is spontaneous when the half-reaction (Cu /Cu) with the larger reduction (+0.34V) acts as the oxidizing agent. In this case, the other half-reaction (Zn /Zn) proceeds as an oxidation. The halfcell potential for this reduction is +0.76 V as it represents the reverse of the half-cell reduction potential as listed in Table 7.2. The sum of the oxidation and reduction half reactions is +0.34V + 0.76 V = +1.10 V. Thus for the galvanic Zn/Cu cell is +1.10V. 

The foregoing example illustrates how equilibrium constants for overall cell reactions can be determined electrochemically. Although the example dealt with redox equilibrium, related procedures can be used to measure the solubility product constants of sparingly soluble ionic compounds or the ionization constants of weak acids and bases. Suppose that the solubility product constant of AgCl is to be determined by means of an electrochemical cell. One half-cell contains solid AgCl and Ag metal in equilibrium with a known concentration of CP (aq) (established with 0.00100 M NaCl, for example) so that an unknown but definite concentration of Kg aq) is present. A silver electrode is used so that the half-cell reaction involved is either the reduction of Ag (aq) or the oxidation of Ag. This is, in effect, an Ag" Ag half-cell whose potential is to be determined. The second half-cell can be any whose potential is accurately known, and its choice is a matter of convenience. In the following example, the second half-cell is a standard H30" H2 half-cell. 

Redox potential is defined by the half cell reduction potential that is created by redox couples that are primarily due to GSH, NAD+ and nicotinamide dinucleotide phosphate. These couples are in ratios of the oxidized to reduced form of the molecules (NAD /NAD, NADP /NADPH, and GSSG/2GSH). The redox couples can be independent, as well linked to each other to form related couples. The redox environment is a reflection of these couples. These ratios can be measured by the Nemst equation, similar to a voltaic cell. 

Redox Electrodes Electrodes of the first and second kind develop a potential as the result of a redox reaction in which the metallic electrode undergoes a change in its oxidation state. Metallic electrodes also can serve simply as a source of, or a sink for, electrons in other redox reactions. Such electrodes are called redox electrodes. The Pt cathode in Example 11.1 is an example of a redox electrode because its potential is determined by the concentrations of Ee + and Ee + in the indicator half-cell. Note that the potential of a redox electrode generally responds to the concentration of more than one ion, limiting their usefulness for direct potentiometry. 

This method involves very simple and inexpensive equipment that could be set up m any laboratory [9, 10] The equipment consists of a 250-mL beaker (used as an external half-cell), two platinum foil electrodes, a glass tube with asbestos fiber sealed m the bottom (used as an internal half-cell), a microburet, a stirrer, and a portable potentiometer The asbestos fiber may be substituted with a membrane This method has been used to determine the fluoride ion concentration in many binary and complex fluondes and has been applied to unbuffered solutions from Willard-Winter distillation, to lon-exchange eluant, and to pyrohydrolysis distil lates obtained from oxygen-flask or tube combustions The solution concentrations range from 0 1 to 5 X 10 M This method is based on complexing by fluonde ions of one of the oxidation states of the redox couple, and the potential difference measured is that between the two half-cells Initially, each cell contains the same ratio of cerium(IV) and cerium(tll) ions... 

As a result, the electromotive force (EMF) of the cell is zero In the presence of fluoride ions, cerium(IV) forms a complex with fluoride ions that lowers the cerium(IV)-cerium(IIl) redox potential The inner half-cell is smaller, and so only 5 mL of cerium(IV)-cenum (III) solution is added To the external half-cell, 50 mL of the solution is added, but the EMF of the cell is still zero When 10 mL of the unknown fluonde solution is added to the inner half-cell, 100 mL of distilled water IS added to the external half-cell The solution in the external half-cell is mixed thoroughly by turning on the stirrer, and 0 5 M sodium fluonde solution is added from the microburet until the null point is reached The quantity of known fluonde m the titrant will be 10 times the quantity of the unknown fluoride sample, and so the microburet readings must be corrected prior to actual calculations... 

Three kinds of equilibrium potentials are distinguishable. A metal-ion potential exists if a metal and its ions are present in balanced phases, e.g., zinc and zinc ions at the anode of the Daniell element. A redox potential can be found if both phases exchange electrons and the electron exchange is in equilibrium for example, the normal hydrogen half-cell with an electron transfer between hydrogen and protons at the platinum electrode. In the case where a couple of different ions are present, of which only one can cross the phase boundary — a situation which may exist at a semiperme-able membrane — one obtains a so called membrane potential. Well-known examples are the sodium/potassium ion pumps in human cells. 

The elemental reaction used to describe a redox reaction is the half reaction, usually written as a reduction, as in the following case for the reduction of oxygen atoms in O2 (oxidation state 0) to H2O (oxidation state —2). The half-cell potential, E°, is given in volts after the reaction ... 

If the equilibrium half-cell potentials for two redox reactions are different, electrons will be transferred from the reduced species in the... 

It is very often necessary to characterize the redox properties of a given system with unknown activity coefficients in a state far from standard conditions. For this purpose, formal (solution with unit concentrations of all the species appearing in the Nernst equation its value depends on the overall composition of the solution. If the solution also contains additional species that do not appear in the Nernst equation (indifferent electrolyte, buffer components, etc.), their concentrations must be precisely specified in the formal potential data. The formal potential, denoted as E0, is best characterized by an expression in parentheses, giving both the half-cell reaction and the composition of the medium, for example E0,(Zn2+ + 2e = Zn, 10-3M H2S04). 

Quinhydrone, a solid-state associate of quinone and hydroquinone, decomposes in solution to its components. The quinhy drone electrode is an example of more complex organic redox electrodes whose potential is affected by the pH of the solution. If the quinone molecule is denoted as Ox and the hydroquinone molecule as H2Red, then the actual half-cell reaction... 

The overall cell potential is +0.96 V, showing that the redox reaction is indeed spontaneous. The standard reduction potential for the half cell Ag2S(s) + 2e - 2Ag(s) + S2 (aq) was obtained from the American Society for Metals (ASM) Handbook, available on the internet. 

As discussed in detail in Section 7.4, the energy liberated by a redox reaction depends on the redox potential of the electron-donating half-cell reaction, relative to the electron-accepting reaction. In the calculation results, we can trace the redox... 

Fig. 22.6. Redox potentials (mV) of various half-cell reactions during mixing of fluid from a subsea hydrothermal vent with seawater, as a function of the temperature of the mixture. Since the model is calculated assuming 02(aq) and H2(aq) remain in equilibrium, the potential for electron acceptance by dioxygen is the same as that for donation by dihydrogen. Dotted line shows currently recognized upper temperature limit (121 °C) for microbial life in hydrothermal systems. A redox reaction is favored thermodynamically when the redox potential for the electron-donating half-cell reaction falls below that of the accepting half-reaction.

The first difference between these two batteries is the voltage they produce a watch battery produces about 3 V and a lead-acid cell about 2 V. The obvious cause of the difference in emf are the different half-cells. The electrode potential E is the energy, expressed as a voltage, when a redox couple is at equilibrium. 

The standard electrode potential of an element is defined as its electrical potential when it is in contact with a molar solution of its ions. For redox systems, the standard redox potential is that developed by a solution containing molar concentrations of both ionic forms. Any half-cell will be able to oxidize (i.e. accept electrons from) any other half-cell which has a lower electrode potential (Table 4.1). 

Redox half-reactions are often written for brevity as, for example, Li+ + e - Li. with the state symbols omitted. The electrode system represented by the half-reaction may also be written as Li+ /Li. The standard redox potentials for ion-ion redox systems can be determined by setting up the relevant half-cell and measuring the potential at 298 K relative to a standard hydrogen electrode. For example, the standard redox potential for the half-reactions... 

The oxidation-reduction potential or redox potential ( h) is a measure of the tendency of a solution to be oxidizing or reducing. Oxidation and reduction are basically electrical processes that are readily measiued by an electrode potential. All measurements are referred to die standard hydrogen electrode, the potential of which is taken as 0.00 V at 298 K, the H2 pressure as 101325 N/m (1 atm) and activities of H2 and as unity. When the half-cell reaction is written as an oxidation reaction ... 

Many redox systems are suitable for use as volumetric reagents for quantitative analysis provided that (i) both states within the oxidized and reduced forms of the redox-active titrant comprise a fast nemstian couple, (ii) all redox states are soluble in the solutions employed, and (iii) the separation between the standard electrode potential for each of the constituent half cells is 0.35/n V (where n is the number of electrons in the titrant couple). 

Let us revisit the electrochemical cell shown earlier in Figure 3.1. In this figure, two redox electrodes are immersed in solutions of their respective ions, with the half cells being connected by a salt bridge. If we were to connect an infinite-resistance voltmeter between the cells, then it would be possible to perform potentiometric experiments such as those described in the previous chapter. One electrode would be positive with respect to the other, with the separation in potential between the two electrodes being the emf - but only if the measurement was performed at equilibrium. (As before, we take the word equilibrium to imply that no charge flows.)... 

Electrode potential, E The energy, expressed as a voltage, of a redox couple at equilibrium. E is the potential of the electrode when measured relative to a standard (ultimately the SHE). E depends on temperature, activity and solvent. By convention, the half cell must first be written as a reduction, and the potential is then designated as positive if the reaction proceeds spontaneously with respect to the SHE. Otherwise, E is negative. 

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. 

Overcharge tests were carried out in LiCo02 cathode half-cells that contained these additives, and a redox shuttle effect was observed between 4.20 and 4.30 V, close to the redox potentials of these additives. The same shuttling effect was observed even after 2 months of storage for these cells, indicating the stability and redox reversibility of these additives. A closer examination of the capacity retention revealed that 4-bromo-l,2-dimethoxybenzene seemed to have the best shuttle-voltage performance for the 4.0 V lithium cell used." The stability of these additives against reductive decomposition was also tested by the authors on metallic lithium as well as on carbonaceous anodes, and no deterioration was detected. 

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