Redox half cell

Note that here we have squared the activity of the iodide ion since the balanced redox half-cell reaction is I2 + 2e 21 . This redox stoichiometry also explains... 

In this present chapter, we will be looking at a slightly more complicated situation, i.e. one in which the contents of two redox half cells are not separated but are allowed to mix. Because mixing occurs, redox chemistry can occur, i.e. electron-transfer reactions are not forbidden. Any electrochemical equilibrium attained is thus a genuine thermodynamic equilibrium and is not frustrated . 

Oxidation-reduction potential is sensitive to pH. Eh decreases with an increase in the pH and increases with a decrease in the pH, if H+ ion or OH- ion is involved in the redox half-cells. 

The zinc-zinc ion, the hydrogen, and the silver-silver chloride electrodes are typical of three common types of aqueous half-cell. A fourth type is the so-called redox half-cell, which involves, for example, both ferrous and ferric ions in solution. A non-reacting electrode, usually of platinum, facilitates the following half-cell process ... 

The metal/ion half-cell generates a potential by the exchange of metal ions between the metal and the electrolyte solution. In contrast, a redox half-cell is based upon an exchange of electrons between die metal and the electrolyte solution. So actually fliere are two sets of standard potential tables, one for metal/ion half-cells (Table 7.2) and one for redox halfcells (Table 7.3). The half-cell potential is of course independent of the interphase area. 

Also known as the standard hydrogen electrode (SHE), it is a redox reference electrode which forms the basis of the thermodynamic scale of oxidation-reduction potentials. The potential of the NHE is defined as zero and based oti equilibrium of the following redox half-cell reaction, typically on a Pt surface 2H+(aq) + 2e H2(g). The activities of both the reduced form and the oxidized form are maintained at unity. That implies that the pressure of hydrogen gas is 1 atm and the concentration of hydrogen ions in the solution is 1 M. 

Suppose now that we connect together two such redox half-cells, corresponding to the reactions... 

The standard hydrogen electrode (SHE), also referred to as normal hydrogen electrode, is the universal reference for reporting relative half-cell potentials. The SHE could be used as either an anode or a cathode depending upon the nature of the half-cell it is used with. SHE is based on the redox half-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... 

Together, the oxidized and reduced forms of the substance are referred to as a redox couple.) Such a sample half-cell is connected to a reference half-cell... 

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... 

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. 

It is so universally applied that it may be found in combination with metal oxide cathodes (e.g., HgO, AgO, NiOOH, Mn02), with catalytically active oxygen electrodes, and with inert cathodes using aqueous halide or ferricyanide solutions as active materials ("zinc-flow" or "redox" batteries). The cell (battery) sizes vary from small button cells for hearing aids or watches up to kilowatt-hour modules for electric vehicles (electrotraction). Primary and storage batteries exist in all categories except that of flow-batteries, where only storage types are found. Acidic, neutral, and alkaline electrolytes are used as well. The (simplified) half-cell reaction for the zinc electrode is the same in all electrolytes ... 

Figure 9.3 The lead storage battery. The key to obtaining electrical energy from a redox chemical reaction is to physically separate the two half-cell reactions so that electrons are transferred from the anode through an external circuit to the cathode. In the process, electrical work is accomplished.

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... 

Defining a reference value for the SHE makes it possible to determine E ° values of all other redox half-reactions. As an example. Figure 19-14 shows a cell in which a standard hydrogen electrode is connected to a copper electrode in contact with a 1.00 M solution of C U . Measurements on this cell show that the SHE is at higher electrical potential than the copper electrode, indicating that electrons flow from the SHE to the Cu... 

As an example of the numerous combinations of the hydrogen electrode with other electrodes, attention is drawn to Figure 6.10 (C) of a galvanic cell. The two redox couples or the half-cells to make up the cell and reactions taking place at the electrodes or the electrode reactions constituting the cell are presented below ... 

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 electrochemical cell can again be of the regenerative or electrosynthetic type, as with the photogalvanic cells described above. In the regenerative photovoltaic cell, the electron donor (D) and acceptor (A) (see Fig. 5.62) are two redox forms of one reversible redox couple, e.g. Fe(CN)6-/4 , I2/I , Br2/Br , S2 /S2, etc. the cell reaction is cyclic (AG = 0, cf. Eq. (5.10.24) since =A and D = A ). On the other hand, in the electrosynthetic cell, the half-cell reactions are irreversible and the products (D+ and A ) accumulate in the electrolyte. The most carefully studied reaction of this type is photoelectrolysis of water (D+ = 02 and A = H2)- Other photoelectrosynthetic studies include the preparation of S2O8-, the reduction of C02 to formic acid, N2 to NH3, etc. 

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. 

This process of creating ATP, known as electron transport phosphorylation, then, involves two half-cell reactions, one at the electron donation site and the other where the electrons are accepted from the transport chain. Taking aerobic sulfide oxidation as an example, the donating species H2S(aq) gives up electrons, two at a time, to a series of redox complexes. With the loss of each pair of electrons, the sulfide oxidizes first to S°, then thiosulfate, sulfite, and finally sulfate. 

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... 

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