Oxidation-reduction cell potential

The corresponding oxidation half-cell potential is the negative of this value, i.e., 7iV/ri = +0.76 V Some standard reduction half-cell potentials are given in Table 15-1. 

Communicating Write two or three sentences describing the processes that take place in a voltaic cell and account for the direction of electron flow. Use the words cathode, anode, oxidation, reduction, and potential difference in your sentences. 

Oxidation-reduction cells were also studied by C. R. Alder Wright and C. Thompson and W. D. Bancroft, who arranged oxidising and reducing agents in order of strength on the basis of the potentials. 

The standard potential for the anodic reaction is 1.19 V, close to that of 1.228 V for water oxidation. In order to minimize the oxygen production from water oxidation, the cell is operated at a high potential that requires either platinum-coated or lead dioxide anodes. Various mechanisms have been proposed for the formation of perchlorates at the anode, including the discharge of chlorate ion to chlorate radical (87—89), the formation of active oxygen and subsequent formation of perchlorate (90), and the mass-transfer-controUed reaction of chlorate with adsorbed oxygen at the anode (91—93). Sodium dichromate is added to the electrolyte ia platinum anode cells to inhibit the reduction of perchlorates at the cathode. Sodium fluoride is used in the lead dioxide anode cells to improve current efficiency. 

Because of lithium s low density and high standard potential difference (good oxidation reduction characteristics), cells using lithium at the anode have a very high energy density relative to lead, nickel and even zinc. Its high cost limits use to the more sophisticated and expensive electronic equipment. 

Oxidation-reduction potential Because of the interest in bacterial corrosion under anaerobic conditions, the oxidation-reduction situation in the soil was suggested as an indication of expected corrosion rates. The work of Starkey and Wight , McVey , and others led to the development and testing of the so-called redox probe. The probe with platinum electrodes and copper sulphate reference cells has been described as difficult to clean. Hence, results are difficult to reproduce. At the present time this procedure does not seem adapted to use in field tests. Of more importance is the fact that the data obtained by the redox method simply indicate anaerobic situations in the soil. Such data would be effective in predicting anaerobic corrosion by sulphate-reducing bacteria, but would fail to give any information regarding other types of corrosion. 

The complete reaction may be regarded as composed of two oxidation-reduction electrodes, a Ox, a Red, and frOx , b Red, combined together into a cell at equilibrium, the potentials of both electrodes are the same ... 

In theory, any oxidation-reduction reaction can be set up in a cell to do electrical work. The amount of reversible1 work is easily calculated. If, during the discharge of a cell, a quantity of electricity Q flows through the external circuit at a constant potential, the amount of electrical work, n e, produced is given by... 

The aluminum-air fuel cell is used as a reserve battery in remote locations. In this cell aluminum reacts with the oxygen in air in basic solution, (a) Write the oxidation and reduction half-reactions for this cell, (b) Calculate the standard cell potential. See Box 12.1. 

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 ... 

A table giving the cell potentials of all possible redox reactions would be immense. Instead, chemists use the fact that any redox reaction can be broken into two distinct half-reactions, an oxidation and a reduction. They assign a potential to every half-reaction and tabulate E ° values for all half-reactions. The standard cell potential for any redox reaction can then be obtained by combining the potentials for its two half-reactions. 

The Zn /Zn reduction potential is more negative than the H3 O /H2 reduction potential (-0.76 V vs. 0 V), so zinc is the anode in this cell. Zinc is oxidized and hydronium ions are reduced, causing electrons to flow from the more negative zinc electrode to the less negative SHE. Again, we reverse the direction of the half-reaction with the more negative potential and find E by subtracting the half-cell potentials ... 

Analytical methods based upon oxidation/reduction reactions include oxidation/reduction titrimetry, potentiometry, coulometry, electrogravimetry and voltammetry. Faradaic oxidation/reduction equilibria are conveniently studied by measuring the potentials of electrochemical cells in which the two half-reactions making up the equilibrium are participants. Electrochemical cells, which are galvanic or electrolytic, reversible or irreversible, consist of two conductors called electrodes, each of which is immersed in an electrolyte solution. In most of the cells, the two electrodes are different and must be separated (by a salt bridge) to avoid direct reaction between the reactants. 

In Table 7-1 the relative tendencies of certain elements to react were listed qualitatively. We can give a quantitative measure of relative tendency to react, called standard reduction potential, as shown in Table 14-2. In this table, the standard half-cell potential for each half-reaction, as a reduction, is tabulated in order with the highest potential first. If we turn these half-reactions around, we change the signs of the potentials and we get oxidation potentials. We thus have half-reactions including both elementary metals and elementary nonmetals in the same table, as well as many half-reactions that do... 

Then add the copper half-cell reduction to the zinc half-cell oxidation and add the half-cell potentials ... 

Equilibrium considerations other than those of binding are those of oxidation/reduction potentials to which we drew attention in Section 1.14 considering the elements in the sea. Inside cells certain oxidation/reductions also equilibrate rapidly, especially those of transition metal ions with thiols and -S-S- bonds, while most non-metal oxidation/reduction changes between C/H/N/O compounds are slow and kinetically controlled (see Chapter 2). In the case of fast redox reactions oxidation/reduction potentials are fixed constants. 

A We write down the oxidation half-equation with the method of Chapter 5, and obtain the reduction half-equation from Table 21-1, along with the reduction half-cell potential. 

Since the overall cell potential is positive, this step is spontaneous. The next step involves oxidation of Fe2+(aq) back to Fe3+(aq), (i.e. the reverse of reaction (i) and the reduction of H202(aq) to H20(l) in acidic solution, for which the half-reaction is... 

Since oxidation occurs at the anode and reduction at the cathode, the standard cell potential can be calculated from the standard reduction potentials of the two half-reactions involved in the overall reaction by using the equation ... 

Because one half-reaction must involve oxidation, one of the half-reactions shown in the table of reduction potentials must be reversed to indicate the oxidation. If the half-reaction is reversed, the sign of the standard reduction potential must be reversed. However, this is not necessary to calculate the standard cell potential. 

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