Uses of Standard Electrode Potentials

The most important application of electrode potentials is the prediction of the spontaneity of redox reactions. Standard electrode potentials can be used to determine the spontaneity of redox reactions in general, whether or not the reactions can take place in electrochemical cells. 

Suppose we ask the question At standard conditions, will Cu ions oxidize metaUic Zn to Zn ions, or will Zn ions oxidize metaUic copper to Cu One of the two possible reactions is spontaneous, and the reverse reaction is nonspontaneous. We must determine which one is spontaneous. We already know the answer to this question from experimental results (see Section 21 -9), but let us demonstrate the procedure for predicting the spontaneous reaction from known values of standard electrode potentials. 

Choose the appropriate half-reactions from a table of standard reduction potentials. 

Write the equation for the half-reaction with the more positive (or less negative) value reduction, along with its potential. 

Write the equation for the other half-reaction as an oxidation and write its oxidation potential-, to do this, reverse the tabulated reduction half-reaction and change the sign of . (Reversing a half-reaction or a complete reaction also changes the sign of its potential.) 

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The Tl -Tl relationship is therefore a dominant feature of thallium chemistry. The standard reduction potentials at 25 °C and unit activity of H+ are TIVtI = —0.336 V, T1 /T1 = +0.72 V, and Tl /Tli = +1.25V. Estimates have also been made for the couples T1 /T1 = +0.33 V and Tl /Tl = 2.22 V. The generally valid limitations concerning the use of standard electrode potentials to predict the redox chemistry of real systems are especially important in the case of thallium factors such as complex formation in the presence of coordinating anions or neutral ligands and pH dependence due to hydrolysis do affect the actual or formal redox potentials. For example, redox potentials have been measmed for TICI/TICI3 =+0.77 V in IM HCl and T10H/T1(0H)3 = —0.05 V in alkaline soluhon. These formal potentials differ from the standard value for Tiin/Tii = +1.25 V. The difference can be attributed to the substanhal difference between the complex forming abilities of Tl and Tl , which will be discussed in detail later. The... 

How is an oxidation/reduction titration curve generated through the use of standard electrode potentials for the analyte species and the volumetric titrant ... 

Standard Electrode Potentials 21-15 Uses of Standard Electrode Potentials... 

An important use of standard electrode potentials is the calculation of the potential of a galvanic cell or the potential required to operate an electrolytic cell. The.se calculated polenlials (sometimes called thermodynamic polenlials) are theoretical in the sense that they refer lo cells in which there is no current. " Additional factors must be taken into account when there is a current in the cell. Furthermore, these potentials do not lake into account junction potentials within the cell. Normally, junction potentials can be made small enough H) be neglected without serious error. 

The Zinc-Copper Cell 21-10 The Copper-Silver Cell Standard Electrode Potentials 21-11 The Standard Hydrogen Electrode 21-12 The Zinc-SHE Cell 21-13 The Copper-SHE Cell 21-14 Standard Electrode Potentials 21-15 Uses of Standard Electrode Potentials 21-16 Standard Electrode Potentials for Other Half-Reactions 21-17 Corrosion 21-18 Corrosion Protection... 

USES OF STANDARD ELECTRODE POTENTIALS Table 21-2 Standard Reduction Potentials in Aqueous Solution at 25°C... 

To generalise, the use of standard electrode potentials can tell us if a reaction is thermodynamically viable. However if a process is predicted to be possible, standard electrode potentials tell us nothing about the likely rate of reaction. 

See p. 435 for discussion of standard electrode potentials and their use. It is convenliona] to write the halfreactions as (oxidized form) = (reduced form). Since... 

Nowadays, tables of standard electrode potentials are used instead of the electromotive series. They include electrode reactions not only of metals but also of other substances [Table 3.1 for detailed tables, see the books of Lewis and Rendall (1923) and Bard et al. (1985)]. 

The review of Martynova (18) covers solubilities of a variety of salts and oxides up to 10 kbar and 700 C and also available steam-water distribution coefficients. That of Lietzke (19) reviews measurements of standard electrode potentials and ionic activity coefficients using Harned cells up to 175-200 C. The review of Mesmer, Sweeton, Hitch and Baes (20) covers a range of protolytic dissociation reactions up to 300°C at SVP. Apart from the work on Fe304 solubility by Sweeton and Baes (23), the only references to hydrolysis and complexing reactions by transition metals above 100 C were to aluminium hydrolysis (20) and nickel hydrolysis (24) both to 150 C. Nikolaeva (24) was one of several at the conference who discussed the problems arising when hydrolysis and complexing occur simultaneously. There appear to be no experimental studies of solution phase redox equilibria above 100°C. 

As it has been shown that the Gibbs function for formation of an individual ion has no operational meanings [12], no way exists to determine such a quantity experimentally. However, for the purposes of tabulation and calculation, it is possible to separate AfGm of an electrolyte arbitrarily into two or more parts, which correspond to the number of ions formed, in a way analogous to that used in tables of standard electrode potentials. In both cases, the standard Gibbs function for formation of aqueous H" " is defined to be zero at every temperature ... 

Thus, the IUPAC decision supports the zinc-minus-copper-plus table of standard electrode potentials. The first thing to do, therefore, when consulting a table of standard electrode potentials is to examine the E° values of the zinc and copper electrodes. If the values are -0.76 and +0.34 V, respectively, the table can be used. If, however, the values are +0.76 and -0.34 V, the convention contravenes the IUPAC decision. To use such a table, one can retain all the magnitudes of the E° values, but change all the signs of the E° values the table will then be in accord with the international convention (Table 7.23). 

Lacking a table of standard electrode potentials, or one that is adequate, what guidelines can be used to identify oxidizing and reducing agents, and to estimate their relative strengths Here are a few. 

It is a relatively simple process to set up a scale of redox potentials in a non-aqueous medium using the standard hydrogen electrode in that medium as the fundamental reference electrode. Thus in liquid ammonia, which is a well studied non-aqueous solvent and for which there exists a considerable amount of thermodynamic information,31 the scale of standard electrode potentials is referred to the standard hydrogen electrode in liquid ammonia (equation 25), which is assigned the value of zero volts, and in which the H+ exists as a solvated species, i.e. NH4+. 

It should be emphasized that many of the potentials listed in tables of standard electrode potentials are values calculated from thermodynamic data rather than obtained directly from cell emf data. As such they are valuable for calculating equilibrium constants of reactions, but caution should be exercised in using them to predict the behavior of electrodes. A steady value for an electrode potential does not necessarily represent the thermodynamic or equilibrium value. 

Referring to a list of standard electrode potentials, such as in Table 8.3, one speaks of an electrochemical series, and the metals lower down in the se-ries(with positive electrode potentials) are called noble metals. Any combination of half-reactions in an electrochemical cell, which gives a nonzero E value, can be used as a galvanic cell (i.e., a battery). If the reaction is driven by an applied external potential, we speak of an electrolytic cell. Reduction takes place at the cathode and oxidation at the anode. The reduction reactions in Table 8.3 are ordered with increasing potential or pe values. The oxidant in reactions with latter pe (or E°) can oxidize a reductant at a lower pe (or ) and vice versa for example, combining half-reactions we obtain an overall redox reaction ... 

Numerous applications of standard electrode potentials have been made in various aspects of electrochemistry and analytical chemistry, as well as in thermodynamics. Some of these applications will be considered here, and others will be mentioned later. Just as standard potentials which cannot be determined directly can be calculated from equilibrium constant and free energy data, so the procedure can be reversed and electrode potentials used for the evaluation, for example, of equilibrium constants which do not permit of direct experimental study. Some of the results are of analjrtical interest, as may be shown by the following illustration. Stannous salts have been employed for the reduction of ferric ions to ferrous ions in acid solution, and it is of interest to know how far this process goes toward completion. Although the solutions undoubtedly contain complex ions, particularly those involving tin, the reaction may be represented, approximately, by... 

As already discussed, the standard hydrogen electrode (SHE) is the chosen reference half-cell upon which tables of standard electrode potentials are based. The potential of this system is zero by definition at all temperatures. Although this reference electrode was often used in early work in electrochemistry, it is almost never seen in chemical laboratories at the present time. It is simply too awkward to use because of the requirement for H2 gas at 1 bar pressure and safety considerations. 

Before we discuss standard electrode potential, we will talk about electromotive force (emf). The electromotive force of a cell is the potential difference between the two electrodes. This can be measured using a voltmeter. The maximum voltage of a cell can be calculated using experimentally determined values called standard electrode potentials. By convention, the standard electrode potentials are usually represented in terms of reduction half-reactions for 1 molar solute concentration. The standard electrode potential values are set under ideal and standard-state conditions (latm pressure and 25°C temperature). From the MCAT point of view, you can assume that the conditions are standard, unless stated otherwise. Table 12-1 shows a list of standard electrode potentials (in aqueous solution) at 25°C. 

In the case of standard electrode potential, it is appropriate to have a standard electrode whose reversible potential is made arbitrarily zero and against which the potentials of other electrodes can be measured. The hydrogen electrode is an accepted standard. It is composed of a rod of platinum covered with platinum black saturated with hydrogen gas at atmospheric pressure. Electrode potential based on this zero are said to refer to the hydrogen scale. However, in experimental work, it is often more suitable to use another standard electrode. Calomel is a common example. It consists of a pool of mercury covered with calomel (mercurous chloride) and immersed in a solution of potassium chloride. 

Substitute into AG° = —nFE H. Use a table of standard electrode potentials to obtain Fceu- The cell reaction equals the sum of the half-reactions after they have been multiplied by factors so that the electrons cancel in the summation. Note that n is the number of moles of electrons involved in each half-reaction. 

The electromotive force (emf), or cell potential, is the maximum voltage of a voltaic cell. It can be directly related to the maximum work that can be done by the cell. A standard electrode potential, or reduction potential, refers to the potential of an electrode in which molar concentrations and gas pressures (in atmospheres) have unit values. A table of standard electrode potentials is useful for establishing the direction of spontaneity of an oxidation-reduction reaction and for calculating the standard emf of a cell. 

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