Standard potential table

Use standard potentials (Table 18.1) torn compare the relative strengths of ditferent oxidizing agents different reducing agents. (Example 18.2 Problems 9-18) 9,10,16,18... 

In making a transition to a quantitative discussion of the electrochemistry of the alkah metals, we begin with a discussion of standard potentials. Table 1 provides a list of standard potentials for half-reactions that take the generic form found in Eq. (1)... 

Note that standard potential tables may list some reactions involving OH- instead of H+ ions. The conversion is relatively simple. For example, for the second equilibrium written above, one can also... 

By convention (cf. 33k) the free energy of formation of every element in its standard state is taken as zero, and, as seen above, the same applies to the hydrogen ion. The standard free energy of formation of liquid water at 26 C is — 56.70 kcal. mole " (Table XXIV), and that of the iodide ion is derived from its standard potential (Table XXXIX) as 1 X 23,070 X ( — 0.536) cal., i.e., — 12.37 kcal. g. ion"" It follows, therefore, that... 

If we know the form of the complex, we could write a new half-reaction involving the acid anion and determine an E value for this reaction, keeping the acid and all other species at unit activity. However, the complexes are frequently of unknown composition. So we define the formal potential and designate this as E°. This is the standard potential of a redox couple with the oxidized and reduced forms at 1 M concentrations and with the solution conditions specified. For example, the formal potential of the Ce +/Ce + couple in 1 M HCl is 1.28 V. The Nemst equation is written as usual, using the formal potential in place of the standard potential. Table C.5 lists some formal potentials. 

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. 

Complexing agents replace water molecules in the hydration sphere of a dissolved metal ion. Their presence changes the stoichiometry of the dissolution reaction of the metal and generally leads to a lower standard potential (Table 2.15) facilitating corrosion. 

The curves in Fig. 7 are calculated from the values of the standard potentials (Table I) for two solvents at two temperatures. It is shown that in the presence of metallic titanium the proportion of Ti increases greatly at high concentrations of titanium salts. An increase in concentration has the same effect as the presence of strong donors of chloride ions. 

If we know the standard potential of a carefiilly-chosen cell (which may be determined experimentally or calculated on the basis of the electrode standard potential tables), then we can calculate the mean activity coefficient at any concentration. 

A problem that has fascinated surface chemists is whether, through suitable measurements, one can determine absolute half-cell potentials. If some one standard half-cell potential can be determined on an absolute basis, then all others are known through the table of standard potentials. Thus, if we know E for... 

Source Values are compiled from the following sources Bard, A. J. Parsons, R. Jordon, J., eds. Standard Potentials in Aqueous Solutions. Dekker New York, 1985 Milazzo, G. Carol , S. Sharma, V. K. Tables of Standard Electrode Potentials. Wiley London, 1978 Swift, E. H. Butler, E. A. Quantitative Measurements and Chemical Equilibria. Freeman New York, 1972. 

The thermodynamics of electrochemical reactions can be understood by considering the standard electrode potential, the potential of a reaction under standard conditions of temperature and pressure where all reactants and products are at unit activity. Table 1 Hsts a variety of standard electrode potentials. The standard potential is expressed relative to the standard hydrogen reference electrode potential in units of volts. A given reaction tends to proceed in the anodic direction, ie, toward the oxidation reaction, if the potential of the reaction is positive with respect to the standard potential. Conversely, a movement of the potential in the negative direction away from the standard potential encourages a cathodic or reduction reaction. 

Table 2-1 Conversion factors and standard potentials for electrochemical metal-metal ion reactions...

A comprehensive list of standard potentials is found in Ref. 7. Table 2-3 gives a few values for redox reactions. Since most metal ions react with OH ions to form solid corrosion products giving protective surface films, it is appropriate to represent the corrosion behavior of metals in aqueous solutions in terms of pH and Ufj. Figure 2-2 shows a Pourbaix diagram for the system Fe/HjO. The boundary lines correspond to the equilibria ... 

A. J. Bard, R. Parsons and J. Jordan Standard Potentials in Aqueous Solution, Marcel Dekker, New York, 1985, 834 pp. G. Milazzo and S. CarOli, Tables of Standard Electrode Potentials, Wiley, New York, 1978, 421 pp. 

Standard half-cell voltages are ordinarily obtained from a list of standard potentials such as those in Table 18.1 (page 487). The potentials listed are the standard voltages for reduction half-reactions, that is,... 

To obtain the standard voltage for an oxidation half-reaction, all you have to do is change the sign of the standard potential listed in Table 18.1. For example, knowing that... 

Using standard potentials listed in Table 18.1, decide whether at standard concentrations... 

H+], calculation of, 192, see also Hydrogen ion Haber, Fritz, 151 Haber process, 140, 150 Hafnium, oxidation number, 414 Haldane, J. B. S., 436 Half-cell potentials effect of concentration, 213 measuring, 210 standard, 210 table of, 211, 452 Half-cell reactions, 201 Half-life, 416 Half-reaction, 201 balancing, 218 potentials, 452 Halides... 

When the activity of the ion M"+ is equal to unity (approximately true for a 1M solution), the electrode potential E is equal to the standard potential Ee. Some important standard electrode potentials referred to the standard hydrogen electrode at 25 °C (in aqueous solution) are collected in Table 2.5.5... 

Several significant electrode potentials of interest in aqueous batteries are listed in Table 2 these include the oxidation of carbon, and oxygen evolution/reduction reactions in acid and alkaline electrolytes. For example, for the oxidation of carbon in alkaline electrolyte, E° at 25 °C is -0.780 V vs. SHE or -0.682 V (vs. Hg/HgO reference electrode) in 0.1 molL IC0 2 at pH [14]. Based on the standard potentials for carbon in aqueous electrolytes, it is thermodynamically stable in water and other aqueous solutions at a pH less than about 13, provided no oxidizing agents are present. 

Table 2. Standard potentials for reactions of carbon materials in batteries containing aqueous electrolytes...

E° values have been measured for many reactions and tabulated as standard half-cell potentials. Table 9.3 summarizes half-cell potentials as standard reduction potentials for a select set of reactions.aa In the tabulations, E° for... 

In Figure 2 the solubility and speciation of plutonium have been calculated, using stability data for the hydroxy and carbonate complexes in Table III and standard potentials from Table IV, for the waters indicted in Figure 2. Here, the various carbonate concentrations would correspond to an open system in equilibrium with air (b) and closed systems with a total carbonate concentration of 30 mg/liter (c,e) and 485 mg/liter (d,f), respectively. The two redox potentials would roughly correspond to water in equilibrium wit air (a-d cf 50) and systems buffered by an Fe(III)(s)/Fe(II)(s)-equilibrium (e,f), respectively. Thus, the natural span of carbonate concentrations and redox conditions is illustrated. 

The electrical double layer at pc-Zn/fyO interfaces has been studied in many works,154 190 613-629 but the situation is somewhat ambiguous and complex. The polycrystalline Zn electrode was found to be ideally polarizable for sufficiently wide negative polarizations.622"627 With pc-Zn/H20, the value of Eg was found at -1.15 V (SCE)615 628 (Table 14). The values of nun are in reasonable agreement with the data of Caswell et al.623,624 Practically the same value of Eff was obtained by the scrape method in NaC104 + HjO solution (pH = 7.0).190 Later it was shown154,259,625,628 that the determination of Eo=0 by direct observation of Emin on C,E curves in dilute surface-inactive electrolyte solutions is not possible in the case of Zn because Zn belongs to the group of metals for which E -o is close to the reversible standard potential in aqueous solution. 

TABLE 12.1 Standard Potentials Species i at 25°C Reduction half-reaction E° (V)... 

FIGURE 12.9 The variation of standard potentials through the main groups of the periodic table. Note that the most negative values are in the s block and that the most positive values are close to fluorine. 

In some cases, we find that available tables of data do not contain the standard potential that we need but do contain closely related values for the same element for instance, we might require the standard potential of the Ce4+/Ce couple, whereas we know only the values for the Ce3+/Ce and Ce4+/Ce3+ couples. In such cases, the potential of a couple cannot be determined by adding or subtracting the standard potentials directly. Instead, we calculate the values of AG° for each half-reaction and combine them into the AC° for the desired half-reaction. We then convert that value of AG° into the corresponding standard potential by using Eq. 2. 

STRATEGY First, write the balanced equation for the cell reaction and the corresponding expression for Q, and note the value of n. Then determine E° from the standard potentials in Table 12.1 or Appendix 2B. Determine the value of Q for the stated conditions. Calculate the emf by substituting these values into the Nernst equation, Eq. 6. At 25.00°C, RT/1 = 0.025 693 V. 

To reverse this half-reaction and bring about the oxidation of water, we need an applied potential difference of at least 0.82 V. Suppose the added salt is sodium chloride. When Cl ions are present at 1 mol-L 1 in water, is it possible that they, and not the water, will be oxidized From Table 12.1, the standard potential for the reduction of chlorine is Cl.36 V ... 

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