Standard Potentials and pH

In HCIO4, HCl and H2SO4 media the pH is sufficiently low to prevent most (but not aU) of these reactions and so the formal potential is not too far removed from the standard electrode potential. 

At higher pH (IM K2C2O4, pH 5) hydrolysis and polymerisation will occur and this is reflected in the huge discrepancy between the standard and formal potentials. In addition, both Fe and Fe + can form couples with the oxalate anion (see Problem 1.17) and these effects also contribute to the large observed difference. In H3PO4 there is likely complexation of both Fe and Fe + with the anions. 

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Table 2.1 states the redox relations at standard conditions. Extended information on the distribution of the redox pairs — still under equilibrium conditions but under varying redox potential and pH — is given in a Pourbaix diagram. Figure 2.4 is an example of such a diagram for the binary sulfur and oxygen system in water at 1 atm and 25°C with the sum of the concentrations of... 

Usually, the thermodynamic properties involved in the electrochemical reactions of Si element are presented in the Pourbaix diagram showing the lines where one species is transformed into another as a function of the potential and pH values. In the case of Si, this presentation is not necessary. In fact, the standard... 

Fig. 3.3.3 Change of the Cd potential and pH of the Cd(OH)2 suspension with the formation of CdS particles under the standard conditions. (From Ref. 2.)...

The thermodynamic information is normally summarized in a Pourbaix diagram7. These diagrams are constructed from the relevant standard electrode potential values and equilibrium constants and show, for a given metal and as a function of pH, which is the most stable species at a particular potential and pH value. The ionic activity in solution affects the position of the boundaries between immunity, corrosion, and passivation zones. Normally ionic activity values of 10 6 are employed for boundary definition above this value corrosion is assumed to occur. Pourbaix diagrams for many metals are to be found in Ref. 7. 

Here Egl is the standard potential of the glass electrode. This quantity varies from specimen to specimen, it depends also on the age and on the pretreatment of the electrode. Within one set of measurements it can be regarded as constant. If we adapt the usual calibration process, described below, it is not necessary to measure the standard potential and to deduct the potential of the calomel electrode from the results, as the pH can be read directly from the pH-meter. 

The water stability boundaries and the locus of measured Eh and pH measurements in natural waters, as reported by Baas-Becking et al. (I960), are shown in Fig. 11.3 (see also Fig. 11.4). It has been observed that frequently the Eh values measured with a Pt electrode differ significantly from values computed from Gibbs free energies or standard potentials and solution concentrations. When they exist, there are two important reasons for such differences. These include (1) misbehavior of the Pt or other indicator electrode (2) the irreversibility or slow kinetics of most redox couple reactions and resultant di.sequilibrium between and among different redox couples in the same water and (3) the common existence of mixed potentials in natural waters (see below). 

The actual potential of a cell (e) depends not only on the standard potential (e°) but also on the concentrations of the reactants and products in solution or their pressures in the gas phase (the things that can vary). Since the concentration of its components affects the potential of a cell, we can use a voltaic cell to measure a concentration if we know its standard potential and we measure its actual potential for the reaction. The pH meter is based on this principle. For either a cell or a half-cell, the potential is given by the Nernst equation. The Nernst equation for the general reaction... 

Figure 1. Representative redox couples of various organic pollutants and iron species. Given are standard potentials at pH = 7 and molar concentrations of the reactants but at standard environmental concentrations of the major anions involved [HCO3-] = [C1-] = 10-3 M [Bi ] = 10-5 M (s) = solid. Adapted from (44).

A standard potential vs. pH diagram correlates the thermodynamics of the aquocobalamin (4+)-Bi2r (23)-Bi2s (40") system (see Fig. 5). The interconversion between the different oxidation states of B -derivatives can usually be monitored effectively by UV-vis spectroscopy, and the relevant data were obtained from potentiostatic measurements, which were followed by UV-vis spectroscopy [90,94], Within the pH range - 1 to 11 and applied potentials = 0.5 V and -1.2 V vs. SCE, seven solution cobalamins are thermodynamically predominant spanning a range of the three formal oxidation states of Bi2 [90]. 

The zero (isopotential) control provides the flexibility to standardize at a point other than pH 7 and then make a slope adjustment without affecting the standardization point. It provides greater accuracy by allowing the two-point calibration to take place over a narrower pH range. The zero control provides a potential to offset the ideal standardization potential and thus provide zero millivolts at a point other than pH 7. The zero control is first adjusted to pH 10.0 when in the standby mode which separates the electrodes from the meter. Then the electrodes are standardized in a pH 10.0... 

The potentiometric determination of an analyte s concentration is one of the most common quantitative analytical techniques. Perhaps the most frequently employed, routine quantitative measurement is the potentiometric determination of a solution s pH, a technique considered in more detail in the following discussion. Other areas in which potentiometric applications are important include clinical chemistry, environmental chemistry, and potentiometric titrations. Before considering these applications, however, we must first examine more closely the relationship between cell potential and the analyte s concentration, as well as methods for standardizing potentiometric measurements. 

Two methods are used to measure pH electrometric and chemical indicator (1 7). The most common is electrometric and uses the commercial pH meter with a glass electrode. This procedure is based on the measurement of the difference between the pH of an unknown or test solution and that of a standard solution. The instmment measures the emf developed between the glass electrode and a reference electrode of constant potential. The difference in emf when the electrodes are removed from the standard solution and placed in the test solution is converted to a difference in pH. Electrodes based on metal—metal oxides, eg, antimony—antimony oxide (see Antimony AND ANTIMONY ALLOYS Antimony COMPOUNDS), have also found use as pH sensors (8), especially for industrial appHcations where superior mechanical stabiUty is needed (see Sensors). However, because of the presence of the metallic element, these electrodes suffer from interferences by oxidation—reduction systems in the test solution. 

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