Indicator electrode potential

Before the equivalence point, the indicator electrode potential is fairly steady near °(Fe3+ Fe2+) = 0.77 V.4 After the equivalence point, the indicator electrode potential levels off near °tCe41 Ce3+) = 1.70 V. At the equivalence point, there is a rapid rise in voltage. 

Equation 2.16 shows that potentiometry is a valuable method for the determination of equilibrium constants, ffowever, it should be borne in mind that the system should be in equilibrium. Some other conditions, which are described below, also need to be fulhlled for use of potentiometry in any application. The basic measurement system must include an indicator electrode that is capable of monitoring the activity of the species of interest, and a reference electrode that gives a constant, known half-cell potential to which the measured indicator electrode potential can be referred. The voltage resulting from the combination of these two electrodes must be measured in a manner that minimises the amount of current drawn by the measuring system. This condition includes that the impedance of the measuring device should be much higher than that of the electrode. 

For this system the indicator-electrode potential is represented by either halfreaction ... 

Potentiometric titration curves normally are represented by a plot of the indicator-electrode potential as a function of volume of titrant, as indicated in Fig. 4.2. However, there are some advantages if the data are plotted as the first derivative of the indicator potential with respect to volume of titrant (or even as the second derivative). Such titration curves also are indicated in Figure 4.2, and illustrate that a more definite endpoint indication is provided by both differential curves than by the integrated form of the titration curve. Furthermore, titration by repetitive constant-volume increments allows the endpoint to be determined without a plot of the titration curve the endpoint coincides with the condition when the differential potentiometric response per volume increment is a maximum. Likewise, the endpoint can be determined by using the second derivative the latter has distinct advantages in that there is some indication of the approach of the endpoint as the second derivative approaches a positive maximum just prior to the equivalence point before passing through zero. Such a second-derivative response is particularly attractive for automated titration systems that stop at the equivalence point. 

Figure 2. (a). In situ XANES spectra for PtMo catalyst collected at the Mo k-edge (20,000 eV) at the indicated electrode potentials, (b). XANES Calibration curve constructed using the change in edge energy as a function of oxidation state of Mo. 

The first term in this equation, Ejnj, contains the information that we are looking for—the concentration of the analyte. To make a potentiometric determination of an analyte, then, we must measure a cell potential, correct this potential for the reference and junction potentials, and compute the analyte concentration from the indicator electrode potential. Strictly, the potential of a galvanic cell is related to the activity of the analyte. Only through proper calibration of the electrode system with solutions of known concentration can we determine the concentration of the analyte. 

To make potential measurements, a complete cell consisting of two half-cells must be set up, as was described in Chapter 12. One half-cell usually is comprised of the test solution and an electrode whose potential is determined by the analyte we wish to measiue. This electrode is the indicator electrode. The other half-cell is any arbitrary half-cell whose potential is not dependent on the analyte. This halfcell electrode is designated the reference electrode. Its potential is constant, and the measured cell voltage reflects the indicator electrode potential relative to that of the reference electrode. Since the reference electrode potential is constant, any changes in potential of the indicator electrode will be reflected by an equal change in the cell voltage. 

We create a voltaic cell with the indicator and reference electrodes. We measure the voltage of the cell, giving a reading of the indicator electrode potential relative to the reference electrode. We can relate this to the analyte activity or concentration using the Nemst equation. 

Fig. 10. ATR spectra of the CN stretching vibration of thiocyanate at an Ag/electrolyte interface measured at indicated electrode potentials (V) vs. Ag/AgCl. The angle of incidence of radiation is 70°. Electrolyte (a), (b) 0.1 M NaC104 + 0.01 M NaSCN (c) 0.1 M NaCl + 0.01 M NaSCN. Thickness of Ag electrode (a), (c) 18 nm (b) 17 nm. Measurement order is top to bottom in each case.

The sign convention for potentiometry is consistent with the convention described in Chapter 22 for standard electrode potentials. In this convention, the in-dicalor electrode is the right-hand electrode and the reference elecirode is on the left, - For direct poten-tiomctric measurements, the potential of a cell is then expressed in terms of the indicator electrode potential, the reference elecirode potential, and the junction potential as shown in liquation 23-1. 

Compare this equation with Eqs. (15.7) and (15.15). By convention, the reference electrode is connected to the negative terminal of the potentiometer (the readout device). The common reference electrodes used in potentiometry are the SCE and the silver/silver chloride electrode, which have been described. Their potentials are fixed and known over a wide temperature range. Some values for these electrode potentials are given in Table 15.3. The total cell potential is measured experimentally, the reference potential is known, and therefore the variable indicator electrode potential can be calculated and related to the concentration of the analyte through the Nemst equation. In practice, the concentration of the unknown analyte is determined after calibration of the potentiometer with suitable standard solutions. The choice of reference electrode depends on the application. For example, the Ag/AgCl electrode cannot be used in solutions containing species such as halides or sulfides that will precipitate or otherwise react with silver. 

In a direct potentiometric measurement, we use an electrode such as a silver wire to measure [Ag ] or a pH electrode to measure [H" ] or a calcium ion-selective electrode to measure [Ca " ]. There is inherent inaccuracy in most direct potentiometric measurements because there is usually a liquid-liquid junction with an unknown voltage difference making the intended indicator electrode potential uncertain. For example. Figure 15-5 shows a 4% standard deviation among 14 measurements by direct potentiometry. Part of the variation could be attributed to differences in the indicator (ion-selective) electrodes and part could be from varying liquid junction potentials. 

Cell voltage = E+ - (reference electrode) where E+ is the indicator electrode potential Before equivalence point Analyte is in excess use analyte Nemst equation to find indicator electrode potential. 

After equivalence point Titrant is in excess use titrant Nemst equation to find indicator electrode potential. 

FIGURE 3.23. Equilibrium potential drop across the polymer/electrolyte interface for a PMPy film in aqueous NaCl electrolytes of concentrations at the indicated electrode potentials (in V). Experimental data from Fig. 3.22 calculated values based on Eqn. 18. 

A titration curve is drawn showing the variation of indicator electrode potential with the quantity of titrant solution added, such a curve having the general form illustrated in Fig. 1. Here the sharp change of potential is evident. 

Fig. 20.41 Energy-dispersive analysis of X-ray emission (EDAX) spectra of the PPy/ROSO -coated gold electrode of Fig. 20.40. For the analyses the electrode was withdrawn from the electrolyte (KCl) at the indicated electrode potentials. The elements corresponding to the X-ray emission peaks from the polymer are indicated by their chemical symbols.

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