Other standard electrodes

The standard hydrogen electrode is not the most convenient of electrodes to set up, because of the necessity of bubbling hydrogen over the platinum electrode. Several other electrodes are commonly used as secondary standard electrodes. One of these is the standard silver-silver chloride electrode, in which a silver electrode is in contact with solid silver chloride, which is a highly insoluble salt. The whole is immersed in potassium chloride solution in which the chloride ion concentration is 1 m. This electrode can be represented as 

We can set up a cell involving this electrode and the hydrogen electrode, 

The standard electrode potential for the silver-silver chloride electrode is thus 0.2224 V. 

The emT at 25° C is 0.3338 V, which is thus the standard electrode potential F . If a saturated solution of KCI is used with the calomel electrode, the standard electrode potential is 0.2415 V. 

I Another electrode comnnonly used as a secondary standard is the glass electrode, 

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The SHE. The H" " H2 couple is the basis of the primary standard around which the whole edifice of electrode potentials rests. We call the H H2 couple, under standard conditions, the standard hydrogen electrode (SHE). More precisely, we say that hydrogen gas at standard pressure, in equilibrium with an aqueous solution of the proton at unity activity at 298 K has a defined value of of 0 at all temperatures. Note that all other standard electrode potentials are temperature-dependent. The SHE is shown schematically in Figure 3.3, while values of Eq r are tabulated in Appendix 3. 

In practice, experimental redox potentials are reported relative to a standard electrode. If the standard is the NHE, one subtracts 4.36 V from the absolute reduction potential (the cost of the free electron) or adds 4.36 V to the absolute oxidation potential (the return from the removed electron) in order to determine the relative potential. Adjustment to other standard electrodes is straightforward, since their potentials relative to the NHE are well known. 

We then construct a standard cell consisting of a standard hydrogen electrode and some other standard electrode (half-cell). Because the defined electrode potential of the SHE contributes exactly 0 volt to the sum, the voltage of the overall cell then lets us determine the standard electrode potential of the other half-cell. This is its potential with respect to the standard hydrogen electrode, measured at 25°C when the concentration of each ion in the solution is 1 M and the pressure of any gas involved is 1 atm. 

We can develop a series of standard electrode potentials by measuring the potentials of other standard electrodes versus the SHE in the way we described for the standard Zn-SHE and standard Cu-SHE voltaic cells. Many electrodes involve metals or nonmetals in contact with their ions. We saw (Section 21-12) that the standard Zn electrode behaves as the anode versus the SHE and that the standard oxidation potential for the Zn half-cell is 0.763 volt. 

Addition of Free Enthalpies and Calculation of Standard Electrode Potentials from Other Standard Electrode Potentials... 

This electrode is given a standard electrode potential of 0.00 V. All other standard electrode potentials are measured relative to this value, state symbol a symbol used in a chemical equation that describes the state of each reactant and product (s) for solid, (1) for liquid, (g) for gas and (aq) for substances in aqueous solution. 

The standard electrode potential of magnesium is given, along with the potentials of other metals, in Table 4.17 and the steady-state potentials of magnesium in various solutions are listed in Table 4.18. ... 

It is apparent that since the electrode potential of a metal/solution interface can only be evaluated from the e.m.f. of a cell, the reference electrode used for that purpose must be specified precisely, e.g. the criterion for the cathodic protection of steel is —0-85 V (vs. Cu/CuSOg, sat.), but this can be expressed as a potential with respect to the standard hydrogen electrode (S.H.E.), i.e. -0-55 V (vs. S.H.E.) or with respect to any other reference electrode. Potentials of reference electrodes are given in Table 21.7. 

When metals are arranged in the order of their standard electrode potentials, the so-called electrochemical series of the metals is obtained. The greater the negative value of the potential, the greater is the tendency of the metal to pass into the ionic state. A metal will normally displace any other metal below it in the series from solutions of its salts. Thus magnesium, aluminium, zinc, or iron will displace copper from solutions of its salts lead will displace copper, mercury, or silver copper will displace silver. 

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)]. 

One distingnishes practical and standard reference electrodes. A standard RE is an electrode system of particnlar confignration, the potential of which, nnder specified conditions, is conventionally taken as zero in tfie corresponding scale of potentials (i.e., as the point of reference nsed in finding tfie potentials of otfier electrodes). Practical REs are electrode systems having a snfficiently stable and reproducible value of potential which are nsed in the laboratory to measure the potentials of other electrodes. The potential of a practical reference electrode may difier from the conventional zero potential of the standard electrode, in which case the potential of the test electrode is converted to this scale by calculation. 

In the example just studied, the electrolysis of HC1 solution, the ions that transport the current (H+ and Cl-) are also the ones that are discharged at the electrodes. In other cases, however, the main ionic transporters of current may not be of the same species as the ions that are discharged. An excellent example is the electrolysis of CuS04 solution between platinum electrodes. A one molal CuS04 solution is quite acid so that the positive current transporters are both Cu2+ and H+ ions. The main negative transporter is the S04 ions. The solution contains, however, a small concentration of OH- ions. In order to determine which ions will be discharged at the electrodes, it is necessary to consider standard electrode potentials of the concerned species ... 

The relationships of the type (3.1.54) and (3.1.57) imply that the standard electrode potentials can be derived directly from the thermodynamic data (and vice versa). The values of the standard chemical potentials are identified with the values of the standard Gibbs energies of formation, tabulated, for example, by the US National Bureau of Standards. On the other hand, the experimental approach to the determination of standard electrode potentials is based on the cells of the type (3.1.41) whose EMFs are extrapolated to zero ionic strength. 

Although a few amperometric pH sensors are reported [32], most pH electrodes are potentiometric sensors. Among various potentiometric pH sensors, conventional glass pH electrodes are widely used and the pH value measured using a glass electrode is often considered as a gold standard in the development and calibration of other novel pH sensors in vivo and in vitro [33], Other pH electrodes, such as metal/metal oxide and ISFETs have received more and more attention in recent years due to their robustness, fast response, all-solid format and capability for miniaturization. Potentiometric microelectrodes for pH measurements will be the focus of this chapter. 

QB For this cell because the electrodes are identical, the standard electrode potentials are numerically equal and subtracting one from the other leads to the value c°dl = 0.000 V. However, because the ion concentrations differ, there is a potential difference between the two half cells (non-zero nonstandard voltage for the cell). [Pb2+] = 0.100 M in the cathode compartment. The anode compartment contains a saturated solution of Pbl2. We use the Nemst equation (with n = 2) to determine [Pb2+] in the saturated solution. 

The standard electrode potential [1] of an electrochemical reaction is commonly measured with respect to the standard hydrogen electrode (SHE) [2], and the corresponding values have been compiled in tables. The choice of this reference is completely arbitrary, and it is natural to look for an absolute standard such as the vacuum level, which is commonly used in other branches of physics and chemistry. To see how this can be done, let us first consider two metals, I and II, of different chemical composition and different work functions 4>i and 4>ii-When the two metals are brought into contact, their Fermi levels must become equal. Hence electrons flow from the metal with the lower work function to that with the higher one, so that a small dipole layer is established at the contact, which gives rise to a difference in the outer potentials of the two phases (see Fig. 2.2). No work is required to transfer an electron from metal I to metal II, since the two systems are in equilibrium. This enables us calculate the outer potential difference between the two metals in the following way. We first take an electron from the Fermi level Ep of metal I to a point in the vacuum just outside metal I. The work required for this is the work function i of metal I. 

Then, knowing F H2, it was relatively easy to determine values of electrode potentials for any other couple. With this methodology, they devised the standard electrode potentials ° scale (often called the E nought scale , or the hydrogen scale ). 

Additionally, other reference electrodes are used which are easier to maintain at standard conditions. These include the silver/silver chloride electrode and the saturated calomel electrode (SCE). The voltage difference between the working electrode and the reference electrode is proportional to the electrochemical potential difference between them. This is written... 

The standard electrode potential of an element is defined as its electrical potential when it is in contact with a molar solution of its ions. For redox systems, the standard redox potential is that developed by a solution containing molar concentrations of both ionic forms. Any half-cell will be able to oxidize (i.e. accept electrons from) any other half-cell which has a lower electrode potential (Table 4.1). 

It is impossible to measure the potential of a half-cell directly and a reference half-cell must be used to complete the circuit. The hydrogen electrode (Figure 4.3) is the standard reference electrode against which all other halfcells are measured and is arbitrarily attributed a standard electrode potential of zero at pH 0. Because it is difficult to prepare and inconvenient to use, the... 

Nevertheless, since in thermodynamics the values of the standard potentials are normally referred to the NHE electrode, when one wishes to obtain thermodynamic data from electrochemical experiments it is usual to transform the redox potentials obtained employing other reference electrodes into values with respect to the NHE, according to the scales given in Chapter 3, Section 1.2. 

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