Electrode half-cell potential

Of the two half-cell electrode potentials, only the cell potential E can be measured experimentally. Therefore, it is not possible to measure the absolute values of any single half-cell electrode potential. To solve this problem, Nemst suggested that the potential of... 

The half-cell electrode potential as a function of the pH of the solution is given as ... 

A DC potential may develop at the electrode metal/solution interphase. The absolute potential of this interphase (half-cell electrode potential) cannot be measured—it must be considered unknown. However, the potential difference between two electrodes can be measured with an ordinary voltmeter connected to the two metal wires from the electrodes. If file metals were different, then they could generate a potential difference of 1 V or more. However, here we presume that the same electrode material is used and that the measured potential difference is small. We will discuss the case for three different electrode materials important in biological work platinum, silver coated with silver chloride (AgCl), and carbon. To the extent that both electrodes are equal, we have a symmetrical (bipolar) system, and the voltage—current dependence should not be dependent on polarity. 

Iron is similar to aluminum in that a protective oxide forms in nearly neutral solutions. However, for iron the field of oxide stability is substantially greater at elevated pH, and iron is far more resistant to alkaline solutions compared with aluminum. Contributing to the overall resistance of iron, is the generally more noble half-cell electrode potential for the anodic dissolution reactions which lower the driving force for corrosion reactions. It is, however, apparent from Fig. 2b that this resistance disappears in more acidic solutions [2]. 

Vi V 2 Standard half-cell electrode potentials (Table 17.1) for metals 1 and 2 (Reaction 17.17)... 

The half-cell reduction potential of the standard hydrogen electrode (SHE) was set arbitrarily to 0.000... V by international agreement. Since it is impossible to determine the potential of a single half-cell without comparing it to another, an arbitrary standard was established. 

The student then added 10. mL of 1.0 M NaCl solution to an empty beaker. He then added one drop of 1.0 M AgN03 to the beaker and stirred well. Since there is an abundance of Cl in the beaker, and the amount of NaCl (10. mL) is magnitudes greater than the one drop of 1.0 M AgN03, it can be assumed that most of the Ag+(aq) will combine with the CT(aq) and that the concentration of the Cl (aq) will remain essentially 1.0 M. A silver metal electrode was immersed in this solution and connected through the salt bridge to the Zn I Zn2+ half-cell. The potential difference was measured as 0.91 volts. 

Because you can measure potential differences, but not individual reduction potentials, all values in the table are relative. Each half-cell reduction potential is given relative to the reduction potential of the standard hydrogen electrode, which has been assigned a value of zero. The design of this electrode is shown in Figure 11.12. 

Measnrements of Ea are usually made with a platinum electrode placed in the soil solntion together with a reference half cell electrode of known potential. The platinnm electrode transfers electrons to and from the soil solution withont reacting with it. Reducing half reactions in the soil tend to transfer electrons to the platinum electrode and oxidizing half reactions to remove them. At eqnilibrinm no electrons flow and the electric potential difference between the half cell comprising the platinnm electrode and the soil solntion and the half cell comprising the reference electrode is recorded. 

Although potentials of half-cell electrode reactions cannot be measured, their consideration is extremely useful. For example, consider the two cells in series, as indicated in reaction (X),... 

However, electrochemical cells are most conveniently considered as two individual half-reactions, whereby each is written as a reduction in the form indicated by Eqs. (2.1)—(2.5). When this is done and values of the appropriate quantities are inserted, a potential can be calculated for each half-cell electrode system. Then that half-cell reaction with the more positive potential will be the positive terminal in a galvanic cell and the electromotive force of that cell will be represented by the algebraic difference between the potential of the more positive half-cell and the potential of the less positive half-cell ... 

Reference half-cells The fact that individual half-cell potentials are not directly measurable does not prevent us from defining and working with them. Although we cannot determine the absolute value of a half-cell potential, we can still measure its value in relation to the potentials of other half cells. In particular, if we adopt a reference half-cell whose potential is arbitrarily defined as zero, and measure the potentials of various other electrode systems against this reference cell, we are in effect measuring the half-cell potentials on a scale that is relative to the potential of the reference cell. 

In Equation (18b), the activity quotient is separated into the terms relating to the silver electrode and the hydrogen electrode. We assume that both electrodes (Ag+/Ag and H+/H2) operate under the standard condition (i.e. the H+/H2 electrode of our cell happens to constitute the SHE). This means that the equilibrium voltage of the cell of Figure 3.1.6 is identical with the half-cell equilibrium potential E°(Ag+l Ag) = 0.80 V. Furthermore, we note that the activity of the element silver is per definition unity. As the stoichiometric number of electrons transferred is one, the Nemst equation for the Ag+/Ag electrode can be formulated in the following convenient and standard way ... 

In thermodynamics, only energy differences are measurable absolute energies are not. Therefore, energies (or enthalpies or free energies) are defined relative to a reference state for which these quantities are arbitrarily set at 0 by international agreement. The same reasoning applies to half-cells Because only differences are measured, we are free to define a reference reduction potential for a particular half-cell and measure other half-cell reduction potentials relative to it. The convention used is to define %° for the half-cell reduction of Hilg) to H30 (d (j ) to be 0 at all temperatures, when the gas pressure at the electrode is 1 atm and the -iiO aq) concentration in solution is I M (Fig. 17.3). 

The foregoing example illustrates how equilibrium constants for overall cell reactions can be determined electrochemically. Although the example dealt with redox equilibrium, related procedures can be used to measure the solubility product constants of sparingly soluble ionic compounds or the ionization constants of weak acids and bases. Suppose that the solubility product constant of AgCl is to be determined by means of an electrochemical cell. One half-cell contains solid AgCl and Ag metal in equilibrium with a known concentration of CP (aq) (established with 0.00100 M NaCl, for example) so that an unknown but definite concentration of Kg aq) is present. A silver electrode is used so that the half-cell reaction involved is either the reduction of Ag (aq) or the oxidation of Ag. This is, in effect, an Ag" Ag half-cell whose potential is to be determined. The second half-cell can be any whose potential is accurately known, and its choice is a matter of convenience. In the following example, the second half-cell is a standard H30" H2 half-cell. 

Similarly, the external reference electrode portion of the galvanic cell is also composed of two potential-generating processes (1) the half-cell reaction potential of the reference electrode (e.g., Ag/AgCl, Hg/Hg2Cl2), and (2) the... 

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. 

Now we can construct a voltaic cell consisting of this reference half-cell and another half-cell whose potential we want to determine. With E eference defined as zero, the overall Eceii allows us to find the unknown standard electrode potential. 

Most sensors are potentiometric and are based on concentration cells. A typical arangement is shown schematically in Figure 17. The solid electrolyte, b, is sandwiched between two electrodes, c. One is exposed to the test gas and the other either to a known concentration of gas or some reference half cell. The potential... 

The overall charge transfer of any corrosion reaction must be the sum of two electrode potentials that are established in the half-cells. Each potential results from the change of the... 

Table 2.4 Equilibrium Potential Values for Commonly Used Reference Electrodes Reference Electrode Half-Cell Reaction Potential V vs. NHE...

The Ha/H standard potential is equal to zero. By substituting the hydrogen partial pressure values in the Nernst equation one can obtain the potential of a hydrogen electrode under the conditions described in the problem. The general Nernst equation for the half-cell electrode reaction is ... 

It is important to record the equipment used. Different half cells have different offsets . The silver/silver chloride half cell gives potentials that are a function of the chloride concentration in them. This is usually about 130 mV more positive than a copper/saturated copper sulphate electrode. This can be compensated for internally in the logging equipment if used or during reporting if the ASTM criteria are being used (see below). 

Based on the individual half-cell reaction potentials, the theoretical electrochemical potential offered by a single Zn/Br cell should be approximately 1.828 V. This value is the Nemstian potential under zero current flow. However, the presence of internal inefficiencies and various resistance contributions seen in practice are expected to result in slightly lower cell voltage values. Another important performance metric for Zn/Br systems is current density, which is the amount of current passing through a unit area of an electrode surface. The current density, in turn, has a direct influence on the electrode capacity (i.e. energy per unit area) as well as the operating efficiency of the overall system. 

Ostwald had what appeared to be a very elegant concept. It involved the measurement of a single electrode potential. The method of measurement was in good accordance with his philosophical views and with the chemistry of the times, and it would, in his opinion, yield an absolute potential. An absolute potential was a sharp contrast to the relative potential obtained by referring a measured half-cell to another single electrode reaction arbitrarily set at zero. Ostwald s measurements of half-cell potentials could be directly related to heats of ionization (1 ). In his opinion, an absolute half-cell redox potential would allow the establishment of an electromotive series which would be analogous to the absolute temperature scale. 

In order to determine the standard electrode potential for a metal, the galvanic cell is designed so that a half-cell is formed by a piece of metal immersed in a solution that contains 1.00 M of ions of that metal, and one half-cell with potential convention defined to be exactly zero volts. This electrode is called the standard hydrogen electrode (SHE) and it consists of a platinum electrode over which Hj gas at 1-atm of pressure is bubbled, immersed in a solution that contains 1.00 M of hydronium ion at 25°C. 

Additionally, electrochemical polarization is a measure of the overpotential and represents a deviation of the electrochemical state of half-cell electrodes induced by an applied external potential. Therefore, the driving force for electrochemical polarization is the overpotential. 

The silver-silver chloride half-cell electrode develops a potential proportional to the chloride concentration, whether it is sodium chloride, potassium chloride, ammonium chloride, or some other chloride salt, and remains constant as long as the chloride concentration remains constant. 

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