Standard half-cell potentials definition

Because the standard-state half-cell potential, , is measured relative to the zero potential of the hydrogen half-cell, = El, and the definition of " given by equation 7.27 is substituted into equation 7.22 to give... 

Erhe andELHE are equilibrium half-cell potentials, or electrode potentials, which depend in sign on the definitions of positive and negative electricity and assignment of Eu + = 0 at standard condi-... 

The cathode consists of platinum which is an inert conductor in contact with the 1 M ions surrounded by hydrogen gas at 1 atm. Such an electrode is called a standard hydrogen electrode which per definition has a half cell potential (symbolised at 298 K by the symbol of s°) of 0.00 volt. The figure below shows the principle in the build up of the standard hydrogen electrode. 

The standard hydrogen electrode has per definition a half cell potential of 0.0 Volt at a Tf concentration of 1,0 M. 

The most widely used reference electrode, due to its ease of preparation and constancy of potential, is the calomel electrode. A calomel half-cell is one in which mercury and calomel [mercury(I) chloride] are covered with potassium chloride solution of definite concentration this may be 0.1 M, 1M, or saturated. These electrodes are referred to as the decimolar, the molar and the saturated calomel electrode (S.C.E.) and have the potentials, relative to the standard hydrogen electrode at 25 °C, of 0.3358,0.2824 and 0.2444 volt. Of these electrodes the S.C.E. is most commonly used, largely because of the suppressive effect of saturated potassium chloride solution on liquid junction potentials. However, this electrode suffers from the drawback that its potential varies rapidly with alteration in temperature owing to changes in the solubility of potassium chloride, and restoration of a stable potential may be slow owing to the disturbance of the calomel-potassium chloride equilibrium. The potentials of the decimolar and molar electrodes are less affected by change in temperature and are to be preferred in cases where accurate values of electrode potentials are required. The electrode reaction is... 

It is very often necessary to characterize the redox properties of a given system with unknown activity coefficients in a state far from standard conditions. For this purpose, formal (solution with unit concentrations of all the species appearing in the Nernst equation its value depends on the overall composition of the solution. If the solution also contains additional species that do not appear in the Nernst equation (indifferent electrolyte, buffer components, etc.), their concentrations must be precisely specified in the formal potential data. The formal potential, denoted as E0, is best characterized by an expression in parentheses, giving both the half-cell reaction and the composition of the medium, for example E0,(Zn2+ + 2e = Zn, 10-3M H2S04). 

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 ... 

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. 

As already discussed, the standard hydrogen electrode (SHE) is the chosen reference half-cell upon which tables of standard electrode potentials are based. The potential of this system is zero by definition at all temperatures. Although this reference electrode was often used in early work in electrochemistry, it is almost never seen in chemical laboratories at the present time. It is simply too awkward to use because of the requirement for H2 gas at 1 bar pressure and safety considerations. 

Table 19.1 lists standard reduction potentials for a number of half-cell reactions. By definition, the SHE has an E° value of 0.00 V. Above the SHE the negative standard reduction potentials increase, and below it the positive standard reduction potentials increase. It is important to know the following points about the table ... 

The total voltage developed under standard conditions is -1-0.76 V. But the voltage of the SHE is 0 by definition therefore the standard reduction potential of the Zn half-cell is ... 

Equation (5) or (11) can be applied directly to half-cell reactions such as (6) and (7) and the resulting potentials obtained will be identical to those obtained from the overall reactions (9) and (10) because of the definition of the SHE as the universal standard. A selection of standard potentials of half-cell reactions is shown in Table 1 [5]. By international convention, electrode reactions in thermodynamic tables are always written as reduction reactions, so the more noble metals have a positive standard potential. Lists such as that in Table 1 are also called electromotive force series or tables of standard reduction potentials. 

By definition, the potential of a redox couple vs SHE is the emf of the Active electrochemical cell, whereby the working electrode is in the half-cell involving the redox couple in question. The counter-electrode is the standard hydrogen electrode at the same temperature. The terminals are made of the same metals and the sum of the possible ionic junction voltages are considered to be equal to zero. In this case we therefore have E,she =... 

As we shall see in Section 21-14, this arbitrary definition will provide a basis on which to describe the relative potential (and hence the relative spontaneity) of any half-ceU process. We can 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 volts to the sum, the voltage of the overall cell is then attributed entirely to the other half-cell. 

Various standard electrode potentials can be calculated for a half-cell after the definition of standard electrode potential. Currently, most of standard electrode potentials can be obtained from the text. 

Suppose that we want to find the standard potential for the half-cell Hg " I Hg. By definition, this potential is 8 for the cell... 

Definitions. Define briefly (a) difference of potential, (b) electromotive force, (c) salt bridge, (d) anode, (e) positive electrode, (f) reference half-cell, (g) standard electrode potential, (h) decomposition potential, (i) overvoltage, (j) sacrificial anode. 

There are several potential benchmarks in common use, but the most ancient is the half-cell in which hydrogen gas is bubbled over a platinum electrode immersed in a solution having a known concentration of hydrogen ions. This historically important reference electrode is called the standard hydrogen electrode (SHE) if a standard solution of acid is used. By definition, the equilibrium potential of this electrode is zero at any temperature. However, the SHE can be somewhat inconvenient to use because of the need to supply hydrogen gas. Therefore, other reference electrodes are much preferred for practical considerations. 

The most acceptable method of obtaining standard electrode potentials is by comparing tbe electrode potential of metals with the standard hydrogen electrode. Since the SHE has zero electrode potential at all temperatures by definition, the electrode potential of a metal is numerically equal to the emf of the cell formed by SHE and the metal electrode. In other words, the emf of the cell represents the electrode potential of the half cell formed by the metal with respect to the standard hydrogen electrode. In such a cell, reaction on the hydrogen electrode is oxidation and reaction on the other electrode is reduction. Such a cell can be expressed as ... 

An important conclusion to be made here is related to the definition of the standard electrode potential given in the lUPAC manual [1] The standard potential of an electrochemical reaction, abbreviated as standard potential, is defined as the standard potential of a hypothetical cell, in which the electrode (half-cell) on the left of the cell diagram is the SHE and the electrode at the right is the electrode in question. Note that E of a half-reaction (or a total electrochemical reaction) as an intensive variable does not depend on the number of electrons nsed in the half-reaction. 

Reference electrode An electrochemical half-cell with a stable and known potential. The most common is the saturated calomel electrode (SCE,-1-0.242 V vs NHE). The primary standard reference electrode is the normal hydrogen electrode (NHE, 0.0000 V by definition). 

Many half-reactions of interest to biochemists involve protons. As in the definition of AG °, biochemists define the standard state for oxidation-reduction reactions as pH 7 and express reduction potential as E °, the standard reduction potential at pH 7. The standard reduction potentials given in Table 13-7 and used throughout this book are values for E ° and are therefore valid only for systems at neutral pH Each value represents the potential difference when the conjugate redox pair, at 1 m concentrations and pH 7, is connected with the standard (pH 0) hydrogen electrode. Notice in Table 13-7 that when the conjugate pair 2ET/H2 at pH 7 is connected with the standard hydrogen electrode (pH 0), electrons tend to flow from the pH 7 cell to the standard (pH 0) cell the measured E ° for the 2ET/H2 pair is -0.414 V... 

You have probably worked with tables of standard reduction potentials before. These tables provide the reduction potentials of various substances. It describes an oxidized species s ability to gain electrons in a reduction half-reaction (like copper in the voltaic cell example). According to this definition, we can use a value from the table to represent the E°red in the expression above, but how do you find the E°ox ... 

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