The standard electromotive force of a cell

In addition, the cell reaction must also be reversible. The concept of equilibrium requires that, for a given system, the same state is obtained when equilibrium is approached from any direction. For a cell, the chemical reaction that takes place when the potential drop along the slide wire of the potentiometer is slightly greater than the emf of the cell must be the same as, but in the opposite direction to, that which takes place when the potential drop is slightly less than the emf of the cell. 

When we write the cell reaction as v,B, = 0, the change of the Gibbs energy for the change of state can be written as 

for the cells represented in Equations (12.35) and (12.36), Equation (12.41) becomes 

The determination of the value of SB may be illustrated by the use of Equation (12.42). We choose the standard state of silver and silver chloride to be the solid state at the experimental temperature and pressure, so that the activities are unity. We also express the activities of the hydrogen and chloride ions in terms of their molalities and activity coefficients. Equation (12.41) then becomes 

SB is the value of the left-hand side of Equation (12.44) in the limit of zero molality. 

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The valne (E°) for the standard electromotive force of a cell in which hydrogen nnder standard conditions is oxidized to hydroninm ions (solvated protons) at the left-hand electrode. This value is used as a standard to measure electrode potentials. 

According to its definition, the standard (reduction) potential of the A/A couple is the standard electromotive force of a cell in which an A/A electrode (where the activities of A and A are made unity) is opposed to an NHE (normal hydrogen electrode) whose potential is assigned to zero by convention. 

The equilibrium constant for a redox reaction can be found from the standard electromotive force of a cell. 

The decrease in free energy of the system in a spontaneous redox reaction is equal to the electrical work done by the system on the surroundings, or AG = nFE. The equilibrium constant for a redox reaction can be found from the standard electromotive force of a cell. 10. The Nernst equation gives the relationship between the cell emf and the concentrations of the reactants and products under non-standard-state conditions. Batteries, which consist of one or more galvanic cells, are used widely as self-contained power sources. Some of the better-known batteries are the dry cell, such as the Leclanche cell, the mercury battery, and the lead storage battery used in automobiles. Fuel cells produce electrical energy from a continuous supply of reactants. 

The electrode potential (reduction potential) for a redox couple is defined as the couple s potential measured with respect to the standard hydrogen electrode, which is set equal to zero (see hydrogen electrode later). This potential, by convention, is the electromotive force of a cell, where the standard hydrogen electrode is the reference electrode (left electrode) and the given half-cell is the indicator electrode (right electrode). The reduction potential for a given redox couple is given by the Nernst equation ... 

Electrode potential - The electromotive force of a cell in which the electrode on the left is the standard hydrogen electrode and that on the right is the electrode in question. [2]... 

No valid method exists for determining the absolute potential of an electrode. All potential measurements require a second electrode, and are therefore relative potentials. It is then necessary to choose one particular electrode to be arbitrarily assigned the zero position on the potential scale. By convention, the standard hydrogen electrode (SHE) is defined to have a potential of exactly 0 V, and is the reference point from which the potentials of all other electrodes are stated (Sec. 2.3). The electromotive force of a cell is the potential difference between two electrodes and is independent of the particular reference-electrode scale used. 

E = the electromotive force of a cell containing the unknown solution = the electromotive force of a cell containing a standard reference buffer solution of known or defined pH, that is, pH,... 

Before we discuss standard electrode potential, we will talk about electromotive force (emf). The electromotive force of a cell is the potential difference between the two electrodes. This can be measured using a voltmeter. The maximum voltage of a cell can be calculated using experimentally determined values called standard electrode potentials. By convention, the standard electrode potentials are usually represented in terms of reduction half-reactions for 1 molar solute concentration. The standard electrode potential values are set under ideal and standard-state conditions (latm pressure and 25°C temperature). From the MCAT point of view, you can assume that the conditions are standard, unless stated otherwise. Table 12-1 shows a list of standard electrode potentials (in aqueous solution) at 25°C. 

In equation 7-3 the term (RT/a ln K can be replaced by E , which is known as the standard electromotive force of the cell. E° is seen to be the value of the cell electromotive force that would be obtained if all of the substances in the reaction were at unit activity. On the other hand, if the system were at equilibrium, then the value of E as well as AG would be zero. 

The electrode potential of any single redox couple is defined as the electromotive force of a cell consisting of the standard hydrogen electrode and the electrode in question, written in the following manner ... 

Electromotive force n. This force is defined as that which causes a flow of current. The electromotive force of a cell is measured by the maximum difference of potential between its plates. The electromagnetic unit of potential difference is that against which 1 erg of work is done in the transfer of electromagnetic unit quantity. The volt is that potential difference against which 1 J of work is done in the transfer of 1 C. One volt is equivalent to 10 electromagnetic units of potential. The international volt is the electrical potential which when steadily applied to a conductor whose resistance is one international ohm will cause a current of one international ampere to flow. The international volt = 1.00033 absolute volts. The electromotive force of a Weston standard cell is 1.0183 int. volts at 20°C. Dimensions T ],... 

A second source of standard free energies comes from the measurement of the electromotive force of a galvanic cell. Electrochemistry is the subject of other articles (A2.4 and B1.28). so only the basics of a reversible chemical cell will be presented here. For example, consider the cell conventionally written as... 

The existence of a contact potential between two different metals was recognized over a century ago by Volta, who ascribed the origin of the electromotive force of galvanic cells to it. This point of view receded somewhat into the background in the later decades of last century, but is now re-established, as will be seen in 5. It is not very easy to demonstrate the existence of this contact potential and its actual value depends very much on the cleanliness of the surface indeed without very careful cleaning of the surface, and removal of surface films, which requires a high standard of vacuum technique, the true value for the clean metal can scarcely be obtained at all. 

The primary medium effect of an electrolyte can also be calculated from the standard potentials of a galvanic cell. The difference of the standard electromotive forces E and E " of the galvanic cells... 

In Eqs. (122) and (123), M(Hg) is an alkali metal amalgam electrode, MX the solvated halide of the alkali metal M at concentration c in a solvent S, and AgX(s)/Ag(s) a silver halide-silver electrode. Equation (124) is the general expression for the electromotive force " of a galvanic cell without liquid junction in which an arbitrary cell reaction 0)1 Yi + 0)2Y2 + coiYi + , takes place between k components in v phases. In Eq. (124) n is the number of moles of electrons transported during this process from the anode to the cathode through the outer circuit, F the Faraday number, and the chemical potential of component Yi in phase p. Cells with liquid junctions require the electromotive force E in Eq. (124) to be replaced by the quantity E — Ej), where Ey> is the diffusion potential due to the liquid junction. The standard potential E° for the cell investigated by Eq. (122) is given by the relationship... 

A measurement of the electromotive force of a suitable cell can therefore be used for the purpose of determining activity coefficients and also the standard free energy change in the cell reactions. 

Figure 19.5 shows the arrangement used to measure the standard electrode potential of a zinc half-cell. The electromotive force is -0.76 V. 

Table 6.11 lists, to the right of the arrows, reducing agents or disposition to electron loss or disposition to oxidation in order of increasing strength. Such a list is more popularly called the electromotive force, or emf, series. The maximum potential difference which can be measured for a given cell is called the electromotive force (abbreviated emf) and represented by the symbol Ecell. It may be recounted that the emf values reported in Table 6.11 are for those cells under specified standard conditions in which all the concentrations are 1 M and pressures are 1 atm. The emf of such a cell is said to be its standard electromotive force, and is given by the symbol E ell. 

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