Electrochemistry electromotive force

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... 

As we have seen, acidity and basicity are intimately connected with electron transfer. When the electron transfer involves an integral number of electrons it is customary to refer to the process as a redox reaction. This is not the place for a thorough discussion of the thermodynamics of electrochemistry that may be found in any good textbook of physical chemistry. Rather, we shall investigate the applications of electromotive force (emf) of interest to the inorganic chemist. Nevertheless, a very brief review of the conventions and thermodynamics of electrode potentials and half-reactions will be presented. 

An electrochemical cell generates a potential difference E. (The symbol E, commonly used in electrochemistry, refers to electromotive force, an archaic term for potential difference.) The electrical work done when n moles of electrons is passed by the cell can be found using Eq. (15-1), w = -nFE. It can be shown that the electrical work done by an electrochemical cell, at constant temperature and pressure, is equal to the change in Gibbs free energy of the cell components,... 

This equation, called the Thomson formula, is given mainly in old textbooks on electrochemistry as an approximate equation which expresses the relation between the EMF of a cell and the heat of reaction. As will be seen from the explanation of electrolysis below, the same formula may be also used for a rough estimate of the decomposition voltage of a compound, if no other data, except the reaction heat, are available. It is worth mentioning that the Thompson formula is precisely valid only with such systems the electromotive force of which does not change with temperature. 

Many writers on electrochemistry, perhaps most, refer to the electrical potential of a phase, particularly of an electrode, as if it was common knowledge what this means. This seems scarcely satisfactory, but as most electrochemical phenomena are concerned with the sum of two or more phase boundary potentials, the indefiniteness does not usually matter, in the final discussion of electromotive forces. 

In a short chapter in a book of this nature it is quite futile to attempt any comprehensive discussion of electromotive force. Only the fundamental notions can be indicated The student is therefore referred to the textbooks dealing specifically with Electrochemistry, especially Le Blanc s Electrochemistry (English or German edition), Lehfeldt s Electrochemistry (in this series of textbooks), and the recent work by Allmand, Applied Electrochemistry This latter book is of special significance in showing the scientific application of pure electrochemistry to industrial problems As regards laboratory experimental methods, i e the measurement of emf, details may be found in any of the textbooks on practical physical chemistry VOT.. II. 113 10... 

Much use of transference numbers has been made in the development of electrochemistry. The chief methods for their determination are (a) the Hittorf method, (6) die moving boundary method and (c) the electromotive force method. Of these the first two will be considered in this chapter. 

A very important branch of electrochemistry is concerned with the electromotive force (emf) or voltage developed in electrochemical cells. Investigations on electrochemical cells provide valuable information of various kinds. For example, they lead to thermodynamic quantities such as enthalpies and Gibbs energies for a variety of chemical reactions, including many reactions important in biology. Also, they allow us to obtain activity coefficients for ions in solution. This chapter deals with the general principles of electrochemical ceils and with some of their more important applications.. ... 

Volta s experiments and conclusions up to this time have been described in detail, since some of them (e.g. the electromotive forces due to the contact of two liquids) anticipated some published later by Davy and Faraday, and his failure to distinguish clearly between a potential difference established by contact and a current maintained by chemical action also reappeared in Faraday s work. Volta s earlier experiments were now eclipsed by his construction of the pile , the first primary battery, which was the real foundation of the science of electrochemistry and of all later developments of the useful applications of electricity. 

Nernst s first outstanding work was his theory of the production of the electromotive force of voltaic cells (1888-9), which is considered later (see p. 705), and his later work on electrochemistry was important. He devised a Wheatstone bridge method of measuring dielectric constants which was largely used, e.g. by Philip. Nernst first showed in detail that solvents of high dielectric constant promote the ionisation of substances, emphasising, however, that solvation of the ions may also have an effect. The first statement was independently briefly suggested somewhat later by J. J. Thomson. A. Sach-anov emphasised that electrolytic dissociation is conditioned not only by the dielectric constant of the solvent but also by the electroaffinities of the ions of the solute and by the formation of solvates and complex ions. ... 

An investigation of why hydroxide makes the Tollens silver mirror test for aldehydes more sensitive has focused on thermodynamic versus kinetic factors. Electrochemistry tends to rule out the former the electromotive force (emf) of an appropriate cell changes little with pH. Exploring the kinetics, single electron transfer processes were confirmed by addition of a radical trap (TEMPO), which slowed the reaction. Rate measurements point to the rate of the formation of the anion of the gm-diol (i.e. the hydrate anion) as the key parameter affected by added hydroxide, a factor that also explains how the rapidity of the test varies with the structure of the aldehyde. 

The electromotive series is a list of the elements in accordance with their electrode potentials. The measurement of what is commonly known as the "single electrode potential", the "half-reaction potential" or the "half-cell electromotive force" by means of a potentiometer requires a second electrode, a reference electrode, to complete the circuit. If the potential of the reference electrode is taken as zero, the measured E.M.P. will be equal to the potential of the unknown electrode on this scale. W. Ostwald prepared the first table of electrode potentials in 1887 with the dropping mercury electrode as a reference electrode. W. Nernst selected in 1889 the Normal Hydrogen Electrode as a reference electrode. G.N. Lewis and M. Randall published in 1923 their table of single electrode potentials with the Standard Hydrogen Electrode (SHE) as the reference electrode. The Commission of Electrochemistry of the I.U.P.A.C. meeting at Stockholm in 1953 defined the "electrode potential" of a half-cell with the SHE as the reference electrode. 

The Status of the Hydrogen Electrode. Probably no area of electrochemistry is more greatly neglected in current texts than the history of the choice of the hydrogen electrode as the reference standard for electromotive force measurements. Since all tables of potentials of oxidation-reduction half-reactions are based on the half-cell reaction 35H2=H +e , it would seem that the selection of this reaction as the standard should warrant more attention. If the selection is treated at all, it is usually dismissed as an arbitrary choice, which it is, with no reference made to the people and events involved in establishing this fundamental reference point for the EMF scale. One possible exception may be noted ( ). The referenced edition of this work is perhaps the best previously existing source on this topic. However, the subsequent edition omits the subject entirely. 

The concepts of EP and related open circuit potential and voltage (OCP and OCV), and electromotive force (emf) are well described in many textbooks on physical chemistry or electrochemistry. However, proving that the measured OCV (OCP) is indeed EP requires an approach that is specific to the particular electrochemical process being analyzed. For the EASP electrodes, a good confirmation of the OCP = EP relationship is the stability of the OCV versus lithium-metal (Li I Li+) reference electrode over time and independence upon lithium ion s concentration. According to the lUPAC [25] rule, the electrochemical circuit for such measurements has the form... 

Many experiments with voltaic cells ( piles ) were made by M. Berthelot. He tried to compare affinity and electromotive force by finding the e.m.f. of a cell which caused electrolysis with evolution of bubbles of gas, and attempted to connect this with the heat of reaction. He later recognised the significance of entropy changes, and experimented with liquid cells and liquid contact potentials, oxidation and reduction, acid-base neutralisation, etc., also the effect of superposition of an alternating current. Berthelot was not really at home in electrochemistry and his work is hardly ever mentioned. Experiments with several types of cells made by Hittorf gave appreciable differences between the chemical and electrical energies. 

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