Practical reversibility, thermodynamics

The dissolution reaction is Pt - Pt2+ + 2e and the value of its reversible thermodynamic potential is 1.2 V on the normal hydrogen scale. The evolution of O2 in acid solution at a current density of, say, 100 mA cm, needs an overpotential on platinum of nearly 1.0 V, i.e., the electrode potential would be >2.0 V. It follows feat at these very anodic potentials platinum would tend to dissolve, although its dissolution would be slowed down by fee fact feat it forms an oxide film at fee potentials concerned. Nevertheless, fee facts stated show feat fee alleged stability of Pt may be more limited than is often thought. This is an important practical conclusion because dissolved Pt from an anode may deposit on fee cathode of fee cell, and instead of having fee surface one started wife as fee cathode, it becomes in fact what is on its surface, platinum. 

The skills developed to produce the equilibrium diagram Figure A.l, are now applied anew. Neither hydrogen nor carbon monoxide occur as free substances in nature, where they are immediately oxidized. They must be made and stored, at thermodynamic and economic cost. The reversible thermodynamics are assessed below, using as the basis of calculation a notional, electrochemical, equilibrium, steam reformer. Figure A.4, for comparison with the alternative practical and irreversible combustion-driven reformers. 

In practice there are several limitations to such measurement. Obviously it implies that both members of the half-reaction are sufficiently stable for a cell to be realized. This is a serious difficulty in organic chemistry owing to usual great reactivities of the species formed upon electron transfers. For the most frequent cases it is then impossible to rely on reversible thermodynamic transformations to determine experimental values of standard reduction potentials. However, these important figures, or at least very precisely approximated values, can be obtained from current intensity potential curves or transient electrochemical methods as is discussed in a later section. 

Since all actual processes occur at finite rates, they cannot proceed with strict thermodynamic reversibility. However, a process may in practice be carried out in such a manner that thermodynamic equations apply to a desired accuracy. Under these circumstances, one might term the process reversible. Practical reversibility is not an absolute term it includes certain attitudes and expectations an observer has toward the process. 

Adhesion is an extremely important concept in both practical and theoretical terms. Unfortunately, there is no completely satisfactory definition of the term that fulfills the needs of both the theoretical surface chemist and the practicing technologist. So far in this book, the term adhesion has been encountered as applied in the ideal or theoretical sense—referring to the reversible thermodynamic process of separating unit area of two phases that originally had a common interface. That aspect of the term was defined in Chapter 2 and will not be repeated here, except where necessary for clarity. Some comments about the reality of that concept will be in order, however. 

The above parameters are derived assuming that the adsorption-desorption process is reversible and that there are no emulsifier-emulsifier interactions in the bulk solution or at the interface. In practice, a thermodynamic interpretation of these parameters may therefore be invalid for many real systems because these assumptions are not met. Nevertheless, these parameters still provide a useful means of characterizing and comparing the interfacial properties of emulsifiers in terms of experimentally measurable quantities. 

The usual situation, true for the first three cases, is that in which the reactant and product solids are mutually insoluble. Langmuir [146] pointed out that such reactions undoubtedly occur at the linear interface between the two solid phases. The rate of reaction will thus be small when either solid phase is practically absent. Moreover, since both forward and reverse rates will depend on the amount of this common solid-solid interface, its extent cancels out at equilibrium, in harmony with the thermodynamic conclusion that for the reactions such as Eqs. VII-24 to VII-27 the equilibrium constant is given simply by the gas pressure and does not involve the amounts of the two solid phases. 

Further reductions in reservoir pressure move the shock front downstream until it reaches the outlet of the no22le E. If the reservoir pressure is reduced further, the shock front is displaced to the end of the tube, and is replaced by an obflque shock, F, no pressure change, G, or an expansion fan, H, at the tube exit. Flow is now thermodynamically reversible all the way to the tube exit and is supersonic in the tube. In practice, frictional losses limit the length of the tube in which supersonic flow can be obtained to no more than 100 pipe diameters. 

In the thermodynamic treatment of electrode potentials, the assumption was made that the reactions were reversible, which implies that the reactions occur infinitely slowly. This is never the case in practice. When a battery deUvers current, the electrode reactions depart from reversible behavior and the battery voltage decreases from its open circuit or equiUbrium voltage E. Thus the voltage during battery use or discharge E is lower than the voltage measured under open circuit or reversible conditions E by a quantity called the polari2ation Tj. 

Retention Rejection and Reflection Retention and rejection are used almost interchangeably. A third term, reflection, includes a measure of solute-solvent coupling, and is the term used in irreversible thermodynamic descriptions of membrane separations. It is important in only a few practical cases. Rejection is the term of trade in reverse osmosis (RO) and NF, and retention is usually used in UF and MF. 

These considerations show the essentially thermodynamic nature of and it follows that only those metals that form reversible -i-ze = A/systems, and that are immersed in solutions containing their cations, take up potentials that conform to the thermodynamic Nernst equation. It is evident, therefore, that the e.m.f. series of metals has little relevance in relation to the actual potential of a metal in a practical environment, and although metals such as silver, mercury, copper, tin, cadmium, zinc, etc. when immersed in solutions of their cations do form reversible systems, they are unlikely to be in contact with environments containing unit activities of their cations. Furthermore, although silver when immersed in a solution of Ag ions will take up the reversible potential of the Ag /Ag equilibrium, similar considerations do not apply to the NaVNa equilibrium since in this case the sodium will react with the water with the evolution of hydrogen gas, i.e. two exchange processes will occur, resulting in an extreme case of a corrosion reaction. 

A hypothetical separation of a homogeneous mixture, carried out in a thermodynamically reversible manner, would require the theoretical minimum expenditure of energy. In practice, however, separations of such mixtures need 50 to 100 times this minimum. Thus, there is significant opportunity for improvement of separations by creating ways to reduce energy consumption without a commensurate increase in capital and operating costs. 

Considerable practical importance attaches to the fact that the data in Table 6.11 refer to electrode potentials which are thermodynamically reversible. There are electrode processes which are highly irreversible so that the order of ionic displacement indicated by the electromotive series becomes distorted. One condition under which this situation arises is when the dissolving metal passes into the solution as a complex anion, which dissociates to a very small extent and maintains a very low concentration of metallic cations in the solution. This mechanism explains why copper metal dissolves in potassium cyanide solution with the evolution of hydrogen. The copper in the solution is present almost entirely as cuprocyanide anions [Cu(CN)4]3, the dissociation of which by the process... 

The electrode processes that are reversible provide values for the equilibrium emfs of cells, which are related to the thermodynamic functions. The condition of reversibility is practically obtained by balancing cell emf against an external emf until only an unappreciable current passes through the cell, in order that the cell reactions proceed very slowly. It may, however, be pointed out that for many of the applications of electrometallurgy, it is clearly necessary to consider more rapid reaction rates. In that situation there is necessarily a departure from the equilibrium condition. Either the cell reactions occur spontaneously to produce electric energy, or an external source of electric energy is used to implement chemical reactions (electrolyses). 

Although from the thermodynamic point of view one can speak only about the reversibility of a process (cf. Section 3.1.4), in electrochemistry the term reversible electrode has come to stay. By this term we understand an electrode at which the equilibrium of a given reversible process is established with a rate satisfying the requirements of a given application. If equilibrium is established slowly between the metal and the solution, or is not established at all in the given time period, the electrode will in practice not attain a defined potential and cannot be used to measure individual thermodynamic quantities such as the reaction affinity, ion activity in solution, etc. A special case that is encountered most often is that of electrodes exhibiting a mixed potential, where the measured potential depends on the kinetics of several electrode reactions (see Section 5.8.4). 

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