Electrolytic-solution pressure

Schofield Phil. Mag. March, 1926) has recently verified this relation by direct experiment. In order to appreciate the significance of this result, it is necessary to consider in more detail the electrical potential difference V and the manner in which it arises. Instead of regarding the phenomenon from the point of view of the Gibbs equation, it has been, until recently, more usual to discuss the subject of electro-capillarity from the conceptions developed by Helmholtz and Lippmann. These views, together with the theory of electrolytic solution pressure advanced by Nemst, are not in reality incompatible with the principles of adsorption at interfaces as laid down by Gibbs. 

The two electrode potentials Fi and F3 are according to the Nemst conception of electrolytic solution pressure given by the expression... 

Vanadium predpitates the metal from solutions of salts of gold, silver, platinum, and iridium, and reduces solutions of mercuric chloride, cupric chloride and ferric chloride to mercurous chloride, cuprous chloride, and ferrous chloride, respectively. In these reactions the vanadium passes into solution as the tetravalent ion. No precipitation or reduction ensues, however, when vanadium is added to solutions of divalent salts of zinc, cadmium, nickel, and lead. From these reactions it has been estimated that the electrolytic potential of the change, vanadium (metal)—>-tetravalent ions, is about —0 3 to —0 4 volt, which is approximately equal to the electrolytic solution pressure of copper. This figure is a little uncertain through the difficulty of securing pure vanadium.5... 

An apparently unresolved issue from the above analysis is the breakdown of the analogy, say, between carbon and metal electrodes for the latter (e.g., ZnlZn2+), it is easy to define an electrolytic solution pressure and an osmotic pressure, as the pioneers of electrochemistry (Nernst, Ostwald, Arrhenius, and Le Blanc) have done (see books.google.com) more than a century ago but I am aware of no such discussion (e.g., CIC+) on carbons in which the exact nature of C+ would have been addressed in those terms. 

With base metals (e. g. alkali metals, iron, zinc etc.) in which the valence electrons in the atoms are more loosely bound the electrolytic solution pressure... 

On the other (copper) electrode the electrolytic solution pressure is lower than the osmotic pressure of the cations in the solution and therefore cupric ions from tho solution are deposited, thus giving the metal a positive charge, while the solution becomes negative due to the excess of anions (SO ). Both kinds of charges Cu++ and SO - are attracted and form again the electrical double layer. In this case, however, the double layer has an opposite effect than at the zinc electrode as it facilitates the transfer of the cupric ions from the electrode to the solution and prevents them being transferred in the opposite direction. Equilibrium will be attained, when the electrostatic forces of the double layer and the solution pressure of copper together will counterbalance the osmotic pressure of the cupric ions in the solution. 

Such electrodes differ from metalic ones only in the fact that the solution due to the negative charge of the ions formed gets negatively charged against the electrode, when the electrolytic solution pressure of the element is greater than the osmotic pressuro of its anions in the solution. 

C is a constant which is characteristic of the metal of which the electrode is composed, and is sometimes called the electrolytic solution pressure. Its numerical value is equal to the ionic concentration of a solution against which the metal would have no difference of potential. This quantity is of the greatest importance for the electrochemical behaviour of the metal. It cannot be determined, however, by measurement with concentration cells, for the solution pressure C disappears from the sum of the various potential differences in equation (5). The calculation of C from the total e.m.f. would be possible if we could choose the ionic concentration of one solution, say Cg, so that the potential difference 3 would be zero. Numerous experiments have actually been carried out with the object of constructing an electrode which would have the absolute potential zero against the solution. These experiments, although in themselves interesting and important, are based on special electrochemical hypotheses and not on purely thermodynamical principles. They are therefore beyond the scope of this book. ... 

In the same publications (1888-9), Nernst also gave, for the first time, an equation for the potential between a metal and a solution of its ions. The mechanism was explained in terms of an electrolytic solution pressure P of a metal, giving its tendency to form ions in a solution of its ions of osmotic pressure ... 

Jones, J.R.E. 1939. The relation between the electrolytic solution pressures of the metals and their toxicity to the stickleback (Gasterosteus aculeatus L.). J. Exp. Biol. 16 425 37. 

In the case of a metal dipping into a solution containing its ions, the tendency of the metal ions to dissolve is thus determined by their solution pressure which Nernst called the electrolytic solution pressure, P, of the metal. The tendency of the metal ions to deposit is measured by their osmotic pressure, p. 

P > p The electrolytic solution pressure of the metal is greater than the osmotic pressure of the ions, so that positive metal ions will pass into the solution. As a result the metal is left with a negative charge, while the solution becomes positively charged. There is thus set up across the interface an electric field which attracts positive ions towards the metal and tends to prevent any more passing into solution (Figure 1.2(a)). The ions will continue to dissolve and therefore the electric field to increase in intensity until equilibrium is reached, i.e. until the inequality of P and p, which causes the solution to occur, is balanced by the electric field. 

P = p The osmotic pressure of the ions is equal to the electrolytic solution pressure of the metal. 

Nemst calculated the potential difference required to bring about equilibrium between the metal and the solution in the following way. He determined the net work obtainable by the solution of metal ions by means of a three-stage expansion process in which the metal ions were withdrawn from the metal at the electrolyte solution pressure P, expanded isothermally to the osmotic pressure p, and condensed at this pressure into the solution. The net work obtained in this process is... 

Here, x denotes film thickness and x is that corresponding to F . An equation similar to Eq. X-42 is given by Zorin et al. [188]. Also, film pressure may be estimated from potential changes [189]. Equation X-43 has been used to calculate contact angles in dilute electrolyte solutions on quartz results are in accord with DLVO theory (see Section VI-4B) [190]. Finally, the x term may be especially important in the case of liquid-liquid-solid systems [191]. 

Simonson J M and Mesmer R E 1994 Electrolyte solutions at high temperatures and pressures Solution Calorimetry, Experimental Thermodynamics yo IV, ed K N Marsh and PAG O Hare (Oxford Blackwell)... 

The freezing points of electrolyte solutions, like their vapor pressures, are lower than those of nonelectrolytes at the same concentration. Sodium chloride and calcium chloride are used to lower the melting point of ice on highways their aqueous solutions can have freezing points as low as —21 and — 55°C, respectively. 

Chapters 7 to 9 apply the thermodynamic relationships to mixtures, to phase equilibria, and to chemical equilibrium. In Chapter 7, both nonelectrolyte and electrolyte solutions are described, including the properties of ideal mixtures. The Debye-Hiickel theory is developed and applied to the electrolyte solutions. Thermal properties and osmotic pressure are also described. In Chapter 8, the principles of phase equilibria of pure substances and of mixtures are presented. The phase rule, Clapeyron equation, and phase diagrams are used extensively in the description of representative systems. Chapter 9 uses thermodynamics to describe chemical equilibrium. The equilibrium constant and its relationship to pressure, temperature, and activity is developed, as are the basic equations that apply to electrochemical cells. Examples are given that demonstrate the use of thermodynamics in predicting equilibrium conditions and cell voltages. 

The interface between a liquid metal and an electrolyte solution can be vibrated by applying an oscillating external pressure variation. An electric signal picked up at the oscillating interface as a function of the applied electrode potential can be registered. It shows a particular dependency, at the signal vanishes [83Mly]. (Data obtained with this method are labelled VE). 

Home, R. A., Day, A. F., and Young, R. P. (1969). Ionic diffusion under high pressure in porous solid materials permeated with aqueous, electrolytic solution. /. Phys. Chem. 73,2782-2783. 

The situations would be totally different when the two surfaces are put in electrolyte solutions. This is because of formation of the electrical double layers due to the existence of ions in the gap between solid surfaces. The electrical double layers interact with each other, which gives rise to a repulsive pressure between the two planar surfaces as... 

The simplest device for measuring ECC at mercury is Gouy s capillary electrometer (Eig. 10.5). Under the effect of a mercury column of height h, mercury is forced into the slightly conical capillary K. In the capillary, the mercury meniscus is in contact with electrolyte solution E. The radius of the mercury meniscus is practically equal to the capillary radius at that point. The meniscus exerts a capillary pressure Pk = directed upward which is balanced by the pressure = ftpegg of... 

Four different electrokinetic processes are known. Two of them, electroosmosis and electrophoresis, were described in 1809 by Ferdinand Friedrich Renss, a professor at the University of Moscow. The schematic of a cell appropriate for realizing and studying electroosmosis is shown in Fig. 31.1a. An electrolyte solution in a U-shaped cell is divided in two parts by a porous diaphragm. Auxiliary electrodes are placed in each of the half-cells to set up an electric held in the solution. Under the inhuence of this held, the solution starts to how through the diaphragm in the direction of one of the electrodes. The how continues until a hydrostahc pressure differential (height of liquid column) has been built up between the two cell parts which is such as to compensate the electroosmotic force. 

In 1861, Georg Hermann Quincke described a phenomenon that is the converse of electroosmosis When an electrolyte solution is forced through a porous diaphragm by means of an external hydrostatic pressure P (Fig. 31.1ft), a potential difference called the streaming potential arises between indicator electrodes placed on different sides of the diaphragm. Exactly in the same sense, in 1880, Friedrich Ernst Dorn described a phenomenon that is the converse of electrophoresis During... 

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