Solubility electrolytes, mixtures

For most organic chemicals the solubility is reported at a defined temperature in distilled water. For substances which dissociate (e.g., phenols, carboxylic acids and amines) it is essential to report the pH of the determination because the extent of dissociation affects the solubility. It is common to maintain the desired pH by buffering with an appropriate electrolyte mixture. This raises the complication that the presence of electrolytes modifies the water structure and changes the solubility. The effect is usually salting-out. For example, many hydrocarbons have solubilities in seawater about 75% of their solubilities in distilled water. Care must thus be taken to interpret and use reported data properly when electrolytes are present. 

Gamsjager, H. Schindler, P. "Solubilities and Activity Coefficients of H2S in Electrolyte Mixtures," Helv. Chim. Acta,1969, 52, 1395-1402. 

The Kirkwood-Buff formalism can be also used to derive the composition dependence of the Henry constant for a sparingly soluble gas dissolved in a mixed solvent containing water-r electrolyte [27]. The obtained equation requires information about the molar volume and the mean activity coefficient of the electrolyte in the binary (water-H electrolyte) mixture. Several expressions for the mean activity coefficient of the electrolyte were tested and it was concluded that the accuracy in... 

The solubility product principle enables simple calculations to be made of the effect of other species on the solubility of a given substance and may be used to determine the species that will precipitate in an electrolyte mixture. One simple result of applying the solubility product principle is the common ion effect. This is the effect caused by the addition of an ionic species that has an ion in common with the species of interest. Since the solubility of a species is given by the product of the concentration of its ions, when the concentration of one type of ion increases, the concentration of the other must decline, or the overall concentration of that compound must decline. We can illustrate this simply by using our previous example of silver chloride. The solubility product of silver chloride at 25°C is 1.56 x lO". This means that at saturation we can dissolve 1.25 x 10 mol of AgCl/lOOOg of water. If, however, we were to start with a solution that has a coneentration of 1 X 10 molal NaCl (hence 1 x 10 molal CP) the solubility product equation can be written in the form... 

Electrolyte Mixtures. The calculation of the solubility of mixtures of strong electrolytes requires knowledge of the thermodynamic solubility product for all species that can precipitate and requires using an activity coefficient calculation method that takes into account ionic interactions. These techniques are well described in Zemaitis et al. (1986), however, we will discuss a simple case in this section. 

A perusal of recent literature shows an increasing interest in technical applications and applied research based on non-aqueous electrolyte properties. The assortment of solvents with widely varying properties, an almost unlimited number of solvent mixtures and soluble electrolyte compounds provides flexibility in tackling a given problem. The unique properties of non-aqueous solutions can be the key in solving special technical problems. 

Electrolyte mixtures come in various forms and add another dimension to the normal complexities of nonelectrolyte solutions entireiy new species can form in water, some of which are not obvious components can precipitate soluble components can affect the vapor pressure of the solution very significantly. In industrial applications, the solutions are often highly concentrated and encounter high pressures and temperatures. Therefore expertise in electrolyte systems has become increasingly aitical in oil and gas exploration and production, as well as in the more traditional chemical industry operations. A variety of correiations are available that can solve the problems that are encountered in industry, but which ones work best This comprehensive handbook not oniy provides easy access to available data but also presents comparative studies of various correlations up to extreme conditions. 

Pabalan, R.T. Pitzer, K.S. (1987). Thermodynamics of concentrated electrolyte mixtures and the prediction of mineral solubilities to high temperatures for mixtures in the system Na-K-Mg-Cl-S04-0H-H20, Geochim, Cosmochim. Acta, Vol. 51, No. 9, pp. 2429-2443, ISSN 0016-7037... 

The presence of electrolytes in the liquid affects gas solubility. Henry s constant for a pure solvent (Ho) can be corrected with the so-called salting-out factors (h) that are ion-specific parameters. Schumpe [6] has suggested the following equation for electrolyte mixtures ... 

In electrolytes based on solvent mixtures both solvent compounds may react to form films of scarcely soluble materials. PC/THF mixtures yield alkoxides and alkylcarbonates [188] EC/ether blends mainly yield alkylcarbonates, which are thought to be the reason for smaller lithium loss during cycling [188]. PC based electrolytes with LiAsF6and LiC104 form films containing alkylcarbonates which allow the access of other molecules, such... 

Sometimes it is important to know under what conditions a precipitate will form. For example, if we are analyzing a mixture of ions, we may want to precipitate only one type of ion to separate it from the mixture. In Section 9.5, we saw how to predict the direction in which a reaction will take place by comparing the values of J, the reaction quotient, and K, the equilibrium constant. Exactly the same techniques can be used to decide whether a precipitate is likely to form when two electrolyte solutions are mixed. In this case, the equilibrium constant is the solubility product, Ksp, and the reaction quotient is denoted Qsp. Precipitation occurs when Qsp is greater than Ksp (Fig. 11.17). 

The auxiliary electrolyte is generally an alkali metal or an alkaline earth metal halide or a mixture of these. Such halides have high decomposition potentials, relatively low vapor pressures at the operating bath temperatures, good electrolytic conductivities, and high solubilities for metal salts, or in other words, for the functional component of the electrolyte that acts as the source of the metal in the electrolytic process. Between the alkali metal halides and the alkaline earth metal halides, the former are preferred because the latter are difficult to obtain in a pure anhydrous state. In situations where a metal oxide is used as the functional electrolyte, fluorides are preferable as auxiliary electrolytes because they have high solubilities for oxide compounds. The physical properties of some of the salts used as electrolytes are given in Table 6.17. 

The quality of the refined metal, and the current efficiency strongly depend on the soluble vanadium in the bath and the quality of the anode feed. As the amount of vanadium in the anode decreases, the current efficiency and the purity of the refined product also decrease. A laboratory preparation of the metal with a purity of better than 99.5%, containing low levels of nitrogen (30-50 ppm) and of oxygen (400-1000 ppm) has been possible. The purity obtainable with potassium chloride-lithium chloride-vanadium dichloride and with sodium chloride-calcium chloride-vanadium dichloride mixtures is better than that obtainable with other molten salt mixtures. The major impurities are iron and chromium. Aluminum also gets dissolved in the melt due to chemical and electrochemical reactions but its concentrations in the electrolyte and in the final product have been found to be quite low. The average current efficiency of the process is about 70%, with a metal recovery of 80 to 85%. 

When a 60 MW turbine at Hinkley A power station disintegrated in 1969 from stress corrosion cracking of a low pressure turbine disc (consequences shown in Plate 1) it was considered that Na H solutions were most probably involved (84) and it was soon found that if NaOH were the sole electrolyte present its maximum concentration (based on vapour pressure depression) was sufficient to have caused the cracking. However, it was also found that in mixtures it was only the free NaOH which led to rapid stress corrosion cracking. Considerations of acid gas solubility and solution thermodynamics showed that at the CO2 and acetate levels present it was most unlikely that free NaOH was present in sufficient quantity to be responsible for the Hinkley failure (85). 

---