Mean activity coefficient chloride salt

Only mean activity coefficients can be experimentally determined for salts, not activity coefficients for single ions. The Maclnnes Convention is one method for obtaining single ion activity coefficients and states that because of the similar size and mobility of the potassium and chloride ions ... 

Calculate the mean activity coefficient of thallous chloride, the solubility of which has been measured in water and in the presence of various concentrations of potassium chloride soiutions at 25 °C, as given in Tabie P.l. The soiubiiity of this salt in pure water is 1.607 x 10 mol kg". (Constantinescu)... 

Figure 4. The Henry constant of oxygen in aqueous solutions of sodium chloride at 25 °C (O) experimental data (a) the Henry constant calculated with eq 24 using for the mean activity coefficient of dissolved salt the Debye—Hiickel equation (b) the Henry constant calculated with eq 24 using for the mean activity coefficient of dissolved salt the extended Debye—Hiickel equation (c) the Henry constant calculated with eq 24 using for the mean activity coefficient of dissolved salt the Bromley equation (d) the Henry constant calculated with eq 15.

Although sodium chloride exists in entirely different phases (liquid and crystal) within this two-phase region, the magnitude of its chemical potential must still be the same in both phases. (This means, incidentally, that the activity coefficient of salt in the saturated solution is >1 if pure substances are taken to be the standard states.) The free energy of the saturated solution differs substantially from that of the salt crystals, however. The free energy of the liquid phase is... 

In both of the above examples, we used an anionic buffer (MOPS or cacodylate). The buffer anions have only repulsive interactions with RNA and can be grouped with chloride ions when calculating mean ion activities. Thus, we apply mean ionic activity coefficients measured with KC1 solutions to solutions in which K+ ions are contributed both by KC1 and K-buffer salts. We strongly advise against the use of cationic buffers such as Tris, because of its idiosyncratic interactions with nucleic acids as compared to group I ions, and particularly against mixing KC1 with Tris buffer, which creates a cationic mixture of unknown activity. 

The Mean Molal Activity Coefficient Calculation for the Macro-Component Salts. The mean molal activity coefficients for the macro-component chloride and sulfate salts are given in Figures 2a and 2b, respectively. Because of the unavailability of measured activity coefficient data in brines, validation of the results presented in Figures 2a and 2b was not possible. 

The solubility of silver chloride in pure water is 1.314 X 10 molal, and the mean ionic activity coefficient is then 0.9985 [Neuman, J. Am. Chem. Soc., 54, 2195 (1932)]. The heat of solution of the salt is 15,740 cal. mole". Taking the entropy of solid silver chloride as 22.97 e.u. mole ", and using the results of the preceding exercise, calculate the standard free energy and heat of formation and the entropy of the Cl" ion at 25 C. 

Values of q> and y for many aqueous electrolytes at 25°C are reported in the books by Hamed and Owen [1], by Robinson and Stokes [2], in subsequent reports from the US National Bureau of Standards (now NIST) [3-8], and have been reevaluated more recently by Partanen and coworkers [9,10]. Values of the mean ionic activity coefficients of representative aqueous electrolytes at 25°C at several molalities m are shown in Table 7.1, in order for the trends with the natures of the ions making up these electrolytes to be seen, and some of these data are shown in Figure 7.1. In all the cases there is a decrease in y at low concentrations but the values tend to increase again at higher molalities. For the univalent cation chlorides, for instance, at 1 m the trend of the y values is HCl > LiCl > NaCl > KCl NH Cl > RbCl > CsCl. For the acids and lithium, sodium, and potassium salts, the trend at this molality among the anions is NOj" < CF < Br < CIO < F (KCIO is insoluble), but for rubidium and cesium the trend is F < Br (not shown) < CF. These trends are commented on in Section 7.2. The values of 1 2 and 2 1 electrolytes are smaller than those for 1 1 electrolytes and the diminution becomes larger as the charge numbers of the ions increase. 

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