Experimental values electrolytes

An electrolysis experiment is performed to determine the value of the Faraday constant (number of coulombs per mole of electrons). In this experiment, 28.8 g of gold is plated out from a AuCN solution by running an electrolytic cell for two hours with a current of 2.00 A. What is the experimental value obtained for the Faraday constant ... 

In this manner, the surface excess of ions can be found from the experimental values of the interfacial tension determined for a number of electrolyte concentrations. These measurements require high precision and are often experimentally difficult. Thus, it is preferable to determine the surface excess from the dependence of the differential capacity on the concentration. By differentiating Eq. (4.2.30) with respect to EA and using Eqs (4.2.24) and (4.2.25) in turn we obtain the Gibbs-Lippmann equation... 

At low electrolyte concentrations, up to about a 10 3 M solution, the Gouy-Chapman theory agrees quite well with experimental values of... 

Figure 2. The rate constants k2 for some electrolytes following the parabolic rate law, vg = k2

It can therefore be seen that the agreement between the calculated and experimental values of Rei is actually quite good, with the probable eause of the calculated value being too large resulting from uneven electrolyte thickness. 

ACTIVITY COEFFICIENTS OF SOME STRONG ELECTROLYTES Experimental Values... 

The differences in the solvation abilities of ions by various solvents are seen, in principle, when the corresponding values of As ivG° of the ions are compared. However, such differences are brought out better by a consideration of the standard molar Gibbs energies of transfer, AtG° of the ions from a reference solvent into the solvents in question (see further section 2.6.1). In view of the extensive information shown in Table 2.4, it is natural that water is selected as the reference solvent. The TATB reference electrolyte is again employed to split experimental values of AtG° of electrolytes into the values for individual ions. Tables of such values have been published [5-7], but are outside the scope of this text. The notion of the standard molar Gibbs energy of transfer is not limited to electrolytes or ions and can be applied to other kinds of solutes as well. This is further discussed in connection with solubilities in section 2.7. 

The parameters can be evaluated from data on mixtures of electrolytes by calculating the differences between the experimental value of and the value calculated with the appropriate values for all pure electrolyte terms and terms but zero values for and j/ terms. For the activity coefficient of MX in a MX-NX mixture one has Eq. (6.42) and equivalent expressions for other cases. 

Because the conventional standard partial molal properties of a hydrogen ion is zero, it follows that the conventional standard partial molal properties of a generic anion A are identical to the experimental values of the corresponding acid electrolyte. Moreover, based on equation 8.104, the standard partial molal properties of a generic cation C can be calculated, once the experimental values for aqueous electrolytes H +A and are... 

A new theory of electrolyte solutions is described. This theory is based on a Debye-Hiickel model and modified to allow for the mutual polarization of ions. From a general solution of the linearized Poisson-Boltzmann equation, an expression is derived for the activity coefficient of a central polarized ion in an ionic atmosphere of non-spherical symmetry that reduces to the Debye-Hiickel limiting laws at infinite dilution. A method for the simultaneous charging of an ion and its ionic cloud is developed to allow for ionic polarization. Comparison of the calculated activity coefficients with experimental values shows that the characteristic shapes of the log y vs. concentration curves are well represented by the theory up to moderately high concentrations. Some consequences in relation to the structure of electrolyte solutions are discussed. 

Figure 13. Calculated charge as a function of potential for the water-copper surface in contact with an electrolyte, compared with experimental values. From Ref. 36, by permission.

Coagulation rates have been measured as a function of electrolyte concentration for a number of sols96 196 204-206, and the predicted linear relationship between log W and log c in the slow-coagulation region seems to be well confirmed. In addition, the experimental values of d log W d log c, although somewhat variable, are of the right order of magnitude compared with theoretical slopes. 

Thus A represents the specific conductance of a hypothetical solution of the electrolyte containing 1 equivalent cm-3. An experimental value of A may be determined by measurement of R, IIA (the cell constant), and N, and use of the relationship... 

The calculated results in the absence of electrolyte will be now compared with the experimental results obtained regarding a lamellar lyotropic liquid crystal SDS (sodium dodecyl sulfate)/pentanol/water/dodecane swollen in a mixture of dodecane and pentanol.24 The weight fraction water/surfactant was 1.552 from the dilution line in the phase diagram, we calculated that the initial concentration of pentanol in the oil-free system was 29 wt % and the concentration of pentanol in the dodecane-based diluant was 8 wt %. The experimental values for the repeat distance were obtained from the X-ray diffraction spectrum (Figure 2 in ref 24) for various dodecane concentrations. 

Fig. 1. Experimental values for the osmotic pressure as a function of separation distance from Ref. [13] (stars water circles 1 M KCl triangles 1 M KBr) are compared with calculations based on the simple equations ((1), (3) and (9)), with parameters reported in Ref. [13] (note that in Ref. [13] it was suggested that hydration interaction increases upon addition of salt) Ah = 1.6 x 10s N/m2, H = 2.1 A, H = 9.2 x KT21 J. b = 39 A, Kc = 5.8 x 10 2(1 J, An = 1.06 A 2, >.fl = 6.0 A (Line 1). Even for H = 0 and the rest of the parameters as before (Line 2), the repulsion at short separations is weaker than in the experiment. This points out that either hydration and/or undulation forces must increase upon addition of electrolyte.

Figure 2. Experimental values of the stability ratio of protein-covered latex particle as a function of electrolyte concentration, at pH = 10.0, reported by Lopez-Leon et al.,1 compared to those calculated from the polarization-based hydration model, for the following parameter values NA = 1.2 x 1018 sites/m2, NB = 1.62 x 1018 sites/m2, A, = 0.9 x lO 20 J, KH = lCL6 M, Aon = 8.95 x 10 8 M, KNh = 0.021 M, (p/e )Na = 1.8 D (1) Ka = 0.76 M, (p/e)ci = 2.3D (2)Kno = 0.62M,(p/e )no3 = -1.8D stars, NaN03 squares, NaCL...

Figure 3. Stability ratios at low concentrations of NaSCN at pH = 10. The calculations have been carried out for Na = Nb = 0.5 x 1018 sites/m2 and various dissociation constants Ans based on the present model (thick lines) and on the DLVO theory, with different values for Ans (dotted lines), predict almost identical results but varies much more rapidly with electrolyte concentration than the experimental values reported in ref 1 (circles). K = A(>i = 1C1 1 M, (/ /r ), — 1.8 D, n = 0.8 M, (p/e )scn = -0.8 D, Ah = 0.1 x 10 201 J. (1) polarization model, ANa = 12.5 x 10-6M DLVO,ANa = 2.5 x 10-eM. (2)polarization model, AKa = 23.1 x 10 6 M DLVO, ANa = 5.83 x 10-6 M. (3) polarization model, Axa = 54.7 x 10 M DLVO, An, = 12.6 x 10-6 M. (4) polarization model, ANa = 164 x 10-6 M DLVO, Axa = 42 x 10 6 M.

Parsons and Zobel plot — In several theories for the electric - double layer in the absence of specific adsorption, the interfacial -> capacity C per unit area can formally be decomposed into two capacities in series, one of which is the Gouy-Chapman (- Gouy, - Chapman) capacity CGC 1/C = 1 /CH + 1 /CGC. The capacity Ch is assumed to be independent of the electrolyte concentrations, and has been called the inner-layer, the - Helmholtz, or Stern layer capacity by various authors. In the early work by Stern, Ch was attributed to an inner solvent layer on the electrode surface, into which the ions cannot penetrate more recent theories account for an extended boundary region. In a Parsons and Zobel plot, Ch is determined by plotting experimental values for 1/C vs. 1/Cgc- Specific adsorption results in significant deviations from a straight line, which invalidates this procedure. 

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