Activity Coefficients of LiCl in Ethanol−Water ... - ACS Publications

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Ind. Eng. Chem. Res. 2008, 47, 2056-2062

Activity Coefficients of LiCl in Ethanol-Water Mixtures at 298.15 K Felipe Herna´ ndez-Luis* Departamento de Quı´mica Fı´sica, UniVersidad de La Laguna, Tenerife, Espan˜ a

He´ ctor R. Galleguillos, Teo´ filo A. Graber, and Marı´a E. Taboada Departamento de Ingenierı´a Quı´mica, UniVersidad de Antofagasta, Antofagasta, Chile

The electromotive force, E, of the cell containing two ion-selective electrodes (ISE) Na-ISE | LiCl(m), ethanol (Y), H2O(100-Y) | Cl-ISE has been measured at a temperature of 298.15 K as a function of the weight percentage Y of ethanol in a mixed solvent. Y was varied between 0 and 80% in 20-unit steps, and the molality of the electrolyte (m) was between 0.02 and 4 mol kg-1, approximately. This electrode system was previously checked and calibrated by measuring the activity coefficients of LiCl in 0, 20, and 40% methanol-water and comparing them with literature values. The values of the standard electromotive force, E0 (molal scale), were determined using routine methods of extrapolation (Debye-Hu¨ckel and Pitzer equations). The results obtained produced good internal consistency within the normal limits of experimental error encountered in these types of measurements. Once E0 was determined, the mean ionic activity coefficients for LiCl, the free energy of transfer from the water to the alcohol-water mixture, and the primary LiCl hydration number were calculated. The variation of these magnitudes with the composition of the mixture is discussed in terms of the ionsolvent and ion-ion interactions and their changes with the properties of the medium. 1. Introduction It is necessary to know the activity coefficients (and related magnitudes) of electrolytes both in aqueous and aqueousorganic media for many practical purposes. Two methods have been typically used for these types of determination, including: isopiestic vapor pressure measurements and emf measurements of galvanic cells. Much has been written about the advantages and disadvantages of these two methods. Although there has not been complete agreement, it appears that the potentiometric method is simpler and more rapid than the isopiestic method, although the first has a definite limitation in that reversible electrodes must be available for the ions of the electrolyte wished to be studied. The potentiometric method1-11 is more suited than the isopiestic method in the case of mixed aqueousorganic solvents because of the nonquantifiable presence of solvents in the vapor phase. Nevertheless, in spite of the major advancement in the ISE,12-20 many ions (anions and cations) still exist for which specific electrodes have not been found or have not been simple enough to fabricate and use. Thus, as far as we know, there is no ISE for Li+ that is commercially available, as are glass electrodes for Na+ or liquid membrane electrodes for K+. Various authors have exploited the quasi-Nerstian response shown by some glass electrodes toward different monovalent ions. For example, Covington and Thain21 compared different types of selective electrodes and described the precautions and safeguards necessary to obtain reliable results. Among electrodes studied by these authors were those termed monovalent cationresponsive glass electrodes, which could respond to one or another ion, but only when there were no other ions present that were capable of causing interference. Jiang et. al22-24 obtained good results with selective electrodes for K+ and Clin obtaining the activity coefficients of RbCl or CsCl in different mixed media. * To whom correspondence should be addressed. Tel.: 34 922 318471. Fax: 34 922 318514. E-mail address: [email protected].

In the present study, we have used a bi-ISE system formed by a Na+-selective glass electrode and another with a solid Cl-selective solid membrane electrode to obtain the activity coefficients for LiCl in water-ethanol mixtures. This was preceded by checking and calibration of this electrode system by measuring the activity coefficients of LiCl in 0, 20, and 40% methanol-water mixtures and comparing them with results obtained from the literature.25 2. Experimental Anhydrous lithium chloride (LiCl) (Fluka, 99%) was dried in vacuo at 370 K for 2 days before use. Afterward, it was stored over silica gel in a desiccator and used without further purification. Both methanol and ethanol (Merck pro analisi, 99.8%) were used as supplied. As the alcohol was always used as a mixture with water, correction for the very small water content of the original product was considered unnecessary. All of the mixed solvents as well as the working solutions were prepared by weighing. For each set of experiments (corresponding to a wt % of alcohol), working solutions were obtained by adding successive solid weighed LiCl to a solution previously prepared of each alcohol and bidistilled water (κ < 1 10-6 S cm-1). The LiCl was contained in a hermetically sealed plastic capsule during weighing to avoid its absorbing humidity from the air. Na-ISE (mod. 6.0508.100) and Cl-ISE (mod. 6.0502.120) were from the Metrohm company. A double-walled Metrohm glass vessel cell was used to hold the electrodes and the solution. The temperature in the cell was maintained constant within 25.00 ( 0.05 °C by circulating water from a Hetofrig model 04 PT thermostat-cryostat. A Guildline model 9540 platinum resistance thermometer was used to monitor the temperature. The solutions were continuously stirred using a magnetic stirrer to avoid concentration gradients in the test solutions. The emf measurements were carried out using a 614 Keithley Electrometer, which had an inner impedance greater than 5 × 1013 Ω and whose resolution was (0.1 mV. To obtain more

10.1021/ie070704i CCC: $40.75 © 2008 American Chemical Society Published on Web 02/19/2008

Ind. Eng. Chem. Res., Vol. 47, No. 6, 2008 2057 Table 1. Experimental Electromotive Force at Different LiCl Molalities in Methanol-Water Mixture at 298.15 Ka 0% Methanol

20% Methanol

40% Methanol

m

E

γb

m

E

γc

m

E

γc

0.1348 0.3033 0.5466 0.6095 0.6567 0.7000 0.7307 0.8282 0.8694 0.9516 1.0203 1.1199 1.3154 1.9277 2.3107 3.3124 4.7511 5.3955

-22.55 -58.81 -87.35 -92.57 -96.17 -99.39 -101.44 -107.72 -110.14 -115.00 -118.05 -123.03 -132.20 -155.85 -168.62 -198.02 -232.92 -248.16

0.774 0.743 0.740 0.743 0.745 0.748 0.750 0.758 0.762 0.770 0.777 0.788 0.813 0.911 0.987 1.250 1.857 2.256

0.01982 0.04710 0.1159 0.2026 0.3471 0.3725 0.4289 0.5078 0.8272 1.0122 1.4450 2.0080 2.9320 3.8177 4.4729 5.0478

37.27 -0.46 -39.73 -64.64 -88.63 -91.66 -98.01 -106.01 -130.51 -141.11 -160.01 -181.22 -209.80 -234.03 -250.64 -264.89

0.857 0.809 0.757 0.730 0.712 0.710 0.709 0.708 0.723 0.740 0.792 0.887 1.110 1.422 1.736 2.090

0.1391 0.3013 0.4687 0.5763 0.7655 0.8617 0.9885 1.3051 1.4741 1.7674 2.0462 2.5941 3.1418 3.7282 3.9214

-70.21 -102.52 -122.83 -132.41 -146.47 -152.31 -159.40 -174.83 -181.84 -193.02 -202.61 -219.53 -234.99 -251.70 -256.82

0.684 0.639 0.624 0.623 0.627 0.632 0.641 0.672 0.692 0.733 0.779 0.890 1.031 1.222 1.295

k ) 54.44 mV σ ) 0.55 mV r ) 0.99995

k ) 54.15 σ ) 0.32 mV r ) 0.99999

k ) 53.88 mV σ ) 0.45 mV r ) 0.99997

〈k〉 ) 54.16 ( 0.16 mV a

Units, m in mol

kg-1;

b

c

E in mV. Ref 32. Ref 25.

precise emf readings, the 2 V analog output of the electrometer was connected to a Keithley model 197A Microvolt DMM with an input greater than 1 × 1012 Ω and resolution of ( 0.01 mV. The correct functioning of this group was periodically checked by measuring potentials using a certified Leeds & Northrup Co. Weston standard cell. The Cl-ISE was connected to a low input (grounded) from the electrometer, because it has a lower impedance than the sodium glass electrode. Depending on the total ionic strength studied, it was observed that, after 40-50 min, the variation of the potential with time was very small. The reading at this moment was considered representative of the cell in equilibrium. 3. Results Ionic mean activity coefficient values for LiCl were determined from the emf measurements of the following galvanic cells without transference:

Na-ISE | LiCl(m), alcohol (Y), H2O (100-Y) | Cl-ISE

(I)

combined with that of the reference

Na-ISE | LiCl(mr), alcohol (Y), H2O (100-Y) | Cl-ISE (II) where m and mr are the molality of LiCl in the working and reference solutions, respectively, and Y is the wt % of the alcohol (methanol or ethanol) in the mixed solvent. The basis for using these ion-selective electrodes and obtaining correct thermodynamic properties is the use of a reference cell in conjunction with the working cell.26-30 The subtraction of the potentials from both cells resulted in the elimination of all of the possible non-thermodynamic potentials present. Thus, applying the equation of Nernst-Nikolsky to cells I and II and subtracting, the following expression is obtained:

∆E ) EI - EII ) -2k log

mγ mrγr

(1)

E is the emf of the cell, k ) (ln10)RT/F is the theoretical Nernst slope, and γ and γr are the mean ionic activity coefficients of LiCl at molality m and mr, respectively. The election of mr is arbitrary, and it does not have an influence on

the results. The summarize, although we are using a nonthermodynamic cell, the values of the potentials obtained, and thus the activity coefficients, are indeed thermodynamic. In no instance have we made extra-thermodynamic assumptions. In some of our previous articles,18-20,31 we have disregarded these potentials non-thermodynamic, and the results have also been good, and thus its effect on the thermodynamic magnitudes obtained does not appear to be very large. In Tables 1 and 2, the E values as a function of the LiCl molalities are shown for each alcohol-water mixture. Because the activity coefficients of the LiCl in water32 and in watermethanol mixtures25 are precisely known, the values of E shown in Table 1 were used to check and calibrate the electrode system used. As can be observed from Table 1, a very good linear relationship (r > 0.9999) is obtained when plotting ∆E versus log(mγ/mrγr). The mean value of the slopes is k ) 54.16 ( 0.16, with the standard deviations of fit being between 0.32 and 0.55 mV. This sub-Nerstian value of k is typical when working with ion-selective electrodes. In this calculation, it was assumed that kLi = kCl = k = (kLi + kCl)/2. This was previously shown when checking the response and reproducibility of each of the electrodes following the instructions of the manufacturer. To use eq 1 to calculate the activity coefficients of the LiCl and the water-ethanol mixtures, it is first necessary to know the values of γr. This is a very important point because it affects the accuracy of the activity coefficients and the other thermodynamic functions subsequently calculated. The determination of γr was carried out following the habitual method of Hitchcock33 and using the extended equation of Debye-Hu¨ckel,1,2 as well as with the Pitzer7,8 equation. For a 1:1 electrolyte, the Debye-Hu¨ckel extended equation can be written as,1,2

Axm + cm - log(1 + 0.002mM) + Ext (2) log γ ) 1 + Baxm where a represented the ion size parameters, c the ion-interaction parameter, M the average molecular mass of the mixed solvent,

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Table 2. Experimental Electromotive Force and Activity Coefficients at Different LiCl in Ethanol-Water Mixtures at 298.15 Ka 20% ethanol m 0.0716 0.1144 0.1904 0.3075 0.3586 0.5011b 0.5445 0.6340 0.7786 0.8756 1.0114 1.1628 1.4006 1.7458 2.4732 3.2385 4.4271

40% ethanol

60% ethanol

80% ethanol

E

γ

m

E

γ

m

E

γ

m

E

γ

-30.78 -51.36 -74.01 -95.63 -102.92 -118.86 -122.81 -130.35 -140.82 -146.92 -154.52 -162.11 -172.74 -186.00 -209.45 -230.38 -258.90

0.785 0.761 0.741 0.726 0.727 0.730 0.731 0.737 0.749 0.759 0.772 0.789 0.821 0.873 1.015 1.209 1.622

0.0367 0.1437 0.2356 0.3807 0.4726 0.5220b 0.6810 0.8374 1.1069 1.3798 1.8718 2.3792 3.2355 4.3195 4.5313

-35.10 -94.27 -116.01 -137.60 -147.62 -152.25 -165.10 -175.38 -190.21 -202.47 -221.33 -237.92 -262.26 -290.06 -295.55

0.768 0.691 0.669 0.655 0.653 0.652 0.657 0.665 0.689 0.717 0.790 0.884 1.090 1.475 1.580

0.0276 0.2261 0.3660 0.4941 0.5881b 0.6882 0.8255 0.9171 0.9904 1.1744 1.2584 1.5134 1.6578 2.1494 2.5207 2.9183 3.5630

-61.17 -146.19 -166.73 -180.25 -188.24 -195.77 -205.03 -210.54 -214.57 -224.05 -227.98 -239.50 -245.52 -263.64 -277.61 -290.85 -311.27

0.711 0.529 0.506 0.499 0.497 0.498 0.506 0.512 0.517 0.533 0.541 0.574 0.596 0.675 0.775 0.887 1.121

0.0775 0.2472 0.3001 0.3863 0.4822 0.6012b 0.8576 0.9716 1.1766 1.4465 1.8100 2.1563 2.7951 3.3546 3.4406

-158.71 -201.31 -208.40 -218.56 -228.15 -238.11 -255.59 -262.34 -273.03 -286.03 -301.94 -315.94 -339.87 -359.35 -362.43

0.479 0.372 0.356 0.343 0.337 0.334 0.340 0.346 0.359 0.384 0.431 0.487 0.625 0.788 0.820

γr ) a/Å ) c/kg mol-1 ) σ/mV )

0.732 5.6 0.1217 0.06

0.656 5.8 0.1180 0.20

0.498 4.1 0.1734 0.09

0.334 3.3 0.2514 0.15

γr ) β0/kg mol-1 ) β1/kg mol-1 ) Cγ/kg2 mol-2 ) σ/mV )

0.729 0.1654 0.6220 0.0004 0.07

0.649 0.1790 0.8911 0.0051 0.19

0.497 0.2617 0.9184 0.0042 0.21

0.335 0.3992 1.1687 -0.0176 0.18

〈γ〉r )

0.730 ( 0.001

0.652 ( 0.003

0.497 ( 0.001

0.334 ( 0.001

a

Units, m in mol kg-1; E in mV. b Reference molality.

and Ext the contribution of the extended terms. A and B represent the Debye-Hu¨ckel constant given by,

A ) 1.8247106d1/2/(rT)3/2kg1/2 mol-1/2

(3)

B ) 50.2901d1/2/(rT)1/2kg1/2 mol-1/2 Å-1

(4)

where d stands for the density, r the relative permittivity (static dielectric constant) of the solvent, and T the temperature. Density and permittivity values were taken from the literature.31 By combining eq 1 and eq 2, the values of a, c, and γr can be optimized for each ethanol-water mixture studied. These values as well as the standard deviation of the fit are shown in Table 2. Consideration or not of the extended terms contribution (Ext) had no significant effect on the fitting. Another way to evaluate the γr values is using the Pitzer equation. Thus, for the 1:1 electrolyte, the Pitzer equation can be written as,7,8

lnγ ) fγ + Bγm + Cγm2

(5)

where

fγ ) -Aφ

[

xm 2 + ln(1 + bxm) b 1 + bxm

]

(6)

Bγ ) 2β0 + 2β1 [1 - (1 + Rxm - R2m/2) exp(-Rxm)] kg mol-1 (7) 2 Rm In these equations, R and b are assumed fixed parameters with values of 2.0 and 1.2 kg1/2mol-1/2, respectively; β0, β1, and Cγ are solute-specific interaction parameters; and Aφ is the Debye-Hu¨ckel constant for the osmotic coefficients defined

by,

Aφ ) 1.4006106d1/2/(rT)3/2kg1/2mol-1/2

(8)

with all symbols having the usual meaning. By combining eq 1 and eqs 5-7, the γr , β0, β1, and Cγ values can be optimized for each ethanol-water mixture studied. These values are summarized in Table 2 together with the standard deviation of the fit. 4. Discussion As can be observed in Table 2, the values of γR obtained with both models are in good agreement. The standard deviations of fit are also comparable, although we must take into account that the Debye-Hu¨ckel model is valid only to about 1 molal, whereas that of the Pitzer model is applicable to a complete range of concentrations. The values of a are within the expected order of magnitude. In all cases, a is greater than the sum of the crystallographic radii (2.76 Å),34 which can be explained by the solvation of the ions. For 20 and 40% ethanol, a > q, (q is the Bjerrum interionic distance35), whereas for 60 and 80% a < q, which suggests that there may be ionic association between the LiCl at these latter percentages.18,20 In Figure 1, the obtained c, β0, β1, and Cγ values of the fits are plotted against the reciprocal of relative permittivity for the different ethanol-water mixtures. It is interesting to note that, both the ionic-interaction parameter of the Debye-Hu¨ckel equation, c, as the parameters of the Pitzer equation, β0 (which can be identified with interactions of like and unlike charged ions) and β1 (which can be identified with the interactions between unlike-charged ions) change linearly with 1/r . This is the typical profile observed for 1:1 electrolytes in mixed solvents.7,8,18,20,36-38 Although in the Pitzer thermodynamic treatment there is no explicit dependence of β0 on the relative permittivity of the solvents, it was found38 that the radial

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Figure 1. Plot of ionic-interaction parameters, c (O), β0 (0), β1 (∆), and C γ (3) for LiCl in ethanol-water mixtures at 298.15 K as a function of the inverse of the relative permittivity. (9, 2, 1): Values in water obtained from Pitzer.7

Figure 2. Plot of log γ vs m1/2 for LiCl in alcohol-water at 298.15 K 0% of alcohol,32 100% of methanol,40 and 100% of ethanol.41

distribution function at hard-core contact (g+-, g++, and g-) was a function of the relative permittivity, which may be the cause of the variation of β0. On the other hand,7,8,38 in the discussion of hard-core effects on osmotic and activity coefficients in terms of the Debye-Hu¨ckel theory, Pitzer has pointed out that β1 would be a function of κ2 (κ being the reciprocal of average radius of ionic atmosphere) that involves the reciprocal of relative permittivity. Thus, the linear dependence of β1 on 1/r observed in this study is reasonable. In relation to the parameter C γ (which represents triple ionic interactions), it is not a simple matter to reach conclusions concerning its variation with 1/r. It is also interesting to note that the Debye-Hu¨ckel ioninteraction parameter c is numerically close to β0 for each ethanol-water mixture such as it occurs in other similar systems.18-20,29-31,39 Substituting the average values for γR, shown in the last row of Table 2, in eq 1 with k ) 54.16 mV, the γ were calculated and appear in this same table for each molality of LiCl and percentage of ethanol. Figure 2 shows log γ versus m1/2 for each water-ethanol mixture studied. The data of Mussini et al.25 are presented for comparative purposes, showing the data on LiCl in watermethanol mixtures. In both cases, the values for the activity coefficient of LiCl in water were taken from Hamer and Wu.32 The corresponding values for pure methanol and ethanol were taken from Safarov,40,41 who obtained them on the basis of vapor pressure measurements.

Figure 3. Plot of log γ vs m1/2 for LiCl, NaCl, and KCl in 40% alcoholwater mixtures at 298.15 K. Methanol-water, (b) Mussini et al.;25 (9) Basili et al.;43 (2) Basili et al.44 ethanol-water, (O) this work; (0) Esteso et al.;45 (2) Mussini et al.45

All the curves show a characteristic trend of the variation of log γ with concentration, which, as is well-known, is governed by two types of short-range interactions,1,2,7,8 including those of ion-ion (ion-pair) and ion-solvent (solvation). A decrease in the log γ of LiCl with a decrease in the static dielectric constant of the solvent mixed (increasing in % of alcohol) is clear, and it must be related to the ion-solvent interactions. In the system containing ethanol, this effect is more pronounced, and, furthermore, a greater shift is observed in the typical minimum toward lower values of m. Figure 3 shows a comparative plot of log γ versus m1/2 for LiCl, NaCl, and KCl in mixtures containing 40% alcohol. The black points represent the measurements made by Mussini et al.,25,42-44 using amalgam electrodes, whereas the white points represent the measurements made in the present study for LiCl. For NaCl, we have used data from a previous study made by our group,45 in which ion-selective electrodes were used for both ions. Although the working ranges were not the same, a clear sequence was observed in the values of log γ: LiCl > NaCl > KCl. Consideration was not made in this analysis for the values corresponding to RbCl and CsCl, because there is a great discrepancy between the values reported by Mussini et al.,42,44 using amalgams, and those of Hu et al.,46,47 using PVC membrane electrodes, which they had fabricated. The standard free energy of transference, ∆Gt0 (or medium effect), is probably one of the magnitudes most commonly used for the understanding of differences in behavior between a dissolved solute in a pure solvent and that of a mixture of solvents. As is well known, this parameter is defined48 as the difference between the standard free energies per mole of electrolyte in a pure solvent (usually water) and another pure or mixed solvent and represents the change in solvation at infinite dilution when the solute is transferred from one solvent to another. Considering the ions to be rigid spheres with radius r+ and r-, in a uniform dielectric medium, ∆Gt0 can be calculated by using the well-known expression of Born,49

∆Gt0 ) -nF(Es0 - Ew0) ) -

(

)(

Ne2 1 1 1 1 + 2 s w r+ r-

)

(9)

where s and w stand for solvent mixture and pure water, respectively, and all of the other symbols have their usual meaning. To experimentally calculate ∆Gt0, it is necessary to know the values of the standard potential of the cell for each percentage studied. These may be obtained beginning with values for E, m, and γ as listed in Table 2, and applying the equation of Nernst-Nikolsky to the cells (I).

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E ) E0* - 2k log mγ

(10)

It should be recalled that E0* includes the asymmetry potential of the electrodes, as asym E0* ) E0 + Easym ) E0 + (asym Na + Cl )

(11)

Because Easym tends to be a small and independent value for the composition of the solvent, we may suppose21,29-31,46,47,50 that (Es0* - Ew0*) = (Es0 - Ew0) ) ∆E0. Prior to the calculation of ∆Gt0, the Born condition for the functional dependency of ∆E0 with 1/r was checked (Figure 4). A notable deviation from the predicted linear behavior was observed above approximately 30% ethanol, which agrees with results reported by various authors.4,25,51,52 In contrast, in some of our previous studies18-20,29,30 this linear dependence was indeed observed, at least to the middle percentages of cosolvent (40-60%), obtaining correlation indexes of above 0.999 and deviations lower than 1-2 mV. It seems, therefore, that the linear dependence depends on many factors including the nature of the solvents and characteristics of the medium, properties of the electrolyte, the electrodes used, precision of the measurements, and other factors. Once ∆E0 has been evaluated, the values for ∆Gt0 are then calculated. These values appear in Table 3 for the methanolwater and ethanol-water together with those obtained from the literature for the first of these systems.21,25,53,54 It can be seen that the agreement is very good, with standard errors of 0.045 and 0.041 kJ mol-1, for 20 and 40% methanol, respectively. This is equivalent to 0.47 and 0.42 mV, that is, approximately half of that suggested by Covington and Thain21 for these types of determinations. Figure 5 shows a plot of standard free energy of transference, ∆Gt0, against the percent alcohol in the mixture. For comparison we have included the corresponding values reported by Mussini et al.25 for the methanol-water system. As can be seen, these values are in very good agreement with the values obtained here. It is remarkable in all cases that the values of ∆Gt0 become more positive with increasing percentages of alcohol, which shows that the transference of LiCl from water to the wateralcohol mixtures studied here is a not a spontaneous process. On the other hand, and as occurs with other electrolytes, the values of ∆Gt0 are greater in the water-ethanol system than in the water-methanol system.31 Obviously, this is due to the properties of the solvent (dielectric constant, polarizability, ability donation electron, etc.), and this would indicate that the LiCl could be somewhat more solvated in the water-methanol system than in the water-ethanol system. To prove the preceding supposition, we can return to the correlation that establishes the dependence of standard electromotive force, expressed in molar concentration scale, Ec0, on the volume fraction of the water in the solvent mixture, φw , according to25,43,44,51,52

∆Ec0 ) Ecs0 - Ecw0 ) nhydr(RT/F)lnφw

(12)

Ec0 ) Em0 + 2k log d

(13)

φw ) (ww/dw)/(ww/dw + wethanol/dethanol)

(14)

where

Em0 is the standard electromotive force in molal scale (up to now we have denoted it E0 for simplicity), nhydr indicates the primary hydration number of the electrolyte, and all of the other

Figure 4. Test of the Born equation for the variation of ∆E0 against the inverse of the relative dielectric constant. Table 3. Values of E0*, ∆E0, and ∆Gt0 for the Methanol-Water and Ethanol-Water Mixtures at 298.15 K %

E0*/mV

〈∆Gt0〉/ standard standard ∆Gt0/ ∆E0/mV kJ mol-1 kJ mol-1 error/kJ mol-1 error/mV

0 -129.68 20 -154.26

0.00 -24.58

40 -180.56

-50.88

0 20 40 60 80 d

-129.68 0.00 -166.17 -36.49 -202.95 -73.27 -246.10 -116.42 -313.63 -183.95

Methanol-water 0.00a 〈0.00〉 〈2.38〉 2.37a b 2.44 2.44c 2.45d 2.21e 4.91a 〈4.92〉 4.97b 4.97c 4.97d 4.76e

0.00 0.045

0.00 0.47

0.041

0.42

Ethanol-water 0.00a 3.52a 7.07a 11.23a 17.75a

a This work. b Feakins and Voice.53 Popovych.54 e Mussini et al.25

c

Covington and Thain.21

Figure 5. Plot of the standard free energy of transference ∆Gt0 vs percent of alcohol for LiCl at 298.15 K, (b) Mussini et al.25

symbols have their usual meaning. Figure 6 is a plot of ∆Ec0 versus (RT/F) lnφw, showing a very good linear relation up to 60%, with a value of nhydr obtained for LiCl of 4.4. A similar analysis carried out for the LiCl in the methanol-water system by Mussini et al.25 gave a value of 5.1, which would confirm our hypothesis. Nevertheless, and as affirmed by Mussini et al.,25 these values are very low compared to those obtained by other methods, as well as being lower than those reported for NaCl.55 There is a low probability that this occurs because the Li+ is smaller than the Na+ and therefore has a larger surface charge density and should thus be more hydrated. The explanation given by Mussini et al.25 is that Feakins and French’s theory51,52 might not give a satisfactory description for some ions, having large hydration as the Li+ ion, when transferred from pure water to a mixed solvent.

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Figure 6. Plot of ∆E0c vs a function of the water volume fraction for ethanol-water mixtures at 298.15 K.

5. Conclusions This work provides reliable activity coefficient data for LiCl in ethanol-water mixtures at 298.15 K. The results were obtained using a galvanic cell without transport, containing a Cl- selective electrode, and an electrode selective for Na+ (at present no electrodes selective for Li+ are commercially available). The correct functioning of this arrangement was confirmed by measuring the activity coefficients of LiCl in water-methanol mixtures and comparing the results with those previously obtained by Mussini et al.25 A good Nerstian response was shown by the electrode system employed. As a second form of verification, the standard free energy of transference, ∆Gt0, was determined for the transference of the LiCl from water to water-methanol mixtures. The values agreed very well with those from the literature, which shows us the suitability of this arrangement for obtaining thermodynamic magnitudes, which would be very useful in various fields of applied chemistry. Acknowledgment The authors thank CONICYT-Chile for financing this research through Fondecyt Project No. 1040299, and likewise to the ORI (Oficina Relaciones Internacionales) of La Laguna University, which made possible the stay of H. R. G. Literature Cited (1) Harned, H. S.; Owen, B. B. The Physical Chemistry of Electrolytic Solutions; Reinhold: New York, 1958. (2) Robinson, R. A.; Stokes, R. H. Electrolyte Solutions; Dover Publications, Inc.: Mineola, New York, 2002. (3) Harned, H. S.; Robinson R. A. Multicomponent Electrolyte Solutions; Pergamon Press: London, 1968. (4) Hills, G. J. Reference Electrodes in Nonaqueous Solution. In Reference Electrodes. Theory and Practice; Ives, D. J. G., Janz, G. J., Eds.; Academic Press: New York, 1961; pp 433-463. (5) Butler, J. N. Experimental Methods: Potentiometric. In ActiVity Coefficients in Electrolyte Solutions; Pytkowicz, R. M., Ed.; CRC Press: Boca Raton, Florida, 1979; pp 81-94. (6) Horvath, A. L. Handbook of Aqueous Electrolyte Solutions; Ellis Horwood Limited: Chichester, U.K., 1985. (7) Pitzer, K. S. Ion Interaction Approach: Theory and Data Correlation. In ActiVity Coefficients in Electrolyte Solutions; Pitzer, K. S., Ed.; CRC Press: Boca Raton, Florida, 1991; pp 75-153. (8) Pitzer, K. S. Thermodynamics; McGraw-Hill, Inc: New York, 1995. (9) Lobo, V. M. M. Handbook of Electrolyte Solutions, Elsevier: Amsterdam, The Netherlands, 1989. (10) Lanier, R. D. The Determination of Activity Coefficients of Sodium Chloride in Aqueous Three-Component Solutions by Cation-Sensitive Glass Electrodes. J. Phys. Chem. 1965, 69, 3992.

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ReceiVed for reView May 17, 2007 ReVised manuscript receiVed December 14, 2007 Accepted December 19, 2007 IE070704I